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HANDBOOK OF HETEROGENEOUS CATALYSIS Second, Completely Revised and Enlarged Edition Volume 1 Gerhard Ertl, Helmuth Knözinger, Ferdi Schüth, Jens Weitkamp (Editors) Wiley-VCH Verlag GmbH& Co. KGaA, Weinheim, Germany, ISBN: 978-3-527-31241-2, 2008

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Introduction 1.1

Principles of Heterogeneous Catalysis1 James A. Dumesic∗ , George W. Huber and Michel Boudart

1.1.1

Introduction

Heterogeneous catalysis is of vital importance to the world’s economy, allowing us to convert raw materials into valuable chemicals and fuels in an economical, efficient, and environmentally benign manner. For example, heterogeneous catalysts have numerous industrial applications in the chemical, food, pharmaceutical, automobile and petrochemical industries [1–5], and it has been estimated that 90% of all chemical processes use heterogeneous catalysts [6]. Heterogeneous catalysis is also finding new applications in emerging areas such as fuel cells [7–9], green chemistry [10–12], nanotechnology [13], and biorefining/biotechnology [14–18]. Indeed, continued research into heterogeneous catalysis is required to allow us to address increasingly complex environmental and energy issues facing our industrialized society. Discussing the principles of heterogeneous catalysis is difficult, because catalysts are used for a wide range of applications, involving a rich range of surface chemistries. Moreover, the field of heterogeneous catalysis is highly interdisciplinary in nature, requiring the cooperation between chemists and physicists, between surface scientists and reaction engineers, between theorists and experimentalists, between spectroscopists and kineticists, and between materials scientists involved with catalyst synthesis and characterization. Furthermore, industrial catalysts are complex materials, with highly optimized chemical compositions, structures, morphologies, and pellet shapes; moreover, the physical and chemical characteristics of these materials may depend on 1 A list of abbreviations/acronyms used in the text is provided at the end of the chapter. ∗ Corresponding author.

hidden or unknown variables. Accordingly, principles of heterogeneous catalysis are typically formulated from studies of model catalysts in ideal reactors with simplified reactants under mild pressure conditions (e.g., 1 bar), rather than from catalytic performance data obtained with commercial catalysts in complex reactors using mixed feed streams under industrial reaction conditions. The principles derived from these more simplified studies advance the science of heterogeneous catalysis, and they guide the researcher, inventor, and innovator of new catalysts and catalytic processes. 1.1.2

Definitions of Catalysis and Turnover

The definition of a catalyst has been discussed many times [19]. For example, a catalyst is a material that converts reactants into products, through a series of elementary steps, in which the catalyst participates while being regenerated to its original form at the end of each cycle during its lifetime. A catalyst changes the kinetics of the reaction, but does not change the thermodynamics. Another definition is that a catalyst is a substance that transforms reactants into products, through an uninterrupted and repeated cycle of elementary steps in which the catalyst participates while being regenerated to its original form at the end of each cycle during its lifetime [20]. The main advantage of using a heterogeneous catalyst is that, being a solid material, it is easy to separate from the gas and/or liquid reactants and products of the overall catalytic reaction. The heart of a heterogeneous catalyst involves the active sites (or active centers) at the surface of the solid. The catalyst is typically a high-surface area material (e.g., 10–1000 m2 g−1 ), and it is usually desirable to maximize the number of active sites per reactor volume. Identifying the reaction intermediates – and hence the

References see page 14

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2

1.1 Principles of Heterogeneous Catalysis

mechanism – for a heterogeneous catalytic reaction is often difficult, because many of these intermediates are difficult to detect using conventional methods (e.g., gas chromatography or mass spectrometry) because they do not desorb at significant rates from the surface of the catalyst (especially for gas-phase reactions). Heterogeneous catalysts typically contain different types of surface sites, because crystalline solids exhibit crystalline anisotropy. Equilibrated single crystals expose different faces with different atomic structures so as to minimize total surface energy. It would be surprising, in fact, if different crystallographic planes exposing sites with different coordination environments possessed identical properties for chemisorption and catalytic reactions. Moreover, most catalytic solids are polycrystalline. Furthermore, in order to achieve high surface areas, most catalysts contain particles with sizes in the nanometer length scale. The surfaces of these nanoscopic particles contain sites associated with terraces, edges, kinks, and vacancies [21]. If the catalyst contains more than one component (as is generally the case), the surface composition may be different from that of the bulk and differently so for each exposed crystallographic plane. Solids normally contain defects of electronic or atomic nature; in addition, they contain impurities which are either known or unknown in the bulk, but are mostly unknown at the surface. Finally, the surface atomic structure and composition may change with time-onstream as the catalytic reaction proceeds. In short, it is normal to expect that a catalytic surface exposes a variety of surface sites, in contrast to displaying a single type of active site. Indeed, it is so normal today to expect such complexity that it seems surprising that, in 1925, when Taylor formulated his principle of active sites or active centers, the report created so much attention and remains one of the most often cited in heterogeneous catalysis [22]. The relative importance of surface structure – as influenced by crystalline anisotropy, surface defects, and surface composition – underlines the difficulty of identifying the active sites, either simple or complex, that are responsible for turning over the catalytic cycle. The identification and counting of active sites in heterogeneous catalysis became the ‘‘Holy Grail’’ of heterogeneous catalysis in 1925, and the situation remains the same today. The activity of a catalyst is defined by the number of revolutions of the catalytic cycle per unit time, given in units of turnover rate (TOR) or turnover frequency (TOF). In cases where the rate is not uniform within the catalytic reactor or within the catalyst pellets, it is useful to report the rate as a site time yield (STY), defined as the overall rate of the catalytic reaction within the reactor normalized by the total number of active sites within the reactor, again in units of reciprocal time. Catalysis by solid materials

has been observed quantitatively at temperatures as low as 78 K and as high as 1500 K; at pressures between 10−9 and 103 bar; with reactants in the gas phase or in polar or non-polar solvents; with or without assistance of photons, radiation or electron transfer at electrodes; with pure metals as unreactive as gold and as reactive as sodium; with multicomponent and multiphase inorganic compounds and acidic organic polymers; and at STYs as low as 10−5 s−1 (one turnover per day) and as high as 109 s−1 (gas kinetic collision rate at 10 bar). TOFs of commonly used heterogeneous catalysts are commonly on the order of one per second. The life of the catalyst can be defined as the number of turnovers observed before the catalyst ceases to operate at an acceptable rate. Clearly, this number must be larger than unity, otherwise the substance used is not a catalyst but a reagent. Catalyst life can either be short, as in catalytic cracking of oil, or very long, corresponding to as many as 109 turnovers in ammonia synthesis. 1.1.3

Steps in a Heterogeneous Catalytic Reaction

During an overall catalytic reaction, the reactants and products undergo a series of steps over the catalyst, including: 1. Diffusion of the reactants through a boundary layer surrounding the catalyst particle. 2. Intraparticle diffusion of the reactants into the catalyst pores to the active sites. 3. Adsorption of the reactants onto active sites. 4. Surface reactions involving formation or conversion of various adsorbed intermediates, possibly including surface diffusion steps. 5. Desorption of products from catalyst sites. 6. Intraparticle diffusion of the products through the catalyst pores. 7. Diffusion of the products across the boundary layer surrounding the catalyst particle. Accordingly, different regimes of catalytic rate control can exist, including: (i) film diffusion control (Steps 1 and 7); (ii) pore diffusion control (Steps 2 and 6); and (iii) intrinsic reaction kinetics control (Steps 3 to 5) of catalyst performance. In addition to mass transfer effects, heat transfer effects can also occur in heterogeneous catalysis for highly exothermic or endothermic reactions (especially in combustion or steam reforming). Figure 1 shows a general effect of temperature on the reaction rate for a heterogeneous catalyst. At low temperatures, diffusion through the film and pores is fast compared to rates of surface reactions, and the overall reaction rate is controlled by the intrinsic reaction

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1.1.4 Desired Characteristics of a Catalyst

of catalysts, and we refer the reader to other articles for further discussion on transport effects in heterogeneous catalysis [23–31].

Slope = 3-5 kJ/mol Slope = Ea/2R Film diffusion controlled regime

ln (rate)

3

1.1.4 Pore diffusion controlled regime

Slope = Ea/R

Desired Characteristics of a Catalyst

The following list provides several of the key attributes of a good catalyst: Intrinsic regime Increasing Temperature

1/ Temperature General effects of temperature on catalytic activity. The intrinsic activation energy is equal to Ea , and R is the gas constant.

Fig. 1

kinetics. As the temperature is increased, the rates of surface reactions typically increase more rapidly than the rates of diffusion, and the overall rate of the catalytic process becomes controlled by intraparticle diffusion. The apparent activation energy in this regime is equal to the intrinsic activation energy divided by two. As the temperature is increased further, mass transfer through the external boundary layer becomes the controlling step. The onset of diffusion limited regimes can be altered by changing the reactor design, the catalyst pore structure, the catalyst particle size, and the distribution of the active sites in the catalyst particles. Values of various dimensionless groups can be calculated to estimate the extents to which transport phenomena may control catalytic performance for specific operating conditions [23–31]; however, these calculations are most reliable for cases where the intrinsic reaction kinetics are known. In these cases, it is possible to make catalysts with structures designed to provide adequate rates of diffusion and yet offering high surfaces areas, leading to high rates of reaction per reactor volume, such as the design of specific pore size distributions (e.g., bimodal distributions containing large pores leading to high accessibility of the active sites within the interior of the catalytic pellet, and small pores that branch from the larger pores leading to high surface areas), the formulation of unique pellet shapes (that lead to high accessibility of the active sites but do not cause large pressure drops through the catalytic reactor), and the synthesis of catalyst pellets containing a spatial distribution of the active material within the catalyst pellet [32]. In some cases, transport effects can be used to improve the selectivity of a catalyst, such as in the case of shape-selective catalysis in zeolites [33–36]. In the following sections, we focus on various factors controlling the intrinsic reaction kinetics

• The catalyst should exhibit good selectivity for production of the desired products and minimal production of undesirable byproducts. • The catalyst should achieve adequate rates of reaction at the desired reaction conditions of the process (remembering that achieving good selectivity is usually more important than achieving high catalytic activity). • The catalyst should show stable performance at reaction conditions for long periods of time, or it should be possible to regenerate good catalyst performance by appropriate treatment of the deactivated catalyst after short periods. • The catalyst should have good accessibility of reactants and products to the active sites such that high rates can be achieved per reactor volume. The first three key attributes of a good catalyst are influenced primarily by the interactions of the catalyst surface with the reactants, products, and intermediates of the catalytic process. In addition, other species may form on the catalyst surface (e.g., hydrogen-deficient carbonaceous deposits denoted as coke) that are not directly part of the reaction scheme (or mechanism) for the overall catalytic process. The principle of Sabatier states that a good heterogeneous catalyst is a material that exhibits an intermediate strength of interaction with the reactants, products, and intermediates of the catalytic process [37, 38]. Interactions of the catalyst surface with the various adsorbed species of the reaction mechanism that are too weak lead to high activation energies for surface reactions and thus low catalytic activity, whereas interactions of the catalyst with adsorbed species that are too strong lead to excessive blocking of surface sites by these adsorbed species, again leading to low catalytic activity. The principle of Sabatier is elegant in its simplicity and generality, but it is deceptively difficult to use in practice. In particular, this principle applies to a catalyst in its working state, and the nature of the catalyst surface can be expected to be dependent on the nature of the catalytic reaction conditions. For example, one may begin the catalytic reaction with the heterogeneous catalyst References see page 14

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4

1.1 Principles of Heterogeneous Catalysis

in a given oxidation state (e.g., containing zero-valent metal particles following treatment of the catalyst in H2 at elevated temperature); however, the nature of the surface can be changed dramatically upon interaction with strongly adsorbed species, such as the formation of carbonaceous deposits (coke), and formation of oxides, carbides, nitrides, or sulfides upon interaction with O, C, N, or S species, respectively [39–44]. In this case, the interactions of these oxide, carbide, nitride, or sulfide surfaces with the adsorbed species enter into the reaction mechanism. Of even greater complexity is the fact that a variety of different types of sites are typically present on a catalyst surface (e.g., sites having different coordination and/or chemical composition), and a majority of the observed catalytic activity may be caused by the contributions from a small fraction of the sites present on the catalyst surface. In this case, the adsorbed species interact with these special surface sites (e.g., steps and defect sites on a metal nanoparticle, or sites present at the metal–support interface of a supported metal catalyst). Another factor that complicates catalyst design is that the strengths of interaction of the surface with adsorbed species typically depend on the surface coverages by adsorbed species. For example, the interaction of a transition metal surface with adsorbed CO may be very strong at low surface coverages (e.g., binding energy of nearly 200 kJ mol−1 ), suggesting that these surfaces would be completely covered and thus poisoned by adsorbed CO at moderate pressures and temperatures; however, these surfaces may carry out catalytic reactions in the presence of gaseous CO at these pressures and temperatures because the differential heat of CO adsorption decreases significantly (e.g., to binding energies near 100 kJ mol−1 ) as the surface coverage by adsorbed CO increases [45, 46]. Accordingly, there is a relationship between activity and the interaction of the surface with adsorbed species at the surface coverage regime appropriate for the catalytic reaction conditions. The aforementioned complications caused by the presence of different types of sites on the surface, and the effects on the surface binding energies caused by changes in surface coverages, clearly make it difficult to interpret the performance of a heterogeneous catalyst in quantitative detail. Tools are certainly available to address these complications, such as kinetic Monte Carlo simulations combined with results from density functional theory (DFT) calculations [47–50]. Yet, from a different point of view, the presence of different types of sites and the effects of surface coverage may well contribute to the robustness of the heterogeneous catalyst for operation over a wide range of reaction conditions. In general, the presence of different types of sites and the effects of surface coverage both contribute to surface non-uniformity (different types

of sites producing a prior non-uniformity, and effects of surface coverage causing induced non-uniformity). At a selected set of reaction conditions, an optimal set of surface binding energies exists that satisfy the principle of Sabatier (as discussed below). Accordingly, the performance of a heterogeneous catalyst with a non-uniform surface will be dominated by the subset of the sites having surface binding energies closest to the optimal values. At higher temperatures, other sites having stronger binding energies with adsorbed species will become the dominant contributors to the observed catalytic activity, whereas sites having weaker binding energies with adsorbed species will control catalytic activity at lower temperatures. Thus, while the effects of surface non-uniformity make it more difficult to predict the performance of a heterogeneous catalyst from a molecular-level understanding, these effects may serve to broaden the range of reaction conditions over which the catalyst can operate effectively. In this respect, our desire to design catalysts having very high selectivity is guided by the synthesis of uniform catalysts, where each site has the optimal properties for production of the desired reaction product. This strategy leads to the idea of highly selective, single-site catalysis as discussed by Thomas et al. [51]. In contrast, the design of catalysts that operate over a wide range of reaction conditions is guided by the synthesis of non-uniform catalysts, such that different subsets of sites control catalyst performance at different reaction conditions. The disadvantage of using non-uniform catalysts, however, is that different sites may display different selectivities for the production of various products, and control over catalytic selectivity may thus be limited [22]. 1.1.5

Reaction Schemes and Adsorbed Species

We now explore further the principle of Sabatier using a specific example: water-gas shift over a metal catalyst −−− −− → (e.g., Cu). This reaction (CO + H2 O ← − CO2 + H2 ) is of importance for the production of H2 from steamreforming of fossil fuels, and for controlling the CO : H2 ratio in synthesis gas mixtures used in methanol and Fischer–Tropsch synthesis processes. For this example, we consider the reaction scheme shown in Fig. 2, where * represents a surface site. The stoichiometric numbers, σi,1 and σi,2 , indicate the number of times that step i occurs to give the overall reaction for reaction schemes 1 and 2, respectively. In this sequence of steps, the water-gas shift reaction can take place via the formation of carboxyl species (COOH) or through the formation of formate species (HCOO) [52]. In the absence of a catalyst, the rate of watergas shift via this mechanism is negligible, because the

ch1.1

1.1.5 Reaction Schemes and Adsorbed Species

si,1

si,2

1.

CO + ∗

CO∗

1

1

2.

H2O +∗

H2O∗

1

1

3.

H2O ∗ + ∗

OH ∗+H ∗

2

1

4.

CO ∗ + OH ∗

COOH ∗+ ∗

1

0

5.

COOH ∗ + OH ∗

CO2∗ + H2O ∗

1

0

6.

CO2∗

CO2 + ∗

1

1

7.

2H ∗

H2 + 2∗

1

1

8.

CO ∗ + OH ∗

HCOO ∗∗

0

1

9.

HCOO∗∗

CO2∗ + H ∗

0

1

overall

CO + H2O

CO2 + H2

Assumed reaction mechanism for water-gas shift reaction over Cu. Adapted from Ref. [52].

Fig. 2

reaction intermediates (e.g., OH∗ , H∗ , COOH∗ , HCOO∗ ) are at very low concentrations in the gas phase. For example, the enthalpy change for step 3 in the gas phase is approximately 500 kJ mol−1 . However, adsorption of the reaction intermediates onto the catalyst surface allows these steps to take place with small enthalpy changes. In the case of copper, the binding energies of H and OH are approximately 250 and 280 kJ mol−1 on Cu(111), such that the enthalpy change for step 3 on the catalyst surface is now slightly exothermic. According to the principle of Sabatier, a good catalyst is a material that adsorbs reaction intermediates with intermediate strength. However, we now must distinguish between reactive intermediates and spectator species on the catalyst surface. In the above reaction scheme, we see that the watergas shift reaction can take place through adsorbed carboxyl species or formate species. Results from DFT calculations indicate that adsorbed formate species have lower energy compared to adsorbed carboxyl species on copper surfaces, suggesting that path 2 for water-gas shift (σi,2 ) would be favored versus path 1 (σi,1 ) based on thermodynamic arguments. However, the activation energy for step 4 is considerably lower than that for step 8, and the primary path for water-gas shift over copper involves the formation and subsequent reaction of adsorbed carboxyl species. Accordingly, the most stable species on the catalyst surface are not necessarily the most reactive species. This idea leads us to distinguish between a most abundant surface intermediate (MASI) and a most abundant reactive intermediate (MARI). In certain cases, the MASI and the MARI may be the same species, but in other cases (such as in this case of water-gas shift on copper), the MASI is a spectator species that

5

does not participate in the overall reaction. In this latter case, the spectator species inhibits the overall reaction by blocking surface sites, and it serves no useful role in the overall reaction scheme. For purposes of elucidating catalytic reaction schemes it is essential to distinguish between reactive intermediates and spectator species. This distinction is of paramount importance in spectroscopic studies of adsorbed species on catalyst surfaces, where the detection of a specific adsorbed species using a spectroscopic method (e.g., the detection by infra-red (IR) studies of adsorbed ethylidyne species on platinum surfaces during ethylene hydrogenation [53]) does not guarantee that this species is a reactive intermediate. Instead, these spectroscopic studies must be conducted under dynamic conditions (e.g., so-called operando measurements, where spectroscopic and reaction kinetics data are collected simultaneously) to determine that the time constant for the formation or disappearance of the surface species is the same as the time constant for the overall catalytic reaction [54, 55]. The overall catalytic reaction is given by a linear combination of elementary steps, and the enthalpy change for the overall reaction, H , is given by:
σi Hi (1) H = i

where Hi are the enthalpy changes for elementary steps i. From the principle of Sabatier, it is now clear that the overall value of H should be composed of approximately equal contributions from each of the values of Hi , giving rise to a relatively flat potential energy diagram of energy versus reaction coordinate in moving from reactants, through adsorbed intermediates, to products. Specifically, any value of Hj that is very negative must be balanced by a value of Hk that is very positive, such that the surface will become highly covered (and poisoned) by the adsorbed species produced in step j , and the activation energy for step k will be high. Both of these effects lead to low catalytic activity. We note that the reaction mechanism can certainly contain steps with positive values of Hi , because the intermediates produced in such a step can be consumed by following steps having negative values of Hi . This situation is termed ‘‘kinetic coupling’’, where the conversion of an unfavorable step is increased by its combination with a favorable step that consumes the unstable products of the first step. The highest value of Hi for a surface reaction that can be tolerated can be estimated from transition state theory. The value of Hi,max depends on the overall rate of the reaction (TOF), the surface coverage by the surface species that reacts in this step (θA ), a frequency References see page 14

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6

1.1 Principles of Heterogeneous Catalysis

factor ν (of the order of 1013 s−1 ), and the temperature T , as given by:   Hi,max θA (2) TOF = ν exp − RT For a reaction operating at 500 K with a TOF of 1 s−1 , the maximum value of Hi that can be tolerated for a species with high surface coverage (θA approaching unity) corresponds to 125 kJ mol−1 , which is still a rather high value. In practice, the highest value of Hi that could be tolerated would be lower than this value of 125 kJ mol−1 , because the surface coverage by the reactive intermediate would typically be lower than unity and the above analysis assumes that the activation energy for the reverse of step i (i.e., the exothermic direction for this step) is equal to zero. This situation where the overall enthalpy change is shared fairly equally between the various steps of the reaction scheme is a necessary condition for high catalytic activity, but it is not a sufficient condition, because we have not yet considered the transition states for the various elementary steps. The aforementioned reaction scheme for water-gas shift involving the formation of carboxyl species contains seven steps, thereby requiring the determination (or estimation) of 13 rate constants to describe the reaction kinetics completely; that is, a forward and reverse rate constant for each step (kfor,i and krev,i ) constrained by the relationship that these rate constants must give the proper equilibrium constant for the overall reaction, Keq , as given below:   kfor,i σi = Keq (3) krev,i i

However, it is a rare case that all of these rate constants are kinetically significant. Thus, while we generally have the desire to know the values for as many rate constants as possible, we typically need to know only the values of a limited number of these rate constants to describe the performance of the catalyst

2.

CO + ∗ H O +∗

3.

H2O ∗+ ∗

1.

4. 5. 6. 7.

Fig. 3

2

CO ∗+OH ∗ COOH ∗+OH ∗ CO2∗ 2H ∗

for the reaction conditions of interest. Unfortunately, at the outset of research on a given catalyst process, we usually do not know which rate constants will be kinetically significant. Accordingly, an important objective of research into a given catalytic process is to identify which steps are kinetically significant, such that further research can focus on altering the nature of the catalyst and the reaction conditions to enhance the rates of these kinetically controlling steps. This situation is illustrated in Fig. 3 for the above case of water-gas shift involving carboxyl species, according to which the rate is controlled by steps 3 and 4, whereas steps 1, 2, 5, 6, and 7 are quasi-equilibrated. The net rate of step 3 in Fig. 3, is twofold faster than the net rates of all other steps, because the stoichiometric number of step 3 is equal to 2 whereas all other stoichiometric numbers are equal to 1. Importantly, the net rate of each step divided by its stoichiometric number is equal to the net rate of the overall catalytic reaction. This equality is due to the principle of kinetic steady state, as stated by Bodenstein (see Ref. [38]), according to which determining the rate of one single reaction (typically the overall reaction) allows one to calculate the net rates of all the other individual reactions. The Bodenstein principle is an important foundation of our thinking about how catalytic cycles turn over. This principle also shows that the notion of a ‘‘slow’’ step in a catalytic cycle at the steady state is a misnomer, because all steps proceed at the same net rate. 1.1.6

Conditions for Catalyst Optimality

It can be shown that the net rate of the overall catalytic reaction is controlled by kinetic parameters which depend only on the properties of the transition states for the kinetically significant steps relative to the reactants (and possibly the products) of the overall reaction [56]. The overall rate is also controlled by an additional kinetic parameter for each surface species that is abundant on

CO∗ H O∗ 2

OH ∗ + H ∗ COOH ∗+ ∗

CO2∗+H2O ∗

CO2 + ∗ H + 2∗ 2

Rates of forward and reverse steps in the water-gas shift reaction on Cu.

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1.1.6 Conditions for Catalyst Optimality

the catalyst surface. Specifically, the net rate of the overall reaction is determined by the kinetic parameters as well as by the fraction of the surface sites, θ ∗ , that is available for the formation of the transition states; the value of θ ∗ is determined by the extent of site blocking by abundant surface species. To illustrate how to determine the optimal activity of a catalyst, we consider an example in which the reaction scheme contains a single rate-controlling step and a single abundant surface species. According to results obtained using DeDonder relations (discussed in Section 5.2.1.10) [56], we may write this reaction scheme in terms of a quasi-equilibrated step involving the transition state for the rate-controlling step, TSi , and a second equilibrated step involving the formation of the most abundant surface species, A∗ , as given below and: −− −− → 1. Reactants + 2∗ − ← − T Si ∗ −− −− → 2. A+∗ − ← −A

(4) (5)

The overall rate of the reaction, rnet , as will be discussed in Section 5.2.1.12, is now given by:   o‡ o‡ S1 H1 ν‡ 1/σ F (ai )θ∗2 (1 − ztot 1 ) exp − rnet = σ1 R RT (6) 

θ∗ = 1 + exp

1 S2o R

−

H2o RT



(7) aA

where F (ai ) is a function of the activities (ai ) of the reactants and/or products of the overall reaction. Neglecting entropy effects, as we change the nature of the catalyst for constant reaction conditions, the primary items in the o‡ above equations that change are H1 and H2o . (Note, we implicitly assume that the reaction mechanism does not change.) Accordingly, the overall rate of the reaction for different catalysts is given by:   o‡ H1 C1 exp − RT (8) rnet =   2 H2o 1 + C2 exp − RT   o‡   S1 ν‡ 1/σ exp (9) C1 = F (ai ) 1 − ztot 1 σ1 R   S2o (10) aA C2 = exp R We next consider that the surface properties of the catalyst are described in terms of some fundamental catalyst property, x. This property x could be a heat

7

of adsorption of one of the reactants [57], the heat of formation of a bulk compound that can be correlated with a heat of adsorption [58], the position of the catalytic element along a horizontal series in the Periodic Table, an electronic property of the catalyst such as Pauling’s d-band character of the metal [59], or the d-band center of the metal [60]. The optimal catalyst can thus characterized by the following relationship:   o‡ o‡ dH1 H1 −C1 exp − RT RT dx drnet =0=   2 o dx H2 1 + C2 exp − RT     o‡ H1 H2o dH2o C2 exp − 2C1 exp − RT RT RT dx + (11)    3 H2o 1 + C2 exp − RT This relationship may be simplified to give:   H2o dH2o 2C2 exp − o‡ dH2o dH1 RT dx   = 2θ =  A H2o dx dx 1 + C2 exp − RT (12) Thus, for the optimal catalyst, the surface coverage by the most abundant surface species is equal to: o‡

dH1 ‡ ω dx = 1 θA = o dH2 2ω2 2 dx

(13)

where the values of ωi are defined as: o‡

dH1 = dx dH2o ω2 = dx

‡ ω1

(14) (15) ‡

In the above derivation, we assume that ω1 and ω2 have the same sign, such that variations in x change the enthalpy of the transition state and the MASI in o‡ the same direction. We also assume that (d 2 H1 )/dx 2 o‡ and (d 2 H2 )/dx 2 are small or zero. This assumption is valid if we are searching for improved catalysts over a small range of x, which typically occurs when testing catalysts. In fact, when we vary x over a large range, then the mechanism of the catalytic reaction would probably change. References see page 14

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1.1 Principles of Heterogeneous Catalysis

Ratenet is maximum at w‡1 = 2qMASIw2

∆H‡tran

∆HMASI

Case 2: w‡1 ≈ w2

∆HMASI

∆H‡tran

x (Fundamental catalyst parameter)

Rate

Case 1: w‡1 >> w2

Intermediate energies

Ratenet is maximum as ∆H‡tran → 0

Rate

Intermediate energies

8

x (Fundamental catalyst parameter)

Case 3: w‡1 0; however, the plot of rate versus x is not symmetric with respect to x. We note that as x

increases for these three cases, the number of vacant sites on the catalyst increases. Cases 2 and 3 clearly illustrate the principle of Sabatier, in which a maximum rate occurs at some moderate level of interaction of the catalyst surface with the intermediates and adsorbed species. While Case 1 appears to contradict ‡ Sabatier’s principle, the situation where ω1  ω2 is highly unlikely. In particular, this situation corresponds to the case where the catalyst interacts more strongly with the transition state than with any of the reactive intermediates. However, if the activation energies for the elementary steps of the mechanism are positive, then the reactants and/or products of the elementary step involving the ratecontrolling transition state are more strongly adsorbed on the surface than is the transition state, leading to the situation described by Case 2 or 3. We may generalize the above expression for catalyst optimality to the case where the surface contains several abundant surfaces species, A∗ , B∗ , C∗ , and D∗ , leading to the following expression: ‡

ω1 = 2(ωA θA + ωB θB + ωC θC + ωD θD )

(16)

In this case, the nature of the optimal catalyst is controlled by the change in the binding energy of the ‡ transition state with respect to x (ω1 ) compared to the changes in the bindings energies of species A, B, C, and

ch1.1

1.1.6 Conditions for Catalyst Optimality

D (ωA , ωB , ωC , ωD ) weighed by their respective surface coverages at the steady state. A bridge between the thermodynamics and kinetics of a reaction is provided by the Brønsted–Evans–Polanyi–Semenov relationship, which states that there is a linear relationship between the activation energy Eact of an elementary step and the heat of reaction if entropy effects are neglected [38]: Eact = E0 + αH

(17)

where H is the enthalpy of reaction, α is the transfer coefficient that varies between zero and one, and E0 is a constant. In other words, if we neglect entropic effects, the activation energy of an elementary step in the exothermic direction is lower when the heat of reaction becomes more favorable (i.e., H becomes more negative). DFT calculations have recently shown that Brønsted–Evans–Polanyi–Semenov relationships are generally upheld in chemical reactions on catalyst surfaces [61–64]. We now consider the following catalytic reaction: A∗ −−−→ B∗ + C∗

(18)

If we use the gas phase A species as the zero energy level (as shown in Fig. 5) we can define the activation energy as: Eact = E0 + α(HBgas + B.E.B + HCgas + B.E.C − HAgas − B.E.A ) = E0 + α(B.E.B + B.E.C − B.E.A ) + αHgas

(19)

where the B.E. terms are the binding energies of the various species on the surface. Agas → Bgas + Cgas

Assume:

dB.E.A dB.E.B dB.E.C = = dx dx dx

Agas

Bgas + Cgas

∆H‡

0 a=0

Eact

B.E.A A* A*

a=1 B* + C*

B* + C*

Schematic potential energy diagram of energy versus reaction coordinate, showing the relationship between the energy of transition state, H= , and changes in the energies of adsorbed species.

Fig. 5

9

We now define the standard enthalpy for the formation of the transition state from gaseous species A as: o‡

H1 = Eact + B.E.A = E0 + α(Hgas ) + α(B.E.B + B.E.C ) + (1 − α)B.E.A

(20)

and we differentiate the enthalpy of the transition state with respect to x, leading to the following result: 

o‡

dH1 dx

=α

dB.E.B dB.E.C + dx dx

 + (1 − α)

dB.E.A dx (21)

This relationship shows how the change of the ‡ transition state enthalpy with respect to x(ω1 ) is related to the changes in the binding energies B.E. of the adsorbed species (ω2 ). If α is equal to zero, then the change of the transition state enthalpy depends only on the change of ‡ the binding energy of A, and ω1 = ω2 . In this case, the transition state is an early transition state that chemically looks similar to A [65]. If α is equal to 1, then the transition state is a late transition state that resembles the products of the reaction, and the change of the transition state enthalpy is related to changes of the binding energies of B and C. If (dB.E.A )/dx = (dB.E.B )/dx = (dB.E.C )/dx, then the late transition state gives rise to the situation ‡ where ω1 = 2ω2 . The above examples show how it is possible to maximize the activity of the catalyst. However, it is often more important to optimize the selectivity of the catalyst. Similar types of analyses can be carried out for these cases. In general, these situations are classified as being series selectivity challenges, such as A → B → C, where B is the desired product, and/or parallel selectivity challenges, such as A → B coupled with A → C. In these cases, if we want to optimize the selectivity of the catalyst, we search for some catalyst property, x, that decreases the enthalpy of the transition state for the desired reaction more than it decreases the enthalpies of the transition states for the undesired reactions. An undesired reaction that leads to progressive blocking of surface sites leads to deactivation of the catalyst with respect to time-on-stream. More generally, various mechanisms exist by which a catalyst can undergo deactivation, such as: (i) poisoning; (ii) thermal degradation (sintering); (iii) leaching of the active site; and (iv) attrition [66]. The first three of these mechanisms are chemical in nature, whereas the last mechanism is physical (e.g., the catalyst pellet breaks apart). Some catalysts do not show any measurable deactivation over periods of years, such as in ammonia synthesis. However, other catalysts lose an important fraction of their activity References see page 14

ch1.1

10

1.1 Principles of Heterogeneous Catalysis

after less than a minute of contact with feed, as in catalytic cracking. In the latter case, if deactivation is caused by coking, the catalyst must be regenerated by continuous regeneration in an oxidizing atmosphere. 1.1.7

Catalyst Design

Given that the performance of a catalyst is controlled by a limited number of kinetic parameters, it is unclear why it is so difficult to design a catalyst from molecularlevel concepts. As noted above, during the early stages of research into a catalytic process, first we do not know which steps in the reaction mechanism are kinetically significant, and which species are most abundant on the catalyst surface under reaction conditions. Second, we do not often know the structure of the active site and its dependence on the nature of the reaction conditions. Third, we do not usually know how the activity and selectivity for the catalytic reaction depend on the structure of the active sites. Fourth, we do not typically know during these early stages the rates of various modes of catalyst deactivation (e.g., sintering, phase changes, deposition of carbonaceous deposits on the surface, etc.), and we do not know whether the catalyst can be regenerated following deactivation. Finally, we must ensure that the texture of the catalyst and the geometry of the reactor are designed in such a way that mass transport of reactants and products to and from the active sites is sufficiently rapid that high rates of reaction per unit volume of reactor can be achieved. Because of these difficulties, the field of heterogeneous catalysis is highly interdisciplinary in nature, and

involves close collaboration between experts in such areas as catalyst synthesis, catalyst characterization, surface spectroscopy, chemical kinetics, chemical reaction engineering and, most recently, in theoretical calculations of catalyst structure and performance using density function theory. These broad studies can be grouped into three levels, as shown in Fig. 6. All studies of heterogeneous catalysis begin at the Materials Level. High-surface area catalytic materials must be synthesized with specific structures and textures, the latter referring to such features as the sizes of the various phase domains and the details of the pore structure. Clearly, the synthesis of catalytic materials must be guided by detailed characterization studies to determine the structures, compositions, and textures of the materials that have been prepared. These characterization studies should be conducted after the catalyst has been subjected to various treatment steps (such as those treatments employed during activation of the catalytic material), and it is most desirable to carry out characterization studies of the catalyst under the actual reaction conditions of the catalytic process. Indeed, the properties of a heterogeneous catalyst are inherently dynamic in nature, and these properties often change dramatically with changes in the reaction conditions (e.g., phase changes, surface reconstructions, changes in surface versus bulk composition, etc.) [67]. The central level of research and development of heterogeneous catalysts involves the quantification of catalyst performance (this is known as the Catalyst Performance Level). These studies can be carried out in a preliminary fashion over a wide range of catalytic materials (e.g., high-throughput studies) to identify promising catalysts

Materials level Characterization studies: catalyst structure, composition, & texture, (ideally under reaction conditions)

Materials synthesis: catalytic materials with specific structures & textures

Catalyst performance level Reaction kinetics studies: activity, selectivity & stability for various reaction conditions

Exploratory studies: promising leads for new catalytic materials & new catalytic reactions

Elucidation level Surface studies: Experimental studies: surface composition and nature interactions of probe of surface sites, (ideally under molecules with reaction conditions) well-defined sites

Fig. 6

Levels of study in heterogeneous catalysis research.

Theoretical studies: stability & reactivity of species on well-defined sites

ch1.1

1.1.8 Catalyst Development

and reaction conditions for further studies. The performance of the catalyst is then documented in greater detail by determining catalytic activity, selectivity, and stability with respect to time-on-stream for various reaction conditions. These measurements must be made at various conversions when multiple reaction pathways exist, because catalytic selectivities in these cases are different, depending on whether the desired products are formed in primary versus secondary reactions, or in series versus parallel pathways. We note here that various definitions of catalytic activity are used, depending on the nature of the study. For practical studies, catalytic activities can be reported as rates per gram of catalyst or per unit surface area. However, for more detailed studies or for research purposes, it is often desirable to report catalytic activities as rates per surface site (i.e., TOFs), with the number of surface sites measured most often by selective adsorption measurements (e.g., adsorption of H2 or CO to titrate metal sites, adsorption of ammonia or pyridine to titrate acid sites). In some cases it is possible to report catalytic activity as rate per active site (also called TOF), when it is possible to distinguish active sites from the larger number of surface sites using special probe molecules (e.g., dissociative adsorption of N2 to titrate sites for ammonia synthesis [68]; selective poisoning by adsorbates that compete with the reactants of the catalytic reaction [69]); or by transient isotopic tracing [70]. For the purposes of catalyst development, it is probably sufficient to work at the Materials Level and the Catalyst Performance Level. However, research into heterogeneous catalysis is dominated by studies conducted at a third level – the Elucidation Level – where the aim is to identify the fundamental building blocks of knowledge which can be assembled to build a molecular-level understanding of catalyst performance in order to guide further investigations to improve catalyst performance. At the Elucidation Level the studies are designed to determine the surface composition and nature of the surface sites on the catalyst [71–74]. Clearly, these investigations must be conducted with the catalyst under controlled conditions (e.g., under ultra-high vacuum, after treatment with H2 , after calcination, etc.) and, where possible, such measurements should be made with the catalyst under reaction conditions. Moreover, the studies may be carried out on real catalytic materials and on more well-defined surfaces (e.g., single crystals, or model samples formed by depositing known amounts of materials onto welldefined supports) [73, 75, 76]. Most measurements at the Elucidation Level involve studies of the interactions of specific probe molecules with the catalyst surface. These probe molecules may be the reactants, intermediates, or products of the catalytic reaction, or they may be more simple species chosen to monitor a specific functionality of the surface. Alternatively, a molecule may be used as

11

a probe because it has an advantageous feature for spectroscopic identification (e.g., CO for infrared studies, a 13 C-containing molecule for NMR studies). These studies of the interaction of probe molecules with surfaces are designed to determine the surface concentrations of different types of surface site, to determine the nature of the adsorbed species formed on the surface sites, and to determine the reactivities of the surface sites by monitoring the adsorbed species on the surface versus time, versus temperature or, most commonly, during a temperature ramp (e.g., temperature-programmed desorption). The third pillar of studies at the Elucidation Level involves the use of DFT calculations to assess the structures, stabilities, and reactivity of species adsorbed onto the surface sites (with the sites being composed of clusters of atoms or as periodic slabs of atoms) [77–81]. These studies are used to help interpret the results obtained from spectroscopic studies of catalyst surfaces (e.g., to predict the vibrational spectra of species adsorbed in different orientations on different sites), to calculate heats of adsorption for various intermediates in a reaction mechanism (e.g., to predict which species are expected to be abundant on the catalyst under reaction conditions), to estimate the energy changes for possible steps in a reaction mechanism (thereby eliminating from further consideration steps with very positive energy changes), and to determine activation energy barriers for steps that are suspected as being kinetically significant in the reaction scheme. Indeed, a key feature of these theoretical studies is the ability to predict how the surface properties are expected to change as the nature of the surface is altered (e.g., by changing the surface structure, or by adding possible promoters). This in turn will provide feedback to the Materials Level with regards to new materials that should be synthesized and which are likely to lead to an improved catalyst performance. In addition, these theoretically based studies provide information about highly reactive intermediates which might be difficult to obtain by direct experimental measurements. Most importantly, studies conducted at the Elucidation Level provide a scientific basis about the working catalysis that may, in future, be used to design different reaction pathways. 1.1.8

Catalyst Development

Catalyst development typically involves testing a large number of catalysts with a feedback loop, as it is currently difficult to design catalysts a priori. In this respect, catalyst development studies involve examining a large number of catalysts, for which recent advances in high-throughput References see page 14

ch1.1

1.1 Principles of Heterogeneous Catalysis

testing have attracted considerable attention [82–87]. Catalyst development through the testing of a wide range of materials was first practiced in 1909 by Mittasch at BASF who, according to Timm [88], issued the following directive to his team who at the time were developing the synthesis of ammonia: • The search for a suitable catalyst necessitates carrying out experiments with a number of elements, together with numerous additives. • The catalytic substances must be tested at high pressures and temperatures, just as in the case of Haber’s experiments. • A very large number of tests will be required. Ten years later, the number of tests conducted had exceeded 10 000, and more than 4000 catalysts had been studied. This extraordinary effort was also extraordinarily successful. What has changed since then, however, is the way in which the systematic search is assisted. Today, armed with an arsenal of principles, concepts, instrumentation and computers, it is possible to identify and to improve new catalytic materials in a much shorter time and with a smaller number of trial samples, especially with the possibility of advanced characterization methods (especially in-situ techniques) and insights from theoretical calculations (e.g., DFT calculations). The practical merit of this ‘‘assisted catalyst design’’ is clear, while its scientific dividend is the possibility of learning as the design proceeds, with the building of a data bank of rate constants and the formulation of more precise models of active sites. With new theoretical insights or principles, quantitative bases of catalyst preparation and reproducibility of catalyst behavior, the future of heterogeneous catalysis still looks very bright. The path to the design of an optimal catalytic process would be clear if the activity, selectivity and stability of the catalyst were to move in the same direction upon an increase in a single process variable, such as temperature. However, this simple behavior is not typically observed, and choices must be made in every instance. For example, while the activity of a catalyst may increase with temperature, its stability usually decreases with temperature. In addition, the relationship between catalytic activity and selectivity is typically very complex, and is not understood in detail until the surface chemistry of the catalytic process has been elucidated. Accordingly, selectivity, stability and activity must be considered together, and trade-offs may have to be negotiated, perhaps by using multi-functional reactors with catalytic distillation or catalytic membranes. Success in heterogeneous catalysis begins with chemistry, but ends with catalytic reaction engineering.

1.1.9

Bridging Gaps in Heterogeneous Catalysis

The above description of research and development into heterogeneous catalysis as being interdisciplinary in nature, involving studies at the levels of materials, catalyst performance and elucidation, can also be cast in the form of building bridges between various types of studies and different types of material. As depicted in Fig. 7, we often talk about bringing together the field of surface science (which traditionally is focused on studies of single crystal surfaces at low pressures) with the field of heterogeneous catalysis (which traditionally is focused on studies of high-surface area catalytic materials surfaces under highpressure reaction conditions). More recently, we have talked about ‘‘bridging the materials gap’’, as we have attempted to use experimental results from studies of welldefined model materials to interpret the performance of more complex, high-surface area catalytic materials. Traditionally, these model materials have been single crystals, cut at various angles to expose surfaces containing different types of sites, such as surfaces with different symmetries and atoms present at terraces, steps, and kinks [89]. More recently, however, these model materials have become highly sophisticated, such as the deposition of nanoparticles with specific sizes and geometries on welldefined support surfaces (e.g., metal nanoparticles supported on thin films of oxides deposited on single crystal metal surfaces, or non-metallic nanoparticles supported directly on single crystal metal surfaces) [73, 75, 76]. We also talk about ‘‘bridging the pressure gap’’, as we attempt to use experimental results from studies conducted at low pressures (less than 10−6 Torr) to interpret the performance of catalysts under high-pressure reaction

High m2

Heterogeneous catalysis

Low P

Pressure gap

Surface science

Materials gap

12

High P

Low m2

Bridging the gap between surface science and heterogeneous catalysis.

Fig. 7

ch1.1

1.1.10 A Philosophical Note

conditions. The origin for this pressure gap comes from the fact that, whereas some spectroscopic techniques can be employed to study the surface and bulk properties of catalysts under high-pressure reaction conditions (e.g., FTIR, Raman, XRD, EXAFS, M¨ossbauer spectroscopy), other spectroscopic and characterization techniques (e.g., XPS, TEM) are most easily conducted with the sample at low pressures (e.g., 10−2 Pa Dipole orientation

∼107

∼105

≥104 ≤103

871

the field strength drops discontinuously to zero. On real surfaces, the electron distribution and also the electric fields vary over distances of a few tenths of a nanometer. Based on pioneering work by Lang and Kohn [9], the detailed electron density profile at a jellium metal surface was studied in the presence of an electric field using density functional theory (DFT) (see, for example, Ref. [10]). Realizing the importance of specific geometric arrangements in adsorption at field emitter surfaces, i.e. ledges and kinks, which are also important in heterogeneous catalysis, Kreuzer introduced quantum chemical concepts [10–12] to provide a thorough understanding of the field-induced charge redistribution and its influence on the surface chemistry and reactivity. As an example, field evaporation of metals in the presence of chemisorbed layers may be associated with the removal of complete metal–adsorbate complexes. Thus a demetallization process of top-layer (kink) atoms due to field penetration is in operation and boosted by adsorption. Field-induced chemisorption of rare gases on a metal [13] and fieldinduced decomposition of chemisorbed species at fields of >1 V nm−1 are further examples. The latter case may be interesting from the viewpoint of heterogeneous catalysis and will be discussed in Section 3.1.3.5.4. B Field Emitter Surfaces and Experimental Methods High electric fields up to 60 V nm−1 can be obtained in a controlled manner at field emitter tips. Figure 1 provides a field ion image along with a ball model of such a tip in nearly 1 : 1 scale. Tips of various (conducting) materials can be prepared with different radii of curvature ranging from ∼10 nm (or even slightly less) to ∼100 nm. It is the nearly hemispherical shape of the apex which makes the tip an appropriate 3D model of a single catalyst grain free of any ceramic (or other non-conducting) support material. For (electrochemical) preparation procedures of tips see, for example, Ref. [14]. In addition to field ion imaging in its traditional mode – field ionization of an image gas (He, Ne, H2 , etc.) at a clean metal surface and acceleration of respective ions towards a screen – further methods have been developed notably with the aim of investigating chemical surface processes. As will be described below, dynamic imaging during an ongoing catalytic surface reaction has demonstrated since the 1990s to provide valuable insight into active site distributions, adsorbate mobilities, spatiotemporal pattern formation, etc., and field effects. In order to elucidate the chemical identity of adsorbed species, socalled atom-probe devices have been developed since the 1960s. These methods are compiled in Table 2. Details can be found in a number of books [14]. References see page 892

872

3.1 Physical Properties

(111)

[010 ]

(101)

(011)

[100]

[10

0] (011)

[010

]

(001)

(111)

(111) (101)

(111)

(a)

(b)

(a) Field ion micrograph of a Rh tip under best image conditions in neon at F = 35 V nm−1 and T = 55 K. Some planes are indicated by their Miller indices along with 001 zone lines. (b) Ball model depicting 3D morphology and atomic arrangements of the tip apex shown in (a).

Fig. 1

Coverage

Short

Impinging material

Medium

Long time

t

Coverage

Time Reaction products

Time

Scheme illustrating field pulses of different repetition frequencies. A systematic variation of the reaction time (time elapsed between two pulses) provides kinetic information.

Fig. 2

Additional comments will be made here on pulsed field desorption mass spectrometry (PFDMS), which has particular importance for fundamental research in heterogeneous catalysis. The general experimental procedure for kinetic measurements is illustrated in Fig. 2. Using short field pulses (width ∼100 ns), timedependent surface processes can be studied. Between any two pulses – this period defines the reaction time tR – gaseous reactants may adsorb on the surface of the field emitter tip and react. Three different pulse repetition rates corresponding to short, medium and long reaction times are depicted. Pulses serve two purposes: they stop the surface reaction process by

field desorption/evaporation of the adsorbed layer and restore a well-defined starting condition (zero coverage, for example) for a new reaction cycle. The longer the time tR – usually tR can be adjusted at any time between 100 µs and 10 s – the more the reaction proceeds and, consequently, the larger are the amounts of products and intermediates. It may occur that product species have short lifetimes at the surface and therefore (partly) escape detection by pulses. The formation of intermediates can, however, straightforwardly be studied in a truly in situ manner, e.g. while dosing reactants and running the device as a flow reactor at pressures below 10−1 Pa. Provided that quantitative field desorption is possible, respective adsorbate coverages can be quantified. PFDMS studies can be performed by avoiding any steady field between pulses. On the other hand, a so-called reaction field FR can be biased to study the field influence from the onset of field evaporation to the onset of field electron emission at reversed polarity. Thus a complete kinetic scheme is accessible for defined values of pressure, temperature and field strength. Instead of field pulses, laser pulses can be used to trigger field desorption in the presence of a steady electric field (PLAP, PIFD). In field ion appearance mass spectrometry (FIAPS), field ions are analyzed for their energy by applying a retarder potential. This amounts to produce an integral energy distribution with an onset potential of δ0 (see Table 2). The energy distribution provides information about the energetics of a surface reaction and allows the discrimination of different reaction pathways. In Fig. 3, the PFDMS method is shown schematically in combination with the video field ion microscopy (FIM) method (b and d). Running the setup as a flow reactor at low reactant pressures allows imaging with

FDS

TFDS

FIAPS

Thermal field desorption spectroscopy

Field ion appearance spectroscopy

PLAP PFID

Pulsed laser atom probe photoninduced field desorption Field desorption spectroscopy

PFDMS

F = const. T = const.

N(i)

F = const.

T

T = const.

F

F = const.

T

T = const.

TR

TD

FR

∆d0

tR

tR

tD

tD

Time

Time

Time

Time

∆d

Integral energy distribution by scanning counter potential δ at constant total ion current

Thermal desorption spectroscopy of ions at constant field

Ramped fields, desorption fields, depending on binding energy

Laser pulse causes T-pulse thermal desorption of ions

Field-dependent field desorption

Field evaporation of solid surface as ions, mass spectrometric analysis; depth profiles of surface composition Field ion source in combination with magnetic sector field or time-of-flight (TOF) mass spectrometers

AP FDMS

FD

Field ionization of noble gases or of reacting gases; atomic resolution with noble gases

FIM

F

Field emission of electrons according to local work function, φ; lateral resolution ∼2 nm

FEM

Field electron microscopy Field ion microscopy Atom probe Field desorption mass spectrometry Pulsed field desorption mass spectrometry

Features

Acronym

Experimental methods

Method

Tab. 2

[26, 27]

[25]

[23, 24]

[21, 22]

[19, 20]

[17] [18]

[16]

[15]

References

3.1.3 Structure and Morphology

873

874

3.1 Physical Properties

[010]

[01 1]

(111) [101] To image recording system

To ToF mass spectrometer

(011)

(101) (113)

[10

(111)

irr or

0]

M

(001) [10

FIM screen

(111)

0]

(317)

(011)

(317) (101) [010]

Tip

(a)

]

1 01

[

[10

1]

(111)

(b) To ToF mass spectrometer To FIM screen

(c)

(d)

(a) Scheme of the probe-hole field ion microscope. Ions hitting the FIM screen provide a micrograph of the tip apex whereas those passing through the probe-hole are mass separated and identified by their flight time. (b) Field ion micrograph of an Rh tip at T = 505 K exposed to an O2 −H2 (1 : 2) mixture at a total pressure of 4.5 × 10−3 Pa. (c) Scheme of the probe-hole principle: the hole in the screen of the field ion device shown in (a) selects a small area of the tip sample during the ongoing reaction. (d) Field ion micrograph during the ongoing H2 + O2 reaction. The micrograph was taken under experimental conditions suitable to investigate the bistability of the reaction (pO2 = 1.5 × 10−3 Pa, pH2 = 2.5 × 10−3 Pa, T = 500 K). The surface composition in (d) is clearly different from that in (c) and the probe-hole device allows local differences to be established. Fig. 3

high spatial resolution. Superimposing short field pulses allows chemical probing on selected surface areas while imaging. In FIAPS, a lens system is added to provide a kinetic energy analysis of the ions [28]. Different applications of high-field methods in heterogeneous catalysis are compiled in Table 3. C Determination of Field Strength Values Conventional field strength determination is based on I –V (Fowler–Nordheim) measurements (I = electron current) combined with field electron energy analysis [29]. F0 values are determined within an accuracy of ±15%. In FIM, the energy analysis of free-space field ionization provides F0 values within 3% accuracy [30, 31]. The field strength can also be calibrated from the best image voltages (BIV) in FIM. For instance, the BIV for Ne is 35 V nm−1 . All these methods determine electric fields,

F0 , at a distance of >1 nm away from the surface. In closer vicinity to the surface ( 7 V nm−1 [96]. The same behavior was seen in a quantum chemical molecular-orbital calculation by Wang and Kreuzer [108]. This led to the remarkable prediction that high-index Ru(CO)x species must be more abundant on stepped Ru surfaces in the absence of electric fields than on Ru field emitter tips under PFDMS conditions. A strongly field-promoted case of subcarbonyl formation was recently reported for Au field emitter surfaces [109] and a comparison with the Ru case is presented in Section 3.1.3.5.4D. We proceed by reporting on the kinetics of subcarbonyl formation for metals such as Ni, Rh, Ru and Co. Reaction time variation (100 µs < tR < 10 s) showed the occurrence of delay times in dicarbonyl formation as one of the common features of all these metal–CO systems. Moreover, during the initial stages, the Me(CO)2 species built up with a time dependence clearly different from the first-order Langmuirian behavior as seen for COad (represented by CO+ and MeCO2+ ). In fact, this excludes the possibility that significant amounts of the dicarbonyl were formed during field desorption of COad . Instead, a consecutive surface reaction must be envisaged, most probably involving Me step sites, i.e. kinks. As field strength variation measurements gave a strong indication for surface mobility of Me(CO)x [100, 106, 107, 110], a reaction model was proposed in which the dicarbonyl formation was associated with Me–Me bond breaking and diffusion into the terrace regions of the surface: [MeCO]kink + COad → Me(CO)2,ad + [ ]kink . Following this scenario, the tricarbonyl and, eventually, the tetracarbonyl species (in the case of Ru) were References see page 892

882

3.1 Physical Properties

Intensity ions/1000 pulses

COad / NiCO

Concentration molecules/site

10

10−4 Ni(CO)2

Ni(CO)3

1

10−5

T = 295 K FDes = 20.6 V nm−1 PCO = 3 × 10−4 Pa 0.1 10−6 10−4

10−3

10−2

10−1

Reaction time / s Variation of the surface concentration of nickel carbonyls, Ni(CO)x , x = 1 − 3, with reaction time tR from a monitored area of 9.5 nm2 (∼150 surface sites), according to [104].

Fig. 6

subsequently formed in the terrace layer by adding on another (or two) COad to the Me(CO)2,ad precursor. For long reaction times, steady amounts of Me(CO)x were measured, thus indicating reversibility of the individual reaction steps. Figure 6 demonstrates the behavior on Ni as a representative example of carbonyl formation on transition metals. Since the Ni(CO)2,3 species are quantitatively field desorbed in this example (field strength variation demonstrates saturated intensities), their mean surface coverage was evaluated to demonstrate the ultimate detection sensitivity of the PFDMS method. From an energetic point of view, Me−Me bond scission requires considerable effort. The occurrence of delay times in Ni(CO)2 and, more generally, in Me(CO)2,ad formation is therefore not surprising. Further support for the given reaction model was obtained on varying the reaction temperature. Any increase in the latter caused an acceleration of the process, i.e. a shortening of the delay times. Evidence for such thermal activation was also provided by infrared (IR) studies in which di- and tricarbonyl species were found to form on supported Rh particles [111]. Direct proof of the COinduced disruption of small Rh crystallites was given by van’t Blik et al. [71] in EXAFS studies. Topographic and morphological alterations caused by subcarbonyl

formation were observed by FIM and are discussed in Section 3.1.3.5.3A. In a recent paper, Vestergaard et al. [112] reported on a CO-induced phase separation for Au/Ni(111) surface alloys. Fast STM measurements at pressures in the millibar region and at 300 K allowed Au nucleation and cluster growth due to Ni removal from steps to be followed. Using DFT calculations, they also proposed a reaction model with Ni(CO)2,3 formation as described above on account of PFDMS data. C Modeling Structures and Interfacial Transport Processes The ability of FEM methods to study surface diffusion with high spatial resolution was recognized as early as the 1950s [113, 114]. For example, valuable kinetic information was obtained for migrating adsorbates such as hydrogen, oxygen, carbon monoxide and others using the ‘‘shadowing’’ FEM technique originally developed by Drechsler [115] and Gomer [116]. The principle of operation was to supply gas to only one side of the emitter shank and to follow the changing emission patterns in the very apex region while or after heating the tip. With the introduction of the probe-hole device, measurements on single-crystal planes became possible in the so-called ‘‘field emission current fluctuation method’’ [117, 118]. From the point of view of catalysis, the chemical identity of structures and diffusing species is of considerable interest. This kind of information is provided by atomprobe mass spectrometry. We give here a few examples for metal–metal oxide systems as these are of considerable relevance in catalysis. PFDMS studies during tungsten oxidation at 1100 K [119] revealed fairly strong variations in the surface composition with increasing oxide layer thickness. A stage-by-stage build-up of various oxide states was found by systematically varying the field pulse amplitude (FP ). At high field strengths mainly WO+ 3 and various suboxides, + /O2+ were detected. Under (n, x = 1, 2) and O WOn+ 3–x these conditions, nucleation and oxide layer growth were largely inhibited. At low field strengths, however, a thick oxide layer built up and, in addition to WO+ 3 , clusters (n = 1, 2; y = up to 5) appeared in the mass of (WO3 )n+ y spectra. The results indicated that WO3 molecules were the major mobile species acting as precursors in the formation of large clusters. These conclusions are left untouched by the observation that ‘‘negative and positive’’ electric fields were found to enhance the formation of oxide crystallites during oxidation [120]. Information on the mobility of MeOy molecules was also obtained in PFDMS studies with Ru metal tips [121, 122]. Again, a variety of oxidic ions (including suboxides and cluster species) were found, with their intensities being strongly dependent on the desorption field strength, i.e. the oxide layer thickness. At a reaction temperature of

3.1.3 Structure and Morphology

700 K, RuO2 and RuO3 molecules appeared to be mobile on top of an oxide layer. The long-range diffusion of these species was proved in experiments with Pt emitter tips covered by a thick Ru layer on their shank. Under continuous supply of oxygen at 850 K, both RuO2 and RuO3 were found to cross the Pt/Ru interface and to migrate into the monitored area of the (oxidized) Pt apex [122]. The initial stages of Rh oxidation were studied by Kellogg [123] in an imaging atom-probe which was directly coupled to a reaction chamber for sample preparation. After oxidation in ∼102 Pa O2 gas at temperatures between 350 and 550 K, various oxygen-containing ions (O+ , RhOn+ , RhO+ 2 ) were detected. From a plot of the cumulative number of oxygen and Rh atoms (with the latter also originating from field evaporation) at different temperatures, information about the layer stoichiometry was obtained. The initial slope of the plot indicated the presence of Rh2 O3 , with a thickness of ∼30 layers. A fairly low activation energy, 16–20 kJ mol−1 , was determined from an Arrhenius plot of the saturated oxygen uptake and was associated with either O2 dissociation or oxygen dissolution into the bulk. The formation of an oxidic surface layer is the most likely reason for the decrease with time of the CO oxidation activity of Rh in catalysts for automotive pollution control. In addition to direct oxidation studies, a number of further attempts have been made to simulate multifunctional catalytic behavior on field emitter tips. For example, Knor and Sotola [124] showed that a realistic modeling of supported metal catalysts is possible. This was demonstrated in FEM/FIM studies with Pd and Mo layers successively deposited on a WOx substrate in UHV conditions. At high temperatures, this metal–metal oxide system was found to display the characteristic features of an SMSI effect (strong metal–support interaction). Islands of molybdenum metal were found to be active in the dissociative trapping of nitrogen molecules. The successive hydrogenation to ammonia occurred on Pd islands after nitrogen spillover from Mo. The occurrence of a similar spillover process from a tungsten surface towards palladium was demonstrated by measuring the characteristic changes in the work function [125]. The hydrogenation of adsorbed nitrogen atoms was also seen on an Ir surface [126]. In a more recent study, the nitrogen spillover process could be demonstrated for a Pd−Mo/Al2 O3 / W surface prepared on a tungsten tip and using alumina, which is one of the important carrier materials in catalysis [127]. Mobile surface species involving spillover phenomena were also reported from studies by means of field ion mass spectrometry [128]. A Pt field emitter could be prepared with isolated areas of zeolite grains. Mass spectrometric methods such as RAPS allowed for distinction between zeolite and platinum surfaces

883

as the origin of field-desorbed ions [129]. During the chemisorption of hydrocarbons, such as cyclohexane, n-heptane and neopentane, zeolites promoted the cleavage of C−H bonds and the dissociative chemisorption of these molecules with subsequent spillover to the Pt surface. D Dynamic Imaging By the beginning of the 1990s, FEM and, slightly later, FIM have been increasingly employed to image catalytic surface reactions on field emitter tips. Using video techniques, real-time information with spatial resolution at the nanoscale became available. In FEM, the local field emission current depends on the local work function, which varies with adsorbate composition. The lateral resolution is of the order of 2 nm and limited by Heisenberg’s uncertainty principle. The maximum resolution of FIM may be better than 1 nm under reaction conditions. Compared with classical FIM work, which aims at revealing the surface structure of metals and alloys, the principle of operation in reactive imaging is to ionize small amounts of reactants or products in the presence of electric fields of the order of 10 V nm−1 . Since ionization is highly localized, the underlying surface structure usually influences the ionization probability of adsorbed species; in some cases morphological and topographical features ‘‘modulate’’ the ion intensities such that reconstruction during reaction becomes visible in real time. The interpretation of video FIM data is not always straightforward, however, and the influence of the imaging field has to be addressed with care. On the other hand, the advantage of reactive imaging with nanoscale resolution is most worthwhile and the possibility to combine FIM with atom-probe techniques provides another clue to the methodological approach. Some of the data accessible in this manner will be presented below. Generally, the kind of information obtained includes (i) identification of catalytically active sites under reaction conditions, (ii) nucleation and dynamic growth of adsorbate islands and (iii) observation of non-linear reaction–diffusion fronts providing coupling between surface planes of 3D metal particles. Most remarkably, self-sustained oscillatory behavior may be observed. Non-linear reaction dynamics in heterogeneous catalysis have attracted increasing interest since Hugo [130] and Beusch et al. [131] first reported on oscillatory states during CO oxidation on supported Pt catalysts. Insight into the underlying mechanisms and spatio-temporal pattern formation was, however, only obtained in studies with crystallographically well-defined single crystals under overall low gas pressures where thermokinetic effects become negligible. A summary of the experimental achievements and the theoretical background is given in References see page 892

884

3.1 Physical Properties

T/K 600 550 500

450

400

350

300

250

Phase diagram of CO + O / [111] - oriented Pt tip Pumping speed = 45 L s−1

10−2

P O2 = 5.3 × 10−2 Pa

Large Pt(110) single crystal

PCO / Pa

tB

10−3

Steady state 1 Low reactivity

Oscillations tA

B

Oscillations

Pt tip Crossing point

10−4 A

Hysteresis

Steady state 2 High reactivity

tA

10−5

tB

1.5

2.0

2.5

3.0

3.5

4.0

4.5

1/T (1000/K) The cross-shaped kinetic phase diagram describing regions of steady state of oscillations and of hysteresis at pO2 = constant, according to [136].

Fig. 7

Ref. [132]. We focus our attention on FEM work in which, in contrast to scanning tunneling microscopy (STM), the whole emitter surface is imaged at once and where the time resolution of the reaction dynamics is limited only by the frequency of the video recording system. The first FEM reports on catalytic surface reactions with oscillatory states were presented independently by van Tol et al. for NO reduction with hydrogen [133] and ammonia on Rh [134] and by Gorodetskii et al. for CO oxidation on Pt emitter tips [135]. As the latter is probably one of the most thoroughly studied reactions in heterogeneous catalysis, we consider it here in more detail and develop some of the basic concepts and achievements. First, adsorption of CO on Pt causes an increase in the work function, , of all Pt single- crystal planes [except Pt(111)] and the (dissociative) adsorption of oxygen causes an even stronger increase in . Since the field emission current is high (low) when  is low (high), local variations in this current reflect changes in the layer composition and 2D spatio-temporal variations monitor the occurrence of a surface reaction. Second, dark areas (oxygen-covered) may be repetitively replaced by bright (CO-covered) and occasionally very bright areas (uncovered). Transitions are not symmetric in time; whereas a CO-covered surface transforms within fractions of a second into an oxygen-covered surface, it takes several seconds for the reverse. Third, a bifurcation

diagram (reactive phase diagram) can be constructed to demonstrate the occurrence of bistability and oscillatory behavior in p –T parameter space. An example of such a bifurcation diagram is shown in Fig. 7, which is based on an FEM brightness analysis of the Pt {011} facet and its surroundings [136]. A cross-shaped feature occurs with lines delimiting steady states of high reactivity (A, ‘‘oxygen side’’) and low reactivity (B, ‘‘CO side’’) from bistable and oscillatory regions below and, respectively, above the cusp point. Interestingly, data obtained with macroscopic Pt (110) single crystals can be extrapolated to fit into the same bifurcation diagram (upper left in Fig. 7). Very similar cross-shaped diagrams are obtained by probing different facets of the Pt tip [137]. Obviously, efficient coupling via fast CO surface diffusion must be in operation if it is assumed that adsorbed oxygen is of rather limited mobility under reaction conditions. The fast transformation of a CO-covered into an oxygencovered surface remains surprising in view of the drastic differences in oxygen sticking probabilities which vary between 0.5 on Pt(110) [138] and 10−3 –10−4 on hexreconstructed Pt(100) [139]. Video FIM investigations of catalytic CO oxidation on Pt were first reported by Gorodetskii et al. [140]. Generally, dynamic imaging in FIM is based on measurements of the local variations in the field ion current of reactant and/or product molecules. Ionization occurs at

3.1.3 Structure and Morphology

critical distances, xc = (I − φ)/eF (I = ionization potential, φ = work function, F = field strength), above the surface and involves quantum mechanical tunneling. The distance, xc , is rather large in the presence of an adsorbed layer. However, according to the resonance model proposed by Kreuzer and Wang [141], the tunneling probability for an electron of the image gas can still be high if its HOMO (highest occupied molecular orbital) shows sufficient overlap with the adsorbate LUMO (lowest unoccupied molecular orbital) lying close to the Fermi level of the metal. The resonance model provides a theoretical explanation for the assumption [140] that the relatively bright regions during the catalytic CO oxidation on Pt tips are due to preferential O2 ionization above Oad -covered regions. The model seems also to be applicable to NO reduction with hydrogen on Pt, where high ionization rates of NO were observed above Oad -covered regions [142]. In order to probe plane activities in CO oxidation, a COad -covered surface was subjected to titration with oxygen [143]. FIM data demonstrate the transformation into an Oad layer to proceed fastest around {011} areas and, more specifically, in {331} planes. Crystallographically equivalent {011} planes always transform first yet nonconcerted before anisotropic wave fronts reach {012}, {122}, {001} and finally {111} planes (after ∼400 s at 300 K). Fronts move slowly on some planes (∼0.8 nm s−1 on {012} at 300 K) and fast in a ‘‘switch-on’’ manner in others (50 nm s−1 on {001}). Their speed increases with increase in temperature. This specificity in plane activities was also observed in the oscillatory regime of the reaction [excluding {111} and showing fast autonomous oscillations on some hundred surface sites of (001)] [144]. It seems likely that the high O2 sticking probability in {011} and their surroundings is jointly responsible for the high activity of these planes. A straightforward evaluation is hampered by the fact that sticking and diffusion coefficients show different dependences on COad and Oad coverages in different planes. Gorodetskii et al. state that Oad sites remain undisturbed while being imaged by O+ 2 at field strengths of ∼15 V nm−1 [143]. Most interestingly, FIM and FEM data seem to fit into the same bifurcation diagram (Fig. 7) if gas pressures are corrected for the ‘‘field compression effect’’ (attraction of O2 molecules due to field-induced polarization). Subsequent work by Suchorski et al. [137] reports that electric fields have negligible influence only at values below 5 V nm−1 . More recent FEM work has addressed fluctuationinduced transitions in the bistable regime of the reaction [145]. Measurements within a specific facet such as Pt {011} show spatially well-correlated transitions between the two branches of the reaction. Focusing on Pt {112}, the amplitude and the spatial coherence of the fluctuations increase the closer the distance to the cusp

885

point of the bifurcation diagram. The spatial correlation decays rapidly at surface steps terminating the facet. It must be concluded that steps either slow CO surface diffusion or act as sites of enhanced chemical reactivity. Interestingly, Suchorski et al. also report on size effects, i.e. the smaller the {112} facet, the further away from the cusp point of the bifurcation diagram the fluctuationinduced transitions occur. The second example of dynamic imaging considered here in more detail concerns water formation from O2 /H2 on Pt and Rh field emitter tips. Both FEM and FIM were employed to reveal specific plane activities by subjecting an oxygen-covered Pt specimen to hydrogen at temperatures between 300 and 395 K [146]. Titration starts in (111) regions, i.e. in {133} terraces, and continues anisotropically to reach {011}, {012}, {013}, {015} until, finally, the central (001) pole. Reaction patterns are similar in FEM and FIM, leading to the conclusion that field effects cannot play a dominant role in the titration. The fact that (111) planes are active in water formation is in agreement with earlier FEM work by Gorodetskii et al. [147] but different from reports by Verheij et al. [148] that suggest the presence of a few yet highly active sites controlling the titration of Oad on macroscopic Pt (111) single crystals. Subsequent STM studies with single crystals of the same orientation at temperatures between 110 and 170 K, i.e. above the formation temperature of water but below the onset of its thermal desorption, demonstrated that chemical waves [149] spread as ringlike features across terraces and steps of the surface. An autocatalytic mechanism was constructed by assuming that water is an intermediate formed according to OHad + Oad → H2 Oad , which immediately reacts to produce two surface hydroxyls via H2 Oad + Oad → 2OHad . Ideally the process continues until the oxygen layer is completely transformed into a water layer. This scenario only applies at low temperatures and is considered immaterial in the temperature range of FEM/FIM titration experiments, where water desorption is fast. Self-sustained kinetic oscillations were observed in FIM experiments at 300 K [146, 150] with periods strongly depending on the hydrogen partial pressure. Similarly to titration experiments [150], reaction fronts were found to start in {331} areas and to move anisotropically across the surface. It is interesting that also the {001} pole turned periodically active although mesoscale photoemission electron microscopy (PEEM) with macroscopic (001) single crystals only exhibited the formation of adsorption islands and pressuredependent concentration waves [150]. Kinetic oscillations on field emitter (001) planes are therefore a clear manifestation of the existence of coupling effects between different adjacent planes. As was emphasized [150], this

886

3.1 Physical Properties

phenomenon is fundamentally different from spillover in supported catalysts. Image formation in FIM is based on oxygen and water ionization. This can be concluded from mass spectrometric analyses of ions while imaging. Both H2 O+ and H3 O+ are observed. Their localized formation in field ionization processes allows the identification of active sites during the H2 −O2 reaction on Pt. The occurrence of H3 O+ ions during catalytic H2 oxidation on Pt is the reason for the occurrence of a new fielddependent reaction channel to produce Oad without face specificity [151]. The reaction of such Oad with hydrogen displays chemical waves which differ from those where Oad was formed by dissociative O2 chemisorption from the gas phase. This has to be expected since this latter process is face-specific whereas the former is not: field

2H2 Oad,f −−−→ H3 O+ + OHad +
ad + e− (Me) 2OHad −−−→ H2 Oad,f + Oad

(1) (2)

H3 O+ is not only formed according to Eq. (1) and a second pathway must be considered: Had + H2 Oad,f −−−→ H3 O+ + e− + 2
ad

(3)

This follows from detailed FIAPS measurements by Sieben et al. [152], according to whom both pathways are associated with different ion appearance energies. H3 O+ is kinetically ‘‘slow’’ in Eq. (1) and ‘‘fast’’ in Eq. (3) due to the different energetics involved (E ≈ 2 eV) and the respective ion intensities can be regarded as reliable probes of the Langmuir–Hinshelwood kinetics of H2 Oad,f formation from H2 −O2 in the presence of an electric field. Although H3 O+ formation is clearly fieldpromoted in FIM experiments, it is interesting that this species may even occur on macroscopic Pt single-crystal surfaces in the absence of an electric field [153]. More recent video FIM work has addressed the H2 −O2 reaction on Rh field emitter tips. A major difference compared with Pt metal is that Rh is much more prone to oxidation. The reaction was found to be structure sensitive, in agreement with earlier FEM reports by Gorodetskii et al. [154], bistable at temperatures between 400 and 500 K and oscillatory at 550 K [155]. Figure 8 presents the respective reaction patterns covering one complete oscillation cycle. Starting on the hydrogen side of the reaction (a), oxygen island formation is seen to occur on the (001) pole after ∼10 s (c). A careful analysis of the nucleation process revealed kink sites in the (001) layer edge to be most active [156]. Subsequently, the Oad layer spreads anisotropically in a back-and-forth manner from (001) to {011} planes such that cross-like patterns appear (d). The {111} planes are invaded along 110

zones via {011} planes. Finally, the whole Rh emitter tip is covered by Oad (e) and the image becomes dim. The oscillatory cycle is closed by a fast (∼100 ms) clean-off reaction which removes the Oad layer in an isotropic manner from the periphery to the (001) pole. This reaction is associated with a ‘‘burst’’ of water formation. Selecting ∼400 atomic sites close to the (001) pole in an atom-probe mass spectrometer (PFDMS) revealed large amounts of H2 O+ /H3 O+ while imaging this fast reaction (see Fig. 9b). Various RhOx ionic species in the same spectrum prove the surface to be oxidized. By contrast, the surface is metallic on the hydrogen side of the reaction (Fig. 9a). Water formation is much less intense under these conditions. The H2 −O2 reaction takes place on the surface of a polyhedral Rh tip. As discussed in Section 3.1.3.5.3A, field-free oxygen adsorption at temperatures above 500 K causes a 3D reshaping of the originally hemispherical Rh specimen (Fig. 4). Major features of the polyhedral morphology are also visible in Fig. 8a–c, i.e. while imaging the hydrogen side of the H2 −O2 reaction. The spatio-temporal patterns associated with Fig. 8 demonstrate oscillations to involve reversible surface oxidation. Generally, dynamic reactions at the tip apex are inevitably coupled to those at the tip shank where the electric field is low. An oscillation mechanism can then be constructed by considering field-promoted diffusion of Had from the shank reservoir to the apex, where fast clean-off occurs according to Oad + 2Had −−−→ H2 O + 3


(4)

Since three empty sites are created per desorbing water molecule (neglecting field-adsorbed intermediate states), an autocatalytic behavior can be constructed. The chemical front during clean-off moves with a speed of ∼0.1 µm s−1 , which is compatible with this scenario. An illuminating example of a field-driven catalytic reaction was provided in FIM/FIAPS studies of H2 −H2 O mixtures on Pt [148]. The kinetics of this reaction differ markedly from the H2 −O2 case. Video FIM patterns along with energy-resolved H3 O+ yields showed a periodic behavior in time and space. To simulate the oscillating H2 −H2 O reaction, a kinetic lattice gas model was constructed comprising H2 Oad , OHad , Oad and Had as adsorbed species occurring during the reaction. Space–time plots for fast and slow H3 O+ [Eqs. (1) and (3)] were calculated and a striking agreement with experiment was found for slow H3 O+ assuming that the high electric field gradient allows rapid diffusion of Had from the shank to the apex of the tip.

3.1.3 Structure and Morphology

887

(111)

]

(101) (111)

[010

(011) [100

]

(001) [10 0

[010

]

]

(111)

(011)

(101) (111) (b)

(c)

(d)

[010

]

(a)

[01

0]

[100 (001) [100 ] ]

(e)

(f)

Series of micrographs showing the most important features of one complete oscillatory cycle. Starting from a surface in the quasi-metallic state (a) an oxide layer is formed first on the topmost plane (b, c). Kink sites of the (001) pole [see square in (b)] react first. The oxide front spreads preferentially along the 100 zones (d) to invade finally the whole visible surface area (e). Experimental conditions: T = 550 K, pH2 = 1.3 × 10−3 Pa, pO2 = 1.0 × 10−3 Pa, F = 9 V nm−1 . (f) Ball model depicting the situation encountered within the square frame of (b) where kink sites are covered with an oxide layer first.

Fig. 8

3.1 Physical Properties

0.15

Ions/1000 pulses

The third major example of dynamic imaging to be presented here is the reduction of NO, which was studied over Rh and Ir using H2 [133, 157–159] or NH3 [134, 159] and NOx over Pt [142, 160, 161]. On all three metals selfsustained oscillatory behavior was observed with distinct surface plane selectivities. Using FEM, van Tol et al. [157] reported reaction fronts moving across the surface of an Rh tip after nucleation in {335} areas at 460 K. Surface defects and grain boundaries were also seen to transmit fronts. Planes such as {001}, {012}, {013} and {113} were, however, omitted by these fronts. The same authors also reported on in-phase oscillations in different {011} planes, which they took as an indication of gas-phase communication. Detailed studies of the plane selectivity and the underlying reaction mechanism showed that planes not participating in oscillations were probably covered by strongly bound Nad . Generally, Oad and Nad species formed by NO dissociation on Rh may react with Had to form H2 O and NH3 . Spectroscopic studies with macroscopic single crystals whose orientation ({335}, for example) was selected according to FEM plane activities showed surface non-linearities to be related to periodic transitions between high-coverage Nad and Oad , with Oad destabilizing Nad such that N2 thermal desorption was accelerated [159]. While the oscillating behavior of the NO−H2 reaction on Rh tips was found to be little influenced by the viewing field in FEM (oscillations continued after switching off and on the field), the same reaction showed a drastic field dependence on Pt tips. Cobden et al. [162] mentioned transient oscillations in FEM, close to the (001) plane, but not on the plane itself. In FIM, however, a quasi-periodic reaction behavior with specific plane reactivities was observed on hemispherical and polyhedrally reconstructed Pt crystal tips [142, 160, 161, 163]. Whereas the reactant partial pressures were similar in both FEM and FIM studies, the temperatures were fairly different, i.e. ∼420 K in Ref. [162] and ∼520 K in Ref. [142]. Catalytic cycles in FIM were of duration ∼0.1 − 0.2 s and were characterized by fast ignition processes (‘‘surface explosions’’) in kinked planes along the 100 zone [161]. Depending on the size and the 3D morphology of the Pt emitter tip, such explosions were observed on {012}, {011} or any other small plane of the 100 zone [161]. After ignition, spatio-temporal pattern evolution led to cross-like features along 100 including the central (001) pole. Autocatalytic water formation was found to be followed by NOad diffusion into empty sites: plane-specific mass spectrometric (PFDMS) data while imaging revealed bursts of NO+ species subsequent to high H2 O+ /H3 O+ production. No NHx ionic species were detected. Self-sustained oscillatory behavior in the NO2 reaction with hydrogen on Pt was observed in both FEM and

H2O+ H3O+

0.1

0.05

O2+

0 20

40

60

80

100

120

140

160

180

m/e

(a) 2 1.5

Ions/1000 pulses

888

H2O+ H3O+

RhO2+ RhO2+

0.3 0.2

RhO+ Rh2+

0.1

RhO22+

Rh+ RhO3+

0 20

(b)

40

60

80

100

120

140

160

180

m/e

(a) Field pulse mass spectrum during the H2 + O2 reaction on the ‘‘hydrogen side’’. About 400 sites in the vicinity of the (001) topmost layer are monitored. Experimental conditions: pO2 = 1.0 × 10−3 Pa, pH2 = 7.6 × 10−3 Pa, T = 450 K, F = 10 V nm−1 . (b) The same reaction on the ‘‘oxygen side’’. Conditions: same as in (a) except for the hydrogen pressure: pH2 = 8.0 × 10−4 Pa. Fig. 9

FIM [160]. Different to the NO + H2 case, kinetic instabilities were found to persist with similar pattern formation while switching the polarity, i.e. inverting the electric field vector. Characteristic times under oscillatory conditions changed due to variations in the gas supply function. Generally, the parameter space (pNO2 , pH2 , T ) to establish oscillatory behavior in either FEM or FIM was found to be much larger than for the NO−H2 reaction. An example of FEM pattern formation during one catalytic cycle at 465 K is shown in Fig. 10. The image brightness first increases in {012} planes before developing cross-like features along 100 . Thus, planes with kinks are most active which is similar to the NO−H2 case. In Fig. 10d the reaction zone begins to collapse while the (001) pole is still active; in Fig. 10e the surface has returned to the initial state of low catalytic activity. The {012} planes play the role of pacemakers in the overall oscillatory mechanism. The relation of FEM patterns with states of different catalytic activity is based on adsorbate-induced work function

3.1.3 Structure and Morphology

889

(011)

(101)

(012)

(102)

(001) (102) (012) (101) (011)

t = 0 ms

(a)

t = 200 ms (b)

(113)

(113)

(001)

(113)

(113)

t = 560 ms (d)

t = 1480 ms (e)

Mean intensity /arb. units

(c)

t = 800 ms

250 200 150 100 50

85 (f)

90

95

Time / s

Sequence of field electron micrographs during the oscillating NO2 + H2 reaction on Pt (pH2 = 1 × 10−3 Pa, pNO2 = 5 × 10−4 Pa, T = 465 K); the reaction ignites in {210} planes (b) and reaction fronts spread anisotropically towards the (001) pole (c–e). (f) Mean FEM intensity vs. time of the oscillating reaction.

Fig. 10

890

3.1 Physical Properties

changes which at 465 K are dominated by Oad . Periodic changes in Oad then reflect rate oscillations of water production. The high dissociative sticking probability of NO2 plays an important role in the fast build-up of high Oad coverages. The reaction with Had then follows essentially according to Eq. (4). 3.1.3.5.4 Field-Induced Surface Phenomena A Field-Induced Chemisorption The redistribution of valence electrons in atoms and molecules due to shifts of orbitals in fields of a few volts per nanometer affects chemical bonding within molecules in addition to bonding to the surface. This redistribution of charge can lead either to an enforced or to a diminished bond strength [10]. A dramatic effect of a field-induced weak chemisorption was observed for noble gas atoms. For example, He was found to adsorb on W field emitters at a temperatures of 200 K [164]. The field-free binding energy between He and W, 6 meV, increases to 128 meV at 37 V nm−1 and to 246 meV at 52 V nm−1 , which is just below the field evaporation field strength of W [13]. These field-dependent binding energies were determined from the temperature dependence of ion intensities using pulsed laser field desorption (PLAP) [165]. The equilibrium position (bond length) of adsorbed He atoms was measured precisely within 10−2 nm by RAPS in combination with electron stimulated field desorption (ESFD) [166]. After fixing the image plane by careful calibration measurements, the W−He equilibrium bond length was found to decrease from its field-free value of 0.32 to 0.2 nm at 52 V nm−1 . Thus, the bond length shrinks to approximately two-thirds of its field-free value while increasing the field strength to close to the field evaporation value of W. In cluster calculations with 14 W atoms [168], simulating the W(111) plane, these experimental results were perfectly reproduced, showing that the field-induced weak chemisorption of He is finally due to an effective involvement of the p-orbitals of helium atoms in creating the surface bond.

B Field-Induced Surface Mobility Electric fields can influence the diffusion behavior on surfaces in such a way that the diffusion is changed from a random walk process to a directional walk, in which adatoms move preferentially in directions of increasing field. Work by Wrigley and Ehrlich [169] and Kellogg [170] demonstrated two mechanisms for surface diffusion: the hopping mechanism of individual particles and the exchange mechanism where the diffusing process consists of energetically favored consecutive exchange steps between adparticles and surface sites. On fcc metals this competition is field dependent. At fields of

several volts per nanometer the exchange mechanism is suppressed and the much slower hopping mechanism becomes operative [171]. C Field Dependence of Surface Binding Energies Binding energies of adsorbed surface molecules can be measured using PFDMS. This was demonstrated for various metal–adsorbate systems such as (CN)2 on Pt [172], CO on Fe [110] and NO on either Pt [92, 173] or Rh [174]. A detailed account of the NO binding energies under fieldfree adsorption conditions was given in Section 3.1.3.5.3B. For evaluating the influence of steady (d.c.) electric fields, the principle of operation was to measure the temperature dependence of either mean lifetimes of adsorbed molecules or equilibrium coverages (maintained by adsorption and thermal desorption between pulses) at different d.c. values, FR , ranging from negative to positive, i.e. from the onset of field electron emission to the onset of field desorption/evaporation. For example, it was found that the field free binding energy, E = 103 kJ mol−1 , of NO molecules on a [111] oriented Rh field emitter is in good agreement with TPD data of macroscopic Rh(111) single-crystal planes [175]. This value increased by about 10% in a negative field of −3 V nm−1 and decreased by 15% in a positive field of +3 V nm−1 . This experimental result was analyzed theoretically by applying DFT. Calculating the binding energies in small Rh−NO and Rh2 −NO clusters for different field strengths provided electron density contour maps which were in good agreement with the experimental trends: the Rh−NO bond was weakened in positive fields whereas it was strengthened in negative fields [176]. D Field-Induced Surface Reactions Two examples will be provided where the catalytic decomposition of surface molecules is either field-enhanced as for NO on Pt(111) or field-retarded as for CH3 OH on Rh. The peculiar behavior of metal subcarbonyls in the presence of steady fields will also be described. PFDMS (in the absence of steady fields) of NO from (111) terraces of a Pt field emitter tip at 300 K only yielded molecular NO+ , indicating the occurrence of molecular adsorption of NO as discussed in Section 3.1.3.5.3B. The field ion mass spectra changed dramatically when the reaction field was varied, as illustrated in Fig. 11 [84]. At a reaction field FR = 4 V nm−1 adsorbed NO molecules become unstable and react. Ions such as N2 O+ , O+ 2 and N+ were observed in the course of this field-induced 2 surface reaction. A thorough theoretical analysis by Kreuzer and Wang [177] has explained this behavior. Due to the particular charge transfer with the 2π electronic level of NO, the Pt−NO bond is weakened and tilted. Additional charge in the 2π ∗ antibonding level leads

3.1.3 Structure and Morphology

N2O+ 100

Ions/s

N2+

O+

10−1

NO+ 10−2

0

2

4

6

8

10

12

Steady field strength / V nm−1 Fig. 11 Field-induced decomposition of NO on Pt, measured on a stepped surface with [111] orientation of the terraces. The reaction field FR is varied and the desorption field, FD , adjusted to FD = 24 V nm−1 . pNO = 6.7 × 10−5 Pa, T = 296 K [84].

to dissociation of the NO molecule. At fields larger than 6 V nm−1 the reduction of the activation barrier in the calculated potential energy surface becomes so significant that NO molecules will snap on top of the adsorbed N atoms, leading to the formation of N2 O as observed experimentally. The relation of these experimental and theoretical studies with heterogeneous catalysis is evident: N2 O is a major product during the catalytic NO reduction on many metal-supported catalysts at low temperatures as, for example, in the exhaust aftertreatment of diesel engines. The mechanism of (unwanted) N2 O formation and its possible role as an intermediate of N2 formation on Pt/SiO2 was discussed in a series of papers by Burch and co-workers [178]. Isotopic kinetic transients and mathematical modeling led these authors to suggest dimeric species, (NO)2 , to be intermediates in the formation of N2 O. Interestingly, studies by de Vooys et al. [179] of NO reduction under electrocatalytic conditions at positive potentials likely to cause electric fields similar in magnitude to those present in Fig. 11 gave evidence for a serial pathway of N2 formation from NO via (NO)2 and N2 O formation. Recent PFDMS measurements with Au tips under mere pulsed-field conditions provided direct proof of (NO)2 dimer formation from NO [180]. Field stabilization of reaction intermediates of the CH3 OH decomposition on Rh and Ru prevents further catalytic reaction [181]. This could be observed in PFDMS

891

measurements, which revealed spectra with ions of methanol, CO, hydrogen and intermediates, CHx O (x = 1 − 3). The reaction was seen to decelerate in positive fields: as the ion intensities of the products decreased, those of the intermediate species increased. This stabilization of CHx O species was counteracted by heat supply, i.e. the overall loss of rate was compensated by an increase in temperature [98–100, 182]. PFDMS studies of the interaction of CO with transition metal tips frequently revealed subcarbonyl formation. In the case of Ru [96], tetracarbonyl species occurred in the mass spectra. Their ion intensities were found to decrease on increasing the steady field, suggesting a stabilization of the early stages of the reaction sequence, in particular adsorbed RuCO and Ru(CO)2 . (see also Section 3.1.3.5.3B). A semi-empirical tight-binding approach demonstrated the occurrence of repulsive interactions between CO ligands in Ru(CO)4,ad at fields larger than 4 V nm−1 to be responsible for this effect. Interestingly, subcarbonyl formation on Au was seen to be promoted by steady fields [109]. Singly and doubly charged ions of Au mono- and dicarbonyl occurred with increasing intensities in the mass spectra while raising the steady field strength, whereas higher homologues – as in the case of Ru or other metals – and molecular CO were absent. This scenario suggests that CO adsorption on Au occurs at steps only. DFT calculations with small Au clusters consistently explained the field stabilization effect by a local charge accumulation on respective surface Au atoms, leading to more strongly bound CO and Au−Au bond weakening [109]. 3.1.3.5.5 Field-Induced Phenomena at Extended Surfaces and Interfaces A Electric Fields at Planar Surfaces High electric fields can also occur at adsorbate-covered planar surfaces. For example, adsorbate islands with different work functions produce patch fields along the line where they touch each other. A single isolated charge at an otherwise non-conducting surface creates a local field of ∼10 V nm−1 , which will decay within a few tenths of a nanometer. In the cavities of zeolite structures, with located charges and counter charges >1 nm apart, more extended fields of >10 V nm−1 were calculated by Pickert et al. [183] and determined experimentally by Angell and Schaffer [184] from the field-dependent C−O stretch frequency (see Section 3.1.3.5.2C). At extended conducting metal surfaces the evaluation of long-range electric fields becomes rather complicated, however. Generally, electronic contributions to binding may be covalent or electrostatic. The covalent contribution References see page 892

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depends mainly on the electron occupation and local density of states at the Fermi level. Electrostatic contributions are due to influences of the Coulomb potential. Long-range charge interactions may be of major importance when considering promoter effects in heterogeneous catalysis [185], as discussed below. Chemisorption and reactivity also depend on the geometric structure of the surface. For example, step sites which are present at ∼12–15% at the surface of a field emitter tip usually cause different binding than terrace sites at extended metal surfaces of low Miller index. Considerable attention has been devoted to the question of adsorbate–adsorbate interactions. Hammer and Norskov [186] provided a summary of theoretical results obtained by using DFT. The various contributions to the interaction energy between (positively charged) alkali metal atoms and molecules such as CO and N2 on Ru(0001) were mapped out to reveal the influence of non-local electrostatic interactions [187]. For example, while the binding energy in the end-on configuration of the N2 molecule was only slightly influenced by Na – different from the CO case [188] – the transition state to dissociation was dominated by its presence. Obviously, the considerable dipole moment of N2 in the transition state – the molecule is in a side-on configuration – caused a strong electrostatic interaction with Na+ (attractive since the electric field vectors have opposite signs). Clearly, this scenario seems to provide a proper theoretical description of the promoter effect of alkali metal metals in ammonia synthesis from H2 and N2 on Ru. In addition to chemical promotion, electrochemical promotion may also be used to tune catalytic properties. The non-faradaic electrochemical promotion of catalytic activity (NEMCA) [189] demonstrates how the activity and selectivity of a catalytic system can be modified by controlling the electric potential [190]. B Experimental Examples of Electric Field Effects The interaction of potassium with co-adsorbed carbon monoxide was studied by Uram et al. [38] on an Ni(111) surface using infrared reflection–absorption spectroscopy (IRAS). Due to the electrostatic field perpendicular to the surface, the C−O vibration showed a Stark shift on the order of 10 cm−1 V−1 nm, which corresponded to the estimated field strength of 2 V nm−1 . At low K coverage, besides a very strong local field effect in the vicinity of Kad , a long-range electrostatic field effect was observed such that up to 25 CO molecules were influenced by a single potassium atom. Investigations of interlayer interaction effects such as physisorbed polar molecules in a second layer on top of chemisorbed and saturated CO layers on metal surfaces showed a site interconversion of chemisorbed CO on

Ni(111) [191]. The terminal CO site (νCO = 1961 cm−1 ) experienced a red shift in the vibrational frequency and gradually transferred into a two-fold bridge site (νCO = 1828 cm−1 ) with increasing amounts of added NH3 . This red shift and site interconversion of terminal CO to bridge-bonded CO were, combined with a work function decrease, a reversible process and driven by the electric field. Many other second-layer molecules, such as CH3 I, CH3 Br and CH3 NH2 , produced the same phenomenon to an extent which was determined by the magnitude of work function changes. In a theoretical model calculation it could be shown that the second-layer molecules – independent of their chemical nature – may give rise to electrostatic fields on the order of 1–5 V nm−1 which affect the bonding of CO to the substrate [191]. Electric field effects were held responsible by Wang and Kreuzer [192] to explain isocyanate (–NCO) formation on transition metal catalysts during the reaction of CO and NO [193, 194]. The proposed mechanism included a shift of adsorbed atomic nitrogen from a three-fold to an ontop position of CO with a subsequent rearrangement to an – NCO surface molecule. Local fields were created by Oad atoms produced by NO dissociation. Conclusion The importance of electric’ fields in chemical systems and in heterogeneous catalysis has been outlined. Experimental methods have been described that display ultimate detection limits of individual atoms and molecules in both imaging and local chemical probing. Different applications have been reported, including studies of surface morphology, surface mobility, reaction mechanisms and kinetic and field effects. Relations of these model studies with real-world catalytic systems have been provided. The application of FEM and FIM to image spatio-temporal patterns with nanoscale resolution has opened interesting perspectives which have not been completely explored so far.

3.1.3.5.6

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3.1.3.6

Gamma Spectroscopies

Hans Niemantsverdriet∗ and Tilman Butz

3.1.3.6.1 Introduction M¨ossbauer spectroscopy and time-differential perturbed angular correlation (TDPAC) belong to the class of techniques which detect solid-state properties mediated by hyperfine interactions via nuclear spectroscopy. In contrast to nuclear magnetic resonance which utilizes r.f. fields, excited nuclear states and γ -ray transitions are involved. M¨ossbauer spectroscopy has been widely applied to the study of iron-containing and – in the case of emission spectroscopy – also to cobalt-containing heterogeneous catalysts and to a lesser extent to catalysts containing tin, iridium, antimony and gold. TDPAC has been applied to the study of molybdenum systems to some extent. This References see page 911 ∗ Corresponding author.

896

3.1 Physical Properties

chapter presents some selected examples for both techniques without any attempt to cover the field completely. .. M ossbauer Spectroscopy M¨ossbauer spectroscopy has yielded extremely useful information on a number of important catalysts, such as the iron catalyst for Fischer–Tropsch and ammonia synthesis and the cobalt-molybdenum catalyst for hydrodesulfurization reactions. The great advantage of the technique is that it uses γ -radiation of high penetrating power, such that the technique can be applied in situ. The technique is limited to those elements that exhibit the M¨ossbauer effect of which the most relevant ones for catalysis are iron, tin, iridium, ruthenium, antimony, platinum and gold. Through the M¨ossbauer effect in iron, one can also obtain information on the state of cobalt. M¨ossbauer spectroscopy provides valuable information on oxidation states, magnetic fields, lattice symmetry and lattice vibrations. Several books on M¨ossbauer spectroscopy [1–3] and reviews on the application of the technique on catalysts [4–9] are available, Millet’s review being the most recent one at the time this chapter was written [9]. 3.1.3.6.2

.. A The M ossbauer Effect Suppose we have two identical atoms, one with its nucleus in the excited state and the other with its nucleus in the ground state. The excited nucleus decays to the ground state by emitting a photon with energy of typically some 10–100 keV. This photon falls on the nucleus of the second atom, which is in the ground state. If both nuclei are free, the second nucleus cannot absorb the photon, because in the event the recoil energy is lost twice, namely on emission and on absorption. If, however, the atoms belong to solid lattices, the recoil energy is taken up by vibrations of the lattice as a whole. Lattice vibrations are quantized in phonons and if the recoil energy is smaller than the energy of the lattice vibrations, a fraction of emission and absorption events (called recoil-free fraction) proceeds without exchange of recoil energy. This pure quantum mechanical phenomenon is the M¨ossbauer effect, named after Rudolf L. M¨ossbauer, who discovered it in 1957 and received the Nobel Prize in 1961 [10]. The intensity of the M¨ossbauer effect is determined by the recoil-free fraction or f factor, which can be considered as a kind of efficiency. It is determined by the lattice vibrations of the solid to which the nucleus belongs, the mass of the nucleus and the photon energy, E0 , and is given by

f = e−kγ x 2

2



position due to lattice vibrations. In the Debye model, the latter can be expressed as a function of temperature and a parameter called the Debye temperature, which indicates the stiffness of the lattice [1–6]. A second condition that must be satisfied in order to observe the M¨ossbauer effect is that one needs nuclei in the excited state as a source for the γ -photons. Figure 1 shows how this is achieved in the case of the most frequently used nucleus for M¨ossbauer spectroscopy, the 57 Fe isotope. In brief, one needs a source which decays to the excited state of the nucleus of interest with a sufficiently long life-time such that experiments are practical. The actual transition used for the M¨ossbauer effect should follow instantaneously. .. B M ossbauer Spectroscopy The nucleus is coupled to its environment through hyperfine interactions, causing the nuclear levels in the absorber to shift as a function of oxidation state and to split if electric field gradients or magnetic fields are present in the solid. In order to detect these effects, the energy of the M¨ossbauer source needs to be varied. This is done by using the Doppler effect: if we move the emitter towards the absorber at a velocity v, the energy of the photon becomes  v (2) E(v) = E0 1 + c

where E(v) is the energy of the γ -quantum emitted by the source, v is the velocity of the source, E0 is the energy difference between the excited state and the ground state of the nucleus and c is the velocity of light. In order to detect shifts and splitting in the nuclear levels due to hyperfine interactions in iron, one needs an energy range of at most 5 × 10−8 eV around E0 , which is achieved with Doppler velocities in the range −10 to +10 mm s−1 . 57Co

Electron capture ~0.6 MeV

9%

123 keV 137 keV

I= 3 2

tn = 10−7 s 14.4 keV

14.4 keV 57Fe

(1)

in which f is the recoil-free fraction, kγ is the wavenumber of the γ -radiation, equal to 2π/λ, and x 2 is the mean squared displacement of atoms from their average

91%

I= 1 2

57Fe

Decay of 57 Co to 57 Fe. The encircled part is the transition .. commonly used for Mossbauer spectroscopy of iron-containing samples [6]. Fig. 1

897

3.1.3 Structure and Morphology

Vmin Vmax

Moving source

Absorber Detector

Transmission

E(v) = Eo (1 + v/c)

Eo

Vmin

Va

Vmin

Vmax

Doppler velocity

Ea

E(v) = Ea

Source

Va

Vmax

Moving source

Absorber

In order to cover all possible transitions in the absorbing nucleus, the energy of the source radiation is modulated by using the .. Doppler effect. For 57 Fe the required velocities fall in the range between −1 and +1 cm s−1 . In Mossbauer emission spectroscopy, the sample under investigation is the source and a single line absorber is used to scan the emission spectrum [6].

Fig. 2

Figure 2 gives a schematic picture of a M¨ossbauer experiment in the transmission mode with a moving single-line source and the absorbing sample in a fixed position. A M¨ossbauer spectrum is a plot of the γ -ray intensity transmitted by the sample against the velocity v of the source. The latter is related to the actual energy by Eq. (2). This is the common mode of operation, called M¨ossbauer absorption spectroscopy, sometimes abbreviated to MAS. It is also possible to fix the 57 Cocontaining source and move the single-line 57 Fe absorber, in order to investigate cobalt-containing catalysts. This technique is called M¨ossbauer emission spectroscopy (MES). Three hyperfine interactions couple the nucleus to its surroundings (Fig. 3): (i) The isomer shift, δ, is the consequence of the Coulomb interaction between the positively charged nucleus and the negatively charged s-electrons. Since the size of the nucleus in the excited state differs from that in the ground state, the Coulomb interaction energies are also different. The isomer shift therefore is a measure for the s-electron density at the nucleus and yields useful information on the oxidation state of the iron in the absorber. Isomer shift values are expressed in velocity units, mm s−1 , and are usually given with respect to the peak position of a reference such as metallic iron. Table 1 lists a few isomer shift values of common iron compounds. The isomer shift contains a contribution from the thermal motion of the individual atoms in the absorber, the

Singlet

Quadrupole doublet d = (v1 + v2)/2 ∆E = v2 − v1

d = v1 1

1 2

v1 0

v1 0

v (mm/s)

d = (v1 + v6)/2 H ∝ v6 − v1 1 23 456

v2

1 23 456

v3 v4 0

v (mm/s)

Magnetic + Quadrupole interaction

Magnetic sextuplet

v1

v2

v5

v6

v / mm s−1

d = (v1 + v2 + v5 + v6)/4 H ∝ v6 − v1 s = (v6 − v5 − v2 + v1)/4

v1 v2

v3 v4 0

v5

v6

v / mm s−1

.. The four most common types of Mossbauer spectra observed in iron-containing catalysts along with the corresponding .. nuclear transitions. Indicated also is how the Mossbauer parameters are derived from the spectra. (From Niemantsverdriet [6].)

Fig. 3

References see page 911

898

3.1 Physical Properties

Tab. 1

.. Mossbauer parameters of common iron compounds

Compound α-Fe2 O3 α-FeOOH Fe3 O4 FeO FeS2 α-Fe (=bcc Fe) θ-Fe3 C SNPa

δ/mm s−1 0.43 0.35 0.30 0.63 1.08 0.28 0.00 0.19 −0.26

EQ /mm s−1 – – – – 0.55 0.60 – – 1.70

ε /mm s−1 −0.10 0.13 – – –

H/T 51.5 38.4 49.2 45.5 – – 33.0 20.8 –

a Sodium

nitroprusside (sometimes used as calibration standard for the isomer shift; commonly metallic iron is used).

second-order Doppler shift, which makes the isomer shift temperature dependent [1–4, 6]. The electric quadrupole splitting, EQ , is caused by the interaction of the electric quadrupole moment of the nucleus with an electric field gradient. The nucleus of iron in the excited state has the shape of an ellipsoid and possesses an electric quadrupole moment. The consequence is that the nucleus can orient itself in two ways in an electric field gradient with slightly different energies and the M¨ossbauer spectrum consists of a doublet (Fig. 3). The origin of the electric field gradient is two-fold: it is caused by asymmetrically distributed electrons in incompletely filled shells of the atom itself and by charges on neighboring ions. The distinction is not always clear, because the lattice symmetry determines the direction of the bonding orbitals in which the valence electrons reside. If the symmetry of the electrons is cubic, the electric field gradient vanishes. (ii) Magnetic hyperfine splitting, the Zeeman effect, arises from the interaction between the nuclear magnetic dipole moment and the magnetic field H at the nucleus. This interaction gives rise to six transitions; the separation between the peaks in the spectrum is proportional to the magnetic field at the nucleus. Of course, all hyperfine interactions can occur simultaneously. In magnetically ordered compounds with a non-vanishing electric field gradient, the shape of the spectrum depends on the relative strengths of the magnetic and the electric quadrupole interaction. In catalysts, the usual situation is that the quadrupole interaction is much smaller than the magnetic interaction. In this case all peaks shift by an energy ε (expressed in mm s−1 ); however, the outer lines shift in the opposite direction as the inner four lines do (see Fig. 3). Examples of this situation are observed in the M¨ossbauer spectra of Fe2 O3 and FeOOH.

The intensity of a M¨ossbauer spectrum depends not only on the recoil-free fractions of the source and the absorber and on the number of absorbing nuclei, but also on the linewidth of the absorption lines and on whether or not saturation effects occur [1–4, 6]. Particularly in the M¨ossbauer spectra of small catalyst particles, one should be aware of the temperature dependence of the absorption area through the recoil-free fraction. If the spectrum contains contributions from surface and bulk phases, the intensity of the former will be greatly underestimated if the spectrum is measured at room temperature [11]. The only way to obtain reliable concentrations of surface and bulk phases is to determine their spectral contributions as a function of temperature and make an extrapolation to 0 K [12]. .. C M ossbauer Spectroscopy in Catalyst Characterization Most applications of M¨ossbauer spectroscopy to catalysts fall into one of the following categories:

• • • • •

identification of phases determination of oxidation states structure information determination of particle size kinetics of bulk transformations.

A typical example of how M¨ossbauer spectroscopy is used in the identification of oxidic, metallic and carbidic phases is provided by a study on titania-supported iron, prepared by impregnating the TiO2 support with a solution of iron nitrate (see Fig. 4) [13]. The top spectrum is that of a freshly impregnated and dried Fe−TiO2 catalyst. It shows a doublet with an isomer shift of 0.37 mm s−1 and a quadrupole splitting of 0.82 mm s−1 , characteristic of a high-spin Fe3+ species. It is difficult to draw conclusions on the type of compound that is present. Reference compounds of iron oxide, Fe2 O3 or Fe3 O4 , possess magnetically split M¨ossbauer spectra (see, for example, the M¨ossbauer parameters of reference compounds in Table 1). The relatively high value of the quadrupole splitting points to a highly asymmetric environment, as surface atoms have. After reduction in H2 at 675 K the catalyst consists mainly of metallic iron, as evidenced by the sextet (δ = 0 mm s−1 , H = 33 T), along with some unreduced iron, which gives rise to two doublet contributions of Fe2+ and Fe3+ in the center. The overall degree of iron reduction, as reflected by the relative area under the bcc iron sextet, is high. One should not consider the relative spectral contributions as concentrations, however, because the three types of iron species may have different recoil-free fractions. When reduced Fe−TiO2 is used as a catalyst for the reaction between CO and H2 to hydrocarbons

3.1.3 Structure and Morphology

Mössbauer spectroscopy of Fe/TiO2 T = 300 K

Fresh catalyst

Reduced in Hz 400 K, 1 h

in Hz 675 K, 18 h After FTS 575 K, 6 h

Exposed to air 300 K −10 −8 −6 −4 −2 0 2 4 6 8 10 Doppler velocity (mm/s)

.. Fig. 4 Mossbauer spectra at room temperature give detailed information on the state of iron in a TiO2 -supported iron catalyst after different treatments: 18 h of reduction in H2 at 675 K, Fischer–Tropsch synthesis and exposure to air. (From van der Kraan et al. [13].)

(the Fischer–Tropsch synthesis) the spectrum changes entirely. All metallic iron converts into an iron carbide, in this case the H¨agg carbide or χ-Fe5 C2 . Apparently the strongly reducing atmosphere has also affected the unreduced iron: all ions are now present as Fe2+ . The bottom spectrum was recorded after exposing the used catalyst to air at room temperature. The spectrum clearly has changed. Although most of the carbide phase is still present, some of the ferrous iron has been oxidized to ferric iron. Hence it is essential that the catalyst be studied under in situ conditions. The conversion of iron catalysts under Fischer–Tropsch conditions has been the subject of several M¨ossbauer studies [14–17]. The example illustrates how M¨ossbauer spectroscopy reveals the identity of iron phases in a catalyst after different treatments. It is typical for many applications of the technique in catalysis: a catalyst is reduced, carburized, sulfided or passivated; after cooling, its M¨ossbauer spectrum is taken at room temperature. However, a complete characterization of phases in a catalyst sometimes requires that spectra are measured at cryogenic temperatures, in particular when catalysts

899

are highly dispersed. Note that limiting oneself to the rules of the nowadays fashionable ‘‘operando’’ community, who specify the acquisition of spectra under catalytic reaction conditions with simultaneous measurements of catalytic conversions, implies that most of the information that M¨ossbauer spectroscopy provides will be lost: spectral intensity degrades severely due to the temperature dependence of the recoil-free fraction (surface phases become virtually invisible) and magnetically split and therefore very informative patterns collapse if the temperature surpasses the Curie point of the superparamagnetic transition temperature in the spectra of small magnetically ordered particles. .. D In Situ M ossbauer Spectroscopy at Cryogenic Temperatures Surface phases have low Debye temperatures. As a result, the recoil-free fraction may be low at room temperature. Thus, measuring at cryogenic temperatures will increase the M¨ossbauer intensity of such samples considerably. However, there are also other reasons which call for low-temperature experiments. An example is given in Fig. 5. The spectrum of a reduced carbon-supported iron catalyst at room temperature shows only a broad single line contribution in the central region of the spectrum, which by itself is difficult to interpret. Spectra of the same sample measured at lower temperatures and/or in applied magnetic fields give much more information, as Fig. 5 shows. The spectrum obtained at liquid nitrogen temperature in a 1-T field allows for a clear interpretation: the majority of the iron gives rise to a magnetic sextet with a magnetic hyperfine field characteristic of bcc iron [18]. Thus, the spectrum at room temperature contains a singlet of superparamagnetic iron, implying that the particles are so small that thermal excitations of energy kT are energetic enough to decouple the magnetization from the lattice. As a result, the magnetization vector of each particle fluctuates rapidly over all directions and the M¨ossbauer transition, which takes place on a time-scale of 10−8 s, feels an average magnetization of zero. The occurrence of superparamagnetism allows one to determine the particle size if an external magnetic field is applied. Figure 5 illustrates how it works. The spectra without an external field show the single peak of superparamagnetic iron. As soon as the magnetic field is applied, magnetic splitting sets in, and its magnitude increases with increasing field strength. The external field orients the magnetization vector of the particles, but thermal excitations let the magnetization vector fluctuate around the direction of the applied field. One measures an average magnetic splitting, given by the Langevin References see page 911

3.1 Physical Properties

Absorption

1.03 T

1.03 T

0.86 T

0.86 T

0.72 T

0.72 T

35

80 K

30 0.55 T

0.55 T

0.01 T

0.01 T

Bobs − B/ T

900

25

20 300 K 80 K

300 K

−8−6−4−2 0 2 4 6 8

15

−8−6−4−2 0 2 4 6 8

Velocity / mm

s−1

1.0

2.0

3.0

B−1 / T−1

(a)

(b)

.. Fig. 5 (a) Mossbauer spectra of a reduced carbon-supported iron catalyst at 80 and 300 K obtained in different applied magnetic fields. The spectra at the bottom, measured without external field, consist mainly of a singlet due to superparamagnetic metallic iron. The application of magnetic fields induces magnetic splitting. (b) Langevin plots according to Eq. (3) for the spectra. The lines extrapolate for 1/H → 0 to the magnetic splitting expected for single domain metallic iron particles; the slopes correspond to a particle diameter of 2.5 ± 0.2 nm. The dashed line is a plot according to the complete Langevin equation and confirms that the use of the high-field approximation is justified. (From Christensen et al. [18].)

equation [19]: Hobs − Hext


kT = H0 1 − µH

for

µH ≥3 kT

(3)

where Hobs is the observed magnetic splitting, Hext the externally applied field [both are vectors, hence Eq. (3) contains the magnitude of their vectorial difference], H0 the bulk magnetic field, µ the magnetic moment of the particle, k Boltzmann’s constant and T the absolute temperature. Equation (3) represents a simplification which is often applicable in practice. Figure 5b illustrates the analysis of the magnetic hyperfine splitting with Eq. (3). A plot of |Hobs − Hext | against 1/Hext gives a straight line with a slope H0 kT /µ, from which µ, the magnetic moment of the particle, follows. As the total magnetic moment is the atomic moment (2.2 Bohr magneton) times the number of atoms in a particle, the latter can thus be calculated and converted into a diameter if we assume that the particles have spherical shape. In this way we find a diameter of 2.5 ± 0.2 nm for the iron particles in Fig. 5 [18]. The usual techniques for the determination of particle sizes of catalysts are electron microscopy, chemisorption, XRD line broadening or profile analysis and magnetic

measurements. The advantage of using M¨ossbauer spectroscopy for this purpose is that one simultaneously characterizes the state of the catalyst. As the state of supported iron catalysts often depends on subtleties in the reduction, the simultaneous determination of particle size and degree of reduction as in the experiments of Fig. 5 is an important advantage of M¨ossbauer spectroscopy. E Kinetics of Solid-State Reactions from Single Velocity Experiments Recording a complete M¨ossbauer spectrum of an iron catalyst typically takes several hours, which is too slow for following reactions in real time. Nevertheless, measuring the intensity of a characteristic peak at constant velocity can monitor processes occurring on the time-scale of minutes to hours, such as the reduction of oxides. Elegant examples have been reported by Delgass and co-workers [15, 20], who studied the conversion of iron and iron nitrides into carbides under Fischer–Tropsch-like conditions. Alternative ways to study the kinetics of bulk transformations would be to monitor changes in weight or in magnetization. Such methods, however, are less specific about the initial and final state of the catalyst than is M¨ossbauer spectroscopy.

3.1.3 Structure and Morphology

.. F In Situ M ossbauer Spectroscopy Under Reaction Conditions In situ characterization becomes an absolute necessity in cases where catalysts change their structure during the start-up of the catalytic reaction. Figure 6 shows the example of a bimetallic Fe−Ir catalyst during synthesis of methanol from CO and H2 . Noble metals such as platinum and iridium are poor CO hydrogenation catalysts, producing mainly methane. Addition of iron, however, increases the activity and shifts the product distribution towards methanol. However, a catalyst containing equal amounts of iron and iridium on silica starts to convert CO and H2 almost exclusively into methane. It typically takes a full day before the product distribution has changed to methanol [21]. What happens with the Fe−Ir catalyst during this activation period? M¨ossbauer spectra taken in situ under high-pressure reaction conditions (Fig. 6) show that the initially reduced Fe−Ir catalyst, consisting of Fe−Ir particles and some unreduced iron believed to be in intimate contact with the support, has changed significantly when the catalyst has reached steady state [7]. The bottom trace in Fig. 6 represents the difference between the working and the initial catalyst. It is characteristic of an iron carbide in the superparamagnetic state. Also, the total absorption area of the spectrum increased, indicating loss of overall dispersion of the iron. Further characterization studies In situ Mössbauer spectroscopy 1:1 FeIr/SiO2 at 523 K

% Transmission

100

After reduction

100 During reaction 48 – 72 h 92 100 Difference spectrum −4

−2

0

2

4

Doppler velocity / mm s−1

.. In situ Mossbauer spectra of a reduced FeIr–SiO2 catalyst at a reaction temperature of 525 K and during CO hydrogenation when the catalyst is in its steady-state methanol-producing state. The bottom spectrum represents the difference between the two upper spectra; it is characteristic of an iron carbide in the superparamagnetic state. (Adapted from Niemantsverdriet and Delgass [7].)

Fig. 6

by M¨ossbauer spectroscopy and EXAFS after cooling to ambient and cryogenic temperatures confirmed the interpretation and added much detail to it concerning the composition of the surface [21]. In brief, during high-pressure CO hydrogenation, the active part of the Fe−Ir catalyst restructures from Fe−Ir alloy particles with a surface enriched in iron to an iridium-rich alloy accompanied by a, probably largely inactive, iron carbide phase. During the reconstruction, the chemical properties of the catalyst surface also change significantly. The interesting notion of the work is that the Fe−Ir catalyst restructures itself by letting excess iron segregate into carbide particles, leading to an FeIr alloy that is well tuned towards methanol formation. .. G M ossbauer Emission Spectroscopy of CobaltContaining Catalysts Catalysts based on molybdenum disulfide, MoS2 , and cobalt or nickel as promoters are used for the hydrodesulfurization (HDS) and hydrodenitrogenation (HDN) of heavy oil fractions [22]. The most direct information on the state of cobalt has come from M¨ossbauer spectroscopy, applied in the emission mode. Such experiments are done with catalysts that contain the radioactive isotope 57 Co as the source and a moving single-line absorber. As the M¨ossbauer spectrum is strictly not that of cobalt, but that of its decay product, iron, the safest approach is therefore to compare the spectra of the Co−Mo catalysts with those of model compounds for which the state of cobalt is known. Figure 7 shows M¨ossbauer spectra of a series of sulfided Co−Mo catalysts as reported by Wivel et al. [23]. The spectra contain essentially three contributions, as indicated by bar diagrams:

(i) catalytically insignificant Co ions inside the Al2 O3 support (ii) the bulk sulfide Co9 S8 , dominant in catalysts of high cobalt content (iii) a formerly unknown state, labeled Co−Mo−S, which is most evident in sulfided Co−Mo catalysts of low cobalt content.

92

98

901

The interesting point about this Co−Mo−S phase is that its presence correlates with the catalytic activity for the desulfurization reaction. Hence this Co−Mo−S phase is either active itself or at least closely associated with the active site. As to the structure of this Co−Mo−S phase, it is important to note that MoS2 has a layer structure of slabs consisting of Mo4+ sandwiched between two layers of S2− ions. As the chemical reactivity of MoS2 is associated with the edges of the slab, it is reasonable to assume References see page 911

3.1 Physical Properties

Co-Al2O3 Co-Mo-S Co9So6

Absorption

Co Mo

Co-Mo-S A

0.09

A

0.27

B

0.53

C

0.75

D

B

Absorption

902

C E

1.19

−6

−4

−2

0

Velocity (mm

2 s−1

4

6

−6

−2

0

2

4

6

−1

)

(a)

−4

Velocity (mm s ) (b)

.. In situ Mossbauer emission spectra of 57 Co: (a) a series of sulfided Co−Mo/Al2 O3 catalysts; (b) MoS2 particles doped with different amounts of cobalt, corresponding to Co/Mo ratios of (A) about 3, (B) 0.05 and (C) 0.25 ppm. The Co−Mo−S phase, active in the HDS reaction, has a spectrum unlike that of any bulk cobalt sulfide and is most clearly observed in the spectra of Co−Mo/Al2 O3 catalysts of low Co content and in the MoS2 particles doped with ppm levels of cobalt. (Adapted from Wivel et al. [23] and Topsøe et al. [24].)

Fig. 7

that edges also form the seat of the catalytic activity. A model system, prepared by impregnating an MoS2 crystal with a dilute solution of cobalt ions, such that the model contains only parts per million of cobalt, appears to have the same M¨ossbauer spectrum as the Co−Mo−S phase (see Fig. 7b) [24]. Scanning Auger spectroscopy and electron microscopy both confirm that cobalt is indeed found around the edge regions. Titration studies using infrared spectroscopy of adsorbed NO and CO, and also structure investigations by EXAFS and theoretical calculations, have lent further support to the structural model for the Co−Mo−S phase. We refer to Topsoe et al. for a review [22]. Recently, the Delft group has reported 57 Co M¨ ossbauer work on Co−Mo catalysts on carbon [25] and alumina [26] under high [i.e. 40 bar (1 bar = 105 Pa)] pressure conditions. Van Berge et al. used M¨ossbauer emission spectroscopy to investigate cobalt Fischer–Tropsch catalysts with respect to stability versus oxidation in water, which has been suggested as a cause of deactivation. Oxidation would imply loss of reduced cobalt and hence loss of activity. The authors found that oxidation can indeed occur, depending on the partial pressure of water in the environment [27].

.. H M ossbauer Spectroscopy of Other Elements Although the majority of M¨ossbauer studies on catalysts deal with iron, other elements exhibiting the M¨ossbauer effect have also been used. Bussiere and coworkers used the M¨ossbauer effect to study the state of tin in supported Pt−Sn [28] and Ir−Sn [29] reforming catalysts and of tin and antimony in mixed Sb−Sn oxides for the selective oxidation of propane [30]. Well known are the investigations by Clausen and Good on supported ruthenium catalysts by means of the difficult 99 Ru isotope [31]. A more recent 99 Ru study was reported by Wagner’s group [32]. Noteworthy is the use of the 125 Te isotope to characterize the state of the tellurium promoter in multicomponent ammoxidation catalysts [33]. Iridium M¨ossbauer spectroscopy has also been applied with success to study supported Ir, Pt−Ir and Fe−Ir alloy catalysts [34–36]. Such experiments are much less straightforward than those with iron: because the 192 Os source decays with a half-life of only 31 h, one can do only a few experiments with one source. Access to a nuclear reactor facility for reactivating the source is essential for doing Ir M¨ossbauer spectroscopy. Moreover, the high energy of the transition, 73 keV, implies that the recoil energy is high and the recoil-free fraction low [Eq. (1)].

3.1.3 Structure and Morphology

Hence the absorber and source need to be cooled to low temperatures, preferably that of liquid helium. Similar considerations hold in the case of gold M¨ossbauer spectroscopy. Nevertheless, with the recent interest in small particles of gold in catalysis, some studies have been carried out, and we review one of these. Supported gold particles a few nanometers in size have been found to display remarkable activity in, for instance, selective and total oxidation reactions. The gold isotope 197 Au exhibits the M¨ ossbauer effect and provides clear information on the oxidation state of the element. Kobayashi et al. [37] reported the evolution of gold species during calcination of Mg(OH)2 -supported Au, prepared by soaking an MgO support with an aqueous solution of HAuCl4 . During the impregnation, the gold deposits as Au(OH)3 and the MgO transforms to Mg(OH)2 . The corresponding M¨ossbauer spectrum in Fig. 8 shows two doublets of Au3+ species. Upon calcination, the Au3+ converts gradually to Au+ and Au0 . The example nicely demonstrates the rich information content of the technique for characterizing gold catalysts. Another interesting example, by Finch et al. [38], describes the application of both 57 Fe and 197 Au M¨ ossbauer spectroscopy on Au−Fe2 O3 catalysts,

AuI AuIII

Au0

AuIII

Transminssion

Before calcination

Component 4 Component 3

Component 2 Component 1

473 K

523 K

573 K

−10

−5

0 Velocity (mm

5

10

s−1)

.. Fig. 8 197 Au Mossbauer spectra of freshly prepared and dried Au catalysts supported on Mg(OH)2 and the samples after calcination at the temperatures indicated. The spectra show that the initially present Au3+ phases in the dried catalyst are converted to Au+ and Au0 upon calcination. (Adapted from Kobayashi et al. [37].)

903

and the same group also reported M¨ossbauer spectra of Au−MgO oxidation catalysts [39]. I Conclusion M¨ossbauer spectroscopy has matured into one of the classical techniques for catalyst characterization, although its application is limited to a relatively small number of elements which exhibit the M¨ossbauer effect. The technique is used to identify phases, determine oxidation states and follow the kinetics of bulk reactions. M¨ossbauer spectra of superparamagnetic iron particles in applied magnetic fields can be used to determine particle sizes. In favorable cases the technique gives information on the structure of catalysts. The great advantage of M¨ossbauer spectroscopy is that its highenergy photons permit measurements on catalysts under in situ conditions. The interested reader is referred to Millet’s review [9] for an overview of new developments in M¨ossbauer spectroscopy which may find application in the field of catalyst characterization in the future. 3.1.3.6.3 Time-Differential Perturbed Angular Correlations (TDPAC) The time-differential observation of the perturbed angular correlation of γ -rays emitted from radioactive nuclei (TDPAC) is another γ -spectroscopic technique which allows the determination of hyperfine interactions such as nuclear magnetic dipole and nuclear electric quadrupole interactions, like the M¨ossbauer effect. However, in contrast to the M¨ossbauer effect, this technique does not allow the determination of the isomer shift (and therefore does not readily yield information on valence states), nor does the size of the effect depend on lattice vibrations (or temperature) via a Lamb–M¨ossbauer factor. Its main advantage is that it can be applied equally well at any temperature and that it requires only about 1010 atoms. Frequently, γ -rays with energies exceeding 100 keV are involved, which renders this technique applicable to in situ studies with practically any sample environment, such as high-pressure or corrosive gases. In the following, a short introduction to TDPAC is given with particular emphasis on spectroscopy utilizing 99 Mo (t 1/2 = 66 h). This isotope is not ideally suited for TDPAC, but it provides access to the Mo site in Mo-based catalysts. Nuclear quadrupole interactions (NQIs) will be discussed exclusively because they serve as fingerprints in the following examples:

(i) frozen molybdate solutions, illustrating the discrimination between monomolybdates and polymolybdates (ii) γ -Al2 O3 impregnated with ammonium heptamolybdate (APM) solutions, addressing the question of the degree of condensation of surface species References see page 911

904

3.1 Physical Properties

Randomly oriented nuclear spins

g1

ORIENTED NUCLEAR SPINS

g1

DET

SELECTION OF AN ALIGNED SUBENSEMBLE OF NUCLEAR SPINS BY VIRTUE OF ANGULAR MOMENTUM CONSERVATION BECAUSE OF THE DETECTION OF g 1

THEREFORE THE EMISSION OF A SUBSEQUENT g 2 FROM THE SAME NUCLEUS IS ANISOTROPIC

COINCIDENCE COUNTRATE OF ORIENTED NUCLEAR SPINS

|I, m 〈

g 1 AND g 2 IS ANISOTROPIC

g1

depends on: SPIN SEQUENCE MULTIPOLARITIES OF g ′s + MIXING RATIOS

|I, m 〈

t DET 1

COINCIDENCE

g2

g1

DET 2

g2 |I, mt 〈

COINCIDENCE

COINCIDENCE

Θ = 90°

Hyperfine interaction leads to nuclear spin precession

Θ = 180°

4 3

Θ = 180°

2 1 0

6 5 4 3

Θ = 190°

2 1 0

COINCIDENCE COUNTS

6 5

COINCIDENCE COUNTS

COINCIDENCE COUNTS

Coincidence countrate depends on the angle Θ between the detectors 6 5 4 3 2 1 0

0 1 2 3 4 5 6 7 8 9

0 1 2 3 4 5 6 7 8 9

0 1 2 3 4 5 6 7 8 9

TIME

TIME

TIME

Anisotropy W (Θ, t ) ~ e−t/t [1+A 22 G22(wt) P 22(cos Θ)]

w: precession frequency

Perturbation function

Principles of γ –γ -TDPAC: (top) illustration of angular correlation; (bottom) schematic experimental set-up for unperturbed (left) and perturbed angular correlations.

Fig. 9

(iii) various forms of MoS2 : single crystals, bulk powders, highly disperse powders, exfoliated, restacked and surface species on γ -Al2 O3 supports. A TDPAC: Theory, Experimental Setup and Data Reduction γ –γ -TDPAC requires excited nuclei which decay via the successive emission of two γ -rays, called a γ –γ cascade (see Fig. 9). Whereas the emission probability of each

of the two γ -quanta individually is isotropic – assuming unoriented nuclear spins – the coincidence count rate between both quanta in general depends on the angle between the two emission directions, i.e. it is anisotropic. The origin of this angular correlation is based on the concept of conservation of angular momentum: the vectors describing the magnitude and orientation of the nuclear spins before and after the emission of the

3.1.3 Structure and Morphology

γ -quantum and the vector describing the direction of emission and the magnitude of the angular momentum ‘‘carried away’’ by the γ -quantum must add up to zero or obey the ‘‘triangle rule’’. As a consequence, the mere observation of a γ -quantum in a certain direction selects a sub-ensemble out of the randomly oriented nuclei because not all nuclei emit in this given direction with the same probability. This sub-ensemble is aligned. In order to profit from this alignment, it suffices to detect the emission of a subsequent γ -quantum, of course in coincidence with the preceding quantum whose detection created the aligned sub-ensemble. This is illustrated schematically in Fig. 9. The coincidence count rate W (θ) can be written as W (θ) = 1 + A2 P2 (cos θ) + · · ·

(4)

where the anisotropy A2 (=size of the angular correlation effect) depends on nuclear decay properties only (nuclear spins involved, multipolarities of the γ -transitions and mixing ratios). It is a constant characteristic of a given cascade. P2 (cos θ) is a Legendre polynomial. The dots indicate that higher terms in the multipole expansion such as A4 P2 (cos θ) can occur. However, they are not relevant for what follows. The connection between the nuclear system and the environment in condensed matter is established via the hyperfine interaction of the intermediate excited state. In order to be efficient, this intermediate state has to be sufficiently long-lived, i.e. it should possess a half-life t1/2 in the range from about 1 ns up to about 1 µs. Shorter half-lives pose severe time resolution problems (apart from the fact that the interaction or ‘‘exploration’’ time is short), whereas larger half-lives render the fulfillment of a ‘‘true’’ coincidence condition from the same nucleus more and more difficult. In what follows, the time-differential observation of the coincidence between γ1 and the delayed γ2 is most appropriate. To accomplish this, γ1 is used to start a ‘‘nanosecond clock’’, whereas γ2 stops it. Every delayed coincidence event is stored in a multichannel analyzer 99

Mo

905

under an address proportional to the time elapsed between γ1 and γ2 . In this way, a time histogram is recorded which will exhibit an exponential decay with the lifetime of the intermediate excited state, i.e. we perform a lifetime experiment. If the nucleus while being in the intermediate excited state interacts with its environment via its nuclear moments, the nuclear spin will precess and therefore ‘‘perturb’’ the angular correlation. This spin precession modulates the exponential decay curve in a characteristic manner which is expressed in the form of a perturbation function G2 (t): W (θ, t) ≈ e−t/τN [1 + A2 G2 (t)P2 (cos θ ) + · · ·]

(5)

This perturbation function contains all relevant information about the hyperfine interaction. In order to be more specific, let us turn to the decay of 99 Mo. The simplified decay scheme is shown in Fig. 10a. 99 Mo decays via β-emission to the start level of the γ –γ cascade in the daughter isotope 99 Tc. This state decays via the emission of a 740-keV γ -quantum to the intermediate excited state with I = 5/2 and τN = 5.39 ns. The de-excitation from this state proceeds in two ways: either a direct transition to the ground state (181-keV emission) or via a 40–141keV cascade. Both cascades are stored simultaneously: the 740–181 cascade and the 740–40–141 triple cascade with the 40-keV transition remaining unobserved. The anisotropy of the first cascade is A2 ≈ +0.10 and of the triple cascade A2 ≈ −0.11. Hence they have to be stored separately and combined subsequently [40]. The intermediate I = 5/2 state splits into three sublevels in the presence of a nuclear quadrupole interaction (see Fig. 10b). Hence the perturbation function G2 (t) contains a superposition of cosine terms with precession frequencies corresponding to the energy differences between the hyperfine split sublevels. Typical examples for relevant perturbation functions are given in Fig. 11. Figure 11a illustrates the case of no interaction (=unperturbed). Figure 11b is typical for liquids in the References see page 911

T ½ = 66 h b− 99Tc

740 keV

EXCITED STATE

740 keV 181 keV I= 181 keV (a)

40 keV 141 keV

5

2

2hwo 3hwo Typically 1 µeV hwo

τN = 5.39 ns 99

Tc GROUNDSTATE (b)

(a) Simplified decay scheme of 99 Mo. (b) Quadrupole splitting of intermediate excited nuclear state, assuming axial symmetry. (Reproduced with permission from Ref. [45].)

Fig. 10

906

3.1 Physical Properties

PURE QUADRUPOLE - INTERACTION FOR l = 5/2 (a)

G2 (t) = 1 4

1

0 0 G2 (t) = exp (–lt)

(b)

l = 1/15 ns−1

1

1

0 (c)

a0 = 0.200 a1 = 0.371 a2 = 0.286 a3 = 0.143

1 h=0 0

1

w = 126 Mrad s−1 0 3 G2 (t) = Σ an cos w n t n=0

(d)

a0 = 0.245 a1 = 0.317 a2 = 0.286 a3 = 0.150

h = 0.6

1

0

1

w1 = 149 Mrad s−1 3 G2 (t) = Σ an exp (–ndwt ) ⋅ cos nwt n=0

(e) 1 h=0

d = 01

0

a0 = 0.200 a1 = 0.371 a2 = 0.286 a3 = 0.143

3 G2 (t) = Σ an exp (–nΓ t ) n=0

a0 = 0.200 a1 = 0.371 a2 = 0.286 a3 = 0.143

1 Γ = 66.7 MRad s−1 0

0

1

w = 126 Mrad s−1 (f)

INTENSITY I (w) [ARBITRARY UNITS]

PERTURBATION FUNCTION G 2 (t)

0 3 G2 (t) = Σ an cos nwt n=0

0

1

0 0

50

100

TIME/ns

0

0.5

1.0

FREQUENCY/GRad s−1

Fig. 11 Typical powder perturbation functions for I = 5/2 and pure quadrupole interaction (left) and their Fourier transforms (right). (Reproduced with permission from Ref. [45].)

fast relaxation regime (here, the direction of the NQI fluctuates so rapidly, that the nucleus has no chance to perform a precession period or a fraction thereof and, hence, the anisotropy decays exponentially and is totally lost after some time). Figure 11c illustrates the case for a static, axially symmetric NQI, whereas Fig. 11d is an example of non-axial symmetry. On the right side in

Fig. 11 the Fourier-transformed perturbation functions are shown: (i) a ‘‘delta function’’ with instrumental linewidth at zero frequency (ii) a Lorentzian (iii) a set of discrete equidistant frequency peaks (iv) a set of discrete, non-equidistant frequency peaks.

3.1.3 Structure and Morphology

This Fourier representation is very useful and is closely related to what would be observed in a nuclear quadrupole resonance experiment. Figure 11e illustrates the effect of a distribution (here assumed Lorentzian) of NQIs, leading to damping in the time domain or to line broadening in the frequency domain. Finally, Fig. 11f shows the perturbation function for a static NQI centered around zero frequency with a Lorentzian frequency distribution of half-width . Note that the anisotropy decays in a multiexponential fashion towards a limiting, non-zero value (called hard core), in contrast to the dynamic relaxation case (ii). The experimental set-up usually consists of more than two detectors in order to improve the efficiency. A standard set-up uses four detectors arranged in a plane at 90◦ intervals. Using each detector as start and stop, 12 coincidences can be stored simultaneously: four combinations with θ = 180◦ and eight combinations with θ = 90◦ . For technical reasons, four out of the eight possible 90◦ coincidences are often discarded. They are combined in such a way that the exponential decay factors and the detector efficiencies cancel and a ‘‘time spectrum’’ A2 G2 (t) is obtained: √ √ 4 W13 W31 W24 W42 − 4 W14 W41 W23 W32 A2 G2 (t) = 2 √ √ 4 W13 W31 W24 W42 + 2 4 W14 W41 W23 W32 (6) where Wij denotes the coincidence count rate between detector i (accepting γ1 ) and detector j (accepting γ2 ) and the labeling of detectors is clockwise (or counterclockwise). The data analysis is performed in two steps. First, the Fourier transform of an experimental time spectrum is performed. This allows the extraction of a first guess of parameters. which are then used for a subsequent least-squares fit analysis with perturbation functions of the type described in Fig. 11 or superpositions thereof. Further information on the theory of TDPAC is given in Ref. [41], and Ref. [42] provides a review with special emphasis on applications in the borderline fields between physics, chemistry and biology. Three examples are now discussed. B Frozen Molybdate Solutions Molybdate solutions at 300 K exhibit weakly perturbed TDPAC spectra of the type in Fig. 11b with very small decay constants λ, irrespective of the pH (i.e. the degree of condensation), due to their rapid tumbling motion. Hence, measurements in the liquid state are not suitable to discriminate between different Mo species. In contrast, frozen solutions (shock-frozen to 77 K) clearly reveal the degree of condensation (see Fig. 12). At high pH, monomolybdates are formed [43], which exhibit a very low NQI. This is a result of the tetrahedral coordination of Mo, which would

907

yield a vanishing NQI for a perfect tetrahedron. Apparently, the MoO2− 4 tetrahedra are only weakly distorted upon freezing. In contrast, condensed Mo−O octahedra are formed at low pH (especially hepta- and octamolybdates). A perfect octahedral environment of Mo would also yield a vanishing NQI. However, upon polymerization the octahedra are strongly distorted leading to a sizeable NQI (see Fig. 12, bottom). Interestingly, the TDPAC spectrum for a frozen heptamolybdate solution is very similar to that of crystalline ammonium heptamolybdate (APM) [44, 45]. In order to produce 99 Mo-labeled solid APM, the solid was neutron irradiated. The radiation damage leads to color center formation in the solid. Interestingly, these color centers are stable when APM is dissolved [46]. A deep-blue solution is obtained which requires stirring under air in order to become colorless. Most likely an oxygen atom is displaced by the prompt γ de-excitation following neutron capture thus forming an oxygen vacancy. It is conceivable that an Mo7 O6− 24 species in solution (possibly solvated) can be stable with an oxygen missing, provided that it is not a peripheral one. Upon depolymerization at pH > 6, the blue color disappears instantaneously. This ‘‘color indicator’’ proved to be very useful for the impregnation studies which will be discussed next. C γ -Al2 O3 Impregnated with APM Solutions The first preparation step for Mo-based heterogeneous supported catalysts is usually the impregnation of a suitable support having a high surface area with an APM solution. We used Al2 O3 PURAL SB from CONDEA with a BET surface area of about 150 m2 g−1 . In the ‘‘pore filling method’’, the amount of solution is kept so low that on the one hand all pores are filled, but on the other hand no supernatant liquid remains. Here, it is particularly difficult to determine the degree of adsorption. Moreover, the pH in the pores certainly shifts upwards during adsorption and it is not clear whether depolymerization takes place prior to adsorption or even in a slow equilibration process between adsorbed polymolybdates and the pore solution at pH > 6.5. Here, in situ TDPAC studies are extremely valuable. Adsorbed heptamolybdates exhibit TDPAC spectra very similar to those of crystalline APM, in contrast to free Mo7 O6− 24 in solution, which yields nearly unperturbed spectra. On adsorption, these molecules are ‘‘immobilized’’. Monomolybdates, however, yield very weakly perturbed spectra both in solution and when adsorbed. Under the assumption that both monomolybdates and polymolybdates are adsorbed and that the fraction that is not adsorbed contains polymolybdates only, it is possible References see page 911

908

3.1 Physical Properties

10 77 K

pH = 11.5

1

5

0

0 pH = 7

1

0

0 pH = 5.5

1

5

INTENSITY/arbitrary units

ANISOTROPY /%

5

0

0 pH = 3.7

1

5

0

−5

0 0

10

20 TIME / ns

30

(a)

0

1 2 3 FREQUENCY/Grad s−1

(b)

Fig. 12 TDPAC spectra for frozen molybdate solutions at various values of pH (a) and their Fourier transforms (b). (Reproduced with permission from Ref. [45].)

to derive the degree of adsorption by performing experiments at 300 and 77 K. Samples were prepared with various loadings ranging from 3 up to 15 wt.% MoO3 , the ‘‘theoretical monolayer coverage’’ corresponding to 12 wt.% MoO3 (denoted by Mo12Al). Figure 13 shows the Fourier-transformed TDPAC spectra for various oxide precursors at 300 and 77 K. The peak at zero frequency is more prominent at low loadings, even at 77 K, suggestive of depolymerization, whereas for Mo12Al essentially polymolybdates are adsorbed. From the small differences between the 300 and 77 K spectra, the degree of adsorption can be extracted, which turned out to be between 75 and 100% [45].

In an attempt to deduce information on the adsorbed species by another method, namely the equilibrium adsorption using excess solution, Luthra and Cheng [47] examined the supernatant by 95 Mo NMR. Since no polymolybdates were found in solution, it was concluded that complete depolymerization took place and that the adsorbed species is monomeric. This conclusion, however, was premature. It is possible to choose equilibrium adsorption conditions such that the supernatant contains monomolybdates only and yet is in equilibrium with adsorbed polymolybdates. This situation has been studied by separate TDPAC experiments on the supernatant and the impregnated support [45]. Interestingly,

3.1.3 Structure and Morphology

Mo 3 Al 77 K

Mo 3 Al 300 K

1

1

0

0 1

Mo 6 Al

Mo 6 Al

0

0

1

1

Mo 9 Al

Mo 9 Al

0

0

1

1

Mo 12 Al

Mo 12 Al

INTENSITY/arbitrary units

1

INTENSITY/arbitrary units

909

0

0 Mo 15 Al

Mo 15 Al

1

1

0

0 0

1

2

3

0

1

2

3

FREQUENCY/Grad s−1 (a)

(b)

Fourier transforms of TDPAC spectra of various impregnated Al2 O3 supports (the number indicates the wt.% MoO3 equivalent loading) at (a) 300 and (b) 77 K. (Reproduced with permission from Ref. [45].)

Fig. 13

this situation can be observed visually when ‘‘blue’’ APM solutions containing color centers (see above) are used for impregnation: first, both the impregnated support and the supernatant are blue; upon equilibration, the supernatant gradually turns colorless, whereas the impregnated support remains blue for several hours. D MoS2 in Various Forms MoS2 consists of hexagonal slabs of Mo, sandwiched between two hexagonal layers of sulfur such that Mo is at the center of a trigonal prism,

i.e. the sulfur layers are exactly superimposed. Therefore, the NQI exhibits axial symmetry. TDPAC spectra of both single crystals and bulk powders exhibit an NQI with axial symmetry and a precession frequency of ω = 111(1) Mrad s−1 . As the particle size decreases, this precession frequency decreases by almost a factor of two for particles of about 50 A˚ diameter. Possible explanations range from size effects, including sulfur to molybdenum ratios >2, to hypothetical octahedrally coordinated Mo, possibly References see page 911

910

3.1 Physical Properties

with protons intercalated. A typical pattern for 325-mesh powder from Aldrich is shown in Fig. 14a. Due to the low precession frequency, only about two-thirds of a full period is observable. When MoS2 (325-mesh) is reacted with n-butyllithium in hexane, an Lix MoS2 intercalation compound is formed. A typical TDPAC spectrum is shown in Fig. 14b. The anisotropy does not decrease in a typical cosine fashion (or superposition of several cosines), but rather in an exponential way, as shown in Fig. 11f, indicating a broad frequency distribution around zero (or very low) frequency. Apparently, the Li distribution in the host MoS2 is not homogeneous. Upon exfoliation (in H2 O, using glycerin as surfactant and ultrasonication), a well-defined suspension is obtained which is stable for several days. The resulting TDPAC spectrum is shown in Fig. 14c. Here, a unique NQI is observed with ω = 56 Mrad s−1 . It is believed that there are single slabs of MoS2 in solution with hydrated Li+ adsorbed to the basal planes. The charge transfer during lithiation is responsible for the lower NQI compared with the bulk MoS2 , similar to observations with intercalation compounds of 2H-TaS2 [48]. Upon restacking the slabs by hydrochloric acid (in the absence of surfactant), spectra almost identical with that for bulk MoS2 are obtained (see Fig. 14d). An interesting situation arises when exfoliated MoS2 is adsorbed on Al2 O3 (Fig. 14e); whereas the main component of the spectrum resembles to a large extent that of bulk MoS2 or restacked MoS2 , there is a highfrequency minority component which indicates Mo atoms with a lower coordination number and possibly another valence state. Such high-frequency components have also frequently been observed in highly disperse MoS2 powders and in surface species obtained by sulfidation of oxidic precursors [49]. They are assumed to be due to coordinatively unsaturated Mo atoms.

10

5

0

10

(b)

5

0

ANISOTROPY / %

10

(c)

5

0

10

(d)

5

0

10

(e)

5

0 −5

Conclusions and Outlook TDPAC spectroscopy on 99 Mo is useful for in situ studies of various stages of heterogeneous catalysts, despite the inherently poor frequency resolution (a consequence of the short τn ). The use of 187 W with excellent frequency resolution [50] looks very promising for the study of tungsten-based catalysts. Since the first edition of this Handbook, progress has been made in several respects. First, new and more efficient spectrometers using six detectors which allow one to record six coincidences under 180◦ and 24 coincidences under 90◦ simultaneously, greatly help in acquiring better statistics [51]. This allowed solid APM to be remeasured [52] and three classes of inequivalent Mo sites to be revealed with population ratios of 4 : 2 : 1, as expected from the crystal structure determination [53]. If it turns out that the strength and symmetry of the

(a)

3.1.3.6.4

0

10

20

30

TIME/ns

TDPAC spectra for MoS2 in various forms: (a) 325-mesh powder (Aldrich); (b) lithiated with n-butyllithium; (c) exfoliated (H2 O, ultrasonication + surfactant); (d) restacked (in HCl); (e) exfoliated and adsorbed on Al2 O3 .

Fig. 14

NQI are characteristic for the type of condensation of octahedra, this ‘‘fingerprint’’ could also be used for other polymolybdates. Second, attempts have been made to investigate Mo2 N, another candidate with potential interest in catalysis [54]. However, the NQI turned out to be so small that no further attempts were made to study supported nanocrystalline Mo2 N. A new direction of research dealt with investigations of InPt–ferrierite-based catalysts [55] and of the structure

References

of extremely nanosized and confined In−O species in ordered porous materials [56]. In both studies, TDPAC was combined with extended X-ray absorption fine structure (EXAFS) studies, a powerful combination indeed. Finally, nuclear resonant scattering using synchrotron light [57, 58] and the emerging new technique of synchrotron-based PAC [59], which uses excitation from the (stable) ground state and thus avoids radioactive parent isotopes, looks very promising, especially in the field of heterogeneous catalysis. References 1. G. K. Wertheim, M¨ossbauer Effect: Principles and Applications, Academic Press, New York, 1964. 2. N. N. Greenwood, T. C. Gibb, M¨ossbauer Spectroscopy, Chapman and Hall, London, 1971. 3. T. E. Cranshaw, B. W. Dale, G. O. Longworth, C. E. Johnson, M¨ossbauer Spectroscopy and its Applications, Cambridge University Press, Cambridge, 1985. 4. J. A. Dumesic, H. Topsøe, Adv. Catal. 1997, 26, 121. 5. H. Topsøe, J. A. Dumesic, S. Mørup, in Applications of M¨ossbauer Spectroscopy, R. L. Cohen (Ed.), Vol. II, Academic Press, New York, 1980, p. 55. 6. J. W. Niemantsverdriet, Spectroscopy in Catalysis, an Introduction, Wiley-VCH, Weinheim, 2007. 7. J. W. Niemantsverdriet, W. N. Delgass, Top. Catal. 1999, 8, 133. 8. F. J. Berry, in Spectroscopic Characterization of Heterogeneous Catalysts, J. L. G. Fierro (Ed.), Part A, Elsevier, Amsterdam, 1990, p. 299. 9. J.-M. M. Millet, Adv. Catal. 2007, 51, 309. 10. (a) R. L. M¨ossbauer, Z. Phys. 1958, 151, 124; (b) R. L. M¨ossbauer, Naturwissenschaften 1958, 45, 538. 11. J. W. Niemantsverdriet, C. F. J. Flipse, A. M. van der Kraan, J. J. van Loef, Appl. Surf. Sci. 1982, 10, 303. 12. J. W. Niemantsverdriet, A. M. van der Kraan, W. N. Delgass, J. Catal. 1984, 89, 138. 13. A. M. van der Kraan, R. C. H. Nonnekens, F. Stoop, J. W. Niemantsverdriet, Appl. Catal. 1986, 27, 285. 14. J. A. Amelse, J. B. Butt, L. H. Schwartz, J. Phys. Chem. 1978, 82, 558. 15. (a) G. B. Raupp, W. N. Delgass, J. Catal. 1979, 58, 337; (b) G. B. Raupp, W. N. Delgass, J. Catal. 1979, 58, 348; (c) G. B. Raupp, W. N. Delgass, J. Catal. 1979, 58, 361. 16. J. W. Niemantsverdriet, A. M. van der Kraan, W. L. van Dijk, H. S. van der Baan, J. Phys. Chem. 1980, 84, 3363. 17. G. LeCaer, J. M. Dubois, M. Pijolat, V. Perrichon, P. Bussiere, J. Phys. Chem. 1982, 86, 4799. 18. P. H. Christensen, S. Mørup, J. W. Niemantsverdriet, J. Phys. Chem. 1985, 89, 4898. 19. P. W. Selwood, Chemisorption and Magnetization, Academic Press, New York, 1975. 20. A. A. Hummel, A. P. Wilson, W. N. Delgass, J. Catal. 1988, 113, 236. 21. L. M. P. van Gruijthuijsen, G. J. Howsmon, W. N. Delgass, D. C. Koningsberger, R. A. van Santen, J. W. Niemantsverdriet, J. Catal. 1997, 170, 331. 22. H. Topsoe, B. S. Clausen, F. E. Massoth, Hydrotreating Catalysis, Springer-Verlag, Berlin, 1996.

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23. C. Wivel, R. Candia, B. S. Clausen, S. Mørup, H. Topsøe, J. Catal. 1981, 68, 453. 24. H. Topsøe, B. S. Clausen, R. Candia, C. Wivel, S. Mørup, J. Catal. 1981, 68, 433. 25. A. I. Dugulan, M. W. J. Craj´e, A. R. Overweg, G. J. Kearley, J. Catal. 2005, 229, 276. 26. A. I. Dugulan, M. W. J. Craj´e, G. J. Kearley, J. Catal. 2004, 222, 281. 27. P. J. van Berge, J. van de Loosdrecht, S. Barradas, A. M. van der Kraan, Catal. Today 2000, 58, 321. 28. R. Bacaud, P. Bussiere, F. Figueras, J. Catal. 1981, 69, 399. 29. K. Lazar, P. Bussiere, M. Guenin, R. Frety, Appl. Catal. 1988, 38, 19. 30. B. Benaichouba, P. Bussiere, J. M. Friedt, J. P. Sanchez, Appl. Catal. 1983, 8, 237. 31. (a) C. A. Clausen, M. L. Good, J. Catal. 1975, 38, 92; (b) C. A. Clausen, M. L. Good, J. Catal. 1977, 46, 58. 32. L. Stievano, S. Calogero, F. E. Wagner, S. Galvagno, C. Milone, J. Phys. Chem. B 1999, 103, 9545. 33. J.-M. M. Millet, H. Roussel, A. Pigamo, J. L. Dubois, J. C. Dumas, Appl. Catal. A 2000, 232, 77. 34. F. E. Wagner, J. A. Sawicki, J. H. Rolston, Hyperf. Int. 1988, 41, 733. 35. H. von Brandis, F. E. Wagner, J. A. Sawicki, K. Marcinkowska, J. H. Rolston, Hyperf. Int. 1990, 57, 2127. 36. F. J. Berry, S. Jobson, Hyperf. Int. 1988, 41, 613. 37. Y. Kobayashi, S. Nasu, S. Tsubota, M. Haruta, Hyperf. Int. 2000, 126, 95. 38. R. M. Finch, N. A. Hodge, G. J. Hutchings, A. Meagher, Q. A. Pankhurst, M. R. H. Siddiqui, F. E. Wagner, R. Whyman, Phys. Chem. Chem. Phys. 1998, 1, 485. 39. K. Blick, T. D. Mitrelias, J. S. J. Hargreaves, G. J. Hutchings, R. W. Joyner, C. J. Kiely, F. E. Wagner, Catal. Lett. 1998, 50, 211. 40. X. Ni, G. Sun, T. Butz, A. Lerf, Chem. Phys. 1988, 123, 455. 41. R. M. Steffen, H. Frauenfelder, in Alpha-, Beta-, Gamma-RaySpectroscopy, K. Siegbahn (Ed.), North-Holland, Amsterdam, 1965, Vol. 2, p. 997. 42. A. Lerf, T. Butz, Angew. Chem. Int. Ed. Engl. 1987, 26, 110. 43. (a) K. H. Tytko, O. Glemser, Adv. Inorg. Chem. Radiochem. 1976, 19, 239; (b) K. H. Tytko, G. Baethe, E. R. Hirschfeld, K. Mehmke, D. Stellhorn, Z. Anorg. Allg. Chem. 1983, 503, 43. 44. T. Butz, A. Lerf, C. Vogdt, A. M. M. Eid, Hyperf. Int. 1983, 15/16, 915. 45. T. Butz, C. Vogdt, A. Lerf, H. Kn¨ozinger, J. Catal. 1989, 116, 31. 46. A. Lerf, C. Vogdt, T. Butz, A. M. M. Eid H. Kn¨ozinger, Hyperf. Int. 1983, 16, 921. 47. W. P. Luthra, W. C. Cheng, J. Catal. 1987, 107, 154. 48. T. Butz, A. Lerf, Rev. Chim. Miner. 1982, 19, 496. 49. P. Mottner, T. Butz, A. Lerf, G. Ledezma, H. Kn¨ozinger, J. Phys. Chem. 1995, 99, 8260. 50. P. Mottner, T. Butz, Chem. Phys. 1990, 149, 199. 51. T. Butz, S. Saibene, Th. Fraenzke, M. Weber, Nucl. Instrum. Methods A 1989, 284, 417. 52. S. K. Das, F. Heinrich, T. Butz, Chem. Phys. 2006, 327, 291. 53. H. D. Evans Jr., B. M. Gatehouse, P. Leverett, J. Chem. Soc., Dalton Trans. 1975, 505. 54. G. Sun, Z. Zhou, S. K. Das and T. Butz, unpublished work. 55. J. M. Ramallo-L´opez, F. G. Requejo, A. G. Bibiloni, M. Renter´ıa, L. Gutierrez, E. E. Mir´o, Z. Naturforsch., Teil A 2000, 55, 327. 56. J. M. Rapallo-L´opez, M. Renteria, E. E. Mir´o, F. G. Requejo, A. Traverse, Phys. Rev. Lett. 2003, 91, 108304.

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3.1.3.7

Solid-State NMR Spectroscopy

Michael Hunger∗ and Wei Wang

Introduction NMR spectroscopy is applicable to any nucleus that possesses a magnetic moment, i.e. a nuclear spin. 1 H, 13 C, 15 N and 31 P nuclei are suitable isotopes allowing the study of a wide variety of reactants and adsorbate complexes interesting for studies in the field of heterogeneous catalysis. Due to their nuclear spin of I = 1/2, sufficient line narrowing is reached by the conventional high-resolution solid-state NMR techniques. Similar conditions exist for 29 Si and 31 P nuclei occurring in the framework of solid catalysts. Isotopes with a nuclear spin I > 1/2, such as 7 Li (I = 3/2), 11 B (I = 3/2), 17 O (I = 5/2), 23 Na (I = 3/2), 27 Al (I = 5/2), 71 Ga (I = 3/2) and 133 Cs (I = 7/2), are additionally characterized by an electric quadrupolar moment. These nuclei are involved in quadrupolar interactions, making their investigation more complicated. In all of the above cases, solid-state NMR spectroscopy is a very useful tool for the investigation of framework atoms, extraframework species and surface sites of solid catalysts and of adsorbate complexes and reaction intermediates formed on these materials. This chapter gives a brief review of the basic principles of solid-state NMR spectroscopy, methods of high-resolution solid-state NMR spectroscopy and special techniques for the preparation of solid catalysts investigated by these methods. As typical applications of solid-state NMR spectroscopy in heterogeneous catalysis, investigations of solid catalysts by 29 Si, 27 Al and 17 O NMR spectroscopy are demonstrated. For the NMR characterization of surface sites on solid catalysts, such as of Brønsted and Lewis acid sites and of base sites, the reader’s attention is directed to Chapter 3.2.4.4. A further objective of the present chapter is the description of in situ NMR experiments that allow one to study the mechanisms of heterogeneously catalyzed reactions. Suitable approaches are the investigation of the H−D exchange between the Brønsted acid sites and reactant molecules by solid-state 1 H and 2 H NMR spectroscopy and the study of the chemical conversion of reactants by solid-state 13 C NMR spectroscopy. Special attention has been devoted to the application of zeolites as solid catalysts. This was done not only because 3.1.3.7.1

∗

Corresponding author.

of their ideal crystalline structures, but also in view of their widespread use in numerous processes of industrial chemistry. Given the extensive application of solid-state NMR spectroscopy in heterogeneous catalysis, no attempt was made to treat the topic exhaustively. More information is available in several review papers and textbooks [1–6]. 3.1.3.7.2 Interactions of Nuclear Spins in Solids Nuclei characterized by a spin I possess a nuclear magnetic moment µ = γ ¯hI , where γ is the magnetogyric ratio and ¯h is Planck’s constant h divided by 2π. When placed in an external magnetic field B0 , the Zeeman interaction quantizes the orientations of the nuclear magnetic moments accompanied by a splitting of the energy level system into 2I + 1 eigenstates with energies [7–11]

E(m) = −mγ ¯hB0

(1)

where m = (I, I − 1, . . . , −I ) is the magnetic quantum number. In the case of nuclear spins I = 1/2, a splitting into two energy levels with m = −1/2 and m = 1/2 occurs. Ensembles of nuclear spins I = 3/2, for example, cause energy level diagrams with four levels (Fig. 1). Transitions between neighboring energy levels can be induced by an electromagnetic radiation, if the radiation frequency ν agrees with the Larmor frequency [7–11]: ν0 =

γB0 2π

(2)

Nowadays, NMR spectroscopy is performed in external magnetic fields with flux densities of up to B0 = 21.1 T, corresponding to the Larmor frequencies of 1 H nuclei of up to ν0 = 900 MHz. The spin ensemble in a sample under study is excited by a single-pulse experiment or a pulse sequence. Following this excitation, the answer of the spin ensemble is recorded in the time domain as free induction decay (FID). The spectrum plotted in the frequency domain is obtained by a Fourier transformation of the FID. Further details of the basic principles of Fourier transform NMR spectroscopy are described in textbooks [7–11]. Nuclear spin ensembles in liquids and gases cause narrow NMR signals since the influence of spatially anisotropic spin interactions on the linewidth is averaged to zero by isotropic motions. In contrast, solid materials such as catalysts and adsorbate complexes give NMR spectra that are strongly affected by the anisotropic spin interactions. The analysis of solid-state NMR signals occurring at characteristic resonance positions and having characteristic lineshapes is a source of structural information. On the other hand, to reach a separation of the NMR signals of nuclei in different local structures

3.1.3 Structure and Morphology

Vzz , occurring at their position. Electric field gradients are caused by non-spherical charge distributions in the environment of the resonating nuclei. Therefore, the quadrupolar interaction is again spatially anisotropic. Analytical expressions for the Hamiltonians of the above-mentioned spin interactions are given in, e.g., Refs. [7–11]. The superposition of the different spin interactions complicates the interpretation and calculation of solidstate NMR spectra. As a measure of the linewidth of solid-state NMR signals it is advantageous to consider the so-called second moment M2 . The full width at halfmaximum ν1/2 of an NMR signal in frequency units is also called static linewidth and amounts to [7–11]

m −3/2 Satellite transition −1/2

Central transition

+1/2 Satellite transition

+3/2 Zeeman interaction

Zeeman + quadrupolar interaction

Energy level diagram of a spin I = 3/2 ensemble for the case of Zeeman interaction only (left) plus quadrupolar interaction (right).

ν1/2 =

Fig. 1

and different chemical surroundings, sophisticated experimental techniques allowing an averaging of the nuclear spin interactions must be applied. The dominating interaction of a nuclear spin ensemble is the Zeeman interaction of the nuclear magnetic moments µ with the external magnetic field B0 . The detailed position and shape of the NMR signal of nuclei in solid materials, however, depend on the interaction of the nuclear spins with their environment. These spin interactions are described by the total Hamiltonian [7–11] ˆ =H ˆ II + H ˆ IS + H ˆ CSA + H ˆQ H

913

(4)

for Gaussian lines. The second moments of NMR signals can be calculated if the type of spin interaction and the local structure of the atoms under study are known. Otherwise, the experimental determination of the second moment is a suitable approach to gain structural information. According to the total Hamiltonian in Eq. (3), the total second moment is given by [3, 7–11] Q,CT

M2 = M2II + M2IS + M2CSA + M2

(5)

with M2II =

NI 3  µ0 2 4 2 1  γI ¯h I (I + 1) rij−6 5 4π NI

(6)

i =j

(3)

ˆ II , H ˆ IS , H ˆ CSA and H ˆ Q are the Hamiltonians of where H the homonuclear (II) and the heteronuclear (IS) magnetic dipole–dipole interaction, the anisotropic chemical shift interaction (CSA) and the quadrupolar interaction (Q), respectively. The dipolar interaction arises from local magnetic fields at the position of the resonating nuclei due to the magnetic moments of the surrounding nuclei. In the case of an interaction of identical nuclear spins I , the homonuclear dipolar interaction occurs. The interaction between different nuclear spins I and S is described by the Hamiltonian of the heteronuclear dipolar interaction. The anisotropic chemical shift interaction arises from a shielding of the external magnetic field by electrons in the environment of the resonating nuclei. The magnetic fields, which are induced by these electrons, oppose B0 . An anisotropic coordination of the resonating atoms to surrounding atoms may cause an anisotropic shielding. Quadrupolar interactions occur for nuclei with a spin I > 1/2, which are additionally characterized by an electric quadrupolar moment eQ. These nuclei interact with the z-component of the electric field gradient,

 1√ 2 ln 2 M2 π

M2IS =

NS NI  4  µ0 2 2 2 2 1  γI γS ¯h S(S + 1) rj−6 k 15 4π NI j =1 k=1




M2CSA =

4 2 2 2 1 2 γI B0  σ 1 + ηCSA 45 3



(7) (8)

92π 2 2 (9) ν 7 qs where I and S are the nuclear spins and NI and NS are the numbers of the resonating and non-resonating nuclei, respectively. γS denotes the magnetogyric ratio of the non-resonating spins S, rij and rj k are internuclear distances and µ0 is the permeability of vacuum. In the case of spins I = 1/2 and assuming that the sample does not contain paramagnetic impurities and that other susceptibility effects can also be neglected, the frequency νcg at the center of gravity of the signal of resonating nuclei is given by [7–11] Q,CT

M2

=

νcg = ν0 (1 − σ ) References see page 930

(10)

914

3.1 Physical Properties

where σ is the shielding constant caused by the anisotropic chemical shielding of the external magnetic field by electrons. The isotropic chemical shift, δ, is defined as δ = σref − σ , where σref is the shielding constant of the nuclei in a reference material. The value of δ is given by [3, 7–11] δ=

1 (δxx + δyy + δzz ) 3

(11)

with the principal values δxx , δyy and δzz of the chemical shift tensor, which are denoted in such a manner that δxx ≥ δyy ≥ δzz is fulfilled. The unit of the chemical shift δ is ‘‘ppm’’, which means parts per million of the resonance frequency v0 . The anisotropy, σ , and the asymmetry parameter, ηCSA , of the chemical shift tensor, used in Eq. (8), are defined by [3, 7–11] σ =

1 (δxx + δyy ) − δzz 2

(12)

and ηCSA =

3 (δxx − δyy ) 2 σ

(13)

In the case of strong quadrupolar interactions, it is practically impossible to excite the whole spectrum in a non-selective manner. Therefore, often the central transition {+1/2 ↔ −1/2} is selectively excited and observed, while the satellite transitions {−3/2 ↔ −1/2} and {+1/2 ↔ +3/2} in Fig. 1 are located outside the spectral range under study. Therefore, the second Q,CT , of quadrupolar interactions used in moment, M2 Eqs. (5) and (9) is that of the central transition only. The frequency at the center of gravity, νcg , of the central transition of quadrupolar nuclei is shifted by the secondorder quadrupolar shift [3, 7–11] νqs = νcg − ν0 (1 − σ ) =−

I (I + 1) − 30

3  ν2
4 Q 1 + 1 η2 ν0 3 Q

3 Vzz eQ 3 = CQCC 2I (2I − 1) h 2I (2I − 1)

(15)

where CQCC is the quadrupolar coupling constant. Finally, the asymmetry parameter, ηQ , of the electric field gradient tensor in Eq. (14) is defined by [3, 7–11] ηQ =

Vyy − Vxx Vzz

3.1.3.7.3 Experimental Techniques of Solid-State NMR Spectroscopy A Magic-Angle Spinning (MAS) The most important high-resolution solid-state NMR technique is fast sample rotation around an axis in the ‘‘magic’’ angle of m = 54.7◦ to the direction of the external magnetic field B0 . This technique is well-known as magic-angle spinning (MAS) [12]. If the sample spinning frequency νrot is large in comparison with the linewidth ν1/2 of the solid-state NMR signal recorded without highresolution techniques, the Hamiltonians in Eq. (3) can be replaced by mean Hamiltonians, which are proportional to the second Legendre polynomial (3 cos2 − 1). Here, denotes the angle between the external magnetic field and the spinning axis. The maximum line narrowing effect is achieved for the angle = m . For an NMR signal, which is broadened by inhomogeneous interactions [13], such as the heteronuclear magnetic dipole-dipole interaction and the chemical shift anisotropy, MAS leads to NMR signals consisting of a narrow central line at the center of gravity, νcg , of the solid-state NMR signal and spinning sidebands at the frequencies [13]

νk = νcg + kνrot (14)

with respect to the line position given in Eq. (10). The second-order quadrupolar shift (νqs ) also affects the value Q,CT of the second moment of the central transition (M2 ) in Eq. (9). The quadrupolar frequency νQ of nuclei with spin I > 1/2, used in Eq. (14), is given by [3, 7–11] νQ =

The principal values Vxx , Vyy and Vzz of the electric field gradient tensor in the molecular frame are denoted in such a manner that |Vxx | ≤ |Vyy | ≤ |Vzz | is fulfilled. The method of moments is suitable for a quantitative evaluation of the line broadening of solid-state NMR signals recorded without application of high-resolution techniques. In the case of high-resolution solid-state NMR spectroscopy, a significant signal narrowing occurs. The quantitative evaluation of the resonance positions and a simulation of the characteristic lineshapes of these narrowed signals is another often used approach for the determination of the spectroscopic parameters defined in Eqs. (5)–(16) [7–11].

(16)

(17)

where k = ±1, ±2, . . . denotes the order of the spinning sidebands. Nowadays, rapid sample spinning is performed with turbines having a gas bearing system for the rotor. Often, the rotors are made from zirconium dioxide. Depending on the rotor diameter, which is commonly between 7 and 2 mm, spinning frequencies between 5 and 35 kHz, respectively, can be achieved. In the case of homogeneous interactions, such as the homonuclear magnetic dipole–dipole interaction between more than two nuclei [13, 14], the residual MAS , of the central line may exhibit a linewidth, ν1/2 significant residual broadening, which can be reduced either by increasing the sample spinning frequency or by a combination of the MAS technique with multiple-pulse sequences leading to the so-called CRAMPS technique

3.1.3 Structure and Morphology

(combined rotation and multiple-pulse spectroscopy) [9, 15, 16]. The residual linewidth of the MAS NMR signals of quadrupolar nuclei with a half-integer nuclear spin I > 1/2 is determined by the second-order quadrupolar broadening of the central transition. This second-order quadrupolar broadening is not completely suppressed by the MAS technique, since the corresponding mean Hamiltonian contains both the second and the fourth Legendre polynomial [17]. Rapid sample spinning around an axis at the magic angle averages the second Legendre polynomial to zero, while the value of the fourth Legendre polynomial is only slightly decreased. The second moment, M2CT,MAS , of the central transition of quadrupolar nuclei under MAS conditions is given by [17] M2CT,MAS =

7 CT 2 M = π 2 νqs 92 2

(18)

That means that application of the MAS technique leads to a reduction √ in the linewidth of the central transition by a factor of 92/7 = 3.6, even at the limit of rapid sample spinning. Often, the analysis of the characteristic lineshapes of the central transition allows a direct determination of νQ and ηQ . On the other hand, the resolution of NMR spectra of quadrupolar nuclei is limited by the second-order quadrupolar line broadening. Since M2CT,MAS strongly depends on ν0 [see Eqs. (18) and (14)], the increase in the external magnetic field is a suitable approach for the improvement of spectral resolution. B Double-Oriented Rotation (DOR) A useful method to remove the second-order quadrupolar line broadening is simultaneous sample spinning around two axes. The corresponding technique, which allows the average of both the second and the fourth Legendre polynomial, is denoted the double-oriented rotation (DOR) technique [18–20]. The experimental device of the DOR technique consists of a large outer rotor reaching a spinning frequency of up to 1.5 kHz and a small inner rotor with a spinning frequency of up to 7 kHz [18]. The angle 1 between the external magnetic field and the rotational axis of the outer rotor corresponds to the magic angle m . The angle 2 between the rotational axes of the inner and the outer rotor amounts to 30.5◦ . The advantage of the DOR technique in comparison with the MAS technique is the total suppression of the second-order quadrupolar broadening. The minimum min , achieved under DOR condiresidual linewidth, ν1/2 tions is determined by the distribution of the isotropic chemical shift, δ, and of the second-order quadrupolar shift, νqs [see Eqs. (11) and (14)]. In addition, the residmin , if the ual linewidth of DOR signals can exceed ν1/2

915

homonuclear magnetic dipole–dipole interaction is not completely averaged by the rotation of the large and, therefore, slowly spinning outer rotor. Investigating, e.g., oxygen atoms in Si−O−Al bonds of zeolite catalysts by 17 O DOR NMR spectroscopy, the above-mentioned distribution of δ and νqs is the limiting factor of the spectral resolution [21]. This problem is independent of the high-resolution technique applied for solid-state NMR spectroscopy of quadrupolar nuclei. A general disadvantage of the DOR technique is its rather complicated handling. In addition, the sample volume of the inner rotor is very small in comparison with the volume of the radiofrequency coil around the large outer rotor. Therefore, additional techniques for the investigation of quadrupolar nuclei by NMR spectroscopy were developed, which are based on a combination of the MAS technique with echo sequences and a separation of spectroscopic parameters in two dimensions. C Multiple-Quantum MAS (MQMAS) Technique In the past decade, the multiple-quantum MAS (MQMAS) technique has been introduced as a new experimental approach for the study of quadrupolar nuclei. The MQMAS technique greatly enhances the resolution of spectra due to nuclei with a half-integer spin I > 1/2 [22, 23]. Basically, this approach combines an excitation of non-observable multiple-quantum transitions {+m, −m} with the experimentally observed single-quantum transition {+1/2, −1/2}. The enhancement of the resolution of solid-state NMR spectra of quadrupolar nuclei stems from the fact that the quadrupolar frequencies for both transitions are correlated. At a specific time of a pulse-sequence, the anisotropic parts of the quadrupolar interactions, responsible for the anisotropic line broadening, are refocused to form an echo. A number of schemes exist that are used to obtain twodimensional MQMAS NMR spectra of quadrupolar nuclei [24]. In a simple form of the experiment, the multiplequantum transitions are excited by a single high-power radiofrequency pulse (Fig. 2a, left) [23]. Subsequently, the multiple-quantum coherence (p = +m, −m) is allowed to evolve for the time t1 . After the evolution period t1 , a second pulse is applied, which converts the multiplequantum coherences into the coherence p = −1, which can be observed during the time t2 (Fig. 2a, right). The signal in the time domain is recorded immediately after the second pulse and the echo is formed at the time t2 = |QA|t1 , where QA is a term denoting a value of the quadrupolar anisotropy. Figure 2b, left, shows a pulse sequence that contains a selective low-power radiofrequency pulse, denoted a z-filter pulse [25]. This pulse References see page 930

916

3.1 Physical Properties

Excitation Conversion pulse pulse

p

Echo

t1

t2

Excitation Conversion pulse pulse z-filter pulse (b)

t1

l:

Echo

t2

Decoupling pulse FID

S: 0

(a) p

(a)

3 2 1 0 −1 −2 −3

p 2 Contact pulse

3 2 1 0 −1 −2 −3

t

p 2

p

Echo

l: p

MQMAS pulse sequences (left) and their coherence pathways (right).

Fig. 2

(b)

0

(19)

where αS and αI are equal to 1 for I = S = 1/2. Otherwise, αS is equal to [S(S + 1) − m(m − 1)]1/2 provided that the high-power radiofrequency pulse applied to the spins S induces transition between the levels with the magnetic spin quantum numbers m and m − 1 [26]. The CP experiment leads to enhancement of the signal intensity of the dilute spins S if the abundant spins I exhibit a higher magnetogyric ratio and/or a higher concentration

p

t

Echo

l: p

p p

p

p p

S: (c)

0

2

p 2

p

4

Nc

Echo

l: Adiabatic pulse

S: (d)

αI γI B1I = αS γS B1S

2t

t

p 2

sequence has the advantage of having a symmetrical coherence transfer pathway (Fig. 2b, right), which significantly diminishes problems with the phasing of the obtained signals. The three-pulse sequence in Fig. 2b is the most widely applied MQMAS experiment for the study of quadrupolar nuclei [24]. The two-dimensional Fourier transformation of the echo decays obtained for different pulse delays t1 leads to a two-dimensional MQMAS spectrum with featured signals lying along the quadrupolar anisotropy axis. The isotropic spectrum can be obtained by a shearing procedure of the two-dimensional spectrum, which is described in detail in Ref. [24]. D Cross-Polarization (CP) Experiment A severe limitation of solid-state NMR spectroscopy is the sensitivity of this method. Cross-polarization (CP) is an effective technique to increase the sensitivity significantly. This technique is advantageously employed in the case of dilute spins in the neighborhood of dipolar coupled abundant spins [9, 26]. The principle of the CP experiment is demonstrated in Fig. 3a. The experiment starts with a π/2 pulse in the channel of the abundant spins I . During a subsequent contact time, τ , magnetization (also denoted spin polarization) is transferred from the abundant spins I to the dilute spins S. The magnetic fields, B1I and B1S , of the contact pulses, which are applied simultaneously to the spins I and S, have to fulfill the Hartmann–Hahn condition [26]:

t1

S:

0 p 2

Adiabatic pulse 2

p

p p

4 p

p p

Nc

Echo

l: Tr /n S: (e)

0

2

4

Nc

Pulse sequences for CP (a), SEDOR (b), REDOR (c), TRAPDOR (d) and REAPDOR (e) experiments.

Fig. 3

than the spins S. The maximum signal enhancement achievable by cross-polarization (ICP ) for spin S = 1/2 nuclei in comparison with the single-pulse excitation (ISP ) amounts to [26] γS ICP = ISP [γI (1 + ε)]

(20)

with ε=

(NS /NI ) (αS /αI )2

(21)

where NI and NS denote the numbers of the spins I and S, respectively, in the sample.

3.1.3 Structure and Morphology

E SEDOR, REDOR, TRAPDOR and REAPDOR Experiments SEDOR (spin echo double resonance) [27, 28], REDOR (rotational echo double resonance) [29, 30], TRAPDOR (transfer of population in double resonance) [31, 32] and REAPDOR (rotational echo adiabatic passage double resonance) experiments [33, 34] are double resonance NMR experiments allowing a more quantitative investigation of the dipolar coupling between spins I and S in solids. The SEDOR experiment is performed without a simultaneous application of high-resolution techniques. A π/2 − τ − π − τ echo sequence is applied to the spins I (Fig. 3b) [27, 28]. The echo refocuses heteronuclear dipolar couplings and the chemical shift anisotropy. During the first pulse delay of the echo sequence, a single π-pulse is applied to the spins S. This pulse inverts the sign of the dipolar coupling, which perturbs the dipolar refocusing process and diminishes the echo intensity for coupled IS spin pairs. By measuring the loss of the echo intensity S(t1 ) as a function of pulse delay t1 , the distance rI S between the coupled spins I and S can be determined via the SEDOR fraction [28]

Sf (t1 ) = with S(t1 ) =

S0 − S(t1 ) S0

 0

π

  cos Dt1 (3 cos2 θ − 1) sin θdθ

(22)

(23)

where D = γI γS ¯h/rI3S . The value of rI S defines the distance between the coupled spins I and S and θ is the angle between the internuclear vector rI S and the external magnetic field B0 . S0 is the echo intensity without application of the π-pulse to the spin S ensemble. The REDOR experiment is in principle a SEDOR experiment in combination with MAS [29, 30]. Spinning of the sample at the magic angle enhances the sensitivity and resolution and allows the measurement of the echo for longer times. A typical REDOR pulse sequence is shown in Fig. 3c. A rotor-synchronized echo sequence is applied to the spins I , which are detected after a time 2τ equaling the even number Nc of rotation periods. For decoupling the dipolar interaction between the spins I and S, π-pulses are applied to the spin S ensemble at every half rotation period Tr . The dipolar coupling is determined by measuring the REDOR fraction, which describes the loss of the echo intensity as a function of the number of rotor periods [35] 1 S0 − S(Nc Tr ) (Nc Tr )2 M2I S = 2 S0 π S(S + 1)

(24)

where M2I S is the second moment of the heteronuclear dipolar interaction given in Eq. (7). This term contains rI S , which is the distance between the coupled spins

917

I and S. Again, S0 is the echo intensity obtained without application of π-pulses to the spin S ensemble. The TRAPDOR and REAPDOR experiments were specifically designed for the study of spins I interacting with quadrupolar nuclei (spin S > 1/2) [31–34]. In the TRAPDOR experiment, a rotor-synchronized echo sequence is applied to the spins I . During the pulse and echo delay, the quadrupolar nuclei are continuously irradiated (Fig. 3d). Continuous irradiation of quadrupolar nuclei in combination with MAS leads to rotationally induced level transitions. These level transitions cause a dipolar dephasing, which is additive for each rotation period. Since it is difficult to calculate this dephasing effect, TRAPDOR is a qualitative experiment only. However, the comparison of spectra obtained by spin I echoes recorded with and without irradiation of the spins S allows to distinguish spins I coupled with spins S and spins I without coupling. The REAPDOR experiment is a variation of the TRAPDOR experiment. The REAPDOR experiment starts with an excitation of the spins I = 1/2 by application of a π/2-pulse. Subsequently, a train of rotor-synchronized π-pulses is applied to the spins I at every half rotation period Tr (Fig. 3e). The π-pulse in the center of this train is omitted. In the first half of the evolution period, the spins I will dephase as a result of the chemical shift anisotropy and the heteronuclear dipolar coupling. In the second half of the evolution period, the magnetization is refocused and the echo is formed. In the center of the evolution period, a so-called adiabatic passage pulse is applied to the spins S. Due to this pulse, the spin states of the quadrupolar spin S nuclei are changed. Therefore, the dipolar dephasing of the spins I , which are coupled with spins S, cannot be refocused in the echo and a decrease of the echo intensity occurs. The optimum duration of the adiabatic passage pulse applied to the spins S is 1/3 to 1/2 of the rotation period Tr [24, 34]. The quantitative evaluation of the REAPDOR fraction is performed similar to the REDOR fraction. 3.1.3.7.4 Preparation of Catalyst Samples for Solid-State NMR Spectroscopy Most solid-state NMR investigations of the framework of solid catalysts performed in the past decades were made using hydrated materials. In this case, the powder sample is filled into the MAS NMR rotor without special treatment. Starting in the early 1980s, solid catalysts in the calcined state were investigated using laboratory-made glass ampoules, which fit as inserts into commercial MAS NMR rotors [36, 37]. Nowadays, commercial glass inserts are offered for all commercial MAS NMR rotor types. Into a 4- and 7-mm glass insert, ca. 25 References see page 930

918

3.1 Physical Properties

and 100 mg, respectively, of catalyst powder can be filled. After calcination of the catalyst material under vacuum and loading with probe molecules or reactants, the glass insert is sealed at the waist and the obtained glass ampoule is inserted into the rotor (see Fig. 1 in Chapter 3.2.4.4). Another frequently applied technique is based on the preparation of the solid catalyst directly inside the sample volume of a MAS NMR rotor. After filling the catalyst material into the MAS NMR rotor, the rotor is put into a fitting at the bottom of a special vacuum equipment. Upon calcination of the catalyst under vacuum and loading with probe molecules or reactants, the MAS NMR rotor is sealed with a gas-tight rotor cap inside the vacuum equipment. A similar system, which allows the calcination of powder material in a horizontal tube inside a heater at temperatures of up to 1000 K, was suggested by Zhang et al. [38] (Fig. 2 in Chapter 3.2.4.4 of this Handbook). After the treatment of the catalyst, the powder material is filled into an MAS NMR rotor at the bottom of the equipment, sealed with a rotor cap from a plug rack and transferred to the NMR spectrometer. The study of the mechanisms of heterogeneously catalyzed reactions by solid-state NMR spectroscopy requires a suitable technique for carrying out experiments under well-defined conditions. For investigations under flow conditions, two different techniques were introduced: (i) the catalytic reaction is performed in an external fixedbed reactor, and subsequently the catalyst is transferred into an MAS NMR rotor after quenching the reaction [39, 40]; (ii) in a true in situ technique, the MAS NMR rotor is used directly as a fixed-bed reactor situated inside a high-temperature MAS NMR probe [41, 42]. The most widely applied technique for in situ MAS NMR investigations of heterogeneously catalyzed reactions Reactant flow

under continuous-flow (CF) conditions is based on the injection of a carrier gas loaded with vapors of the reactants into the spinning MAS NMR rotor via an injection tube [41, 42]. For this purpose, the solid catalyst is pressed to a hollow cylinder using a special tool. An injection tube is inserted into the sample volume up to the bottom of the MAS NMR rotor reactor via an axially placed hole in the rotor cap. Via this tube, the feed is injected into the inner space of this hollow cylinder and flows from the bottom to the top of the sample volume inside the MAS NMR rotor reactor. The product stream leaves the sample volume via an annular gap in the rotor cap. In some applications, the reaction products leaving the CF MAS NMR probe were led to the sampling loop of an on-line gas chromatograph [42]. A recent development is the coupling of in situ CF MAS NMR spectroscopy under flow conditions with in situ UV/visible spectroscopy [43]. The CF MAS NMR–UV/visible probe (Fig. 4) is based on the abovementioned injection technique. In the CF MAS NMR probe, a glass fiber is attached to the bottom of the stator. At the bottom, the rotor is equipped with a quartz window. Via this quartz window and using the glass fiber, the catalyst material in the sample volume can be investigated by a fiber-optic UV/visible spectrometer. The above technique is suitable, e.g., for simultaneous in situ CF MAS NMR and UV/visible studies of the formation of aromatic compounds and carbenium ions on the surface of solid catalysts. NMR Investigations of Catalyst Frameworks For a number of nuclei, such as 1 H (spin I = 1/2, Irel = 1), 7 Li (I = 3/2, Irel = 0.272), 11 B (I = 3/2, Irel = 0.133), 23 Na (I = 3/2, Irel = 9.27 × 10−2 ), 27 Al 3.1.3.7.5

MAS NMR rotor

Catalyst Quartz window

Glass fiber

Product flow

Rf coil

Gas bearing

Fig. 4

UV/ visible beam

Scheme of the in situ CF MAS NMR–UV /visible probe applied to the study of heterogeneously catalyzed reactions [43].

919

3.1.3 Structure and Morphology

(I = 5/2, Irel = 0.207), 29 Si (I = 1/2, Irel = 3.69 × 10−4 ), 31 P (I = 1/2, I −2 51 V (I = 7/2, I rel = 6.65 × 10 ), rel = 71 Ga (I = 3/2, Irel = 5.65 × 10−2 ) and 133 Cs 0.383), (I = 7/2, Irel = 4.82 × 10−2 ) [8], the investigation of solid catalysts by solid-state NMR spectroscopy using samples with isotopes in natural abundance is a standard characterization. The second value in the parentheses above is the relative intensity Irel in comparison with 1 H nuclei (Irel = 1), which have the highest NMR sensitivity. Other nuclei, such as 13 C (I = 1/2, Irel = 1.76 × 10−4 ), 15 N (I = 1/2, Irel = 3.85 × 10−6 ) and 17 O (I = 5/2, Irel = 1.08 × 10−5 ) [8], require an isotopic enrichment, making experiments more expensive. In the following part, examples of the application of solid-state 29 Si, 27 Al and 17 O NMR spectroscopy for the characterization of solid catalysts are discussed. Additionally, the potential of solid-state NMR spectroscopy for the investigation of the mechanisms of heterogeneously catalyzed reactions is demonstrated. The characterization of surface sites of solid catalysts by solid-state 1 H, 13 C, 15 N and 31 P NMR spectroscopy is described in detail in Chapter 3.2.4.4 of this Handbook. A Solid-State 29 Si NMR Spectroscopy of Catalysts The basic structural units of zeolite catalysts are TO4 tetrahedra with silicon atoms at the central T-positions. In the second coordination sphere of these T-atoms, various metal atoms, such as aluminum, boron, gallium, iron and titanium, can be incorporated into the framework. Depending on the number of metal atoms that are incorporated in the second coordination sphere of T-positions, the tetrahedrally coordinated silicon atoms (Q4 ) are characterized by up to five different environments, denoted Si(nT) with n = 0, 1, 2, 3 and 4. In the case of aluminum atoms incorporated at T-positions, each type of Si(nAl) species yields a 29 Si MAS NMR signal in a well-defined range of chemical shifts. Figure 5 summarizes the chemical shifts of these Si(nAl) species. The relative intensities of the 29 Si MAS NMR signals of the Si(nAl) species are a function of the framework composition. Therefore, the framework nSi /nAl ratio of zeolite catalysts can be calculated via [4] 4 

ISi(nAl) nSi n=0 = 4 nAl  0.25nISi(nAl)

(25)

n=0

where ISi(nAl) is the intensity of the Si(nAl) signal, i.e. of the NMR signal caused by silicon atoms with n aluminum atoms in the second coordination sphere. In the same manner as described for 29 Si MAS NMR spectroscopy of

Al O AlOSiOAl O Al

Al O AlOSiOSi O Al

Al O SiOSiOSi O Al

Al O SiOSiOSi O Si

Si O SiOSiOSi O Si

Si(4Al)

Si(3Al)

Si(2Al)

Si(1Al)

Si(0Al)

Si(4Al) Si(3Al) Si(2Al) Si(1Al) Si(0Al) −80

−90

−100

−110

−120

d29 / ppm Si

29 Si

NMR shift range of Si(nAl) species in the framework of zeolite catalysts. The dotted lines for the Si(4Al) species give the chemical shift range occurring for sodalites with different guest compounds [4]. The shift values are referenced to tetramethylsilane (TMS: δ29Si = 0 ppm).

Fig. 5

aluminosilicate-type zeolites, the chemical compositions of the framework of the gallium analogue of zeolite Beta, [Ga]Beta [44], and of zincosilicates can be obtained [45]. The 29 Si MAS NMR spectra of zeolites VPI-7 and VPI-9, for example, show signals of Si(1Zn) and Si(2Zn) units at resonance positions of −88.5 to −95.6 and −77.9 to −81.0 ppm, respectively, referenced to tetramethylsilane (TMS: δ29Si = 0 ppm) [45]. In each zeolite catalyst, terminal hydroxyl groups bound to silicon atoms exist [Q3 , Si(3Si,1OH); Q2 , Si(2Si,2OH)]. These hydroxyl groups are located at the outer surface of zeolite particles or at internal framework defects. It is important to note that Si(1Al) species (δ29Si = −95 to −105 ppm) occur at the same resonance positions as Si(3Si,1OH) species (δ29Si = −100 to −103 ppm) [46]. A quantitative evaluation of the relative intensities of Si(nT) signals therefore requires the determination of the concentration of Q3 silicon species, which can be performed via quantitative 1 H MAS NMR spectroscopy of the SiOH groups in the dehydrated material (see Chapter 3.2.4.4 of this Handbook). A number of zeolite catalysts are characterized by structures with crystallographically non-equivalent T-positions in chemically equivalent environments. Silicon atoms located at such T-positions may cause a splitting of the 29 Si MAS NMR signals occurring at different chemical shifts. For the determination of the 29 Si NMR shifts of silicon References see page 930

3.1 Physical Properties

δ29Si = −5.230 − 0.570α ppm The following correlation between δ29Si cos α/(cos α − 1) was theoretically derived [48] δ29Si = −223.9 cos α/(cos α − 1) + 5n − 7.2 ppm

(26) and

(27)

where n is the number of aluminum atoms in the second coordination sphere of the T-atoms. Comparing data of 29 Si MAS NMR and XRD studies on unloaded zeolite ZSM-5, Fyfe et al. [52] obtained a correlation coefficient of 0.97 in Eq. (26). The comparison of experimental 29 Si MAS NMR spectra with resonance positions calculated via Eq. (26) or (27) allows the examination of structure models developed by Rietveld refinement of powder X-ray patterns. Fyfe and co-workers [53–56] investigated the connectivities between silicon atoms at crystallographically non-equivalent T-positions in zeolites ZSM-5, ZSM12, ZSM-23, ZSM-39 and KZ-2. These investigations were performed by two-dimensional COSY (correlation spectroscopy) and INADEQUATE (incredible natural abundance double quantum transfer experiment) experiments [57]. The COSY experiment utilizes the J -coupling between neighboring nuclear spins. The pulse sequence consists of an excitation of the 29 Si nuclei via a crosspolarization sequence (Fig. 3a), an evolution period, a π/2-pulse inducing the magnetization transfer between neighboring J -coupled spins and an acquisition period. After two-fold Fourier transformation, the 2D spectrum is obtained, which consists of cross-peaks caused by two

silicon atoms connected via oxygen bridges. An important advantage of the INADEQUATE experiment is that it can be performed with materials containing 29 Si nuclei in natural abundance. A recent development is the double-quantum correlation experiment introduced by Brouwer et al. [58]. This experiment utilizes through-space dipolar interactions between 29 Si nuclei in natural abundance in pure silica zeolites to provide information about bonding connectivities and also long-range distances between silicon atoms in zeolite structures. Connectivities and distances of 29 Si nuclei in solid catalysts with other atoms such as 11 B [59], 23 Na [60] and 27 Al nuclei [61, 62] can be investigated by doubleresonance experiments (Fig. 3). {27 Al}−29 Si REAPDOR NMR spectroscopy has been employed to study the local structure of silicon atoms in the framework of the aluminum-substituted molecular sieve ETS-10 [62]. In this catalyst, some of the T-positions were occupied by aluminum atoms, leading to the ETAS-10 material (ETAS-10 = ETS-10 with framework positions substituted by aluminum atoms) with an nSi /nAl ratio of 22 [62]. The structure of ETS-10 comprises corner-sharing SiO4 tetrahedra and TiO6 octahedra linked via oxygen bridges. The titanium atoms are coordinated via oxygen bridges to two titanium atoms and four silicon atoms, Ti(4Si,2Ti). The silicon atoms are coordinated either to four silicon neighbors, Si(4Si), or to three silicon neighbors and one titanium neighbor, Si(3Si,1Ti). The crystallographic nonequivalence of the T-positions leads to eight different titanium and three different silicon sites. The inset in Fig. 6 shows the 29 Si MAS NMR spectrum of ETAS-10 [62]. It consists of four signals at −90.2 (A), 1.0

B

C D

A

0.8 −84

−88

0.6

−92

−96 −100 −104 −108

d29 / ppm Si

S0

atoms at crystallographically non-equivalent T-positions, it is advantageous to use all-silica materials, which consist of Si(OAl) species only. Upon treatment of the zeolite in its ammonium form with water vapor at temperatures of 873–1073 K, dealumination and ultrastabilization of the framework occur, which bring about a dramatic decrease in the 29 Si MAS NMR linewidth [47]. The above-mentioned splitting of 29 Si MAS NMR signals of silicon species at crystallographically nonequivalent T-positions is explained by different bonding geometries of these SiO4 tetrahedra. Empirical correlations and theoretical considerations showed that the 29 Si NMR shifts, δ29Si , of Si(nAl) species are linearly related to the average value of the four Si−O−T bond angles, α, at the silicon atom in the center of the tetrahedron. With linear regression analysis, quantitative relationships between the values of δ29Si and α, sec α, sin(α/2) and cos α/(cos α − 1) were found [4, 47, 48–51]. For Si(4Al) units in sodalite, cancrinite, thomsonite and zeolites A, X, Y and ABW, Eq. (26) with a correlation coefficient of 0.988 was obtained [51]:

S0-S (NcTr)

920

0.4 0.2 0.0 0.0

0.5

1.0

1.5

2.0

2.5

3.0

NcTr / ms Experimentally obtained REAPDOR fraction [S0 − S(Nc Tr )]/S0 (open circles) and theoretically derived REAPDOR curves for the molecular sieve ETAS-10. The inset shows the 29 Si MAS NMR spectrum of ETAS-10 with signals A−D [62].

Fig. 6

3.1.3 Structure and Morphology

−94.0 (B), −96.0 (C) and −104.5 ppm (D). The REAPDOR experiment was performed using the double-resonance pulse sequence shown in Fig. 3e. In this figure, spins I and spins S stand for 29 Si and 27 Al nuclei, respectively. In the center of the pulse sequence, a radiofrequency pulse with a duration of Tr /2 was applied to the spins S. The REAPDOR measurements were restricted to a dephasing period of 3.0 ms due to the rapid spin-spin relaxation of the observed signals. Using the above-mentioned approach, a REAPDOR response could be observed for signal A at −90.2 ppm only. Signals B, C and D showed negligible dephasing. This result and a comparison of the 29 Si MAS NMR spectra of ETAS-10 and ETS-10 [63, 64] led to the assignment of signal A to Si(2Si,1Ti,1Al), signals B and C to Si(3Si,1Ti) and signal D to Si(4Si) species. The experimentally obtained REAPDOR fraction [S0 − S(Nc Tr )]/S0 of ETAS-10 is plotted in Fig. 6 as open circles. In addition, calculated REAPDOR curves for 29 Si−27 Al dipolar coupling constants of D = 195 (top), 185 (middle) and 175 kHz (bottom) are shown. An acceptable agreement between the experimental and the calculated curves occurs for a 29 Si−27 Al dipolar coupling of 185 ± 10 kHz, which corresponds to an Si−Al distance of 323 ± 5 pm. This value falls within the Si−Al distance range of 309–330 pm estimated on the basis of X-ray diffraction data for the zeolite structure [62]. The above-mentioned study demonstrates the capability of double-resonance spin-echo NMR experiments for quantitative investigations of the local structure of solid catalysts. B Solid-State 27 Al NMR Spectroscopy of Catalysts According to Loewenstein’s rule, the formation of Al−O−Al bonds in aluminosilicates and aluminophosphates is forbidden, and only Al(4Si) and Al(4P) species can exist in the corresponding frameworks [4]. Therefore, 27 Al MAS NMR spectra of hydrated zeolite catalysts and aluminophosphates consist, in general, of only one signal of framework aluminum atoms (Alf ) in a range of chemical shifts between ca. 35 and 60 ppm (referenced to a 0.1 M aqueous solution of Al(NO3 )3 in D2 O) [4, 65, 66]. In hydrated aluminosilicates and aluminophosphates, only small deviations from the ideal tetrahedral symmetry of the AlO4 units occur, which lead to weak quadrupolar interactions and weak second-order quadrupolar line broadenings. Extra-framework aluminum (Alnf ) species in hydrated zeolites, formed e.g. as a result of calcination and steaming, are octahedrally coordinated AlO6 species and cause 27 Al NMR signals at −15 to 0 ppm [4, 65, 66]. If extra-framework aluminum exists as polymeric aluminum oxide or oxide hydrates in zeolite cages or

921

channels, a strong second-order quadrupolar line broadening may occur owing to distortions of the octahedral symmetry of these AlO6 species. In some cases, an additional broad signal occurs at 30–50 ppm, which is caused by aluminum atoms in a disturbed tetrahedral coordination or a pentacoordinated state [4, 67]. For the 27 Al MAS NMR signals of tetrahedrally coordinated framework aluminum atoms, no definite relationships were found between the chemical shift, δ27Al and the nSi /nAl ratio or the Si−Al order scheme of the zeolite framework. However, by 27 Al and 29 Si MAS NMR spectroscopy of lithium and sodium halide aluminosilicate sodalites with high-speed sample spinning, Jacobsen et al. [68] obtained the following linear correlation between the 27 Al NMR shift, δ27Al and the 29 Si NMR shift, δ29Si , of aluminum and silicon atoms, respectively: 1.03δ29Si δ27Al = + 151.94 ppm ppm

(28)

Whereas the MAS technique reduces the second-order quadrupolar line broadening by a factor of only ca. 3.6 [see Eq. (18)], application of the DOR technique results in a complete averaging of this interaction (see Section 3.1.3.7.3C). Crystalline aluminophosphates are materials with a high degree of framework ordering. The framework of aluminophosphates is built of AlO4 and PO4 tetrahedra in an alternating arrangement. The 27 Al DOR NMR spectra of these materials consist of narrow signals due to aluminum atoms at crystallographically non-equivalent T-sites [69, 70] or caused by framework aluminum atoms interacting with probe molecules [71, 72]. At the top of Fig. 7, the 27 Al DOR NMR spectrum of the dehydrated aluminophosphate VPI-5 is shown [73]. It consists of two signals at 36 and 33 ppm with an intensity ratio of 2 : 1. Addition of water molecules dramatically alters the local bonding of the framework aluminum atoms in VPI-5, as indicated by the spectrum at the bottom of Fig. 7. The lines at 36 and 33 ppm have disappeared, and two partially resolved signals associated with tetrahedrally coordinated framework aluminum atoms are present at 41 and 40 ppm. In addition, a high-field signal occurs at −18 ppm in the range ascribed to octahedrally coordinated aluminum species. The integrated intensities of these three signals have a ratio of 1 : 1 : 1. The recording of spectra in different magnetic fields allows the determination of the second-order quadrupolar shift as a function of the Larmor frequency. In this way, the isotropic chemical shifts of the above-mentioned signals of aluminum species in VPI-5 were determined as 43.6, 41.6 and −10.4 ppm, respectively [74]. Equations (14) and (15) can be used to calculate the quadrupolar frequency, References see page 930

922

3.1 Physical Properties

36

Dehydrated VPI-5

33

(a) 41

40 Hydrated VPI-5

−18

60 (b)

40

20

d27

AI

0

−20

/ ppm

Fig. 7 27 Al DOR NMR spectra of the hydrated (a) and dehydrated (b) aluminophosphate VPI-5 [73].

νQ , and the quadrupolar coupling constant, CQCC . For the signals at 43.6 and 41.6 ppm, quadrupolar coupling constants of 2.0–2.3 and 1.0–1.2 MHz, respectively, were found [75]. No second-order quadrupolar shift occurred for the signal at −10.4 ppm [74]. Quadrupolar coupling constants of 1.0–1.2 and 2.0–2.3 MHz are characteristic for tetrahedrally coordinated aluminum atoms located at different T-positions in the framework of VPI-5. The latter was explained by the correlation between the CQCC values and the shear strain parameters of AlO4 tetrahedra in a number of crystalline materials [75]. In addition, this correlation could be used to assign the signals at 43.6 and 41.6 ppm to aluminum atoms at two different T-positions in six-membered oxygen rings (Al3 and Al2, respectively). The low quadrupolar interaction of the aluminum species giving rise to the signal at −10.4 ppm indicates the octahedral coordination [74], which agrees well with the characteristic resonance position of this signal. These aluminum species are located at the position Al1 in four-membered oxygen rings. The above-mentioned variation of the 27 Al DOR NMR spectrum of VPI-5 due to hydration is an example for the change in the coordination of framework aluminum atoms in a zeolite framework from AlIV to AlVI . An alternative approach for the study of quadrupole parameters of 27 Al nuclei in crystalline materials is the application of MQMAS. In the last decade, this technique has been widely used to study aluminum species in dealuminated zeolite catalysts [76–79]. Solidstate NMR investigations of calcined and steamed zeolites indicated that a coordination change of aluminum atoms from tetrahedral to octahedral coordination is not always

accompanied by dealumination of the framework [80, 81]. Whereas the 27 Al MAS NMR spectra of steamed zeolites consist of signals around 0 ppm due to the octahedrally coordinated aluminum atoms, which were often assigned to extra-framework aluminum species, adsorption of a strong base, such as ammonia and pyridine, leads to a partial transformation of AlVI to AlIV species [80]. Moreover, Wouters et al. [81] found agreement between the framework nSi /nAl ratios of thermally treated zeolites H-Y determined by 29 Si and 27 Al MAS NMR spectroscopy upon adsorption of ammonia. Woolery et al. [82] assumed a calcination-induced transformation of SiOHAl groups to three-fold-coordinated Lewis acidic framework aluminum atoms (AlIII ) with neighboring SiOH groups. In contrast, Wouters et al. [83, 84] suggested partial hydrolysis of framework Al-O bonds and a generation of framework-connected AlOH groups upon calcination or steaming of zeolite catalysts. In a second step, these framework aluminum atoms coordinate to water molecules, giving rise to octahedrally coordinated framework aluminum species. Subsequent adsorption of ammonia on the material converts the AlVI atoms back to AlIV atoms. The ammonia-induced coordination change of aluminum atoms in zeolites was investigated by high-field 27 Al MAS NMR and 27 Al MQMAS NMR spectroscopy [85]. Figure 8 shows high-field 27 Al MAS NMR spectra of zeolites H-Y steamed at 753 K for 2.5 h at a water vapor pressure of 47.4 kPa (deH-Y/47.4) and 94.0 kPa (deH-Y/94.0). In addition, the latter sample was washed in 0.1 M HCl after steaming. Before recording the spectra on the right-hand side, adsorption of ammonia at a pressure of 5 kPa for 1.5 h on the dehydrated samples was performed on both samples. Upon rehydration, it was found that AlVIa and AlVIb species are involved in a coordination change from AlVI to AlIV . By 27 Al MQMAS NMR spectroscopy, a quadrupolar coupling constant of CQCC = 2–3 MHz (typical for framework aluminum atoms) was determined for the AlVIa and AlVIb atoms. The AlVIc atoms, which were not involved in the ammoniainduced coordination change, cause an 27 Al MQMAS NMR signal with a CQCC value of ca. 5 MHz. This large quadrupolar coupling constant indicates that the corresponding AlVI atoms are due to aluminum species in extra-framework clusters. This assumption is supported by the observation that washing of the sample with dilute HCl leads to preferential removal of the AlVIc atoms (see Fig. 8d) [85]. SEDOR and REDOR experiments are suitable for the determination of the bonding geometries of aluminum atoms in solid catalysts. As an example, van Eck and Veeman [28] investigated the Al-P distance in the aluminophosphate AlPO-5. The 27 Al MAS NMR spectrum of this material consists of a single signal at 38.4 ppm due

3.1.3 Structure and Morphology

deH-Y/47.4

deH-Y/47.4a

AIIVa

AIVIa

AIIVb AIVIb

AIIVa

(a)

AI

AIVIa

deH-Y/94.0a AIIVa

IVa

AIIVb

75

50

AIIVb AI

25

AIVIc

VIb

0

−25

100

d27AI / ppm (c)

AIVIc

(b)

deH-Y/94.0

100

AIIVb

AIVIc

75

50

25

0

−25

d27AI / ppm (d)

Fig. 8 27 Al MAS NMR spectra of dealuminated zeolites deH-Y recorded before (a, c) and after (b, d) adsorption of ammonia in an external magnetic field of B0 = 17.6 T [85].

to tetrahedrally coordinated Al(4P) species. By application of the NMR experiment shown in Fig. 3b with 27 Al nuclei as spins I and 31 P nuclei as spins S, the SEDOR fraction was determined [Eqs. (22) and (23)]. Fitting of the SEDOR curve led to the dipolar coupling constant of D = 405 ± 10 Hz for framework aluminum atoms in the aluminophosphate AlPO-5. The corresponding Al-P distance was 315 ± 3 pm. Kao and Chen [86] investigated zeolite H-Beta dealuminated by ammonium hexafluorosilicate. In this work, the REDOR experiment shown in Fig. 3c was utilized with 27 Al nuclei as spins I and 19 F nuclei as spins S. The studies focused on the REDOR behavior of extra-framework aluminum species occurring in the 27 Al MAS NMR spectrum at 0 ppm. The REDOR fraction obtained could be described by a dipolar coupling constant calculated with an Al−F distance of 180–220 pm. This finding indicates that the aluminum atoms occurring in extra-framework clusters of the dealuminated zeolite H-Beta under study are directly bound to fluorine atoms. C Solid-State 17 O NMR Spectroscopy of Catalysts Due to the low natural abundance of 17 O nuclei of 0.037%, 17 O NMR spectroscopy of solid catalysts requires an isotopic enrichment. This enrichment can be performed by, e.g., treatment of the powder material in a reactor with H17 2 O vapor at 523 K for some hours [21]. Early 17 O NMR spectroscopic studies of zeolites Na-A and Na-Y, dealuminated zeolites H-Y, gallosilicates and aluminophosphates were performed by Timken and co-workers [87, 88] under static conditions. In this case, the spectra of aluminosilicate-type zeolites consist of broad quadrupolar patterns corresponding to a quadrupolar coupling constant, CQCC , of

923

about 4.6–5.2 and 3.1–3.2 MHz due to oxygen atoms in Si−O−Si and Si−O−Al bridges, respectively. The asymmetry parameters (ηQ ) were found to be 0.1 for Si−O−Si and 0.2 for Si−O−Al bridges. The isotropic chemical shifts, δiso , referenced to H17 2 O, were determined as 44–57 ppm for Si−O−Si and 31–45 ppm for Si−O−Al bridges [87]. Simulation of the 17 O NMR spectra of aluminophosphates and gallosilicates gave the following parameters: CQ = 5.6–6.5 ppm, ηQ ≈ 0 and δ = 61–67 ppm for Al−O−P bridges and CQ = 4.0–4.8 ppm, ηQ ≈ 0.3 and δ = 28–29 ppm for Si−O−Ga bridges [88]. These results were verified by MAS and variable-angle spinning (VAS) experiments in different magnetic fields. Recently, Peng et al. [89] performed a selective 17 O NMR study of framework oxygen atoms contributing to the local structure of bridging OH groups (SiOHAl) in zeolite H-Y. By application of a 17 O−1 H REDOR experiment, the signal of the above-mentioned oxygen atoms was identified at an isotropic chemical shift of 28 ppm, with a CQCC value of 6.6 MHz and an asymmetry parameter of 0.8. Nowadays, solid-state 17 O NMR investigations of catalysts are performed applying the DOR and MQMAS technique [21, 90–93]. A prerequisite for obtaining well resolved 17 O DOR and MQMAS NMR spectra is the presence of ordered local structures, such as those occurring in siliceous zeolites and in zeolites with an alternating arrangement of silicon and aluminum atoms at T-positions (nSi /nAl = 1). In these cases, the different signals of 17 O nuclei at crystallographically non-equivalent positions appear as well resolved lines in the 17 O DOR and MQMAS NMR spectra. Studying a siliceous zeolite FER by 17 O MQMAS NMR spectroscopy, Bull et al. [90] found up to 10 different signals of oxygen atoms at crystallographically non-equivalent positions in Si−O−Si bridges. While the isotropic chemical shifts of these signals cover a range of 28.0–43.1 ppm, only a slight difference in the quadrupolar coupling constants being in the range 5.22–5.62 MHz was determined. Pingel et al. [21] investigated zeolites Na-A and NaLSX (nSi /nAl = 1) with 17 O DOR and MQMAS NMR spectroscopy. The significant effect of the DOR technique in comparison with the MAS technique for the study of quadrupolar nuclei is demonstrated in Fig. 9. Whereas the 17 O MAS NMR spectrum of zeolite Na-LSX (n /n = 1) Si Al consists of a single broad signal, application of the DOR technique leads to splitting of the spectrum into three different signals [21]. However, it is well known that the structure of zeolite X is characterized by four crystallographically non-equivalent oxygen positions [94]. Therefore, the peak in the middle position must be a superposition of two signals of oxygen atoms at different positions. References see page 930

924

3.1 Physical Properties

(a)

17

(b)

17

The signals obtained along the δ2 -axis of the twodimensional spectrum in Fig. 10a correspond to the MAS NMR spectrum (projection on top). These signals are affected by second-order quadrupolar interactions. The signals obtained along the δiso -axis (projection at the lefthand side) are comparable to those in the DOR NMR spectrum. In this dimension, the anisotropic secondorder quadrupolar line broadenings are averaged by the multiple-quantum experiment. In Fig. 10b, slices cut parallel to the δ2 -axis at different δiso values are depicted. These slices correspond to the separated signals of the MAS NMR spectrum shown as a projection on top. The simulation of the signals obtained by slices allows the determination of the quadrupolar parameters of the corresponding 17 O nuclei. In this way, quadrupolar coupling constants of CQCC = 3.2–3.6 MHz and asymmetry parameters of η = 0.15–0.4 were obtained for the 17 O nuclei located at the four oxygen positions in zeolite Na-LSX [21]. Data obtained by 17 O MQMAS NMR spectroscopy of 17 O nuclei at crystallographically non-equivalent oxygen positions in zeolites Na-A and Na-LSX yielded a correlation between the isotropic chemical shift, δiso , and the Si−O−Al bond angle, β

O MAS NMR

O DOR NMR

80

60

40

20

0

−20

d17 / ppm O

17 O NMR spectra of zeolite Na-LSX (n /n = 1) recorded Si Al with the MAS (a) and DOR (b) technique in an external magnetic field of B0 = 11.7 T [21].

Fig. 9

The potential of the MQMAS technique is demonstrated in Fig. 10, which shows the triple-quantum (3Q) MAS NMR spectrum of 17 O nuclei in zeolite Na-LSX. In the case of 3QMAS, multiple-quantum transitions {+3/2 ↔ −3/2} are excited by the first pulse of the NMR experiments in Fig. 2. The excitation of triple-quantum transitions leads to spectra with a better signal-to-noise ratio than spectra obtained upon excitation of quintuplequantum transitions {+5/2 ↔ −5/2}, which could also be performed for 17 O nuclei with spin I = 5/2.

δiso = −0.71β + 143.7 ppm

(29)

with a correlation coefficient of r = 0.93 [21]. Equation (29) may be used to calculate Si−O−Al bond angles β in materials where determination of the local structure by X-ray diffraction is not possible.

diso = 41.1 ppm

20 diso / ppm

40

diso = 45.9 ppm diso = 49.4 ppm diso = 53.6 ppm 60

40

20

d2 /ppm

60

60 (a)

40

δ2 / ppm

20 (b)

(a) Sheared 17 O MQMAS NMR spectrum of zeolite Na-LSX (nSi /nAl = 1) recorded in an external magnetic field of B0 = 17.6 T. (b) Slices parallel to the δ2 -axis at isotropic chemical shift values of δiso = 41.1–53.6 ppm [21].

Fig. 10

3.1.3 Structure and Morphology

Solid-state 17 O NMR spectroscopy of non-crystalline solids may be complicated by the distribution of the bond geometries in the local structure of the resonating nuclei and, therefore, of the quadrupolar parameters. Kozhevnikov and co-workers [95, 96] investigated various heteropoly acids (HPA) by 17 O MAS NMR spectroscopy in a magnetic field of B0 = 9.4 T. In the case of a molybdenum-containing HPA, they distinguished 17 O MAS NMR signals of oxygen atoms in terminal Mo=O species at 941 ppm and in two bridged Mo−O−Mo species at 566 and 549 ppm (referenced to H17 2 O). The above-mentioned resonance positions of 17 O nuclei are strongly affected by paramagnetic shifts due to the neighborhood of paramagnetic molybdenum atoms. Oxygen atoms in nanophase MgO were studied by Chadwick et al. [97] via 17 O MAS NMR spectroscopy. The spectra of materials with different particle sizes were described by three different signals: A narrow signal at 47 ppm due to bulk MgO, a narrow signal at 41 ppm of bulk MgO with oxygen atoms bound to residual – CH3 groups and an underlying much broader signal caused by hydroxyl oxygens at the surface of the particles. Van Eck et al. [98] applied spin-echo and triple-quantum MAS NMR experiments for the characterization of 17 O nuclei in silica made via the sol–gel technique in an external magnetic field of B0 = 7.0 T. Using the techniques mentioned above, the signals of oxygen atoms in the Si−O−Si bridges and in the SiOH fragments could be separated. For the former signal, a quadrupolar coupling constant of CQCC = 5.3 MHz, an asymmetry parameter of η = 0.0 and an isotropic chemical shift of δiso = 42 ppm were determined. The SiOH fragments caused signals in a wide range of CQCC values and at isotropic chemical shifts of δiso = 0 ± 20 ppm. The behavior of the latter signals is explained by the abovementioned distribution of the bond geometries in the local structure of the SiOH fragments. 3.1.3.7.6 NMR Investigations of Heterogeneously Catalyzed Reactions A H−D Exchange Between Adsorbate Molecules and Solid Acid Catalysts Investigated by In Situ 1 H and 2 H MAS NMR Spectroscopy Under Batch Conditions H−D isotope exchange between hydroxyl groups of solid acid catalysts and adsorbate molecules can provide valuable information, not only for evaluating the acid strength of the hydroxyl groups, but also for studying the mechanisms of heterogeneously catalyzed reactions on these solid acids [99–108]. Characterization of the surface acidity of solid acids by studying the H−D exchange can be achieved by 1 H MAS NMR spectroscopy, which is described in Chapter 3.2.4.4 of this Handbook. In this section, some applications of in situ 1 H and 2 H MAS NMR spectroscopy

925

for studying the H−D exchange during heterogeneously catalyzed reactions are described. Stepanov et al. [99] investigated the H−D exchange of propane-d8 on acidic zeolite H-ZSM-5 within the temperature range 457–543 K by means of in situ 1 H MAS NMR spectroscopy under batch conditions using samples fused in glass ampoules. The reaction temperature was controlled with a Bruker BVT-200 variable-temperature unit and the temperature inside the sample volume was calibrated with an accuracy of ±2 K [99]. Figure 11 shows the stack plot of the in situ 1 H MAS NMR spectra of propane-d8 adsorbed on zeolite H-ZSM-5 measured in sequence at 519 K. The increasing intensities of propane signals from both methyl (CH3 , 1.0 ppm) and methylene groups (CH2 , 1.45 ppm) indicate the proton transfer from the acidic hydroxyl groups to the deuterated propane molecules. The ratio between the signal intensities of methyl and methylene groups is 3 : 1 in the last spectrum. Therefore, both methyl and methylene groups of propane are involved in the H−D exchange in a nonregiospecific manner. The rates of hydrogen exchange are, however, significantly different for methyl and methylene groups, as indicated by experiments carried out at 457, 473 and 519 K. By plotting the exchange 1H

MAS NMR

CH3

CH2

2.0

1.0

0.0

d1 / ppm H

Stack plot of the 1 H MAS NMR spectra of the methyl (1.0 ppm) and methylene (1.45 ppm) groups of propane-d8 on zeolite H-ZSM-5 at 519 K. The first spectrum (bottom) was recorded at t = 3 min and the last spectrum (top) at t = 5 h. The time between subsequent spectra was 5 min [99].

Fig. 11

References see page 930

926

3.1 Physical Properties

rates as a function of temperature (Arrhenius plot), apparent activation energies of 108 and 117 kJ mol−1 were determined for the hydrogen exchange in the methyl groups and the methylene group of propane, respectively. A pentavalent transition state was proposed to rationalize the experimental observations. The different exchange rates for methyl and methylene groups were explained by the different reactivities for protonation of primary and secondary C−H bonds in alkanes [99]. Haouas et al. [102] applied in situ MAS NMR spectroscopy to investigate the initial stage of propane activation on Al2 O3 -promoted sulfated zirconia (SZA) within the low-temperature range 303–375 K. In this case, the samples were prepared in gas-tight MAS NMR rotors and heated in a variable-temperature MAS NMR probe. Timeresolved 1 H MAS NMR experiments during the reaction of propane-d8 on Al2 O3 -promoted SZA at 355 K showed that the signal of methyl groups increased at first, whereas the signal for methylene groups appeared much later. This finding indicates a regioselective H−D exchange 1H

between the Al2 O3 -promoted SZA catalyst and the methyl hydrogen atoms of propane [102]. Activation energies for the H−D exchange of the methyl group and the D-scrambling of the methylene group were determined as 54 and 78 kJ mol−1 , respectively [102]. The difference of ca. 24 kJ mol−1 may account for the additional energy required to generate the adsorbed 1-propylium cations [102]. The intramolecular hydrogen scrambling of propane was further investigated by 1 H and 2 H NMR spectroscopy. Figure 12 shows the 1 H and 2 H MAS NMR spectra of propane-d8 adsorbed on Al2 O3 -promoted SZA after longer contact times at different temperatures. The 2 H NMR spectrum (Fig. 12a, right) obtained before heating consists of signals of CD3 (0.5 ppm) and CD2 (1.0 ppm) groups in an intensity ratio of 3 : 1. In the 1 H MAS NMR spectrum (Fig. 12a, left), no signals occur. After starting the reaction, a substantial decrease of the 2 H MAS NMR signal intensity can be observed with a CDy H(3−y) :CDx H(2−x) ratio below 3 (Fig. 12, right), while the 1 H MAS NMR spectra have corresponding signals in a 2H

MAS NMR

17

MAS NMR 68

83

32 (f)

20

67

80

33 (e)

90

58

10

42 (d) 97

40

3

60 (c) 93

63 37

7 (b)

75 25

4.0

3.0

2.0

1.0

0.0

d1 / ppm H

−1.0 −2.0

4.0 (a)

1 H (left) and 2 H (right) MAS NMR spectra of adsorbed propane-d 8

3.0

2.0

1.0

0.0

−1.0 −2.0

d2 / ppm H

on Al2 O3 -promoted SZA (a) before reaction and after, (b) 50 h at 303 K, (c) 25 h at 321 K, (d) 25 h at 335 K, (e) 16 h at 344 K and (f ) 8 h at 355 K. The relative intensities of protons and deuterons in the methyl (0.5 ppm) and methylene (1.0 ppm) groups are given on top of each spectrum [102].

Fig. 12

3.1.3 Structure and Morphology

ratio much larger than 3 at all temperatures (Fig. 12, left). This finding agrees with a rapid H−D exchange between methyl groups and the solid catalyst and indicates that an intramolecular H−D scrambling between methyl and methylene hydrogen atoms of propane molecules occurs at higher temperatures. Additionally, an intramolecular 13 C scrambling by a skeletal rearrangement process is favored at higher temperatures, which was monitored by 13 C MAS NMR spectroscopy [102]. Beck et al. [103] studied the H−D exchange between benzene-d6 molecules and hydroxyl groups on various acidic zeolites within the temperature range 333–393 K. Benzene-d6 was loaded on the activated zeolite under shallow-bed conditions via a CAVERN device [104] at liquid nitrogen temperature, which is well below the onset of isotopic exchange. Subsequently, the sample was transferred into an MAS NMR rotor, sealed and rapidly placed in a variable-temperature MAS NMR probe. For the H−D exchange of benzene-d6 on zeolites H,NaY, USY and H-ZSM-5, activation energies of 107, 85 and 60 kJ mol−1 , respectively, were obtained [103]. These results were confirmed by Mildner and Freude [105], who applied a high-temperature MAS NMR probe with a laser heating system. They obtained activation energies of 102 and 93 kJ mol−1 for the H−D exchange of benzene-d6 on zeolites H-Y with degrees of cation-exchange of 85 and 92%, respectively [105]. In agreement with results obtained by theoretical calculations, the experimental activation energies indicate that benzene molecules are adsorbed at hydroxyl groups via van der Waals bonds [103]. Therefore, the H−D exchange between benzened6 and hydroxyl groups on solid acid catalysts does not involve free benzenium cations as intermediates [103]. Xu et al. [106] studied the H−D exchange occurring upon adsorption of acetone-d6 on acidic zeolite H-ZSM-5. By 1 H{27 Al} TRAPDOR NMR spectroscopy, the 1 H MAS NMR signal at 16.4 ppm observed after loading acidic zeolite H-ZS-5 with acetone was assigned to bridging hydroxyl protons involved in a strong hydrogen bond with adsorbed acetone molecules [106]. Upon adsorption of acetone-d6 on zeolite H-ZSM-5 at room temperature, a 1 H MAS NMR signal appeared at 2.4 ppm due to protons of methyl groups in adsorbed acetone molecules. The observed H−D exchange between the Brønsted acid sites and adsorbed acetone-d6 molecules was explained by a monomolecular mechanism involving keto–enol isomerization via a concerted action of bridging hydroxyl protons and neighboring framework oxygen atoms in zeolite H-ZSM-5 [106]. 1 H and 2 H MAS NMR spectroscopy under batch conditions was also applied to study the H−D exchange during n-butene conversion on acidic zeolite H-ferrierite within the temperature range 290–373 K [107] and upon

927

adsorption of isobutane on acidic zeolite USY within the temperature range 343–383 K [108]. B Methanol-to-Olefins (MTO) Reaction on Acidic Zeolite Catalysts: Hydrocarbon Pool Mechanism Studied by In Situ Continuous-flow (CF) MAS NMR Spectroscopy Until now, the detailed mechanism of the catalytic conversion of methanol to olefins (MTO) on solid acid catalysts has been a matter of discussion [109–112]. Recent studies verified that the MTO process under steady-state conditions is dominated by a hydrocarbon pool mechanism [112–114]. Among the in situ NMR techniques successfully applied to study the conversion of methanol on acidic zeolites are the stop-and-go method under batch conditions [115, 116], the pulse-quench method [117, 118] and various flow techniques [119–122]. The in situ MAS NMR technique under continuousflow (CF) conditions allows a direct NMR investigation of the formation and transformation of surface species [41] and a simultaneous gas chromatographic analysis of the reaction products [42] under steady-state conditions. To investigate methanol conversion under steady-state conditions by in situ 13 C CF MAS NMR spectroscopy, a flow of 13 C-enriched methanol (13 CH3 OH) was continuously injected into a spinning MAS NMR rotor reactor filled with the calcined zeolite catalyst. Figure 13 shows the in situ 13 C CF MAS NMR spectra obtained during the conversion of 13 CH3 OH on H-SAPO-34 at reaction temperatures of 548–673 K [121]. The formation of hydrocarbons on the working catalyst was evidenced by the 13 C MAS NMR signals appearing in the alkyl region of 10–40 ppm accompanied by broad signals in the olefinic and aromatic region of 129–134 ppm (Fig. 13a–d). It was found through deconvolution of the spectra that the hydrocarbon pool species on the working catalyst under steady-state conditions consist of a mixture of C6 −C12 aromatics and olefins, such as hexenes, hexadienes and polymethylbenzenes [121]. To study the organic species occluded in H-SAPO-34 after methanol conversion, the working catalyst was further purged with dry carrier gas (nitrogen) at 673 K. The spectrum obtained after purging consists of signals at 22 and 129 ppm (Fig. 13e), indicating the presence of alkylated aromatic compounds occluded in the cages of H-SAPO-34 [121]. In order to identify the catalytic role of the hydrocarbon pool in the MTO process on zeolite H-ZSM-5, in situ 13 C CF MAS NMR spectroscopy was performed with an alternating flow of 13 CH3 OH and 12 CH3 OH [114]. After the conversion of 13 CH3 OH under steady-state conditions, the reactant flow was switched to 12 CH3 OH without changing the reaction parameters. Assuming that the References see page 930

928

3.1 Physical Properties

30 − 32

T = 548 K

switching from 13 CH3 OH to 12 CH3 OH, as depicted in Scheme 1. After the reactant flow had been switched from 13 CH OH to 12 CH OH for 1 h, the 13 C isotope abundance 3 3 of the alkyl groups bound to the compounds of the hydrocarbon pool had decreased by ca. 40% [114]. The 1 H CF MAS NMR spectra, however, did not show any change in the total number of hydrogen atoms contributing to the hydrocarbon pool [114]. This finding provides direct evidence that the hydrocarbon pool does play an active role in the MTO reaction under steady-state conditions, in agreement with the work of Arstad and Kolboe [113] and Haw et al. [112]. Under steady-state conditions, continuous methylation of arenes occurs, which can split off ethene or propene under the reaction conditions in a further step. Subsequently, these alkylbenzenes are methylated again and are, therefore, the reactive compounds of the hydrocarbon pool. The in situ MAS NMR–UV/visible technique, shown in Fig. 4, provides complementary information on the working catalyst inside the MAS NMR rotor reactor [43]. Whereas NMR spectroscopy allows the identification of alkyl signals in more detail, UV/visible spectroscopy offers high sensitivity for detecting the formation of aromatic compounds and unsaturated carbenium ions [43]. This technique was recently applied to study the formation of hydrocarbons and carbocations during the conversion of methanol on a weakly dealuminated zeolite H-ZSM-5 [43]. The 13 C CF MAS NMR spectrum (Fig. 14a, left) recorded during the conversion of 13 C-enriched methanol at 413 K consists of signals at 51 and 61 ppm due to methanol and dimethyl ether, respectively. In the simultaneously recorded UV/visible spectrum (Fig. 14a, right), bands occur at 275, 315 and 375 nm, which indicates the formation of neutral aromatic compounds, mono- and dienylic carbenium ions, respectively [123–125]. The conversion of ethene (13 C isotope in natural abundance) was further studied on the working catalyst after methanol conversion (Fig. 14b). 13 C MAS NMR signals at 14, 23 and 32 ppm (Fig. 14b, left), occurring after the conversion of ethene at 413 K for 1 h, are due to alkyl groups of alkylated cyclic compounds. The simultaneously recorded

18 − 25 11

(a)

T = 573 K

40 132 − 134

(b) 20

T = 623 K 132 − 134

(c) 20

T = 673 K

129

(d)

T = 673 K, purging with N2

22

129

200

150

100

d13c

(e)

50

0

−50

/ ppm

In situ13 C CF MAS NMR spectra of calcined H-SAPO-34 recorded during the conversion of 13 CH3 OH (W/F = 25 g h mol−1 ) at reaction temperatures of 548 (a) to 673 K (d). Spectrum (e) corresponds to spectrum (d), but was recorded after purging the working catalyst with dry nitrogen at 673 K [119].

Fig. 13

alkyl groups of the hydrocarbon pool are involved in the conversion of methanol, e.g. by adding methanol and splitting off products such as ethylene, the 13 C isotope abundance of these groups would decrease strongly after

(n+2)

13CH

3OH

(13CHx)n

H

H

13C

13C

H

H

H (13CHx)n

H

(13CH

x)n

12C

H Scheme 1

13C

H H

2 12CH3OH

H +

H 12C

H

H + 2H2O

H 13C

H

3.1.3 Structure and Morphology

13C

UV / visible

CF MAS NMR 51

275

61

929

315

375

23

13CH OH 3

× 40

300

(a) 61

13CH OH + 3 12CH = 12CH 2 2

32

23 14

375 × 10

300 370

(b)

12CH = 12CH 2 2

34

120

450

14

120 200 160 (c)

24

80

40

0

−40

δ13 / ppm C

250

300

350

400

450

500 550

l / nm

In situ 13 C CF MAS NMR (left) and UV/visible (right) spectra of dealuminated zeolite H-ZSM-5 recorded during conversion of −1 12 12 3 OH under continuous-flow conditions (W/F = 25 g h mol ) at 413 K for 2 h (a), during a subsequent conversion of CH2 = CH2 (W/F = 10 g h mol−1 ) at 413 K for 1 h (b) and during conversion of 12 CH2 =12 CH2 (W/F = 10 g h mol−1 ) at 413 K on a fresh catalyst for 2 h (c) [43]. Asterisks denote spinning sidebands. The narrow peaks at ca. 500 nm in the UV/visible spectra are due to an artifact caused by the equipment. Fig. 14 13 CH

UV/visible spectrum (Fig. 14b, right) shows bands at 300 and 375 nm, which indicate the formation of neutral cyclic compounds and dienylic carbenium ions, respectively [123–125]. As a control experiment, the spectra recorded after the conversion of ethene on a freshly dealuminated zeolite H-ZSM-5 are shown in Fig. 14c. The 13 C MAS NMR spectrum consists of signals at 14, 24 and 34 ppm (Fig. 14c, left). Additionally, broad signals occurred at ca. 120 ppm, i.e. in the chemical shift range of olefinic and aromatic compounds. The UV/visible spectrum consists of bands similar to those in Fig. 14b and an additional weak band at ca. 450 nm. The latter were attributed to benzenium ions or trienylic carbenium ions in an alkylated form as found by 13 C MAS NMR spectroscopy [43]. The result is in agreement with the work of Karge et al. [125], who studied methanol conversion on dealuminated zeolite H-ZSM-5. Extra-framework aluminum species acting as Lewis acid sites were claimed to accelerate the formation of hydrocarbons and carbenium ions [125]. Conclusions Due to the development of new techniques and the further increase in the magnetic field strength available for commercial applications, solid-state NMR spectroscopy has become a routine method for the

3.1.3.7.7

characterization of solid catalysts in recent decades. As an important advantage, solid-state NMR spectroscopy allows the investigation of the local structure of the nuclei under study in both crystalline and amorphous materials. Active sites in solid catalysts and the specific behavior of the frameworks of these materials often depend on local effects, such as framework defects, the substitution of framework atoms, guest compounds, etc. Therefore, solid-state NMR spectroscopy is nowadays a widely applied analytical method delivering information about solid catalysts, which are complementary to those obtained by other analytical methods. Whereas in early solid-state NMR spectroscopic studies in the field of heterogeneous catalysis the characterization of the catalyst framework was the dominant application, an increasing number of recent studies have focused on the investigation of surface sites, i.e. on the determination of their concentration, strength and accessibility. In this case, the advantage of NMR spectroscopy as a quantitative method can be utilized. New developments are the different ex situ and in situ techniques used to clarify the interaction of probe molecules and reactants with active surface sites on solid catalysts. Furthermore, References see page 930

930

3.1 Physical Properties

some of these approaches are suitable for investigating the mechanisms of heterogeneously catalyzed reactions under conditions that are relevant for industrial processes. Following a general trend of modern in situ methods, the combination of in situ solid-state NMR spectroscopy under continuous-flow conditions with other analytical techniques, such as on-line gas chromatography and UV/visible spectroscopy, was recently introduced to obtain complementary information about reactants and intermediates on working catalysts. As demonstrated in this chapter, the above-mentioned techniques of solidstate NMR spectroscopy make this method an important and valuable tool for research in the field of heterogeneous catalysis. References 1. G. Engelhardt, in Handbook of Heterogeneous Catalysis, G. Ertl, H. Knoezinger, J. Weitkamp (Eds.), Vol. 2, VCH, Weinheim, 1st Ed., 1997, p. 525. 2. M. Stoecker, in AdvancedZeolite Science and Applications, J. C. Jansen, M. Stoecker, H. G. Karge, J. Weitkamp (Eds.), Studies in Surface Science and Catalysis, Vol. 85, Elsevier, Amsterdam, 1994, p. 429. 3. M. Hunger, E. Brunner, in Molecular Sieves–Science and Technology, Vol. 4, Characterization I, H. G. Karge, J. Weitkamp (Eds.), Springer-Verlag, Berlin, 2004, p. 201. 4. G. Engelhardt, D. Michel, High-Resolution Solid-State NMR of Silicates and Zeolites, Wiley, Chichester, 1987, 485 pp. 5. J. F. Haw, Adv. Catal. 1998, 42, 115. 6. M. Hunger, J. Weitkamp, in In-Situ Spectroscopy of Catalysts, B. M. Weckhuysen (Ed.), American Scientific Publishers, Stevenson Ranch, CA, 2004, p. 177. 7. C. P. Slichter, Principles of Magnetic Resonance, SpringerVerlag, Berlin, 1996, 655 pp. 8. R. K. Harris, Nuclear Magnetic Resonance Spectroscopy, Longman Scientific and Technical/Wiley, New York, 1989, p. 236. 9. M. Mehring, Principles of High Resolution NMR in Solids, Springer-Verlag, Berlin, 1983, 342 pp. 10. E. O. Stejskal, J. D. Memory, High Resolution NMR in the Solid State, Oxford University Press, Oxford, 1994, 189 pp. 11. A. Abragam, Principles of Nuclear Magnetism, Clarendon Press, Oxford, 1994, 599 pp. 12. E. R. Andrew, A. Bradbury, R. G. Eades, Nature 1958, 182, 1659. 13. M. M. Maricq, J. S. Waugh, J. Chem. Phys. 1979, 70, 3300. 14. E. Brunner, D. Freude, B. C. Gerstein, H. Pfeifer, J. Magn. Reson. 1990, 90, 90. 15. R. G. Pembleton, L. M. Ryan, B. C. Gerstein, Rev. Sci. Instrum. 1977, 48, 1286. 16. G. Scheler, U. Haubenreisser, H. Rosenberger, J. Magn. Reson. 1981, 44, 134. 17. D. Freude, J. Haase, in NMR Basic Principles and Progress, Vol. 29, P. Diehl, E. Fluck, H. Guenther, R. Kosfeld, J. Seelig (Eds.), Springer-Verlag, Berlin, 1993, p. 1. 18. A. Samoson, E. Lippmaa, A. Pines, Mol. Phys. 1988, 65, 1013. 19. A. Samoson, E. Lippmaa, J. Magn. Reson. 1989, 84, 410. 20. A. Samoson, A. Pines, Rev. Sci. Instrum. 1989, 60, 3239.

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3.1.3.8

Vibrational Spectroscopies

.. Gerhard Mestl and Helmut Knozinger∗

Introduction Vibrational spectroscopies are among the most promising and most widely used methods for catalyst characterization, because very detailed structural information can be obtained from vibrational spectra. Also very important is that several vibrational spectroscopies can be applied under in situ conditions and they can very successfully be used for studies of ill-defined high-surface-area porous materials. In situations where X-ray diffraction techniques are not applicable, vibrational spectroscopies can often provide information on phase transitions and changes in compositions of bulk catalyst materials, their crystallinity and on the nature of surface functional groups. In molecular solids, the individual molecules may be distorted and their symmetry altered compared with the free molecules, leading to modifications of their vibrational behavior. In addition, factor group splitting, unit cell coupling and splitting of longitudinal and transverse optical modes (LO/TO splitting) may occur [1–4]. These effects will also depend on the size and morphology of the crystals. The lattice dynamics are described by the concept of phonons with their energy dispersion and changing symmetry properties in k-space [5–7]. Solid-state vibrational spectra are further complicated by the presence of defects [8, 9], which can give rise to the appearance of additional vibrational modes because of the relaxation of the k-selection rule. This selection rule is also weakened for small crystallites [10, 11] for which surface modes may also become detectable. Vibrational spectra of powders must be considered as a superposition of the spectra of all randomly 3.1.3.8.1

∗

Corresponding author.

oriented crystallites in the specimen and they may be further complicated by inhomogeneous particle size distributions. Obviously, the dynamic properties of solids are also affected by temperature and hydrostatic pressure [12]. This chapter outlines the basic principles and opportunities of vibrational spectroscopies that can be used for the structural characterization of catalytic materials. Representative examples for the vibrational characterization of selected materials such as bulk oxides (including simple binary oxides, multicomponent materials and zeolites and molecular sieves) and supported catalysts are discussed. The vibrational analysis of surface groups, particularly of hydroxy groups, is also addressed. In many cases, surface hydroxy groups are simply formed by dissociative chemisorption of water molecules, thus reducing the surface free energy of, for example, oxides. Hydroxy groups may also be constituents of the solid-state structure, while at the same time accessible to gas-phase molecules, as for example in zeolites. It is for this reason that the vibrational characterization of surface hydroxy groups is included. Basic Principles of Vibrational Spectroscopies Principally, photons, electrons and neutrons can be used as probes. A wide variety of techniques has been developed for photon vibrational spectroscopies, which have to be selected according to the surface and optical properties of the system to be studied. A great advantage of photon spectroscopies (photons as probe and information carrier) is the fact that they can be applied even in the presence of gas phase, in contrast to electron spectroscopies. Because of its relative simplicity and wide applicability, infrared transmission–absorption spectroscopy [13–18a] and diffuse reflectance spectroscopy [13, 17–21] are most frequently used today in catalysis research. Several variants of laser Raman spectroscopy (LRS) [16–18, 22–31] are finding increasing applications, whereas internal reflection [attenuated total reflection (ATR)] [32, 33] and optoacoustic spectroscopies (OAS) [17, 34, 35] are being used less frequently. Sum frequency generation (SFG) has been developed to study surfaces and interfaces [36–39]. It is a highly sensitive technique (applicable to single-crystal surfaces) with surface specificity, because the selection rules permit finite SFG signals only for systems which are lacking inversion symmetry (such as surfaces and interfaces). Electron spectroscopies [electron energy loss spectroscopy (EELS)] have hardly any application to rough surfaces of electrically non-conducting, polydisperse materials and cannot be used in the presence of gas phases. Inelastic electron tunneling spectroscopy (IETS) is a sophisticated technique which can provide high-resolution spectra of model samples. The experiment, however, requires relatively costly equipment and is far less widely 3.1.3.8.2

3.1.3 Structure and Morphology

applicable than, for example, infrared transmission or Raman spectroscopy. Inelastic neutron scattering is discussed in detail in Chapter 3.1.3.9. A recent review was published by Albers and Parker [39a]. A Infrared Transmission–Absorption Spectroscopy The KBr technique is routine for IR transmission–absorption spectroscopy of powder samples. However, for in situ investigations, this technique is not applicable and pressed self-supporting wafers have to be used in this case. The applicability of the transmission technique is determined by the properties of the solid powder under study. Thus, samples which exhibit only weak bulk absorption and which have average particle sizes d which are smaller than the wavelength of the infrared radiation in the region of interest will be optimally suited for the transmission mode. The particle size condition (λ > d) which determines the wavelength range of suitably low scattering losses is usually met in the mid- and far-infrared regions, whereas scattering losses become strongly involved in the near-infrared region. However, most samples show strong bulk absorption in the low-wavenumber region (roughly 770 K) and low-temperature (80 K) spectroscopy has recently been designed by Zeilinger et al. [56]. Stair [56a] described a fluidized bed reactor for UV–Raman studies of working catalysts. Some major problems which may be encountered in laser Raman spectroscopy include: (i) sample sensitivity to heating effects of the laser beam, which may become very severe when colored samples are studied (ii) low sensitivity of the technique (iii) background fluorescence, a problem which is sometimes so severe on oxide surfaces that weak Raman signals remain undetectable. The heating effects of the laser beam can be reduced by simply applying low laser power levels (400 A), in the Raman spectrum at 685 and 100 cm−1 , which were suggested to arise from relaxation of selection rules due to surface effects [149]. Everall et al. [155] were able to correlate linearly the crystallite size with the relative intensity of the band at 110 cm−1 . Siegel and coworkers [156, 157] also studied nanophase TiO2 (Fig. 2) and found that the Raman spectrum was strongly References see page 965

3.1 Physical Properties

influenced by defect structures within the material. They were able to correlate Raman band shifts with the degree of oxygen deficiency in TiO2−x (Fig. 3). 19 ). Anatase crystallizes in the space group I 41 /amd (D4h −1 Raman bands at 144, 197, 399, 519 and 639 cm were reported [158–160]. Factor group analysis predicts six Raman-active and three IR-active modes for anatase [161]. Phase transitions between low- and high-pressure TiO2 modifications have been studied by Raman spectroscopy [162–165]. The temperature-induced anatase to rutile transformation at 1070 K has been followed by Raman spectroscopy [154, 166] and Balachandran and Eror [154] showed that the rutile phase can be stabilized at temperatures below 723 K. d Zirconium Oxides Zirconia (ZrO2 ) crystallizes in two metastable modifications (cubic and tetragonal phases) and in a stable monoclinic phase. The metastable cubic zirconia modification is usually stabilized by addition of Ti4+ , Y3+ , Ce3+ , Cr3+ , Mg2+ or Ca2+ . This cubic modification exhibits Raman bands at 490 (w) and 610 cm−1 [167, 168]. The intensity of the prominent Raman band at 610 cm−1 increased with increasing Y3+ concentration [169]. 15 ) exhibits prominent Raman The tetragonal phase (D4h bands at 147 (B1g ), 261 (Eg ), 320 (B1g ), 464 (Ag ) and 642 (Eg ) cm−1 [168, 170, 171], whereas the monoclinic

A

B

100

300

500

440

430

420

154

150

146

142 0.12

0.08

0.02

0

x Dependence of peak positions of TiO2 Raman bands on the degree of oxygen deficiency x (adapted from Ref. [157]).

Fig. 3

5 ) shows bands at 174, 186, 219, 305, 334, modification (C2h 347, 380, 476, 502, 537, 559, 616 and 638 cm−1 . Raman spectra of these two zirconia modifications are compared in Fig. 4. The phase transitions between tetragonal and monoclinic ZrO2 have been followed by in situ Raman spectroscopy [167, 168, 170, 171] and the influence of pressure was also demonstrated [172]. The pressure dependence of band positions may be calibrated for the determination of the hydrostatic pressure in a reactor.

e Magnesium Oxide MgO crystallizes in the NaCl structure and frequently forms cubes terminated by well-defined (100) faces. For symmetry reasons, firstorder Raman scattering is not allowed, but second-order Raman spectra were reported [173]. MgO microcrystals having diameters smaller than 100 nm exhibited a firstorder Raman spectrum, which was probably induced by symmetry reduction caused by distortions in the surface region of the particles [174].

700

Raman shift /cm−1

Raman spectra of nanophase TiO2 : (A) as prepared; (B) after annealing in air at 970 K (this sample is considered to be stoichiometric) (adapted from Ref. [157]).

Fig. 2

450

Wavenumbers /cm−1

940

f Molybdenum Oxides MoO3 crystallizes in the space 16 ) with four formula units in the unit cell, group Pbnm (D26 which has dimensions of 396.28 pm (a axis), 1385.5 pm (b axis) and 369.64 pm (c axis) [175]. The orthorhombic unit cell extends over one layer and two half-layers above and

3.1.3 Structure and Morphology

Intensity

A

B

0

200

400

600

800

∆nR/cm−1 Raman spectra of (A) tetragonal and (B) monoclinic zirconia.

Fig. 4

below the latter. The MoO3 layers are separated by a van der Waals gap of about 700 pm. The Mo atoms in MoO3 have adjacent oxygen neighbors at distances between 167 and 196 pm, whereas the distances of the remaining two oxygen atoms that complete a distorted octahedron are as long as 225 and 233 pm. The four close oxygen neighbors form a distorted tetrahedral arrangement around the metal center (Fig. 5). The structure of MoO3 can therefore be considered as being built up from MoO4

941

tetrahedra, each sharing two corner oxygen atoms with two neighboring tetrahedra to form chains running in the c direction. Crystals of MoO3 typically grow in the form of needles or platelets with this direction as the principal axis. Figure 6a shows a Raman spectrum of polycrystalline MoO3 , which exhibits characteristic bands at 94 (B1g ), 114 (B2g ), 125 (B3g ), 157 (Ag , B1g ), 195 (B2g ), 217 (Ag ), 244 (B3g ), 289 (B2g , B3g ), 335 (B1g , Ag ), 366 sh (Ag ), 376 (B1g ), 469 (Ag , B1g ), 663 (B2g , B3g ), 816 (Ag , B1g ) and 993 (Ag , B1g ) cm−1 . The assignment of the bands was based on results of a single-crystal study and valence force field (VFF) analysis by Py and coworkers [177, 178], in which the structure of MoO3 proposed by Kihlborg [175] (see Fig. 5) was used. Small frequency shifts and variations of relative intensities were recognized as compared with single-crystal spectra in the spectra of the polycrystalline material in the spectral region of the rigid chains below 200 cm−1 . These effects are probably a consequence of the microcrystalline nature of the MoO3 sample that gave the spectrum of Fig. 6a. The VFF calculations [178] indicated that the frequencies of rigid-chain modes decrease with decreasing interactions between the chains within the half-layers of the structure, which may be related to perturbations of the layer structure. Interestingly, analogous frequency shifts in the rigid-chain mode regime were also detected when MoO3 crystallites were disintegrated mechanically in a ball-mill [179]. References see page 965

167 pm 225 pm 104° 195 pm

104°

98°

149° 173 pm 195 pm b

233 pm c a

Coordination around the Mo centers in MoO3 and section of the layered MoO3 structure: chains of MoO4 tetrahedra along the c axis form half-layers in the ac plane. Two half-layers are stapled along the b axis (adapted from Ref. [175]).

Fig. 5

279

3.1 Physical Properties

100

990

888

660

954

724

472

370

75

469

500

94 110 120 142 156 193 234

993

663

335 376 366

94

125 114 157 195 217 244

289

Intensity/a.u.

334

818

816

942

900

100

500

900

Raman shift/cm−1 (a)

(b)

(a) Raman spectrum of polycrystalline MoO3 (spectrum recorded at room temperature). (b) Raman spectrum of oxygen-deficient MoO3−x obtained by heating MoO3 in vacuum (0.1 Pa) at 648 K for 10 h (spectrum recorded at 648 K) (adapted from Ref. [176]).

Fig. 6

Relative band intensities were also shown to depend on the crystallite orientation. For example, Ozkan et al. [180] found increasing intensity of the Raman band at 816 cm−1 (symmetric stretch of terminal oxygen atoms) relative to that of the band at 993 cm−1 (antisymmetric stretch of terminal oxygen atoms) when the sample contained increasing amounts of crystallites with high basal-to-etch plane [i.e. (010): (100)] ratios. The thermal expansion of MoO3 is anisotropic. As a consequence, the Raman scattering tensor becomes temperature dependent [176]. This effect can also influence the relative Raman band intensities in addition to the population-induced intensity changes, when Raman spectra are recorded at elevated temperatures. When MoO3 is thermally treated in vacuum at elevated temperatures, non-stoichiometric MoO3−x is formed. Figure 6b shows a Raman spectrum as an example of a sample that had been heated in vacuum to 648 K. The blue color of MoO3−x results from the polaron conductance (intervalence transitions) of these suboxides [181] and leads to an increased absorption coefficient. This and the smaller Raman scattering cross-section of suboxides should be the reason for the lower signal-to-noise ratio of the spectrum of the suboxide (Fig. 6b) as compared with that of the stoichiometric MoO3 (Fig. 6a). The non-stoichiometry leads to negative frequency shifts by several wavenumbers in the regime of the translational (c direction) rigid-chain modes, suggesting perturbations of the structure that result in vibrationally less interacting chains, consistent with the VFF calculations

of Py and Maschke [178]. The variations of Mo−Mo distances in suboxides as reported by Kihlborg [182] are also in line with this observation. The spectrum in Fig. 6b shows an intense background over the entire frequency range up to 1000 cm−1 . The loss of oxygen from the MoO3 lattice leads to ordered crystallographic shear structures [182, 183] via extended shear defects [184], thus destroying the translational symmetry. The observed background intensity may therefore arise from a multitude of combination modes of acoustic and acoustic plus optical modes [176]. Three new Raman bands in the Mo−O stretching region, namely at 724 (vw), 888 and 954 (vw) cm−1 (see Fig. 6b) indicate significant modifications of the oxygen coordination spheres around the Mo centers in the MoO3−x sample. Oxygen-deficient MoO3 contains tetrahedra along the shear planes and the distortion of the coordination polyhedra is decreasing from undisturbed parts of the MoO3 lattice towards the shear planes [182]. The appearance of the bands at 954 and 888 cm−1 in MoO3−x (Fig. 6b) suggests the presence of octahedra being less distorted than in stoichiometric MoO3 and of tetrahedral coordination spheres, respectively. The band at 888 cm−1 falls in the typical frequency range, in which Mo−O stretching modes of tetrahedrally coordinated species are observed (Table 2). More severely reduced suboxides have block or columnar structures of ReO3 type and contain even pentagonal bipyramids [185, 186], which should lead to more significant modifications of their Raman spectra. MoO2 has the rutile structure and gives rise to Raman

3.1.3 Structure and Morphology

Tab. 2

943

Vibrational bands (v/cm−1 ) of solid molybdenum–oxygen compounds

Species

(NH4 )4 [Mo8 O26 ] · 4H2 O (NH4 )6 [Mo7 O24 ] · 4H2 O

(NH4 )2 [Mo2 O7 ] H4 [SiMo12 O40 ]

Mo=O stretch

Mo−O−Mo asym. stretch

Mo−O−Mo sym. stretch

Mo=O bending

Mo−O−Mo deformation

Ref.

967–947

846

600

468

236

[188]

918 912 934

670 860

574 526 630

365 338 450

218 152 245

[189]

908 890 880

840

550

220 115

937 910 982 962 915SiO

720

415 363 340 308 456 365 467 393 371 334 416

(MoO6 )− (in Ba2 CaMoO6 )

892 880 805

680 620 504

812 650 875

(MoO4 )2− (in K2 MoO4 )

650

291

[190]

290 246 210 160

[191–193]

[194]

320

[195]

311 364 354 327

[196] [197]

820 MoOF4 MoOCl4

1039 1008

CaMoO4 Na2 MoO4 Al2 MoO4

882 892 1006

bands at 203, 228, 345, 363, 461, 495, 571, 589 and 744 cm−1 [187]. Molybdenum–oxygen vibrations of several molybdenum–oxygen compounds which have relevance as catalysts or as reference compounds for the interpretation of spectra of supported catalysts (see Section 3.1.3.8.4) are summarized in Table 2. g Vanadium Oxides Tetrahedral VO3− 4 , as present in aqueous solution, has the most prominent Raman band at 827 cm−1 . This band is associated with the symmetric stretching mode [199, 200]. For solids with isolated VO3− 4 units, this band lies in the range 829–845 cm−1 [198]. The laser Raman spectra of [V10 O28 ]6− in aqueous solution and in some solids show two prominent bands at 994 and 970 cm−1 [198]. The idealized overall symmetry of the polyanion is C2h , the structure is centrosymmetric and built up by an array of VO6 octahedra [201] which are edge-sharing and strongly distorted [202]. The IR and Raman spectra of magnesium orthovanadate (Mg3 V2 O7 ) and magnesium metavanadate (MgV2 O6 ) have been reported [203].

390

[198] [198] [24]

13 [204, 205]. CornerV2 O5 crystallizes in space group D2h sharing VO4 tetrahedra form layers in the ac plane; one short V=O bond is oriented along the b direction. The results of a factor group analysis of V2 O5 [206] are summarized in Table 3.

h Tungsten Oxides The irreducible representation of the tetrahedral WO2− 4 ion contains four fundamental modes, v1 (A1 ) + v2 (E) + v3 (F2 ) + v4 (F2 ), and all are Raman active. These modes are observed at 931 (A1 ), 324 (E), 833 (F2 ) and 324 cm−1 (F2 ). Two of these correspond to the stretching modes (931 cm−1 A1 symmetric stretch, 833 cm−1 F2 antisymmetric stretch) while the two bending modes are degenerate and appear at 324 cm−1 . Crystalline Na2 WO4 has a spinel structure with a tetrahedral site symmetry [207] and the Raman bands are observed at 928 (A1 ), 313 (E), 813 (F2 ) and 373 cm−1 (F2 ). Distortions from the ideal tetrahedral structure would lead to removal of degeneracy giving rise to splitting of bands. For example, calcium tungstate has a scheelite References see page 965

944

3.1 Physical Properties Highest Raman frequencies (v/cm−1 ) for tetrahedral and octahedral tungsten–oxygen compounds (adapted from Ref. [213])

Raman frequencies (v/cm−1 ) of V2 O5 (D13 )a (adapted 2h from Ref. [206])

Tab. 4

ag

b1g

Octahedral compounds

996

996

Tab. 3

530 483 406 306

202

105

105

cryst

b3g

703

703

285

287 231

146

147

406 311

199

a

b2g

Assignment v(OM) v(OM3 ) (high) v(OM3 ) (low) δ(OM2 ) δ(OM) δ(OM3 ) δ(OM) δ(OM2 ) δ(OM2 )n δ(OM2 )n V−V mode

= 3au + 3b1u + 6b2u + 6b3u + 7ag + 7b1g + 3b2g + 4b3g .

structure [208] where the WO2− 4 group is slightly distorted from the ideal tetrahedron along a line perpendicular to the two opposite edges and the site symmetry is S4 . The correlation between Td (tetrahedral) and S4 can be represented as: A1 → A, E → A + B and F2 → B + E, which shows the removal of degeneracy. Accordingly, Raman bands for CaWO4 are observed at 922 (A), 838 (B), 794 (E), 403 (B), 334 (A), 210 (A) and 118 cm−1 (E). Al2 (WO4 )3 possesses both distorted and undistorted tetrahedra [209]. Octahedrally coordinated polyhedra of the WO6 type exhibit only three Raman-active modes. One of a few cases of an ideal octahedron reported in the literature is Li6 WO6 [210]. The three fundamental Raman-active modes occur at 740 (A1g ), 430 (Eg ) and 360 cm−1 (F2g ). Crystalline WO3 possesses the cubic ReO3 structure [211], the simplest three-dimensional structure formed from vertex-sharing octahedral groups. At high temperatures, the structure is undistorted, whereas at moderate temperature the lattice is distorted. The abundance of peaks in the Raman spectrum is a consequence of the appreciable distortion of the WO3 structure from the ideal octahedral arrangement of ReO3 . The major vibrational bands are located at 808, 714 and 276 cm−1 and have been assigned to the W−O stretching mode, the W−O bending mode and the W−O−W deformation mode, respectively [206]. Other minor bands are at 608, 327, 248, 218, 185 and 136 cm−1 . Generally, the position of the highest Raman band reflects the highest bond order (shortest W−O bond) present in the tungsten oxide structure [204, 206]. Thus, regular tetrahedral groups (WO4 ) would show Raman bands at higher frequencies than octahedral (WO6 ) groups. However, distortions in the structure would change the bond orders and shift the Raman bands, resulting in much overlap. Table 4 represents the highest

Li6 WO6 WO3 FeWO4 CoWO4 MnWO4 NiWO4 CuWO4 H2 WO4 (NH4 )6 H2 (W3 O10 )4

vmax /cm−1

740 805 856 886 886 893 908 951 980

Tetrahedral compounds CaWO4 Cs2 WO4 Na2 WO4 BaWO4 Fe2 (WO4 )3 Al2 (WO4 )3 (WO4 )2−

vmax /cm−1

913 920 928 940 1026 1060 931

wavenumber Raman bands of several tungsten oxygen compounds. It can be seen from Table 4 that, whereas octahedra have bands below 910 cm−1 , the bands are above 980 cm−1 in tetrahedral structures. The extreme case of ammonium tungstate possesses WO6 octahedra of significant distortion and the symmetric stretch appears at 980 cm−1 . A structure that possesses both octahedral and tetrahedral tungsten oxide groups is Na2 W2 O7 [212, 214]. The tungsten–oxygen octahedra share corner oxygen atoms with adjacent octahedra and are also attached to the tungsten-oxygen tetrahedra; the tetrahedra share oxygen atoms with two different adjacent tungsten–oxygen octahedra in the chain. The Raman bands [212] observed at 957 and 940 cm−1 are assigned to W−O stretching modes of the tetrahedra, whereas those at 853, 763 and 538 cm−1 are associated with the W−O stretching of the octahedra. It is interesting that the Raman bands due to WO2− tetrahedra dominate in intensity over those 4 due to WO6− 6 octahedra which show weaker bands. Dixit et al. [27] cited literature illustrating the nature of W−O vibrations and how the frequencies change with respect to normal coordinates and molecular geometries and upon the phase and environmental condition surrounding the tungsten–oxygen bonds. i Rhenium Oxides Raman bands of several rhenium– oxygen compounds are summarized in Table 5. j Niobium Oxides Niobium oxide exists in different polymorphic forms and the phase transformations of niobium oxide depend strongly on the heat treatment [219]. Niobium oxide can be modified by cation substitution into the crystalline lattice to form different kinds of niobium compounds, such as perovskite structures [220], layered structures [221] and Nb6 O8− 19 clusters [222]. Most of the niobium compounds contain an octahedral NbO6 structure unit with different extents of distortion present

3.1.3 Structure and Morphology

Tab. 5

945

Raman band positions (v/cm−1 ) of selected rhenium–oxygen compounds

ReO− 4 (aq.) [215] 971 916 332 332

NaReO4 [216]

ReO3 F [217]

ReO3 Cl [217]

ReO3 Br [217]

α-Li6 ReO6 [208]

Re2 O7 (gas) [208]

963 928 980 331 331

1009 980

1001 961

997 963

680 505

1009 972

403 321

293 344

350 332

360

341 322 456 185

in the structures. Niobium oxide compounds containing an NbO4 tetrahedral structure unit are extremely rare. The major Raman frequencies of niobium compounds with corresponding structures as classified by Jehng and Wachs [223] are presented in Table 6. species does not exist in The tetrahedral NbO3− 4 solution because of the high electronegativity and small radius of the oxygen atom. For the regular tetrahedral structure, Blasse [224] observed major bands at ca. 816 (v1 ), ca. 650 (v3 ), ca. 420 (v4 ) and ca. 340 cm−1 (v2 ). In the slightly distorted octahedral structures, Raman bands at ca. 870 cm−1 (antisymmetric stretch of the Nb−O−Nb linkage), 620–630 and 520–580 cm−1 (symmetric stretch) corresponding to slightly different Nb−O bond distances and ca. 430 and 375 cm−1 (bending modes) occur. In highly distorted octahedral NbO6 structures, the major Raman bands appear in the 850–1000 cm−1 region. For the heptaniobate and octaniobate structures, the Raman band positions have been shown to be very similar to those in NbO6 structures and to depend on the extent of distortion. Raman spectral correlations of niobium oxide structures and Relationship between structures of niobium–oxygen compounds and Raman frequencies (v/cm−1 ) (adapted from Ref. [223])

Tab. 6

Symmetry

Raman bands/cm−1

Compound

Tetrahedral

790–830

Distorted octahedral (C2v ) (NbO7 and NbO8 )

500–700

YNbO4 YbNbO4 Nb2 O5 (amorphous, TT, T and H)

Distorted octahedral (C4h )

850–1000

LiNbO3 NaNbO3 KNbO3 Nb2 O5 (H) AlNbO4 K8 Nb6 O19

Assignment

vs (Re−O) vas (Re−O) δs (O−Re−O) δas (O−Re−O) vs (Re−O−Re) δ(Re−O−Re)

niobium oxide solution chemistry have been extensively investigated [225]. k Chromium Oxides Chromium–oxygen compounds prefer a tetrahedral coordination of the Cr6+ cation. Various types of tetrahedrally coordinated chromium oxide compounds are known [226–228]. Tetrahedrally coordinated chromium can exist as a monomer (CrO2− 4 ), 2− 2− ), trimer (Cr O ) and tetramer Cr O dimer (Cr2 O2− 3 10 4 13 , 7 both in aqueous solution and in crystalline lattices. The vibrational assignments for these structures are presented in Table 7. The chromate ion has regular tetrahedral (Td ) symmetry in aqueous solution, but in a crystalline phase the symmetry is lower than Td (usually Cs site symmetry). The isolated chromate ion exhibits four Raman-active fundamentals corresponding to the Td irreducible representation. These are at 846 (symmetric stretch), 904 (antisymmetric stretch) and 371 and 348 cm−1 (bending mode). In general, Cr6+ oxide compounds possess strong Raman bands in the 800–1000 cm−1 region assigned to Cr−O stretching modes and bands of moderate intensity in the 300–400 cm−1 region associated with Cr−O bending modes. The dichromate ion contains one oxygen bound to two CrO3 groups, whereas in the trimer two terminal CrO3 groups are linked together by a CrO4 group [227]. The tetramer possesses two terminal CrO3 units connected by a Cr2 O7 unit and crystalline CrO3 has a chain structure of CrO4 tetrahedra with two terminal Cr=O bonds and two bridging Cr−O bonds for each Cr6+ cation [227]. Substitution for the oxygen atoms by monovalent anions (Cl− , Br− ) successively reduces the symmetry from Td [229]. α-Cr2 O3 gives rise to one strong Cr−O stretching vibration at 543 cm−1 [28, 231]. l Antimony Oxides Antimony trioxide (Sb2 O3 ) exists in two crystalline forms, namely the cubic senarmontite, which is stable below 840 K, and the orthorhombic References see page 965

946

3.1 Physical Properties

Tab. 7

Raman band positions (v/cm−1 ) of selected chromium–oxygen compounds

CrO2− 4 (aq.) [226]

CsCrO3 Bra [229]

CrO2 Cl2 a [230]

K2 Cr2 O7 [228]

K2 Cr3 O10 [228]

Cs2 Cr4 O3 [228]

CrO3 [228]

Assignment

978 961 945

1001 975

945

980 945 930

vas (CrO2 ) vs (CrO2 ) vas (CrO4 )/vas (CrO3 )

902 770

903 818 761

560

562 518 378

904 840 831 719 560 512 485 378 343 316

994 984 884

847

955 947 933 908

357 368 348

242 369 360

370

358

217

262 230 212

211 a Only

247 224

vs (CrO4 )/vs (CrO3 ) vas (CrOCr)

563 497 404

vs (CrOCr) δ(CrO2 ) δ(CrO4 )/δ(CrO3 )

375 338

ρ(CrO2 OCrO) δ(OCrO)

208

ρ(CrO2 ⊥OCrO)

chromium–oxygen vibrations are listed.

valentinite, which exists between 840 K and the melting point [232]. Cubic Sb2 O3 is built of discrete Sb4 O6 molecules and crystallizes in a face-centered cubic lattice of the space group F3 dm (Oh7 ). It is isomorphous with arsenite As4 O6 [233]. Symmetry considerations show that the symmetry of the free molecule remains unchanged in the crystal site, so that the correlation of its internal modes with those in the crystal is straightforward. A center of symmetry is introduced in the crystal and in-phase and out-of-phase coupling of two molecules in the unit cell lead to an IR-active and a Raman-active component, derived from each normal mode of the discrete molecule [234]. Valentinite crystallizes in the orthorhombic space group 10 and is built up by infinite polymeric Sb−O−Sb chains D2h running along the c axis [235]. The Raman spectrum of orthorhombic Sb2 O3 in the internal mode region 400–800 cm−1 is dominated by the spectral features of the polymeric Sb−O chains, whereas the lattice mode region below 400 cm−1 is controlled by the crystal space group and unit cell coupling [232]. The Raman spectrum of an Sb2 O3 sample that largely consists of the cubic senarmontite with only small contributions of valentinite is shown in Fig. 7a. The major bands at 82 (B2 ), 118 (E), 189 (B2 ), 254 (A1 ), 355 (E), 373 (B2 ), 450 (A1 ) and 712 (B2 ) cm−1 are characteristic of the cubic phase and have been assigned by Beattie et al. [234] on the basis of a single-crystal investigation. The weak

bands observed in Fig. 7a at 141, 189, 219, 295, 500 and 596 cm−1 can be attributed to the minority valentinite phase in the sample [232]. A comparison of the Raman spectrum shown in Fig. 7a with the IR data reported by Cody et al. [232] indicates that the bands at 684, 712 and 904 cm−1 of the cubic phase and the band at 596 cm−1 of the orthorhombic phase should be IR-active and Ramanforbidden modes, respectively. These may have become Raman active in the disperse, poorly crystalline powder sample due to a weakening of the spectroscopic exclusion rule [236]. The oxidation of Sb2 O3 to Sb2 O4 and phase transitions in both oxides were studied by Raman and IR spectroscopy, although no band assignments were given by Cody et al. [232]. The low-temperature α-Sb2 O4 modification crystallizes in the space group C29 /Pna21 [237]. It consists of corrugated sheets of Sb5+ −O octahedra which are linked by sharing corners. Adjacent sheets are interconnected by Sb3+ ions. The high-temperature β-Sb2 O4 modification crystallizes in the monoclinic structure (space group C2/c) [238]. The major difference between the two phases is the coordination of the Sb3+ ions. Figure 7b shows a Raman spectrum of the orthorhombic α modification, which exhibits bands at 72, 91, 117, 142, 199, 221 (sh), 239, 255 (sh), 262, 283 and 353 cm−1 in the regime of external modes and at 400, 413 (sh), 463, 614, 652, 714, 758 and 825 cm−1 in the regime of

254

100

712

900

Raman (a)

100

758 825

463

413

614 652 714

400

262 221 255

117 91

500

×10

283 353

72 142

904

239

596

450

219 295 355

82 118 141

373

189

500

684

Intensity / a.u.

×10

Fig. 7

947

199

3.1.3 Structure and Morphology

500

900

shift/cm−1 (b)

Raman spectra of (a) Sb2 O3 (senarmontite) and (b) Sb2 O4 (adapted from Ref. [239]).

internal modes. The two bands at 614 and 652 cm−1 should be IR active and Raman forbidden [232]. Their appearance in the Raman spectrum shown in Fig. 7b is probably explained by the poor crystallinity of the powder sample [236]. The temperature dependence of the Raman spectra and the oxygen-exchange properties of Sb2 O3 and Sb2 O4 have been reported by Mestl et al. [239]. B Zeolites and Other Crystalline Porous Materials Vibrational spectroscopies can provide important information on the structure and dynamics of these crystalline solids. The vibrational spectra are characteristic of the nature of the zeolite lattice framework and effects of the SiO2 : Al2 O3 ratios and of exchangeable cations on the framework structure can be investigated. Moreover, vibrational spectra are sensitive to isomorphous substitutions of framework atoms. Several review articles dealing with vibrational properties and lattice dynamics of zeolitic materials have been published [143, 240, 241]. a Synthesis of Zeolites A detailed understanding at a molecular level of the steps occurring in solution during zeolite synthesis are of utmost importance for optimal control over the synthesis process. Raman spectroscopy offers the advantage that it permits vibrational spectra to be recorded from solution species and also from amorphous and crystalline phases that are present during zeolite synthesis. Hence several workers have used Raman spectroscopy in attempts to elucidate the

synthesis mechanisms of various zeolites [242–248]. The aluminosilicate gel used for the synthesis of zeolite A [242, 245], even though amorphous, was shown to have a structure consisting of predominantly four-membered rings connected in a random fashion. It is considerably depolymerized, consisting of Si atoms with one and two non-bonded oxygen atoms. For the transformation of this gel to zeolite A to proceed, it is essential to have −1 Al(OH)− 4 ions (band at 621 cm ) in solution. During the nucleation period, the gel reorganizes its structure by interaction with these Al(OH)− 4 anions and forms nuclei of zeolite A. During the nucleation period of zeolite Y, the solid amorphous phase consists of predominantly sixmembered aluminosilicate rings, which act as building blocks for the formation of zeolite Y [246]. Raman spectroscopy also showed that the presence of polymeric, highly condensed silicate units is essential if zeolite Y crystallization is to take place. The synthesis of zeolite ZSM-5 was examined by monitoring the Raman spectra of the tetrapropylammonium template ion [248]. b Framework Vibrations in Zeolites Because of the enormous importance of zeolite materials as adsorbents and catalysts, their IR spectra and, less frequently, their Raman spectra have been used extensively to characterize these samples. The early IR studies of zeolites have been reviewed [240, 241]. References see page 965

3.1 Physical Properties

314 320

290

and compared with frequencies of other pure SiO2 polymorphs. The IR spectrum of Al-free faujasite is shown in Fig. 8 and the corresponding Raman spectrum is reproduced in Fig. 9. Although both spectra clearly contain bands in the symmetric (700–850 cm−1 ) and antisymmetric (1000–1200 cm−1 ) stretching regions, the total number of detectable bands still remains below the predicted number of IR-active (12) and Ramanactive (24) bands [143]. The strong Raman bands observed near 500 cm−1 in Fig. 9 belong to mixed stretching and bending modes (see below). The effect of the incorporation of Al3+ cations in the framework of faujasite at different Si : Al ratios is shown in Fig. 10. A comparison of the IR spectra of zeolites X and Y having Si : Al ratios of 1.12 and 2.98 with the spectrum of Al-free faujasite in Fig. 8 shows that the presence of aluminum does not give rise to the formation of additional bands. Because of the random distribution of Al3+ cations within the framework, it has a line-broadening effect. Also, the band positions are shifted to lower frequencies as the Si : Al ratio decreases, this effect being most pronounced

200

613

529

460

400

484

437

682

350

Transmittance

400

600

683

1178 1209

837

484 530

613

791

1636

Wavenumbers/cm−1

1061

461

The zeolite framework is composed of the network formed by TO4 tetrahedra sharing corners. As opposed to pure crystalline silica forms (such as silicalite or Alfree faujasite), a certain percentage of the TO4 tetrahedra contain network-forming cations which isomorphously substitute the Si atoms. For example, Al-substituted silicalites are known as zeolite ZSM-5 and the incorporation of AIO4 tetrahedra into the faujasite structure yields zeolite Y. As a consequence, certain similarities in the vibrational structures of zeolites and their pure SiO2 parent structures are expected. The vibrational spectrum of a zeolite can be described as the sum of several contributions, the first of which is given by the framework of the zeolite. The second contribution originates from the charge-balancing cations, which are located in channels and cages formed by the framework and which are required to compensate the negative charge resulting from the substitution of lower charged cations (e.g. Al3+ ) for the framework silicon. The third contribution to the vibrational spectra of zeolites has the same principal origin and results when protons are used for charge compensation. Hydroxy groups are formed in this case and their vibrational behavior is discussed in Section 3.1.3.8.5. The charge-compensating cations vibrate against the framework with frequencies typically below 400 cm−1 and are called translational modes in the notation of solid-state vibrational spectroscopy. Translational (and rotational) modes are classified as external vibrations, as opposed to internal vibrations, which are characteristic of covalently bonded structures. In contrast to this convention, the terms internal and external have been used in zeolite spectroscopy to describe vibrations within tetrahedral building units and between them, respectively. Consequently, van Santen and Vogel [143] suggested the terms intratetrahedral and intertetrahedral for the internal framework vibrations and denote the translational modes of charge-compensating cations as external vibrations in the sense of solid-state spectroscopy. As already discussed for the pure SiO2 polymorphs (Section 3.1.3.8.3Ab), the vibrational spectra of zeolites are also dominated by three frequency regimes, namely 1000–1200, 700–850 and 250–500 cm−1 . Based on the concept of localized and delocalized vibrational modes, van Santen and coworkers [143, 144] attributed these frequency regimes to localized and delocalized antisymmetric stretching vibrations, localized and delocalized symmetric stretching vibrations and bending vibrations, respectively. The intratetrahedral modes are evidently of the localized type, whereas intertetrahedral modes are delocalized. In addition, torsional and mixed modes are located at frequencies below 200 cm−1 . The characteristic IR and Raman frequencies of silicalite in the three major regimes are summarized in Table 1

Transmittance

948

1080

500

1000

1500

Wavenumbers/cm−1 Room temperature Fourier transform IR spectrum of aluminum-free faujasite (adapted from Ref. [143]).

Fig. 8

3.1.3 Structure and Morphology

949

on T−O−T angles in the framework structure. Dutta et al. [253] noted the sensitivity of the position of this band on the T−O−T angle. The two sharp bands observed in the Raman spectrum of Al-free siliceous faujasite at 488 and 510 cm−1 (Fig. 9) correspond to T−O−T angles of 141◦ and 147◦ , respectively [249]. As shown in Fig. 11, these bands broaden significantly when the Si : Al ratio is decreased, their relative intensities change and the position of the more intense component shifts to slightly higher frequencies. The incorporation of heteroatoms in the framework of silicon-rich zeolites with the MFI structure having a 1000

510

500

488

∆nR/cm−1 Raman spectrum of aluminum-free faujasite (adapted from Ref. [249]).

Fig. 9

E 507

795

A 460

1145

975

461

732

F

D

1021

B

800

400

C

514

1200

Wavenumbers/cm−1

for the antisymmetric stretching modes [250, 251]. The frequency of these vibrational modes depends primarily on the polarity and length of the Al−O−Si bonds, since the mass of Al is nearly identical with that of Si. In fact, the average strength of the Al−O bond is lower than that of the Si−O bond, as indicated by the longer average Al−O bond distance of 0.173 nm compared with the Si−O bond length of 0.162 nm. Similar frequency shifts were observed by Dutta and coworkers [249, 252] in the Raman spectra of zeolites A and Y. As already mentioned, the strongest Raman bands observed near 500 cm−1 belong to mixed stretching and bending modes and should therefore depend strongly

516

B

467

Fig. 10 Room temperature IR spectra of faujasite zeolites with different Si : Al ratios: (A) zeolite Y(Si : Al = 2.98) and (B) zeolite X(Si : Al = 1.12) (adapted from Ref. [250]).

A 400

500

600

∆nR/cm−1 Fig. 11 Raman spectra of faujasite zeolites with varying Si : Al ratios in the region between 400 and 600 cm−1 . The Si : Al ratios are (A) 1.0; (B) 1.3; (C) 2.6; (D) 3.3; (E) 4.5; and (F) ∞ (adapted from Ref. [249]). References see page 965

950

3.1 Physical Properties

small number of heteroatoms can be described as a solid solution of, for example, [AlO4 ] in [SiO4 ] in the case of Al-substituted silicalite (ZSM-5). This implies that the general features of the vibrational spectra of silicalites (see Table 1) should not be altered substantially by the substitution of Si for Al, as has been observed [143, 254, 255]. This should also be the case when other elements such as Ga, Ge, Be, Fe or Ti are isomorphously substituted for Si in silicalite, although frequency shifts can occur due to the changing T−O bond polarities and length and because of the different masses of the heteroatoms. For example, the symmetric stretching mode was observed at 788 cm−1 for silicon sodalite with a band half-width of 30 cm−1 ; the symmetric modes in aluminum sodalite were found at 741, 718 and 672 cm−1 with approximately the same half-widths, at 700, 635 and 590 cm−1 in gallium aluminum sodalite and at 630, 585 and 555 cm−1 in gallium sodalite. The band half-width in the latter two materials is significantly larger [256]. A band at 960 cm−1 was observed in defective silicalite samples which could be associated with an O3 Si−OH stretching vibration in fully hydroxylated nanocavities (hydroxylated nests) [254, 257, 258]. Titanium silicalite has been investigated by IR and Raman spectroscopy and again only minor changes were observed for the framework vibrations, although the antisymmetric and symmetric stretching vibrations turned out to be sensitive to the presence of titanium [138, 255]. The most prominent extra feature observed in the IR spectra of Ti silicalite is represented by a band at ca. 960 cm−1 [138, 255, 259–261]. This band was interpreted as superposition of O3 Si−OH stretching modes of groups present in hydroxylated nests and of a titanium-sensitive mode by Scarano et al. [255]. Astorino et al. [138] associated this band with an antisymmetric stretching mode of a Si−O−Ti bridge. The strongest Raman band near 500 cm−1 was considered to be the most useful vibrational mode for the identification of MFI structures [138], since it is associated with a structure-sensitive symmetric stretching–bending mode of T−O−T bridges. Vibrational spectroscopic investigations of titanium silicate molecular sieves including infrared and visible and UV-excited Raman spectroscopies were recently reviewed by Ratnasamy et al. [262]. The adsorption of water and ammonia has shown that the Ti4+ cations in Ti silicalite function as Lewis acid centers [138, 255]. During the adsorption, the Ti4+ cation moves from the tetrahedral position toward a more ‘‘external’’ relaxed location, where the SiO−Ti bonds become more polar and elongated relative to the usual situation. The coordination of NH3 gives rise to an expansion of the overall coordination of the Ti4+ cation to five- or six-fold [138, 255]. In the presence of H2 O,

the SiO−Ti bonds may reversibly open with formation of SiOH and TiOH groups [255]. IR and Raman data for B- and Fe-substituted silicalites have also been reported [255]. As a general conclusion, frequency shifts of framework vibrations among the various zeolites appear to be dominated by variations in T−O−T angles and there appears to be little reason to assign particular frequencies to localized ring vibrations of the lattice [143]. c Effect of Non-Framework Cations Charge-compensating, exchangeable cations within the channel and cage systems of zeolite structures give rise to the external (translational) vibrations which are located in the farinfrared region [263], but they also have an influence on the position of framework bands [143]. The antisymmetric stretching mode at 1000 cm−1 in zeolite A shifted to 1080 cm−1 on replacement of Na+ by H+ [264]. The band also shifted slightly on substitution of K+ , Li+ and divalent cations for Na+ , while much smaller shifts were reported for the symmetric stretching bands at 675 and 550 cm−1 and for the bending modes at 465 cm−1 . Raman spectra of cation-exchanged zeolite A suggested that Li+ distorts the framework, whereas Na+ , K+ and Tl+ do not alter the zeolite structure and the cations occupy the same sites [265]. When Na+ was partially exchanged for H+ in X and Y zeolites, the symmetric stretching bands (700–800 cm−1 ) shifted to higher frequencies by 10–30 cm−1 and bands between 450 and 700 cm−1 shifted to lower frequencies by 5–20 cm−1 [266]. Datka et al. [267] reported a linear increase in the frequencies of symmetric and antisymmetric stretching and of Si−O−Si bending vibrations with Sanderson’s average compound electronegativity in faujasite- and ZSM-5-type zeolites, when the Si : Al ratio was increased and when alkaline earth metal cations and protons were substituted for Na+ and K+ . Van Santen and Vogel [143] gave an interpretation of the effect of non-framework cations on the framework vibrations. They assumed that the T−O bond is weakened when a cation is coordinated to a given oxygen atom because the electron density of T−O bonds of this oxygen atom will be affected. The hybridization of the oxygen atom, which is largely sp in character for Si−O−Si bond angles close to 180◦ , will acquire a higher percentage of sp2 character, causing the bond angle to decrease. This will cause the coupling between the Si−O oscillators to decrease and the frequencies of the antisymmetric stretchings to decrease and those of symmetric stretchings to increase. The latter effect is not in accord with the experimental results. However, the larger cations of lower electronegativity are coordinated to more than one anion. When they are replaced by a

3.1.3 Structure and Morphology

d Aluminum Phosphate Molecular Sieves Vibrational investigations of phosphate-based molecular sieves are still relatively scarce. The vibrational spectra of αAIPO4 (berlinite), the aluminophosphate analogue of α-quartz, have been interpreted [273, 274]. The Raman spectrum [275] and the mid-IR spectrum [276] of VPI5 have been reported. The mid-IR-spectra of VPI-5 and AlPO4 -H2 are compared in Fig. 12. These two materials are constructed from the same building block and, therefore, they show similar framework spectra in the mid-IR region. The IR spectrum of SAPO-37, which has the faujasite structure, also shows a similar signature [277]. Although additional bands are present due to the different composition of the material, the frequency regions known for zeolites (see above) can be recognized, which are associated with antisymmetric stretching vibrations (950–1300 cm−1 ) and symmetric stretching vibrations (650–800 cm−1 ) of Al−O−P bonds and with lattice deformation modes (350–500 cm−1 ). Raman spectroscopy has proved extremely useful for investigations into the vibrational structure of materials of the AlPO family [278, 279]. Figure 13 shows the Raman spectra of AlPO4 -5, AlPO4 -8 and VPI-5 and Table 8 summarizes the observed Raman lines of these molecular sieves and of AlPO4 -11. A factor group analysis for these materials in comparison with that for α-quartz and berlinite was also reported [279]. Topology-dependent shifts of the strong bands at 1100–1200 cm−1 and obvious differences of the Raman signatures in the 350–650 cm−1 region of the spectra shown in Fig. 13 can be recognized.

B

Transmittance

proton being attached to only one framework oxygen atom, several oxygen atoms are set free and will show the reverse effect, this eventually resulting in an overall decrease in frequencies for the symmetric and an increase for the antisymmetric stretching vibrations. The external translational modes of exchangeable cations are typically located in the frequency range below 250 cm−1 . The application of far-IR spectroscopy in the investigation of the translational vibrations of extra-framework cations in zeolites has been pioneered by Ozin and coworkers [263, 268, 269]. For faujasite zeolites they demonstrated [263, 268] that the combined use of the frequencies and intensities of site-specific metal cation far-IR modes permits secure metal cation vibrational assignments for sites I, I , II, III and III , distinction of different cations and the establishment of cation occupancies and locations in the zeolite framework. The solid-state ion exchange can be monitored by far-IR spectroscopy and the resulting cation distributions can be determined [270]. Raman spectroscopy has also been applied successfully to studies of the translational modes of extra-framework cations in faujasitic zeolites [271, 272].

951

A

1500

1080

660

Wavenumbers/cm−1 Fig. 12 Infrared spectra of the structural vibrational region for (A) AlPO4 -H2 and (B) VPI-5 (adapted from Ref. [276]).

(c)

400

1200

(b)

400

1200

(a)

400

1200

∆nR/cm−1 Raman spectra of (a) calcined AlPO4 -5, (b) AIPO4 -8 and (c) VPI-5 (adapted from Ref. [279]).

Fig. 13

Although the detailed band assignments given by Holmes et al. [279] deviate from those mentioned above and discussed for zeolites in Section 3.1.3.8.3Bb, the data nevertheless indicate that Raman spectroscopy may well References see page 965

952

3.1 Physical Properties

Raman frequencies (v/cm−1 ) of aluminophosphate molecular sieves (strongest bands underlined) (adapted from Ref. [279])

Tab. 8

AIPO4 -11

AIPO4 -5

AIPO4 -8

VPI-5

1240 1140 1121 1066 1048

1236 1141 1125

1227 1137 1115

1217 1167 1113 1081

567 494 461 410 404 378 262

566 499 462

1050 640 610 483 438

126

398 361 267 247 221 201 143 122

377 300 281 227 207 177 119

85

61

60

213 178

605 492 485 433 375 305 255 217 204 184 115 105 60

be further developed as a diagnostic tool for probing framework structures of molecular sieves. It is well known that aluminosilicates [280, 281] and molecular sieves [279] often exhibit a strong background fluorescence. Demuth et al. [282], in Raman studies of SAPO-5, demonstrated that the luminescence excited at 532 nm could effectively be quenched when pseudooctahedrally coordinated Cr3+ species were deliberately produced in the one-dimensional channel system. The photoluminescence quenching mechanism was assumed to involve an energy transfer from the molecular sieve photoexcited states to the metastable 2E state of Cr3+ , giving rise to the well-known 2E → 4 A2 luminescence at 696 nm of Cr3+ . C Multicomponent Oxides Multicomponent oxides are widely used in selective oxidation (see Chapters 14.11.1 and 14.11.8) and ammoxidation (see Chapter 14.11.8) reactions and vibrational spectroscopies including IR and Raman spectroscopy have been used for their characterization [26]. Raman spectroscopy has been shown to be an extremely powerful tool for in situ investigations into working oxidation catalysts [283]. This section concentrates on two catalyst systems as examples, namely molybdates, with special emphasis on bismuth molybdates and V−P−O catalysts. a Bismuth Molybdates A variety of bismuth molybdates, namely α-Bi2 Mo3 O12 , β-Bi2 Mo2 O9 and γ -Bi2 MoO6 , are

known and they have been shown to be present under catalytic conditions. Raman spectroscopy enables the syntheses of these phases to be followed, including the precipitation in aqueous solution as a function of temperature, pH and rate of mixing, the aging of the precipitate and its calcination [284]. The series of spectra shown in Fig. 14 represent the evolution of Bi−Mo−O species at various time intervals as the precipitation proceeds and the formation of the solid oxides during drying and calcination for α-Bi2 Mo3 O12 and γ -Bi2 MoO6 . The formation routes of the different mixed oxides could be associated with the state of polymerization of Mo−O and Bi−O species, this being strongly determined by the solution pH. Many of the reported IR and Raman studies concern the identification of the different bismuth molybdate phases present in the catalytic material. This is not an easy task since the catalysts are typically multiphasic and the lattices of individual phases are highly distorted so that a clear distinction between octahedral and tetrahedral coordinations is often difficult. A thorough vibrational analysis of α-Bi2 Mo3 O12 , β-Bi2 Mo2 O9 and γ -Bi2 MoO6 , was reported by Matsuura et al. [285] and the frequencies observed by these and other authors have been compiled by Mehicic and Grasselli [26]. Figure 15 shows the IR and Raman spectra of the three phases mentioned above. The Raman frequencies observed for bismuth molybdates with widely varying Bi : Mo atomic ratios (0.67 ≤ Bi : Mo ≤ 14) were tabulated and show that many of these compositions were multiphasic [286]. An assignment of each of the reported Mo−O bond lengths in each of the various identified phases to its corresponding observed Mo−O Raman stretching frequency was possible based on an empirical stretching frequency–bond length relationship [28] that was established for several metal–oxide systems including molybdates [287] and bismuth oxides [288]. Several Raman studies that were devoted to the elucidation of the nature of sites responsible for the activity and/or selectivity of bismuth molybdate catalysts have used 18 O isotope labeling. The active participation of lattice oxygen in the partial oxidation of propene and its incorporation into the reaction products could be demonstrated by the 18 O labeling technique [289, 290]. Using Raman spectroscopy, Ono and Ogata [291] determined the location of 18 O tracers incorporated in the lattices of γ -Bi2 MoO6 and α-Bi2 Mo3 O12 after exchange. With the γ -phase, all types of lattice oxygens could be exchanged with gas-phase 18 O2 , whereas with the α-phase, the lattice oxygens that exchanged were those Mo tetrahedra with adjacent Bi ions. Oxygens situated next to Bi ion vacancies did exchange initially. In the catalytic oxidation of propene on the α-phase, oxygen incorporation also occurred at the terminal oxygen vacancies.

3.1.3 Structure and Morphology

953

Calcined Calcined Dried 30 h Dried

12 h

30 h

8h

8h 5h 5h 4h 4h 3h

3.5 h

2h

3h

1h 1h 10 min 200

600

1000

200

Raman (a)

Fig. 14

600

1000

shift/cm−1 (b)

Raman spectra recorded during the formation of (a) γ -Bi2 MoO6 and (b) Bi2 MoO12 (adapted from Ref. [284]).

1000 (a)

600

200

0

1000 (b)

600

200

Raman shift/cm−1

0

1000

600

200

0

(c)

Infrared (top) and Raman (bottom) spectra of (a) Bi2 MoO6 at high temperature, (b) Bi2 Mo2 O9 and (c) Bi2 Mo3 O12 (adapted from Ref. [285]).

Fig. 15

954

3.1 Physical Properties

Grasselli and coworkers [292, 293] monitored the oxygen loss and the incorporation of 18 O during redox cycles applied to β-Bi2 Mo2 O9 by in situ Raman spectroscopy. From the spectra shown in Fig. 16, it was concluded that β-Bi2 Mo2 O9 disproportionated into α-Bi2 Mo3 O12 and γ -Bi2 MoO6 , the latter phase being enriched on the surface. By identifying those positions in the structures at which 18 O was incorporated, it was proposed that oxygens bridging Bi and Mo atoms were responsible for hydrogen abstraction, whereas those connected exclusively to Mo atoms insert into the allylic intermediate. Based on these investigations, the possible structure of the multifunctional site shown in Fig. 17 was proposed. The observed disproportionation of γ -Bi2 MoO9 is consistent with the fact that bismuth molybdate catalysts are typically multiphasic; the various phases present may act synergistically. A synergistic cooperation of the

A a

a

a +g

a

g a

a B

g

800

400

∆nR/cm−1 Raman spectra of (A) β-Bi2 Mo2 O9 after treatment in air at 873 K and (B) after reduction with propene and reoxidation with air at 753 K (α refers to Bi2 Mo3 O12 and γ to Bi2 MoO6 ) (adapted from Ref. [292]).

Fig. 16

O O

O Bi

Bi

O

O

O

O O

O Mo

O

O O

Fig. 17 Schematic representation of the structure of multifunctional active sites on the surface of γ -Bi2 MoO6 (adapted from Ref. [293]).

α- and γ -phases has been demonstrated by IR and Raman spectroscopy [294]. In situ Raman experiments have indicated [59] that the β-phase disproportionated under reducing conditions, as reported by Grasselli et al. [295], and in oxygen provided that the material had not been calcined above 773 K. If calcined at temperatures >773 K, however, the β-phase was stable and a mixture of α-, β- and γ -phases did not show any change in phase composition that was detectable by Raman spectroscopy over 24 h, under conditions of propene oxidation. b Synergy Effects in Selective Oxidation Synergy effects in the selective oxidation of C4 hydrocarbons to maleic anhydride have been reported for mixtures of molybdates such as MnMoO4 and CdMoO4 with MoO3 [180]. Using in situ Raman spectroscopy in combination with 18 O isotopic labeling, Ozkan et al. [180] demonstrated that pure MoO3 was strongly reduced by butadiene at 723 K, whereas pure MnMoO4 hardly exhibited any detectable spectral changes. In the mixture with MnMoO4 , MoO3 was much less reduced. Reoxidation of the pure phases with 18 O2 , bridging oxygen atoms in Mo16 O3 (bands at 665 and 815 cm−1 ) were quantitatively substituted by 18 O, whereas oxygens in terminal positions (band at 992 cm−1 in Mo16 O3 ) were only partially exchanged. No isotopic shifts were observed for pure MnMoO4 under the same conditions and an analogous behavior of the two components was found in their mixtures. The different exchange rates of bridging and terminal oxygen atoms in MoO3 was interpreted by a higher rate of insertion of bridging oxygens into the allylic intermediate as compared with terminal oxygen atoms. The major difference exhibited by the mixed oxides was the fact that 16 O was not quantitatively replaced by 18 O in bridging positions and that it yielded less CO2 . It was therefore suggested that bridging oxygen atoms might be responsible for complete oxidation, whereas Mo=O groups provide the oxygen for selective oxidation (this conclusion being in contrast to the mechanism proposed

3.1.3 Structure and Morphology

by Grasselli and coworkers [292, 293]) and that the MnMoO4 phase assists in replenishing oxygen vacancies in MoO3 by spill-over. In physical mixtures of antimony oxides with MoO3 , spill-over oxygen, if present, remained below the limits of detectability by Raman spectroscopy [296]. c V−P−O Catalysts V−P−O catalysts are active in the selective oxidation of n-butane. In situ Raman studies demonstrated [55, 297–299] that the basic (VO)2 P2 O7 phase was stable under working conditions, whereas the δ-VOPO4 phase transformed into α-VOPO4 [55]. The precursor VO(HPO4 ) · 0.5H2 O became totally disordered during catalyst activation in an n-butane–air atmosphere with the simultaneous onset of the generation of maleic anhydride [298]. It was concluded that the creation of the active surface may not depend solely on the nature of the bulk catalyst phase. The active sites were assumed to form specifically under reaction conditions and that their creation requires the presence of the reacting molecules and products that may provide a template to enable surface structures of the catalyst to ‘‘crystallize’’. In situ Raman spectroscopy also showed [299] that Fe and Co promoters had a significant influence on the transformations of the VO(HPO4 ) · 0.5H2 O precursor into the VOPO4 /(VO)2 P2 O7 catalyst. The presence of the promoters influenced the phase composition and the dispersion of VOPO4 phases on the (VO)2 P2 O7 matrix during the activation process. The role of lattice oxygen species in the catalytic oxidation of n-butane to maleic anhydride was investigated by laser Raman and Fourier transform IR spectroscopy using β-VOPO4 labeled with 18 O [300, 301]. Preferential incorporation of 18 O at P−O−V sites was observed and specific active oxygen centers were identified. The investigations mentioned in Sections 3.1.3.8.3Ca and b clearly show the potential of laser Raman spectroscopy as a technique that permits monitoring of phase compositions and transformations of bulk oxide catalysts under reaction conditions. 3.1.3.8.4 Vibrational Spectra of Supported Catalysts Oxides of transition metals such as Mo [302–316], W [28, 307, 316a,b], Cr [28, 308], V [307, 309] and Re [310] supported on a second high-surface-area metal oxide such as alumina, titania, silica or zirconia comprise a class of technologically highly important catalysts which find application in a wide variety of large-scale processes. They are frequently described as monolayer catalysts, indicating that the supported transition metal oxide active phase is highly dispersed over the surface of the support. Structural characterization of such materials is difficult because of their very nature and most structure-sensitive techniques,

955

except perhaps X-ray absorption spectroscopy, have failed to provide detailed structural information for the active phase at a molecular level. Therefore, vibrational spectroscopies play an important role in the elucidation of the structures of supported transition metal oxide species or overlayers, since vibrational spectra reflect the nature of adsorbates having molecular character as well as separate solid phases or compounds. Laser Raman spectroscopy (LRS) is particularly powerful for investigations into the structure of supported oxide catalysts. All characteristic vibrational features of the mentioned transition metal oxides fall into the frequency range below 1100 cm−1 (Section 3.1.3.8.3A) and thus into a regime that is entirely opaque in IR transmission spectroscopy (Section 3.1.3.8.2A) because of the strong absorption of the Me−O vibrations in the support oxides having highly ionic bond character. As a consequence, these support materials (particularly alumina and silica) have very low Raman scattering cross-sections, so that the normal modes of the minority components, namely the transition metal oxides having a relatively high covalent bond character resulting in high Raman scattering cross-sections, can most frequently be detected by LRS with relative ease in the frequency region from 50 to 1100 cm−1 . A major advantage of LRS for the study of these catalytic materials is its potential for in situ studies at elevated temperatures under controlled atmospheres. Since the early work of Villa et al. [311] in 1974, LRS has found steadily increasing application for the characterization of supported transition metal oxide catalysts, including their generation. The results have been reviewed in several excellent articles [24–28, 306, 312, 313]. Therefore, this section only presents some LRS results that were obtained for molybdate-based catalysts, as an example. This system has been chosen as being the best characterized and because of its importance in relation to hydrotreating catalysis. A Supported Molybdate Catalysts Supported molybdate catalysts are typically prepared by impregnation of a suitable support (mostly alumina) from an aqueous solution containing the appropriate concentration of the (NH4 )6 Mo7 O24 precursor. This solution may contain the tetrahedral monomer MoO2− 4 or the heptaanion 6+ is octahedrally coordinated , in which Mo Mo7 O6− 24 by oxygen atoms, depending on the molybdenum concentration and pH. The equilibrium between the two species is 6− + −−− 7MoO2− 4 + 8H −−− Mo7 O24 + 4H2 O

References see page 965

(13)

956

3.1 Physical Properties

Solution spectra of these two anionic species are shown in Fig. 18. The monomeric species has Td symmetry and is therefore expected to give rise to four normal modes: vib = v1 (A1 ) + v2 (E) + v3 (F2 ) + v4 (F2 )

(14)

Assuming that the hydrated species in solution still maintains Td symmetry, these bands are indeed seen in the spectrum in Fig. 18, namely the totally symmetric stretching mode v1 (A1 ) at 896 cm−1 , the weaker antisymmetric stretching mode v3 (F2 ) at 841 cm−1 and the degenerate symmetric [v2 (E)] and antisymmetric [v4 (F2 )] bending modes at 318 cm−1 . The heptaanion Mo7 O6− 24 is built up of edge-sharing, distorted MoO6 octahedra and gives rise to the characteristic solution spectrum shown in Fig. 18. These spectra may be taken as fingerprint spectra for comparison with spectra that are obtained for surface species. The bands of the heptaanion at 840 and 790, 560 and 220 cm−1 , which are associated with antisymmetric and symmetric stretching modes and with deformation modes of Mo−O−Mo bridging groups, are of particular diagnostic importance as their detection is clear evidence for the presence of condensed, polymeric species. Monomeric and condensed anionic molybdate species may adsorb from the impregnation solution on to the surface of the support, provided that the isoelectric point (IEP) of the solid allows a positive surface polarization to develop at the pH conditions of the solution (see

846

318

938

896

359

218

895

A

Chapter 2.4.2). In this case, the adsorbed molybdate anions may be stabilized on the surface by formation of Al−O−Mo bonds during the drying and calcination procedures. However, if the surface polarization is negative, then anion adsorption can hardly take place and the molybdate will precipitate on the support surface during drying and it will decompose during calcination to form microcrystals of MoO3 . Formation of MoO3 will also occur when the total molybdate loading exceeds the monolayer capacity of the support material used. Finally, if the molybdate loading is high and the calcination temperature exceeds approximately 870 K, a solid-state reaction between Al2 O3 and MoO3 with formation of Al2 (MoO4 )3 may occur. Figure 19 shows two series of Raman spectra of molybdena-on-alumina in the impregnated, dried and calcined states of samples that were impregnated at pH 6 and contained 3 and 8 wt.% MoO3 . After impregnation, for the sample with low loading (spectrum A in Fig. 19b), two broad bands at 918 and 338 cm−1 are detected. These bands, although shifted by about 20 cm−1 to higher frequencies, fall in the region of the v1 (A1 ) mode and the degenerate v2 (E) and v4 (F2 ) modes of tetrahedral MoO2− species. In the dried state, these 4 two bands shift to still higher frequencies and become broader. It has been suggested [64] that the spectra were to be attributed to distorted tetrahedral surface species, although the impregnation solution should have contained the heptaanion at pH 6. Hence, decondensation of the polyanion to form the tetrahedral monomer must have occurred. Undoubtedly, this surface monomer must have a symmetry lower than Td . Stencel et al. [314] suggested the surface structures in Scheme 1, dependent on the state of hydration. The spectra in Fig. 19 were recorded under ambient conditions so that the samples should be hydrated. The corresponding structure should thus belong to point group C2v or perhaps C1 . Therefore, the degenerate normal modes characteristic for the Td symmetry must be split and the major bands are expected to shift to higher frequencies because of the rehybridization in the anchored molybdate species. Whereas the frequency shifts are obvious, the splittings are not clearly detected, presumably because of the large width of the bands which do, however, show some structure in spectrum B in Fig 19b. When this 3 wt.% sample is calcined at 773 K,

B

O 400

800

0

Mo O Al Al

O

Raman shift / cm−1 Raman spectra of (A) solution.

Fig. 18

MoO2− 4 and (B)

Mo7 O6− 24 in aqueous

O

Scheme 1

OH H +H2O −H2O

O

H OH O

O

Mo O

Al

Al

950

328

930 899

943

C

B

344

338

918 894

221

921

A

334

344

B

A

219

326

328

939

Intensity / a. u.

C

957

938 900 865

3.1.3 Structure and Morphology

800

400

0

800

400

0

Raman shift/cm−1 (a)

(b)

Raman spectra of alumina-supported molybdate catalysts: (a) 8 wt.% MoO3 on γ -Al2 O3 after impregnation at pH 6 (A) in the wet state, (B) after drying at 393 K and (C) after calcination at 773 K; (b) 3 wt.% MoO3 on γ -Al2 O3 after impregnation at pH 6 (A) in the wet state, (B) after drying at 393 K and (C) after calcination at 773 K (adapted from Ref. [64]).

Fig. 19

the center of the strongest band is shifted to 938 cm−1 and exhibits broad tailing towards lower frequencies. The band in the bending region also becomes asymmetric towards lower frequency, most probably indicating an additional band near 220 cm−1 . This suggests that at least partial condensation of the surface molybdate species with formation of Mo−O−Mo bridges has occurred. In the calcined state, therefore, monomeric and condensed surface molybdate species are most likely present in 3 wt.% MoO3 /γ -Al2 O3 . The sharp lines at 820 and 998 cm−1 are indicative of the formation of a small amount of MoO3 . In the preparation of the more highly loaded 8 wt.% MoO3 /γ -Al2 O3 .catalyst, the impregnation solution is more concentrated and, therefore, the polyanion concentration is enhanced. Spectrum A in Fig. 19a of the impregnated sample clearly shows the deformation band of Mo−O−Mo bridges at 219 cm−1 . This and the position of the strongest band at 939 cm−1 are consistent with their association with a polymeric species. Polymeric species are also detected in the dried state and after calcination. In this latter state again some MoO3 was formed, as indicated by the sharp lines in spectrum C in Fig. 19a. Stencel et al. [314] reported on frequency shifts of the terminal stretching mode as a function of MoO3 loading on alumina between 3 and 25 wt.%. When the Raman spectra were recorded under ambient conditions, i.e. with the samples in a hydrated state, an increase in frequency from 938 to 970 cm−1 was observed with increasing

loading. This behavior was associated with an increasing degree of condensation of the surface molybdate species. The position of the terminal stretching mode was also found to be sensitive to the state of hydration [314–318]. The reversible effect of water on the band position is shown in Fig. 20, which clearly shows that the Mo=O stretching band at 950 cm−1 , which is characteristic of the hydrated state, shifts to 1006 cm−1 after dehydration of the 15 wt.% MoO3 /Al2 O3 sample. This influence of water was interpreted as being due to a weakening of the Mo=O bond strength by the coordination of water molecules. Finally, when catalyst materials containing high MoO3 loadings (beyond the monolayer capacity of the alumina support) were calcined at temperatures higher than 770 K, microcrystalline MoO3 and Al2 (MoO4 )3 were detected by LRS [319, 320]. As shown in Fig. 21, Al2 (MoO4 )3 was observed, as indicated by bands at 390 (not shown), 830 and 1006 cm−1 after calcining an MoO3 /Al2 O3 catalyst at 1170 K. As discussed in Chapter 2.4.7, MoO3 wets alumina because of the strong chemical interaction between these two compounds, which is responsible for the stabilization of the molecular species on the surface of the γ -Al2 O3 support. In contrast, silica has an IEP near pH 2, so that under impregnation conditions the surface polarization is negative. Moreover, the chemical interaction between MoO3 and SiO2 is extremely low because of the acidity References see page 965

3.1 Physical Properties

1030

895

830

1006

1006

958

950

C

700

900

1100

∆nR /cm−1 Raman spectrum of 14 wt.% MoO3 /Al2 O3 after calcination at 1173 K for 2 h (adapted from Ref. [320]).

Fig. 21

820

B

996

Mo/SiO2, 770 K

A

800

1000

∆nR/cm−1 Raman spectra of 15 wt.% MoO3 /Al2 O3 : (A) catalyst exposed to the atmosphere for 3 months; (B) catalyst calcined in 1 bar O2 at 823 K for 1 h, spectra recorded in situ in O2 atmosphere; (C) calcined catalyst after exposure to 14 mbar H2 O vapor at 298 K for 1 h (adapted from Ref. [314]). Fig. 20

1000

Fig. 22

of the silica surface and MoO3 does not wet SiO2 . As a consequence, monolayer formation is not possible on SiO2 supports and the dominant phase produced in the MoO3 /SiO2 system is polycrystalline MoO3 , as shown in Fig. 22. The examples given so far provide clear evidence for the potential of LRS for the molybdenum speciation as a function of preparation and treatment parameters and this conclusion can safely be extended to other supported oxide catalysts [28, 307–310]. B Quantification of Raman Spectra Comparison of relative Raman peak intensities of different species on a catalyst surface will lead to erroneous estimates of their relative abundance if their relative Raman scattering crosssections are not known. Unfortunately, the determination

800

600

400

∼ /cm−1 ∆n R

200

0

Raman spectrum of MoO3 /SiO2 .

of Raman scattering cross-sections is an extremely difficult task even for optically homogeneous samples and appears to be practically impossible for polydisperse materials for reasons that will be outlined below. Baltrus et al. [321] attempted to analyze Raman spectra of MoO3 /Al2 O3 and other supported materials quantitatively by using admixed KNO3 as an internal standard. The technique used for quantification was to mix KNO3 of known weight with different amounts of supported molybdena and tungsten catalysts. The integrated intensity of the symmetric stretching mode of the NO− 3 anion at 1050 cm−1 was then compared with the intensity of the Mo−O−Mo stretching mode at 820 cm−1 and the W−O−W stretching mode at 808 cm−1 . The accuracy of the method was claimed to be ±6% for MoO3 and WO3 at concentrations greater than 0.1%.

3.1.3 Structure and Morphology

The scattering cross-sections of supported species relative to those of the corresponding unsupported oxide were also determined. This was accomplished by using mixtures of, for example, MoO3 and catalysts with different loadings that contained the interaction species only, but no MoO3 . The cross-section of crystalline MoO3 relative to that of surface molybdena species was estimated to be 16 : 1. Likewise, a value of 137 : 1 was reported for the WO3 -supported tungsten system. These values may be considered as order-of-magnitude information and they presumably indicate that the relative amounts of microcrystalline oxide in supported catalysts may be easily overestimated. However, we believe that such values have to be considered with extreme care. Serious sources of uncertainty of the calibration methods mentioned above are the unavoidable inhomogeneities in mixtures of two solid powders typically having different particle sizes and morphologies. An enrichment of one component at the macroscopic surface of a pressed specimen may occur. Moreover, depending on particle size distributions, the laser focus may hit preferentially one or the other component at a given position, a problem that can be eliminated by sample rotation. Encapsulation, particle aggregation or disintegration phenomena may occur during the mechanical mixing of the components. Particle size and size distributions also influence the scattering properties of the sample, so that the scattering volume remains principally unknown and may vary for different samples depending on composition and specimen preparation conditions, such as mixing procedures and pressing. An additional weakness of the method occurs when optical multichannel analyzers are being used for the quantitative measurement of bands in different wavenumber regions having significantly different band shapes and widths. This is due to variations of the diode characteristics. This problem can be largely eliminated when the scanning multichannel technique (SMT) is applied (see Section 3.1.3.8.2D). Absorption caused by sample coloration also affects the band intensities and can be corrected via the Kubelka–Munk formalism as discussed above (see Section 3.1.3.8.2G). In conclusion, quantitative Raman spectroscopy of supported catalyst samples must be considered extremely critically and it may presumably be impossible within reasonable limits of accuracy. C Reduced Molybdate Catalysts Raman cross-sections of Mo suboxides are very small and the samples are typically strongly colored or even black, so that goodquality Raman spectra are difficult to record. Payen et al. [317] reported a Raman study on H2 reduction of alumina-supported molybdena catalysts. They proposed the formation of Mo5+ −OH and Mo6+ −O−Mo5+ and/or

959

Mo6+ −O−Mo4+ species. Reduction in D2 resulted in an isotope shift of a band at 840 cm−1 to 760 cm−1 , suggesting their assignment to Mo5+ −OH and Mo5+ −OD groups. A band at 760 cm−1 that was not sensitive to isotope labeling was associated with the Mo6+ −O−Mox+ group (x = 4, 5). It is remarkable that the Raman spectra did not indicate the formation MoO2 , even under the most severe reduction conditions. D Sulfided Mo/Al2 O3 Catalysts Sulfided Mo/Al2 O3 catalysts are black and consequently their Raman spectra usually suffer from a relatively poor signal-to-noise ratio. Nevertheless, the characteristic Raman bands of MoS2 at approximately 407 and 380 cm−1 have been detected on sulfided Mo/Al2 O3 catalysts [322–326]. In addition, bands in the range 500–550 cm−1 , which were associated with the existence of [S−S]2− groups presumably located at edges of surface MoS2 slabs, have been detected [323, 326]. The band positions were dependent on the slab size of the MoS2 [325] and on the temperature and composition of the gas phase [324]. The latter effect was associated with the hydrogen uptake of MoS2 in an H2 –H2 S atmosphere. Schrader and Cheng [323] and Payen et al. [325] monitored the transformation of the surface molybdate in the oxide precursor into the sulfided state using in situ LRS. Molybdenum oxysulfides were detected as intermediate species prior to the MoS2 formation. 3.1.3.8.5 Vibrational Spectra of Surface and Structural Hydroxy Groups Oxide surfaces are preferentially terminated by hydroxy groups for energetic reasons [327, 328] and structural hydroxy groups must be present in zeolites and molecular sieves when protons are involved as charge-compensating cations. Hydroxy groups play a most important role as surface functional groups in catalysis and surface chemistry as either acidic or basic sites or as reactive surface groups [328]. OH groups represent diatomic surface oscillators, which give rise to typical O−H stretching frequencies. They can therefore be considered as intrinsic surface probes, the vibrational frequencies of which provide information on the coordination of the probe and, hence, on the local surface structure. For most disperse high surface area oxides, the real surface structure and the type of terminating crystallographic planes are often unknown, so that IR spectra in the hydroxy stretching region may be the only source of information on the surface structure of the particles constituting the oxide material under consideration. References see page 965

3.1 Physical Properties

B 4550

7285

Reflectance

A

Wavenumer / cm−1 Fig. 23 Diffuse reflectance spectra in the near-infrared region of an SiO2 surface after dehydroxylation at (A) 473 and (B) 773 K (adapted from Ref. [329]).

A

3660

B

2630 2760

3660 3550 3740

Adsorbed molecular water may complicate the identification of surface OH groups since the OH stretching vibrations of both species occur in the same wavenumber range (3200–3800 cm−1 ). However, the v2 deformation vibration of the water molecule occurs between 1600 and 1650 cm−1 , whereas that of, e.g., surface silanol groups was found near 870 cm−1 [329, 330]. A discrimination between surface OH groups and adsorbed molecular water therefore becomes possible on this basis, as shown for the Al2 O3 /H2 O [331, 332] and TiO2 /H2 O [333] systems. The absence of the deformation band near 1600–1650 cm−1 , however, does not necessarily indicate the complete desorption of molecular water, since the extinction coefficient of the deformation vibration is significantly lower than that of stretching vibrations. The combination bands v2 + v3 of molecular water and (vOH + δOH ) of hydroxy groups permit a distinction between both species. The (v2 + v3 ) mode of water adsorbed on silica surfaces has been observed between 5100 and 5300 cm−1 (depending on the degree of hydration), whereas the (vOH + δOH ) band of surface silanol groups is located at about 4550 cm−1 [329, 334–336]. As an example, Fig. 23 shows the overtone and combination spectra measured in diffuse reflectance of a dehydrated aerosil sample for two dehydroxylation temperatures. The (vOH + δOH ) band of the silanol groups occurs at 4550 cm−1 and the first overtone of the O−H stretching mode v02 is found at 7285 cm−1 (the fundamental vibration v01 is near 3750 cm−1 , see below). A second problem is related to the discrimination between surface and inaccessible internal or bulk hydroxy groups, which exist, for example, in silica particles. As shown in Fig. 24, an Aerosil surface after dehydration at 473 K gives rise to a sharp band at 3740 cm−1 , a broader feature with a maximum at 3660 cm−1 and a shoulder near 3550 cm−1 . The 3660 cm−1 band was

Transmission

960

Wavenumber/cm−1 Fig. 24 Infrared transmission spectra of (A) hydroxy and (B) deuteroxy groups on an SiO2 surface after thermal treatment at 473 K (adapted from Ref. [328]).

ascribed to inaccessible internal or bulk OH groups [337, 338], whereas the band at 3550 cm−1 is assumed to be due to mutually interacting surface silanol groups [337–339]. Deuterium exchange with D2 O quantitatively shifts the 3740 cm−1 band to 2760 cm−1 and the shoulder near 3550 cm−1 to 2630 cm−1 (corresponding to a wavenumber shift by a factor of 0.74), whereas a broad band at 3660 cm−1 is observed in the OH stretching region, which does not find an isotope-shifted analogue in the O−D stretching region (Fig. 24). The corresponding OH groups are therefore inaccessible for deuterium exchange and must be assigned to internal OH groups, which cannot interact with adsorbed species [339]. The assignment of the O−H stretching frequencies of isolated and unperturbed OH groups on the surface of oxide materials deserves detailed consideration. A multiplicity of O−H stretching bands has been observed for many binary oxides. Examples are complied in Table 9. Band assignments have been made on the assumption that the crystallites are terminated by different low-index crystallographic planes, suggesting the coordination of surface OH groups as a frequency-determining factor [370–372]. Tsyganenko and Filimonov [341] proposed a systematic interpretation by correlating the O−H stretching frequencies with the coordination numbers of the hydroxy oxygen atom. Coordination numbers of 1–4

3.1.3 Structure and Morphology

Tab. 9

Vibrational frequencies (v/cm−1 ) of unperturbed hydroxyl groups on binary oxide surfaces

Oxide

vOH /cm−1

BeO

3730, 3620 3735, 3630 3752, 3610 3745 3730, 3610 3740 3745 3746, 3610 3707, 3695 3765, 3650 3700, 3610 3680–3700, 3733 3740–3745, 3760–3780 3785–3800 3745–3750 3673 3640 3725, 3670, 3645 3728, 3707, 3672, 3654, 3636 3715, 3665 3720, 3665–3675 3740, 3690, 3660 3685 3700, 3670 3680–3690, 3650–3670

MgO

CaO

Al2 O3 SiO2 GeO2 SnO2 TiO2 (anatase)

TiO2 (rutile)

α-Cr2 O3 ZnO

ThO2

Scheme 2

2745, 2660

2760, 2670

2735, 2719

2725–2733, 2755–2760 2790–2803

Ref. [340] [341] [342] [343] [344, 345] [346] [340] [347] [348] [349] [341] [334, 350, 351]

2760 2702 2640

[334, 341] [352] [353] [354] [355] [356, 357] [360, 361] [358] [356] [359] [362, 363] [364] [365] [366] [367] [368]

2755, 2730, 2700

3670, 3620 3680, 3620 3770, 3670 3760, 3660 3725, 3662 3742

for hydroxy groups are found in many compounds [373]. For simplicity, the various OH groups having coordination numbers 1, 2, 3, etc., are referred to as type I, II, III, etc., hydroxy groups (Scheme 2). Table 10 represents some assignments given by Tsyganenko and Filimonov [341] on the basis of this concept. The first row gives the oxygen coordination number in the bulk lattice of the respective binary oxide. The maximum coordination number of a surface hydroxy oxygen atom must necessarily be lower by one. Hence only terminal type I hydroxy groups are expected for silica, whereas type I and type IV hydroxy groups can occur on alkaline earth metal oxides.

H O

H O M

vOD /cm−1

2700, 2675

ZrO2 (monoclinic) ZrO2 (tetragonal)

H O M

961

M

M M M

2780, 2703

2758

[369]

Vibrational frequencies of free hydroxyl groups on binary oxides with different crystal structure (adapted from Ref. [341])

Tab. 10

Oxygen coordination number 2 3 4

6

Frequency/cm−1

Oxide

SiO2 TiO2 (anatase) γ -Al2 O3 ZnO CeO2 ZrO2 MgO CaO

Type I

Type II

3750 3725

3670

3800 3710 3770 3750 3700

3740 3675

Type III

Type IV

3700 3622 3640 3670 3630 3610

Comparing the OH stretching frequencies of OH groups of the same type for various oxides clearly indicates References see page 965

962

3.1 Physical Properties

SiOH

GeOH

3785

3695

3715

Sn–OH

3725

Transmission

3775

SiOH

Wavenumbers/cm−1 3700

3500

3700

3500

Wavenumber/cm−1 Infrared transmission spectra in the O−H stretching region of modified Aerosil silica surfaces (adapted from Ref. [376]).

Infrared transmission spectrum in the hydroxy stretching region of γ -Al2 O3 after thermal dehydroxylation at 773 K; spectrum recorded at 80 K (adapted from Ref. [363]).

Fig. 26

Fig. 25

the strong influence of the nature of the metal atom. It is well known that the O−H stretching frequency of hydroxides of various elements depends linearly on the electronegativity of the central atom [374]. The same trends were observed for surface hydroxy groups [375, 376] and the correlation was rationalized by quantum chemical calculations [375, 377–379]. A representative example is shown in Fig. 25 for SiO2 samples that were modified with ≡GeOH and ≡SnOH groups. The respective O−H stretching frequencies for these type I hydroxy groups were found at 3750 cm−1 for SiOH, at 3682 cm−1 for GeOH and at 3665 cm−1 for SnOH [376]. A special and most complicated case is that of alumina surfaces. The transitional phases η- and Comparison of the O−H (O−D) stretching fundamentals (v/cm−1 ) for γ -, η- and δ-Al2 O3

Tab. 11

vOH

vOD

γ -Al2 O3 [379a, 380]

γ -Al2 O3 [380a]

η-Al2 O3 [380b]

η-Al2 O3 [351]

γ -Al2 O3 + δ-Al2 O3 [380c]

3700 3733 3744 3780 3800 2733 2759 2803

3700 – 3745 3760 3785 2725 2760 2790

3710 – 3740 – 3785 2730 2755 2790

3695 – 3730 3775 3795

3670 3685 3727 3775 3790

γ -Al2 O3 have a defect spinel structure in which a certain number of cation positions remains unoccupied. Table 11 summarizes O−H and O−D stretching frequencies of isolated unperturbed hydroxy groups that have been reported for η-, γ - and δ-Al2 O3 by various research groups and Fig. 26 shows a typical hydroxy IR spectrum of γ -Al2 O3 . Considering the (100) and (111) planes, Tsyganenko and Filimonov [341] concluded that only three types of OH groups could occur, namely types I, II and III, with O−H stretching frequencies at 3800, 3740 and 3700 cm−1 (see Table 10). The observation of two additional, although weaker, bands at 3780 and 3733 cm−1 (see Table 11) was accounted for by the fact that Al cations occupy octahedral and tetrahedral sites. The bands in the 3760–3780 cm−1 range were assigned to OH groups in the coordination sphere of tetrahedrally coordinated Al3+ cations (Altet ) and those in the range 3700–3750 cm−1 to OH groups being shared by an octahedrally and a tetrahedrally coordinated Al3+ cation (Aloct and Altet , respectively) [371, 372]. Peri [380, 381] observed a single band at 3800 cm−1 in AlPO4 , in which Al3+ occurs exclusively in tetrahedral coordination, and Morterra and coworkers detected bands at 3730–3740 cm−1 on the surfaces of α-Al2 O3 [372] and of MgAl2 O4 [382] containing only octahedral Al3+ cations. A simple electrostatic model was proposed by Kn¨ozinger and Ratnasamy [351] for the interpretation of the O−H stretching frequencies of hydroxy groups on the surfaces of transition aluminas, which turned out to be consistent with the assignments mentioned above. The model was based on the general concept that the

3.1.3 Structure and Morphology

O−H stretching force constant was largely determined by the coordination of the OH group, with the oxygen coordination of the involved Al3+ cations being an additional determining factor. The low index planes (100), (110) and (111) were considered as terminating the crystallites. Within an OH layer on top of the A layer [351, 383] of a (111) plane, two types of OH configurations can be distinguished, namely a terminal OH group which is coordinated to a single tetrahedral Al3+ cation: type Ia

(Altet )OH

and a bridging OH group which links a tetrahedral and an octahedral cation: type IIa

(Altet )(Aloct )OH

Similarly, in the OH layer on top of a B layer of the (111) plane two OH configurations can occur. Both are bridging groups. The type IIb configuration links two octahedral cations: type IIb

(Aloct )2 OH

whereas the type III OH group is coordinated to three cations in octahedral interstices: type III

(Aloct )3 OH

If potential vacant cation sites are also taken into account, a fifth configuration, type Ib, can also exist: type Ib

(Aloct )OH

Based on electrostatic arguments using Pauling’s electrostatic valence rule [384], the O−H stretching frequency assignments given in Table 12 were proposed. The most important result of the above considerations is that one should expect a maximum of five different OH configurations on the surface of transitional aluminas with the Tab. 12

(111)

(110)

(100)

spinel structure, their occurrence and relative abundance depending on the relative contributions of the different crystal faces. The OH groups in these various configurations bear slightly differing net charges. As a consequence, they should possess different properties, namely bond orders and force constants. The five observed frequencies of the stretching vibrations of isolated surface OH groups should therefore correlate with the estimated net charges of the five possible OH configurations. The band with the highest frequency (3800 cm−1 ) is assigned to configuration Ib, which bears the most negative net charge (–0.5), whereas the band of lowest frequency (3700 cm−1 ) is attributed to configuration III with a net charge of +0.5. The remaining three bands are assigned following decreasing frequency to the corresponding configurations with increasingly positive charge. This assignment is in agreement with the fact that the oscillator frequency ωe of free OH− species is 3839 ± 10 cm−1 and decreases by about 100 cm−1 to 3735 cm−1 for the neutral OH• radical [385]. In these species a shift of the stretching frequency by 100 cm−1 thus corresponds to a charge difference of one elemental charge. The same estimated charge and frequency regions are exhibited by the isolated OH group configurations on alumina surfaces. The simple electrostatic model of the alumina surface has found support from quantum chemical calculations [378, 386]. However, recent molecular dynamics simulations seem to indicate that microcrystals of γ -Al2 O3 are not terminated by well-defined crystallographic planes but rather bear an amorphous outer layer [387, 388]. It was also proposed that more than five chemically distinct OH configurations would exist on this surface [387]. Several more recent experimental and theoretical publications reported on modifications of the electrostatic model and O−H band assignments [389–394]. References see page 965

Possible OH configuration and corresponding O−H stretching frequencies on alumina surfaces

Crystal face

Layer

B B A A A, B C C D

Configuration

(Aloct )3 OH (Aloct )2 OH (Aloct Altet )OH (Altet )OH (Aloct )OH (Aloct )2 OH (Altet )OH (Aloct )OH (Aloct )OH

963

Type

III IIb IIa Ia Ib IIb Ia Ib Ib

Coordination numbers of surface anion Aloct

Altet

3 2 1 – 1 2 – 1 1

– – 1 1 – – 1 – –

νOH /cm−1

3700–3710 3740–3745 3730–3735 3760–3780 3785–3800 3740–3745 3760–3780 3795–3800 3785–3800

964

3.1 Physical Properties

It is interesting that broad bands had been observed at about 3500 and 3300 cm−1 in transition aluminas, which were associated with protons trapped in octahedral and tetrahedral vacancies of the defect spinel structure [395, 396]. When all cation vacancies were saturated with protons in the lattice, the material was described as an ideal ‘‘protospinel’’ with the formula HAl5 O8 or Al2 O3 · 0.2H2 O. Earlier, bands were also observed in Hdoped MgO in the frequency range 3300–3340 cm−1 which were assigned to O−H stretching fundamentals of one proton substituting a divalent cation [397]. Further details about the O−H stretching spectra of silicas can be found in review articles [328, 334]. Amorphous aluminosilicates containing up to approximately 50 wt.% Al2 O3 did not give rise to O−H stretching bands of hydroxy groups associated with Al3+ cations [398, 399]. After heat treatment at 773 K and higher temperatures, only the stretching band near 3750 cm−1 attributed to SiOH groups could be observed. According to Peri [398], a band at 3650 cm−1 may or may not appear in spectra of aluminosilicates depending on the composition and extent of surface hydration. This band infers the presence of bridging OH groups of the type Si(OH)Al that are known to exist in crystalline aluminosilicates. Such structural hydroxy groups form in zeolitic materials when protons are present as charge-compensating cations (Scheme 3). The hydroxy stretching spectra of these groups have been studied extensively and described in several reviews [18, 400–402]. They typically give rise to OH stretching frequencies in the range 3600–3650 cm−1 , as can be seen from the data summarized in Table 13. As an example, Fig. 27 shows an IR spectrum in the O−H stretching region for an HY zeolite. Three bands can readily be distinguished. The weak band at 3740 cm−1 is typically attributed to terminal SiOH groups located on the external surface of the zeolite crystallites, whereas the other two bands at 3625 and 3540 cm−1 are characteristic of bridging zeolite hydroxy groups. Based on their accessibility for adsorbate molecules, the first of these bands could be associated with OH groups located in the supercage, whereas the second band must be attributed to OH groups in the sodalite cages [400, 409, 410]. The low frequency of the latter band is thought to be caused by the perturbation of the OH oscillator by the electrostatic field created by the close oxygen atoms in the sodalite cage.

Si

Scheme 3

H O

Al

Hydroxy stretching frequencies of framework OH groups of zeolites

Tab. 13

Zeolite

Al-HZSM-5

Ga-HZSM-5

Fe-HZSM-5 B-HZSM-5 H95 Na5 Y H90 Na10 Y H70 Na30 Y H40 Na60 Y H20 Na80 Y

4000

nSi /nMe 3+ (lattice)

vOH /cm−1

Ref.

23.8 35 65 24 13.6 14–20 29.5 99 34 35 42 2.90 2.47 2.50 2.50 2.50

3618 3610 3610 3617 3617 3625–3610 3622 3620 3615 3630 3725 3643 3640 3647 3648 3650

[403] [404] [405] [406] [407] [254] [403] [404] [405] [404, 405] [404] [407] [408] [407] [407] [407]

3800

3600

3400

3200

Wavenumbers/cm−1 Fig. 27

Hydroxy stretching IR spectrum of zeolite HY.

This low frequency is not observed in ZSM-5 zeolites with MFI structure. A further general characteristic of O−H stretching frequencies of structural OH groups in aluminosilicates is the decrease in the frequency of OH groups in large cages or channels as the Al content is lowered [411, 412]. The O−H stretching bands of structural OH groups in zeolites are typically rather broad, having bandwidths at half-maximum of 30–40 cm−1 even at liquid N2 temperature [403]. This linewidth must be inhomogeneous and originates from a multitude of different bridging hydroxy types which are present in a zeolite matrix. Reasons for this heterogeneity of OH groups in zeolites [413–415] may be either or a combination of: (i) an inhomogeneous distribution of the trivalent cations in the zeolite lattice [386, 416] (ii) crystallographic inequivalence of the T atoms due to varying T−O distances and T−O−T bond angles [413, 415]. The sensitivity of the O−H stretching

References

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3.1.3.9

Neutron Scattering

Herv´e Jobic∗

3.1.3.9.1 Introduction Neutron scattering is a uniquely sensitive probe for studying the structure and dynamics of materials on the atomic and molecular level. Neutron-based techniques cover an extremely wide range of applications, extending from nuclear physics to biology. Several of these techniques have been applied in catalysis research. The structure of hydrogenous layers and of molecules adsorbed in zeolites has been determined by neutron diffraction. Textural or clustering studies have been performed by small-angle neutron scattering (SANS). The vibrational modes of catalysts and of adsorbed molecules have been measured by inelastic neutron scattering (INS). The rotational and translational motions of molecules have been characterized by quasielastic scattering (QENS). Other neutron techniques such as reflection and magnetic scattering have not been used so far in catalysis. 3.1.3.9.2 Interaction of Neutrons with Matter Neutrons interact with nuclei via very short-range nuclear forces. A neutron has both particle-like and wave-like properties: it has zero charge, spin 1/2, a mass slightly larger than that of a proton and a wavelength λ. The potential interest in neutrons, in contrast with other probes, lies in the fact that its wavelength matches interatomic distances and that its energy covers a wide References see page 981 ∗ Corresponding author.

972

3.1 Physical Properties

range of molecular motions. Therefore, both the statics and the dynamics of a sample can be studied. The neutron wavelength is given by the de Broglie relation λ = h/mv, where h is Planck’s constant and v the neutron velocity. The associated wavevector k has the magnitude k = 2π/λ. This allows one to define the neutron momentum as p = mv = ¯hk and the neutron energy as E = 12 mv 2 = ¯h2 k 2 /2m. The two basic quantities in a neutron scattering experiment are the momentum transfer ¯hQ and the energy transfer ¯hω, given by ¯hQ = ¯h(k0 − k )



¯hω = E0 − E =

(1)



 ¯h2  2 k0 − k 2 2m

(2)

where k0 and k are, respectively, the incident and scattered wavevectors. Elastic scattering corresponds to k = k0 , so that only momentum is transferred (¯hω = 0). When the neutrons exchange also energy with the sample, ¯hω will be positive for k < k0 (energy loss or creation of an excitation) and the scattering is called inelastic. In most neutron scattering experiments, one measures the double-differential cross-section, d2 σ/ddE, which represents the number of neutrons scattered into the solid angle d with energy in the range dE. The amplitude of the scattered wave varies between nuclei (because of different isotopes or spins), so that averages have to be performed for each element. The total cross-section per scatterer is obtained by integrating over energies and solid angles:  σ =

d

dσ = d



 dE

d

d2 σ = 4π b2

ddE

(3)

where b is the scattering length, which can be real or complex. The real part is usually positive and the imaginary part is related to the absorption of neutrons. The scattering lengths vary irregularly from one atom to another or even from one isotope to another, and they are determined by experiment. The scattering cross-section can be split into coherent and incoherent contributions. The coherent elastic scattering has a phase term in exp(iQ · R) and thus takes into account interference effects. The coherent scattering cross-section corresponds to an average over all isotopes and spin states: σcoh = 4π b 2

(4)

The incoherent part, which has no phase relationship, corresponds to the difference between the total and the coherent cross-sections:   (5) σinc = 4π b2 − b 2

and therefore represents the mean square deviation from the mean potential, which can be due to isotopic or spin effects. For each natural element, an average over all stable isotopes has to be performed. The relative incoherent and coherent cross-sections of some elements are given in Table 1. For a sample containing only one type of atom (one isotope) with zero spin, then b = b and the scattering is totally coherent (e.g. 12 C). The absolute values are given in barns (1 barn = 10−28 m2 ), which indicates that the interaction of neutrons with matter is relatively weak. It appears from Table 1 that there is no systematic variation of the cross-section across the Periodic Table. This contrasts with X-ray scattering, where the scattering power increases with the number of electrons in the atom. Hydrogen has the largest neutron cross-section, which explains why the most neutron measurements deal with hydrogenated molecules. In the case of diffraction experiments, large incoherent scattering should be avoided since it increases the background below the Bragg peaks. It is therefore better to use deuterated compounds because the incoherent cross-section of D is much weaker than that of H. For inelastic studies, the energy of the scattered neutrons is analyzed and both coherent and incoherent scattering can be measured. In the case of adsorbed systems, a high surface area is required for the substrate. The best signal will be obtained with hydrogenated adsorbates, but the scattering from deuterated molecules or from molecules which do not contain H atoms can be investigated on high-flux instruments. Experimental The production of intense beams of neutrons requires large reactor- or acceleratorbased neutron sources. A complex system of guides and advanced instruments has to be designed to exploit the full power of these costly facilities. At present, Europe still has the lead in neutron science. The two main sources are (i) the high-flux reactor at the Institut LaueLangevin (ILL), in Grenoble (France) and (ii) the ISIS spallation source at the Rutherford Laboratory (UK). The new reactor FRM-II in Munich (Germany) has recently come into operation and its instrument suite is being developed. Some competition is expected in the future: the USA is building a spallation source at Oak Ridge, TN, and Japan is building a similar facility in Tokai (the power level should rise progressively in both places starting in 2006, full power being attained around 2010). Neutrons issued from the source are slowed in moderators whose temperatures can vary between 25 and 2000 K. This gives ˚ neutron wavelengths ranging from 0.2 to 20 A. The number and diversity of neutron instruments around a neutron source are large. Each instrument is

3.1.3.9.3

3.1.3 Structure and Morphology

973

Coherent, incoherent and absorption cross-sections for some elements, in barns (1 barn = 10−28 m2 ). The absorption cross-section is proportional to the neutron wavelength, ˚ The nuclear spins are given for some isotopes, and also their relative abundance (%) here λ = 1 A.

Tab. 1

Element

H

He

B C

N

Atomic number

1

Spin

1 2

1 (99.985) 2 (0.015)

1

3 (0.00014) 4 (99.9998)

0

2

1 2

5 6 12 (98.9) 13 (1.1)

0

14 (99.63) 15(0.37)

1

19 (100)

1 2 5 2

7

O F

8 9

Al Si P S Cl Ar

13 14 15 16 17 18

V Ni Ru Cd Pt

Mass number (natural abundance)

27 (100)

1 2

1 2

31 (100) 32

1 2

36 (0.337) 38 (0.063) 40 (99.6)

0 0 0

23 28 44 48 78

designed to cover a given range of momentum transfer and/or time-scale. There are two main classes of instruments: elastic scattering instruments, which are used for determining the structure of materials, and inelastic scattering spectrometers, which, by measuring energy transfers, give information on atomic and molecular motions. The beam size at the sample position is typically 4 × 3 cm. The absorption of neutrons by matter is small for most elements so that neutrons are highly penetrating and can be used in complex sample environments such as furnaces, cryostats and pressure cells. The few elements which absorb neutrons are used for protection (e.g. Cd) or for detection (e.g. B). After a neutron experiment, the sample may become activated. The amount of time for it to decay is usually of a few days (e.g. for Ni or Pd), but it is several years for Co. Sample quantities vary from a few grams, for diffraction or quasi-elastic experiments, to several tens of grams,

σcoh

1.7568 1.7583 5.592 1.34 4.42 1.34 3.54 5.551 5.559 4.81 11.01 11.03 5.21 4.232 4.017 1.495 2.163 3.307 1.02 11.531 0.458 77.9 1.5 0.42 0.018 13.3 6.53 3.04 11.58

σinc

σabs

80.26 0.19 80.27 0.19 2.05 0 0 0.0004 1.6 2963 0 0 1.7 426. 0.001 0 0 0 0.034 0 0.5 1.1 0.5 1.1 0 0 0 0 0.0008 0.005 0.008 0.004 0.005 .007 5.2 0.225 0 0 0 5.08 5.2 .07 3.46 0.13

0.13 0.1 0.1 0.29 18.6 0.38 2.89 0.44 0.37 2.82 2.5 1.42 1400 5.72

for inelastic measurements. The neutron cell can be cylindrical or slab-shaped, depending on the neutron beam size and scattering strength. For inelastic and quasielastic studies, the cells are usually made of aluminum, but stainless-steel and quartz containers have been used. For diffraction, the sample is generally contained in thinwalled vanadium cans. At present, the ILL is unique for dynamic studies in that a very wide range of energy transfers are accessible using combined instruments (Fig. 1). The range 1–500 meV (1 meV = 8.065 cm−1 ) can be covered using time-of-flight (TOF) and beryllium filter (BeF) spectrometers, allowing the study of vibrational modes of adsorbed molecules. Depending on their time-scales, the rotational and translational motions can be characterized on TOF, backscattering (BS) and neutron spin-echo (NSE) instruments. The energy transfers involved in QENS are References see page 981

3.1 Physical Properties

Structure Neutron diffraction, which corresponds to coherent scattering, is analogous to X-ray diffraction. There are, however, important differences. Neutrons are scattered by the nuclei and X-rays by the electrons. Furthermore, light nuclei are much easier to locate by neutron diffraction in structures that include heavy atoms. In neutron diffraction, no energy analysis is performed. The observed count rate into a solid angle  is proportional to the differential scattering cross-section, dσ/d. Considering a perfect crystal with N unit cells, v0 being the volume of the unit cell, one obtains for the coherent differential scattering cross-section 3.1.3.9.4

dσcoh (2π)3  δ(Q − τ )|Fhkl |2 =N d v0 τ

(6)

The δ function ensures that scattering only occurs when Q coincides with a reciprocal lattice vector τ , which corresponds to Bragg scattering. The intensities of the reflections from the lattice planes (hkl) of a crystal are governed by the structure factor of the unit cell, Fhkl :  Fhkl = bi exp[2πi(hxi + kyi + lzi )] (7) i

where the summation is over the atoms of the unit cell, at positions (xi , yi , zi ). For a real crystal, the atoms are not located at the sites of a perfect lattice, because of static or dynamic disorder. This implies a higher background and a reduction of intensity of the Bragg peaks which is taken into account by a Debye–Waller factor, exp(−2Wd ). Most of the studies performed on surfaces have dealt with small physisorbed molecules; another application of neutron diffraction is the localization of protons and of molecules adsorbed in zeolites [1]. Amongst recent studies [2, 3], one can mention as an example the location of pyridine in AgNaY zeolite (Fig. 2) [2]. Small-angle neutron scattering (SANS) is similar to small-angle X-ray scattering (SAXS). This technique also corresponds to coherent scattering and it gives structural information about inhomogeneities in the sample, with characteristic lengths ranging from 1 nm to 1 µm. Again, an important difference between the two methods lies in the fact that SAXS is sensitive to inhomogeneities of the electron density whereas SANS detects variations in scattering length densities. For two-phase systems, the scattered intensity is proportional to the square

Time scale/ps 104

102

10−2

1

BeF

10

1

1

10 BS

0.1

0.01

TOF 100

Length scale/Å

usually 100

FPIA QELS

Photographic imaging Electromagnetic wave interaction

0.8–300 0.005–2

[dynamic light scattering, photon correlation spectroscopy, optical beating spectroscopy] Heterodyne method Homodyne method Small-angle neutron scattering Ultra-small-angle neutron scattering Small-angle X-ray scattering

SANS USANS SAXS

Wave interaction

0.001–10

Electromagnetic wave interaction

0.001–31

Ultra-small angle X-ray scattering Laser Doppler velocimetry [laser Doppler anemometry]

USAXS LDV

0.5–10

[Coulter principle, Coulter Counter] Electrical sensing zone Differential mobility analysis

ESZ

Aerodynamics and electromagnetic scattering Volume displacement

Scanning electron microscopy Field emission SEM Transmission electron microscopy X-ray gravitational sedimentation Optical centrifugal sedimentation X-ray centrifugal sedimentation Sedimentation field flow fractionation

SEM FE-SEM TEM XGS OCS XCS SdFFF

Sieving Gas adsorption surface area analysis [Brunauer–Emmett–Teller] Acoustic attenuation spectroscopy [ultrasonic attenuation spectroscopy, ultrasonic spectroscopy] Coupled phase theory Scattering theory Multiple scattering theory Electroacoustic spectroscopy [Electrokinetic sonic amplitude]

– BET

C Shape of Catalyst Fine Powders Apart from size factors such as a circle with the same area as the particle shadow or as the perimeter of the particle shadow, shape parameters such as circularity, convexity and aspect ratio are also important for characterizing the morphology of catalyst powders. These factors can only be determined in diluted systems where isolated catalyst particles are resolved. Usually complementary techniques such as optical ‘‘flow particle image analysis’’ and ‘‘scanning electron microscopy with digital image analysis’’ need to be combined to achieve a representative model of the catalyst powder morphology. Finally, a

DMA

Electrostatic classification Imaging

Applicable size range/µm

0.4–>100 0.005–1 0.02–10

Imaging Sedimentation Sedimentation Sedimentation Sedimentation classification Size exclusion Surface area

0.01–0.5 0.5–100 0.01–>5 0.01–>5 0.03–>1

AAS

Acoustic wave interaction

0.025–>100

ESA

Electroacoustic response

0.1–10

2–>100 no limit

statistically significant number of catalyst particles of the order of several thousand objects should be analyzed. State-of-the-art digital image analysis can be applied to most automatically operating optical tools. These systems are commercially available. A straightforward interpretation is only possible if catalyst particles of uniform nature (only one phase) are present. In other cases, the morphology pattern may serve only as a fingerprint tool. However, the so-called scattergrams can be very characteristic for a given catalyst powder, where a particular morphological parameter is monitored as a function of the size fraction of the catalyst particles.

3.1.3 Structure and Morphology

Some of the morphological parameters are described qualitatively below. The ‘‘area convexity’’, sometimes called the ‘‘solidity’’, of an image of a catalyst particle may be calculated by the object area A, divided by the area which is confined by the so-called convex hull. This area is the sum of the shadow area A plus the area which is circumscribed by the tangent that touches two points of the shadow area and the corresponding segment of the object. The shadow image of a rough surface may show several or many convex segments and the convex hull area is calculated accordingly. The perimeter convexity is the convex hull perimeter divided by the actual object perimeter. The aspect ratio is the minimum diameter divided by the maximum diameter of a catalyst particle. The elongation is defined as 1 – aspect ratio. The circularity is defined as the circle circumference of the equivalent area divided by the perimeter. By definition, a sphere has a circularity 1. A square of the same side length as the sphere diameter has a circularity 0.886 and a triangle of the same height and the same base line has a circularity 0.777. A rectangle with an aspect ratio of 1 : 4 has a circularity of 0.660 and a rectangle with an aspect ratio of 1 : 10 has a circularity of 0.509. D Determination of the Size by Classification In catalyst production, the ‘‘classification process’’ [7] is used to control or limit the shortcomings of the previous processing stages, for example to eliminate oversize particles at the exit of a mill or spray drier. ‘‘Separation’’ generally refers to dissimilar materials, whereas ‘‘classification’’ refers to different grades of the same material [8]. Classification techniques are also applied to characterize catalyst powders with respect to their particle size distribution. The technologies may use liquid suspensions Tab. 3

(‘‘wet’’: water) or a carrier gas (‘‘dry’’: air). The wet technologies include hydrocyclones, centrifuges, wet dynamic classifications, screens/sieving and sedimentation and the dry methods include sieving, static classifiers (cyclones), dynamic classifiers (single-stage, multi-stage), cross-flow classifiers and counter-flow classifiers (elutriators). a Sieving (125 mm to 37 µm) Particle size analysis from about 125 mm to 37 µm by using a stacked series of woven wire or punch plate sieves arranged in decreasing order of aperture size [9] is still one of the most frequently used methods to characterize a powdered catalyst sample. The lower size range starting at about 5 µm can be included in the analysis by using micromesh sieves, but at the expense of the time required to sieve an equivalent mass of powder. Usually, the sieves are identified according to their ASTM mesh size. A 400-mesh sieve relates to a minimum square aperture of 37 µm. Sieving can be performed either dry or wet, with manual or machine agitation and for a set time or until a sufficiently low and constant powder flow-rate is observed through the sieves. Key variables that influence sieving results include particle shape, presence of very fine particles, initial sieve loading, time and method of agitation and cohesiveness of the powder (in dry sieving only). Repeatability can be high, although reproducibility is often poor due to the many variables that provide sources for user error [10]. Several sieving machines are commercially available. Table 3 shows some of the corresponding standards. The results of the sieve analysis should be combined with data from other sizing techniques to provide a composite report accounting for all of the sample material. References see page 997

Standards for sieving analysis

Standard

Title

ISO 565:1990

Test Sieves – Metal Wire Cloth, Perforated Metal Plate and Electroformed Sheet – Nominal Sizes of Openings Test Sieves and Test Sieving – Vocabulary Test Sieving – Part 1: Methods Using Test Sieves of Woven Wire Cloth and Perforated Metal Plate Test Sieves – Technical Requirements and Testing – Part 1: Test Sieves of Metal Wire Cloth

ISO 2395:1990 ISO 2591–1:1988 ISO 3310-1:2000 ISO 3310-1:2000/Cor 1:2004 ISO 3310-2:1999 ISO 3310-3:1990 ISO/CD 9276-6

987

Test Sieves – Technical Requirements and Testing – Part 2: Test Sieves of Perforated Metal Plate Test Sieves – Technical Requirements and Testing – Part 3: Test Sieves of Electroformed Sheets Representation of Results of Particle Size Analysis – Part 6: Descriptive and Quantitative Representation of Particle Shape and Morphology

988

3.1 Physical Properties

Tab. 4

Standards for sedimentation measurements

Standard

Title

ISO 13317-1:2001

Determination of Particle Size Distribution by Gravitational Liquid Sedimentation Methods – Part 1: General Principles and Guidelines Determination of Particle Size Distribution by Gravitational Liquid Sedimentation Methods – Part 2: Fixed Pipette Method Determination of Particle Size Distribution by Gravitational Liquid Sedimentation Methods – Part 3: X-Ray Gravitational Technique Determination of Particle Size Distribution by Centrifugal Liquid Sedimentation Methods – Part 1: General Principles and Guidelines Determination of Particle Size Distribution by Centrifugal Liquid Sedimentation Methods – Part 2: Photocentrifuge Method

ISO 13317-2:2001 ISO 13317-3:2001 ISO 13318-1:2001 ISO 13318-2:2001

b Sedimentation For particles which are too small to be classified by sieving, elutriation and sedimentation are used. The airflow concepts of elutriation, free vortex and forced vortex are used either separately or in combination in the design of today’s air classifiers. Elutriation is the process of separating by washing, in the case of air classifiers using air as the washing medium. Elutriation is generally used to separate the bulk of fines and coarse particles by introducing the feed into an air flow. The air should be humidified to avoid aggregation. The fractionation tube is maintained in continuous vibration. The airflow raises the fine particles against gravity to a fines collector. Being too heavy to be carried upwards, the coarse particles decelerate and fall with gravity, against the flow, into the coarse fraction collector. The cut point can be adjusted by increasing or decreasing airflow velocity within the tube. Air classification that achieves cut sizes of a few micrometers typically follows Stokes’ law: " dp = 18ηv/(ρp − ρL )g (8)

where dp is the theoretical cut-size, η is the air viscosity, ν is the airflow velocity, ρp is the particle density and g is gravitational acceleration. Equation (8) describes the theoretical cut-size based on elutriation. Increasing or decreasing the classifier’s air-flow velocity (v) adjusts the cut-size. Elutriation is a crude method of separation that is seldom used alone. Today, sedimentation methods are used in production rather than for the purpose of characterization of catalyst powders. In particular, the information on the particle morphology gained by this method is very limited. The X-ray sedimentation technique is based on similar principles but the size is recorded as the catalyst particles settle through a liquid medium. The mass fraction in each size class is determined by the absorption of soft X-rays. These instruments are known under the brand ‘‘Sedigraph’’ The particle size measurement range is 0.1–300 µm.

The corresponding standards for sedimentation are summarized in Table 4. E

Determination of the Size by Optical Methods

a Microscopic Methods The interaction of electromagnetic radiation with a catalyst powder specimen can cause absorption, diffraction and scattering and the interaction of a beam of electrons with a catalyst powder sample may cause in addition backscattering of primary electrons and formation of secondary electrons. In addition to scattering and absorption, there is another feature of light that can be used for determining the particle size, namely the frequency. Information from frequency changes or shifts is used in ‘‘dynamic light scattering’’. All the optical methods are widely and frequently used in heterogeneous catalysis to determine particle size, particle size distributions and morphology. In special cases also chemical information can be achieved. The optical microscope (µm to mm) Optical microscopes are the fundamental tools for obtaining an overview of the size and morphology of catalyst powder particles – dry or in optically transparent suspension. Basically, the optical microscope magnifies an image by sending light through the object plane. The condenser lens focuses the light on the sample plane and the objective lenses (magnification from ×10 up to ×2000) magnify the beam, which contains the image, to the projector lens so the image can be analyzed by the observer. Depending on the particle size, visible light is mostly absorbed or scattered and cannot pass through catalyst particles. Therefore, a shadow image is obtained. To study the surface structure of catalyst particles, a type of microscope is used, where the light is reflected from the examined surface. The light is fed through the same objective using a semi-transparent mirror. A stereo microscope allows even better inspection of the surface of catalyst particles and particularly to analyze

3.1.3 Structure and Morphology

whether the shadow of the particle observed by the transmission microscope is in reality composed of strongly bound agglomerates of primary particles. Moreover, the stereo microscope reveals some morphological features of the particles. The stereo, binocular or dissecting microscope uses two eyepieces (or sometimes two complete microscopes) to provide slightly different viewing angles to the left and right eyes. In this way it produces a three-dimensional (3-D) visualization of the sample being examined. The stereo microscope has a useful maximum magnification of only ×100. All three types, the transmission microscope, the reflection mode microscope and the stereo microscope, are commercially available and all three commercial types include a calibration scale in the field observed. The maximum resolution that one can obtain is controlled by the wavelength of the light that is used to probe the catalyst particle; nothing smaller than the wavelength being used can be resolved. Visible light has wavelengths of 400–700 nm; larger than many catalyst powder particles of interest. Significantly better statistical confidence compared with the above-discussed microscopes is obtained if a sample is taken from a dilute suspension of particles and passed through a measurement cell in a flat particle flow, which ensures that all particles lie in the same focal plane of a stroboscopic illumination (about 60 s−1 ) and that are oriented with their largest area facing a charge-coupled device (CCD) camera which takes the images. Each individual particle representation is automatically extracted and analyzed for size and shape by conventional electronic image processing techniques. For particle sizes ranging from about 0.8 to about 300 µm these easy to handle particle analyzers are commercially available. Transmission electron microscope ( 0 with Eb (i) ≈ 0; and (5) I < 0 with Eb (i) ≈ 0. Note that the results of this analysis have to be considered as semiquantitative due to charging problems. This is because an error δ in the charge referencing procedure introduce an error 2δ in the intercept [and therefore in the 2(VM + kq) value]. With reference to Fig. 2, we can discuss the different cases using practical examples. In the Zn Wagner plot for ZnSO4 we can write 2[VM (ZnSO4 ) + kq(ZnSO4 )] = 2[VM (Zn) + kq(Zn)] + I (ZnSO4 )

(14)

with I (ZnSO4 )  0. The Madelung potential in the case of cations is always a negative quantity and for each chemical state |VM (i)|  |VM (m)|. Equation (15) demonstrate that the position of the ZnSO4 in the Wagner plot is essentially determined by the high negative value of the Madelung potential. Note that in the Zn Wagner plot for all the chemical states I < 0. Note also that kq(i) is always a positive quantity for a cation and the highest negative values for I are measured at the lowest binding energies with respect to the metallic state. In the Zn Wagner plot, a few compounds lie very close to the line of the metallic state with slope +3. This means that for these chemical states I ≈ 0 and therefore we can write [VM (i) − VM (m)] = −kq(i)

(16)

An increase in the local charge q(i) leads to a parallel increase in the negative value of the Madelung potential VM (i). The analysis discussed above is also in agreement with the empirical observations reported by O’Keeffe [32]. O’Keeffe showed that the Madelung potential at an ion site scales approximately as the charge on the ion, i.e. VM (i) = −aq(i)

(17)

where the value of the constant a scales with the size of the atom concerned. A final case to be discussed in relation to the Wagner plots is when the Eb values for different chemical states are practically constant and equal to the value for the metallic state. In this case, from the Wagner plot we have I = α (i) − α (m)

(18)

This means that [VM (i) + kq(i)] − VM (m) = [R(i) − R(m)]

(19)

This situation occurs when the initial and the final state effects have a similar value. [Note that Eq. (19) can be written as Eb (i) = Eb (m).] In this case, when the Auger parameter shift is negative [i.e. I (i) < 0] we have |VM (i)|  kq(i). The shift is more negative the more negative is the Madelung potential. According to Eq. (19), this is also related to the ionic character of the chemical bond. In fact, considering Fig. 2, we see that at the value References see page 1038

1036

3.2 Chemical Properties

of the binding energy of the metallic state the negative Auger parameter shifts are in the order ZnTe < ZnS < ZnO < ZnAl2 O4 , in agreement with the order of ionicity of the chemical bonds. [No examples of positive Auger parameter shift (i.e. I > 0, see Eq. (18)) are present in Fig. 2. In this case |VM (i)| ≤ kq(i) and the shift should be more positive the more covalent the compound is.] The same analysis of the Auger parameter shifts and of the initial state effects reported above can also be applied to the cases of atomic-like core–core–valence and core–valence–valence Auger transitions. An example of application of the oxygen Auger parameter in the characterization of oxygen-containing compounds is reported in Ref. [18]. In the case of band-like core–core–valence and core–valence–valence Auger transitions the above approximations are not valid and the analysis should be based on the Cini–Sawatzky theory, from which the relevant initial and final states parameters can be obtained by means of Auger lineshape analysis [16–18]. 3.2.3.1.8 Thermochemical Approach Core-level electron binding energies of alloys and mixed-valence compounds can also be interpreted using a thermodynamic approach based on the equivalent cores approximation. This method bypasses detailed electronic considerations and refers only to macroscopic quantities. This approach was used extensively by Jolly and co-workers (see Ref. [33] and references cited therein) to improve predictions on chemical shifts in inorganic molecules and solids and to estimate thermodynamic energies from XPS data. The general idea is to exchange the core ionized Z ∗ atom with the valence ionized state of the (Z + 1)+ atom. It is assumed that this exchange requires an energy that depends only on the atomic number Z of the core-ionized atom and is independent of its chemical environment. According to this approximation also a (Z + 1) impurity atom dissolved in a Z metal is electronically identical to a fully screened core-ionized Z ∗ atom in the Z metal. With these assumptions, a Born–Haber cycle can be performed following the general procedure outlined by Jolly [33] ˚ and Johansson and Martensson [34]. Applications of the method have been reported in studies of metal bulk and alloys [34], surface core-level shifts [35] and binding energies in adsorbates [36]. The binding energy shift between the metal core level referred to the Fermi M and the corresponding free atom referred level Ec,F to the vacuum level EcA is given by the following equation: Z+1 Z+1 M Z = I1Z+1 + Ecoh − Ecoh + Eimp EcA − Ec,F

(20)

where M denotes the metal, A the atom, c the core, F the Fermi level, Z the atomic number and I1 the first

Z+1 ionization energy. Ecoh is the cohesive energy and Eimp is the solution energy of impurity atom (Z + 1) in Z metal. The dominating contributions to the shifts are the difference in cohesive energy between the (Z + 1) and Z metal and the first ionization energy of the (Z + 1) atom. The connection between electronic and thermodynamic quantities which emerges from such cycles shows that photoelectron spectroscopy can be used to determine such relatively inaccessible thermodynamic quantities as the implantation energy. The application of the equivalent cores approximation in combination with a Born–Haber cycle to the interpretation of core-level binding energies for adsorbates on metal surfaces is also very useful and of direct relevance to heterogeneous catalysis. Tomanek et al. [36] described the adsorption of CO on metal surfaces using the C 1s binding energy. Starting from the initial state, which is a chemisorbed CO molecule, we have first to overcome (i) the adsorption energy Echem (CO/Me) to desorb CO and then the dissociation energy D(CO) to obtain isolated atoms. Next, the gaseous C atom is core ionized. Within the ‘‘equivalent cores’’ model, the final state of the core ionized Z ∗ atom can be replaced by the (Z + 1)+ ion with its core electrons all present, but its highest-lying valence electron removed. Thus, the C∗ 1s atom is replaced by an N+ 2p ion. For a C atom in the adsorbed CO the fully screened final state C* (1s hole, plus a screening electron) is equivalent to N (neutral). Thus, in order to neutralize the N+ ion, the first ionization energy I N of N is released. Next, the dissociation energy D(NO) is released when N combines with oxygen atom. Finally, NO is adsorbed on the surface thereby gaining the adsorption energy (f ) Echem (NO/Me) of an NO molecule. From this cycle, the C 1s binding energy can be calculated as (i)

EbF (C 1s; CO/Me) = Echem (CO/Me) + D(CO) + EbV (C 1s; C) − I N − D(NO) (f )

− Echem (NO/Me)

(21)

For calculating the C 1s binding energy values in CO adsorbed on different metal substrates, Tomanek et al. [36] used the values D(CO) = 11.2 eV, I N = 14.5 eV and D(NO) = 6.5 eV. The chemisorption energies of CO and NO on several transition metal surfaces are reported in Table 1 in their paper. Using the equivalent cores approximation C∗ ≈ N+ and a Born–Haber cycle involving CO2 and NO2 , the binding energy of the carbon free atom EbV (C 1s; C) is estimated to be 295.4 eV [36]. Tomanek et al. [36] extended their analysis to surface species containing nitrogen and the halogens bromine and iodine.

3.2.3 Valence States

Mo/Al2O3 Mon +3d

Intensity/counts s−1

H2 1173 K

0 2 3

5

6 5 4 3

6 5

4 3

6

Oxidic

245

240

n = 0,2,3

2 n = 0,2,3,4 2 3 0 4 3 5 04

H2 973 K

H2 773 K

0 2 3

n=6

6

235

n = 3,4,5,6

230

225

220

Eb /eV Typical Mo 3d curve-fitted spectra for oxidic and reduced Mo/Al2 O3 catalysts. The Mo 3d5/2 –3d3/2 doublets are designated by their assigned oxidation states. (Handbook of Heterogeneous Catalysis, 1st Ed.)

Fig. 3

3.2.3.1.9 Illustrative Examples of Application of XPS and AES-XAES to Heterogeneous Catalysis We report some illustrative examples of the application of XPS and Auger spectroscopy in the characterization of heterogeneous catalysts to detect directly chemical changes which could be directly related to the catalytic activity. Several reviews have been devoted to these subjects [1d, 8, 37, 38] and also to the important related subject of the characterization of model catalysts based on single crystals (see the pertinent chapters in this book and references cited therein). An analysis of Mo 3d XPS spectra of supported Mo catalysts reported by Hercules and co-workers [39, 40] can illustrate very well the potential of XPS in detecting changes in oxidation states of Mo after several treatments. Mo/Al2 O3 catalysts were prepared by incipient wetness impregnation using ammonium paramolybdate solutions. The samples were studied after oxidation in O2 at 773 K and after reduction in H2 in the range 773–1173 K. The different pretreatments are those typically employed in the activation of these catalysts for propene hydrogenation, propane hydrogenolysis and benzene hydrogenation. In every case the samples were pretreated and transferred to the spectrometer analysis chamber without contact with the ambient atmosphere.

1037

In Fig. 3 we report the typical Mo 3d curve-fitted spectra for oxidic and reduced Mo/Al2 O3 catalysts. The 3d5/2 and 3d3/2 doublets are designated by their assigned oxidation states according to Hercules and co-workers [39]. The binding energy values were corrected for charging using the Al 2p line of the support fixed at 74.5 eV. A non-linear least-squares fitting routine was used to curve fit the Mo 3d envelope after background subtraction. The results show that Mo oxidation states ranging from +6 to 0 are produced on reduction in the range 773–1173 K (see Fig. 3). Note that for theoretical reasons each 3d5/2 –3d3/2 doublet has a fixed intensity ratio and a fixed energy separation. The binding energies of Mo 3d5/2 for Mon+ species supported on γ -Al2 O3 determined by Hercules and co-workers were as follows [39, 40]: Mo6+ ≈ 233.1 eV, Mo5+ ≈ 231.8 eV, Mo4+ ≈ 229.9 eV, Mo3+ ≈ 228.8 eV, Mo2+ ≈ 228.2 eV and Mo0 ≈ 227.8 eV. It is important to note that the Mo average oxidation state estimated from XPS for a reduced catalyst agrees with that obtained by measuring the O2 consumption on reoxidation. The catalytic activity for benzene hydrogenation occurs only for reduction temperatures ≥873 K. A correlation of the catalytic activity with the Mo oxidation states determined by XPS indicates that Mo2+ and Mo metal are the active species, with Mo metal the most active species [39]. Hercules and coworkers also found that the catalytic activity for propene hydrogenation and for propane hydrogenolysis increases for Mo oxidation states ≤ +4. In particular, the catalytic activity for propane hydrogenolysis correlated with the quantity of Mo metal [40]. Hercules and co-workers also reported the variation of XPS IMo 3d /IAl 2p intensity ratio as a function of reduction temperature at constant Mo loading (Mo 8 wt.%). A decrease in the intensity ratio with increasing reduction temperature up to about 973 K followed by a leveling off at higher temperatures was found. The decrease in the IMo 3d /IAl 2p intensity ratio could not be attributed to changes in the BET surface area and was interpreted as being mostly due to some recrystallization of the reduced moieties or to changes in the morphology of the reduced species (e.g. tower-like shape vs. platelets for the oxidic catalysts) [39]. In a recent paper, Briand et al. [41] reported on the methodical aspects in the XPS analysis of molybdena catalysts with TiO2 , CeO2 and Al2 O3 supports. They found that supported Mo oxide becomes reduced under X-ray irradiation during extended XPS data acquisition and established with higher accuracy the monolayer capacity of Mo oxide species. The XPS quantitative analysis of supported catalysts was reviewed by Cimino et al. [42]. Starting from the basic quantitative equations of XPS, the development of different models is summarized, including those for References see page 1038

1038

3.2 Chemical Properties

flat surfaces and rough profile surfaces, and also models treating high surface area and porous materials. A few examples of the application of XPS and AES to the characterization of zeolites and metal-exchanged zeolites is also timely (for reviews, see Refs. [8], [20] and [23–25]). Okamoto et al. [43], using XPS, demonstrated that the basic strength of framework oxygen, related to the O 1s binding energy, increases with increasing Al content regardless of the crystal structure and with decreasing electronegativity of the counter cation. The dependence of the core-level binding energies of Si, Al and O bonded in zeolites on the concentration of skeletal Al was also investigated by Jirka [44]. Poncelet and co-workers characterized the Al coordination at the outer surface of dealuminated mordenites [45, 46] and beta zeolites [47] by XPS using a careful analysis of the Al 2p and the Al KLL Auger transitions excited by X-rays. The basis of their analysis was the Al Auger = E (Al 2p) + E (Al KLL). The parameter defined as αAl b k authors using the simple electrostatic model of the Auger parameter (see the discussion above and Ref. [20] for details of the application of the model to zeolite materials) were able to demonstrate three types of Al at the surface of dealuminated mordenites and beta zeolites, namely hexa-, tetra- and tricoordinated. The application of XPS to the characterization of supported clusters and cluster-derived species has been reviewed by Guczi [48]. Small clusters present in general a binding energy that is a function of the cluster nuclearity. Due to the molecular character of the clusters, the core levels present higher binding energies also because the extra-atomic relaxation energy should be reduced compared with the metal bulk [48]. The valence band is also modified and the density of states is very low at the Fermi level. Moretti and Porta [23] demonstrated the sensitivity of the Auger parameter approach to the number, size and polarizability of small Cu clusters entrapped in A zeolites. The effects of the low nuclearity on the core-level binding energy and on the density of states at the Fermi level were confirmed. Similar results were obtained in the case of Pd clusters entrapped in Y zeolites [24]. Initial and final state effects on metal cluster–substrate interactions, as determined by the use of the Auger parameter and Wagner plots, were also reported for copper clusters of different size deposited on ordered alumina ultra-thin films [49] and on Dow Cyclotene and highly oriented pyrolytic graphite [50]. The application of the Auger parameter and Wagner plots to the characterization of oxide/oxide interfaces was reviewed by Gonz`alez-Elipe and Yubero [51]. The electronic and bonding interactions developed at the interface between the oxidic substrate and deposited oxidic phase determine the observed variations of the initial and final state parameters as a function of surface coverage.

It is evident that the use of X-ray excitation to study Auger and photoelectron spectra simultaneously is a rewarding technique since more information can be obtained by studying them together than studying either separately. Future developments will involve high-energy, high-resolution Auger and photoelectron spectroscopy using monochromatic radiation of Cr (Kβ), Cu (Kα) anodes. These sources will permit the study of intense core–core–core Auger transitions for most of the transition metals in the Periodic Table. Then an analysis of the Auger parameters and Wagner plots will add valuable information to characterize heterogeneous catalysts and the surface chemistry of materials in general. A first step in this direction was made by Beamson and co-workers [52]. The use of a high-energy X-ray source, such as monochromated Cu Kα radiation, provides direct access to ‘‘deep’’ core levels, the binding energy of which can therefore be related to the respective core–core–core Auger electron emission and offers also the advantage of a much greater sampling depth than that available in commercial XPS spectrometers. References 1. (a) N. H. Turner, J. A. Schreifels, Anal. Chem. 2000, 72, 99R; (b) D. M. Hercules, J. Chem. Educ. 2004, 81, 1751; (c) J. E. Castle, C. J. Powell, Surf. Interface Anal. 2004, 36, 225; (d) A. M. Venezia, Catal. Today 2003, 77, 359; (e) K. Siegbahn, J. Electron Spectrosc. Relat. Phenom. 2004, 137–140, 3. 2. W. N. Delgass, T. R. Hughes, C. S. Fadley, Catal. Rev. 1970, 4, 179. 3. C. S. Fadley, in Electron Spectroscopy: Theory, Techniques and Applications, C. R. Brundle, A. D. Baker (Eds.), Vol. 2, Academic Press, New York, 1978, pp. 1–156. 4. W. F. Egelhoff Jr., Surf. Sci. Rep. 1987, 6, 253. 5. R. Z. Bachrach, in Synchrotron Radiation Research, R. Z. Bachrach (Ed.), Vol. 1, Plenum Press, New York, 1992, pp. 1–59. 6. C. J. Powell, A. Jablonski, I. S. Tilinin, S. Tanuma, D. R. Penn, J. Electron Spectrosc. Relat. Phenom. 1999, 98–99, 1. 7. U. Gelius, B. Wannberg, P. Baltzer, H. Fellner-Feldegg, J. Larsson, P. M¨unger, G. Vegerfors, J. Electron Spectrosc. Relat. Phenom. 1990, 52, 747. 8. (a) T. L. Barr, in Practical Surface Analysis by Auger and X-ray Photoelectron Spectroscopy, D. Briggs, M. P. Seah (Eds.), Wiley, Chichester, 1983, pp. 283–358; (b) T. L. Barr, in Practical Surface Analysis, 2nd Ed., D. Briggs, M. P. Seah (Eds.), Vol. 1, Wiley, Chichester, 1990, pp. 357–436. 9. C. J. Powell, J. Chem. Edub. 2004, 81, 1734. 10. M. A. Kelly, in Surface Analysis by Auger and X-Ray Photoelectron Spectroscopy, D. Briggs, J. T. Grant (Eds.), IM Publications, Chichester, 2003, pp. 191–210. 11. H. Peisert, T. Chass´e, P. Streubel, A. Meisel, R. Szargan, J. Electron Spectrosc. Relat. Phenom. 1994, 68, 321. 12. D. A. C. Gregory, A. D. Laine, P. S. Fowles, A. Takahashi, P. Weightman, J. Phys.: Condens. Matter 1993, 5, 3843. 13. J. Q. Broughton, P. S. Bagus, J. Electron Spectrosc. Relat. Phenom. 1980, 20, 127.

3.2.3 Valence States 14. (a) J. Q. Broughton, P. S. Bagus, J. Electron Spectrosc. Relat. Phenom. 1980, 20, 261; (b) J. Q. Broughton, P. S. Bagus, J. Electron Spectrosc. Relat. Phenom. 1980, 21, 283. 15. G. Moretti, J. Electron Spectrosc. Relat. Phenom. 1998, 95, 95. 16. (a) P. Weightman, J. Electron Spectrosc. Relat. Phenom. 1994, 68, 127; (b) D. E. Ramaker, in Surface Analysis by Auger and X-Ray Photoelectron Spectroscopy, D. Briggs, J. T. Grant (Eds.), IM Publications, Chichester, 2003, pp. 465–500; (c) P. Weightman, in Surface Analysis by Auger and X-Ray Photoelectron Spectroscopy, D. Briggs, J. T. Grant (Eds.), IM Publications, Chichester, 2003, pp. 797–814; 17. J. C. Fuggle, in Electron Spectroscopy: Theory, Techniques and Applications, C. R. Brundle, A. D. Baker (Eds.), Vol. 4, Academic Press, New York, 1981, pp. 85–152. 18. G. Moretti, J. Electron Spectrosc. Relat. Phenom. 1992, 58, 105. 19. C. D. Wagner, A. Joshi, J. Electron Spectrosc. Relat. Phenom. 1988, 47, 283. 20. F. Filippone, G. Moretti, Appl. Surf. Sci. 1998, 135, 150. 21. G. Moretti, in Surface Analysis by Auger and X-Ray Photoelectron Spectroscopy, D. Briggs, J. T. Grant (Eds.), IM Publications, Chichester, 2003, pp. 501–530. 22. G. Moretti, Surf. Interface Anal. 1991, 17, 352. 23. G. Moretti, P. Porta, Surf. Science 1993, 287/288, 1076. 24. G. Moretti, A. Yu. Stakheev, W. M. H. Sachtler, J. Electron Spectrosc. Relat. Phenom. 1992, 58, R1. 25. G. Moretti, F. Filippone, M. Satta, Surf. Interface Anal. 2001, 31, 249. 26. J. E. Castle, A. M. Salvi, F. Decker, G. Moretti, Surf. Interface Anal. 2002, 33, 533. 27. G. Moretti, A. M. Salvi, M. R. Guascito, F. Langerame, Surf. Interface Anal. 2004, 36, 1402. 28. G. Moretti, P. Porta, Surf. Interface Anal. 1990, 15, 47. 29. C. D. Wagner, L. H. Gale, R. H. Raymond, Anal. Chem. 1979, 51, 466. 30. (a) C. D. Wagner, in Practical Surface Analysis by Auger and X-Ray Photoelectron Spectroscopy, D. Briggs, M. P. Seah (Eds.), Wiley, Chichester, 1983, pp. 477–509 (b) C. D. Wagner, in Practical Surface Analysis by Auger and X-Ray Photoelectron Spectroscopy, D. Briggs, M. P. Seah (Eds.), Wiley, Chichester, 1983, pp. 283–358; (c) Practical Surface Analysi, 2nd Ed., D. Briggs, M. P. Seah (Eds.), Vol. 1, Wiley, Chichester, 1990, pp. 595–634. 31. NIST XPS Database, http://www.nist.gov/srd. 32. M. O’Keeffe, J. Solid State Chem. 1990, 82, 108. 33. W. L. Jolly, in Electron Spectroscopy: Theory, Techniques and Applications, C. R. Brundle, A. D. Baker (Eds.), Vol. 1, Academic Press, New York, 1977, pp. 119–149. 34. B. Johansson, N. M˚artensson, Phys. Rev. B 1980, 21, 4427. 35. A. Flodstr¨om, R. Nyholm, B. Johansson, in Synchrotron Radiation Research, R. Z. Bachrach (Ed.), Vol. 1, Plenum Press, New York, 1992, pp. 199–251. 36. D. Tomanek, P. A. Dowben, M. Grunze, Surf. Sci. 1983, 126, 112. 37. C. Defosse, in Characterization of Heterogeneous Catalysts, F. Delannay (Ed.), Marcel Dekker, New York, 1984, pp. 225–298. 38. W. N. Delgass, in Spectroscopy in Heterogeneous Catalysis, W. N. Delgass, G. L. Haller, R. Kellerman, J. H. Lunsford (Eds.), Academic Press, New York, 1979, pp. 267–322. 39. M. Yamada, J. Yasumaru, M. Houalla, D. M. Hercules, J. Phys. Chem. 1991, 95, 7037. 40. J. Yasumaru, M. Yumada, M. Houalla, D. M. Hercules, in New Frontiers in Catalysis, Proceedinds of the 10th International

41. 42. 43. 44. 45. 46. 47. 48.

49. 50. 51.

52.

1039

Congress on Catalysis, L. Guczi, F. Solymosi, P. T´et´enyi (Eds.), ` Budapest, 1993, pp. 1867–1870. Vol. 2, Akad´emiai Kiado, L. E. Briand, O. P. Tkachenko, M. Guraya, I. E. Wachs, W. Grunert, Surf. Int. Anal. 2004, 36, 238. A. Cimino, D. Gazzoli, M. Valigi, J. Electron Spectrosc. Relat. Phenom. 1999, 104, 1. Y. Okamoto, M. Ogawa, A. Maezawa, T. Imanaka, J. Catal. 1988, 112, 427. I. Jirka, J. Phys. Chem. B 1997, 101, 8113. M. J. Remy, M. J. Genet, P. P. Nott´e, P. F. Lardonois, G. Poncelet, Microporous Mater. 1993, 2, 7. M. J. Remy, M. J. Genet, G. Poncelet, P. F. Lardonois, P. P. Nott´e, J. Phys. Chem. 1992, 96, 2614. F. Collignon, P. A. Jacobs, P. Grobet, G. Poncelet, J. Phys. Chem. B 2001, 105, 6812. L. Guczi, in Metal Clusters in Catalysis, B. C. Gates, L. Guczi, H. Kn¨ozinger (Eds.), Studies in Surface Science and Catalysis, Vol. 29, Elsevier, Amsterdam, 1986, pp. 209–219. Wu, E. Garfunkel, T. E. Madey, J. Vac. Sci. Technol. A 1996, 14, 1662. D.-Q. Yang, E. Sacher, Appl. Surf. Sci. 2002, 195, 187. A. R. Gonz`alez-Elipe, F. Yubero, in Handbook of Surface and Interfaces of Materials, H. S. Nalwa (Ed.), Vol. 2, Academic Press, New York, 2001, p. 147. (a) G. Beamson, S. R. Haines, N. Moslemzadeh, P. Tsakiropoulos, P. Weightman, J. F. Watts, Surf. Interface Anal. 2004, 36, 275; (b) G. Beamson, S. R. Haines, N. Moslemzadeh, P. Tsakiropoulos, J. F. Watts, P. Weightman, K. Williams, J. Electron Spectrosc. Relat. Phenom. 2005, 142, 151.

3.2.3.2

UV–Vis–NIR and EPR Spectroscopies

Zbigniew Sojka∗ , Fran¸cois Bozon-Verduraz, and Michel Che

3.2.3.2.1 Introduction Ultraviolet–visible–near-infrared (UV–Vis–NIR) and electron paramagnetic resonance (EPR) spectroscopies are techniques which both involve transitions related to electrons. However, a major difference is that for UV–Vis–NIR, allowed transitions correspond to the promotion of electrons from a ground term to excited terms without a change of the electron spin, whereas for EPR, allowed transitions occur with a change of the electron spin. Another difference is that the former applies to most systems, whereas the latter is more selective since it concerns paramagnetic species only, i.e. those which contain one or more unpaired electrons. Let us now examine the case of a d1 ion, in order to obtain some insight into the energy involved in such transitions. Figure 1 gives the splitting of the 2 D term of such a free d1 ion upon forming an ML6 complex (Oh symmetry), upon losing a ligand to give a surface References see page 1063 ∗ Corresponding author.

1040

3.2 Chemical Properties

A1 (z2) Eg B1 (x2–y2) 2D

T2g

E (yz, zx)

ms = +1/2 ms = −1/2

B2 (xy)

B Free ion

Complex ion (Oh)

(a)

(b)

Complex ion

External

(C4v) Magnetic field (Compressed square pyramid) (c)

(d)

Survey of the energy levels involved in UV–Vis–NIR (DRS) and EPR spectroscopies for a d1 ion: (a) as a free ion (only the fundamental 2 D term is given); (b) engaged in an ML6 octahedral complex (Oh symmetry); (c) engaged in a surface ML4 L square-pyramidal complex (C4v symmetry); and (d) subjected to an external magnetic field. For the sake of clarity, the spin multiplicity 2S + 1 = 2 given for the 2 D term has been omitted in the symbols in (b) and (c). The vertical arrows represent the spin-allowed electronic transitions in (b) and (c) and the EPR transition in (d). Fig. 1

ML4 L species (C4v symmetry) and finally subjected to an external magnetic field. Figure 1 shows that UV–Vis–NIR, which deals with electronic transitions from the ground term to excited term(s), will detect for the d1 system only one electronic transition (2 T2g → 2 Eg ) for the octahedral complex and three (2 B2g → 2 Eg ; 2 B2g → 2 B1g and 2 B2g → 2 A1g ) for the tetragonal complex. Such d–d transitions, of the order of 104 cm−1 , are usually found in the visible or near-infrared region of the spectrum (Fig. 2). As seen in Fig. 1, an external magnetic field lifts the degeneracy of magnetically degenerate energy levels: this is the Zeeman effect. Upon application of a microwave radiation, EPR spectroscopy allows one to detect transitions, of the order of 1 cm−1 , between magnetically non-degenerate levels. Thus the transitions are far more energetic, by about four orders of magnitude, in UV–Vis–NIR than in EPR. In what follows, the principles and applications of these spectroscopies to the field of heterogeneous catalysis will be reviewed. Transition metal ions (TMI) will be primarily concerned in the examples given hereafter. In fact, they can act both as catalytic sites and as probes of their own interactions with an oxide support, not only during the course of catalyst preparation but also upon further thermal treatment, adsorption and catalytic reaction. Any change in their coordination sphere may affect their optical or magnetic properties and may therefore be followed by spectroscopies, such as UV–Vis–NIR and

EPR, respectively, within the conceptual frameworks of interfacial coordination chemistry [1] and surface organometallic chemistry [2]. Since the last edition of this Handbook, the major achievement has been to apply these spectroscopies to catalysts under their working conditions (vide infra) and to couple them with other techniques/spectroscopies. UV–Vis–NIR Spectroscopy UV–visible and near-infrared specA Introduction troscopies allow the study of electronic transitions between orbitals or bands in the case of atoms, ions or molecules in gaseous, liquid or solid state. In the NIR region there appear also overtones and combinations of stretching and bending vibrations. It is highly recommended to cover both the UV ( 190 nm)–visible (400–800 nm) and NIR (800–2500 nm) ranges because, along with the ligand vibrational fingerprints, NIR may contain a lot of low-energy d–d transitions together with several f–f transitions and metal–metal charge transfers (Fig. 2). Investigations in the field of heterogeneous catalysis concern mainly catalyst preparation and adsorption: (i) the study of precursor solutions, (ii) their interaction with the support, (iii) the chemical changes undergone by precursors upon the various pretreatments (calcination, reduction) leading to the active phase (e.g. oxide, sulfide, metal), (iv) the modification of the latter upon contact with reactants, promoters or poisons and (v) the nature of adsorbed species (intermediate, inert or poison). The samples are mostly examined by transmission spectroscopy (TS) or diffuse reflectance spectroscopy (DRS); specular reflectance (SR) has been little used. TS is employed in the case of solutions, thin films and crystals whereas powders, gels and turbid media are examined by DRS. Several reviews have been devoted to the applications of UV–Vis–NIR spectroscopy to catalysis [3–7]. The aim of this part is to focus on (i) the preparation and characterization of catalysts and (ii) their interaction with probe molecules and reactants.

3.2.3.2.2

B

General Background

a Electronic Transitions The electronic transitions are of two types and involve orbitals or levels localized either on the same metal (M) atom or on two adjacent atoms, respectively [8]. The first class includes the following metal-centered (MC) transitions:

• d → d and (n − 1) d → ns in transition elements • f → f and 4f → 5d in rare earth elements • ns → np in the main group elements.

3.2.3 Valence States

Ultraviolet 200 400 n (cm−1)

50,000

E (eV)

6

Visible

Infrared

800

25,000

2,500

12,500

3

LMCT and MLCT

Near

d

l (nm)

4,000

1.5

MMCT (intervalence)

1041

0.5 Vibration overtones and combinations

d transitions

The ultraviolet (UV), visible (Vis) and near-infrared (NIR) ranges and the domains where the different types of transitions are generally observed (see text).

Fig. 2

The second class involves charge-transfer (CT) transitions (Fig. 2) from an occupied level centered on a donor atom to a vacant level centered on an acceptor. This class includes: • ligand-to-metal (LMCT) and metal-to-ligand (MLCT) charge transfers [9]; • metal-to-metal charge transfers (MMCT), sometimes called intervalence transitions, e.g. Fe2+ → Fe3+ in Fe3 O4 or Fe2+ → Ti4+ in sapphire [8]; • transitions between molecular orbitals (n → π ∗ , π → π ∗ ) in inorganic or organic molecules or ions either free or coordinated to a metal, named intraligand charge transfers or ligand-centered (LC) in the latter case; Transitions of the two classes appear in molecular complexes, which will be examined in more detail below. Most of them are also found in non-molecular solids. b Molecular Complexes In heterogeneous catalysis, molecular complexes are used either in solution (as precursors) or on supports (as catalytically active phases); those involving TMIs are the most important class. Metal-Centered (MC) Transitions (i) For d → d transitions, their number depends on the electronic configuration and on the symmetry of the complex. Let us consider an octahedral complex with a central d1 ion (Fig. 1b). If this complex is exposed to radiation with photons of energy hν = E = E(Eg ) − E(T2g ), an energy absorption occurs. The electron is promoted from the ground term T2g to the excited term Eg and only one transition, referred hereafter to as d–d transition, is observed. As in most spectroscopies, this energy absorption is detected, amplified and plotted as a function of wavelength λ (nm), wavenumber ν˜ (cm−1 ) or, more rarely, energy E (eV) of the photons.

For complexes containing more than one d electron on the central ion, the situation is more complicated because

of repulsive interactions between the electrons and a polyelectronic approach has to be used. For an octahedral complex involving a d2 ion, several configurations can be 2 , t 1 e1 , e2 ). To each of those are associated obtained (t2g 2g g g states or terms (both words can be found in the literature, but the latter is more often used) with given energies and energy absorptions can occur between the ground term and excited terms leading to several transitions, as indicated in Fig. 3. For reasons too involved to be presented here, high-spin octahedral complexes form two categories: those with configurations d1 , d4 , d6 and d9 , on the one hand, and those with configurations d2 , d3 , d7 and d8 , while configuration d5 is a separate case (see below). d → d transitions are of moderate or weak intensity because they are forbidden by the Laporte (orbital) selection rule (allowed transitions obey l = ±1). In addition to the Laporte rule, d–d transitions are submitted to the spin selection rule, which requires that the spin multiplicity of the levels involved be the same (S = 0). These selection rules may be partially or even totally relaxed by various mechanisms such as vibronic coupling, spin–orbit coupling or exchange interaction in the case of polymetallic systems [10, 11]. Figure 3 presents the spin-allowed d–d transitions in the case of octahedral complexes. It shows that for d1 , d4 , d6 and d9 high-spin configurations, one transition is expected; however, the Jahn–Teller effect [10, 12, 13] induces distortions either in the ground term or in the excited term, which turns the octahedral symmetry into square planar and leads to an asymmetric band: Cr2+ (high-spin), Cu2+ and Ti3+ are well-known examples. Figure 3a shows further that for d2 , d3 , d7 and d8 high-spin configurations, three transitions are expected. For a high-spin octahedral complex with d5 configuration, there is no spin-allowed transition possible without changing the value of S. The case of low-spin configurations (Fig. 3b) is rather complicated because the expected transitions are overlapping and/or obscured by CT transitions. References see page 1063

1042

3.2 Chemical Properties

d1

d2

d3

d4

d5

d6

d7

d8

d9

(1)

(3)

(3)

(1)

(0)

(1)

(3)

(3)

(1)

eg Number of transitions (a) High spin

t2g

Excited terms Ground term Number of transitions (b) Low spin

2E g 2

T2g

3

A2g

4

4

3T 1g

4

4

3T 1g

3

4

4T 2g

3

4

3

T2g

3

T1g

A2g

T1g

T1g

T1g T2g

4A 2g

5

5

T2g

5E g

Eg

6

A1g

T2g

T1g

T1g

T2g

A2g

2

T2g

2

Eg

eg (∗)

(∗)

(2)

(∗)

t2g 1

T2g

Excited terms Ground term

5

3

1

T1g

3T 1g

2T 2g

1A 1g

2E g

Spin-allowed d–d transitions in octahedral complexes with various dn configurations: (a) high-spin; (b) low-spin. This diagram may be applied to tetrahedral complexes by removing the g subscript (no inversion center) and by inverting the order of the spectroscopic terms; e.g. for a d1 tetrahedral complex, the transition is 2 E2 → 2 T2 . (∗) these low-spin cases are rather complicated because the expected transitions are overlapping and/or obscured by CT transitions. Fig. 3

d–d transitions give information on oxidation states and symmetry of the TMI and allow the determination of the crystal field (CF) parameter  (or 10Dq). This parameter increases with the oxidation state of the metal ion and depends also on the symmetry of the environment and on the nature of the ligands; common ligands have been ordered by increasing values of  in the spectrochemical series [9, 10, 12, 13]. Further, the covalent character is reflected by the value of β = B/B0 , where B0 and B are the Racah parameters for the free cation and the complexed cation respectively. B0 is available in the literature and B is determined, along with Dq, from Tanabe–Sugano diagrams [10, 12, 13]. Since the complexation lessens the repulsion between d electrons (because of d orbital expansion towards the ligands), β is always 3 eV, colored for 1.5 eV < Eg < 3 eV and brown–black when Eg < 1.5 eV. The filled valence band may often be identified to oxygen 2p levels and the empty conduction band with metal (n − 1)d or ns orbitals and the interband transition is analogous to an LMCT. For d1 to d9 oxides, d–d transitions have been observed. Figure 4 presents a correlation diagram between a molecular complex (MA6 ) and a non-molecular compound (MA) built from Mz+ and Aa− ions. It must be noted that, in the latter case, any departure from stoichiometry has a strong influence on the spectrum and the color of the compound, e.g. NiO, V2 O5 . The influence of defects on the optical spectrum depends on their nature and concentration and on the local symmetry [17, 18]. On the other hand, insulating oxides often used as catalyst supports (SiO2 , Al2 O3 , MgO) are white (Eg > 5 eV) and the spectra of supported TMI are easily obtained. The nature of the information given by DRS depends on the surface-to-volume ratio of the sample. For solids Molecular complex MA6

Antibonding

t1u a1g eg t2g p

Bonding

with a low surface area, the main information concerns bulk properties; for larger surface areas (say >10 m2 g−1 ), the contribution of surface species becomes significant, which can be checked through adsorption of a molecular probe. Information concerning active phase–support interactions may be obtained, especially at low loadings (e.g. a TMI may be isolated, engaged in a solid solution or in a definite compound). In addition, surface ions present coordinative unsaturation (cus) and show distortions from high symmetries (octahedral, tetrahedral). This cus favors electron transfers between surface sites and molecular probes, which allow the characterization of acid–base and redox properties. Octahedral, tetrahedral and, to a lesser extent, square-planar symmetries have been well studied and energy level diagrams are available in several books [10, 12, 13, 19, 20]. Theoretical models concerning lower symmetries (C3v , D3h ), much less investigated, may be found in Refs. [19–23]. d Vibrational Transitions Apart from low-energy electronic transitions, the NIR range contains overtones and combinations of stretching and bending vibrations. It is convenient to record on the same spectrum the vibrational fingerprint of organic or inorganic ligands, in particular adsorbed species and the electronic transitions associated with these ligands (intra-ligand, CT and M-centered transitions). In NIR–DRS, surface groups such as OH, SH and NH are easily detected in addition to the overtones and combinations of vibration of ligands employed in catalyst preparation (NH3 , acetylacetonate, etc.). A detailed examination of OH overtones and combinations in silica can be found in Ref. [24]. References see page 1063

Free ions MZ+ and Aa−

np

(M z +)

ns

(M z +)

Non molecular compound Conduction band eg t2g d orbitals

(n −1) d (M z +) p (A a − ) s (A a − )

Valence band

s

Decreasing M − A length

1043

Decreasing M − A length

Correlation diagram between the energy levels of a molecular octahedral complex and a parent non molecular compound, e.g. Ni(H2 O)2+ 6 and NiO (composed of NiO6 units) (after Ref. [16]).

Fig. 4

1044

3.2 Chemical Properties

C Diffuse Reflectance Spectroscopy (DRS) a Theoretical Background For small particles which are isotropic and not in contact with others, light undergoes single scattering and the Mie theory applies [4, 5, 25]. Although important in colloidal media, this theory cannot be used for powders, where multiple scattering occurs. In this case, one has to use a phenomenological theory with separate absorbance and scattering constants. The Schuster–Kubelka–Munk (SKM) model [3–7, 25] is widely accepted and allows one to obtain quantitatively the absorption spectrum of a solid sample from diffuse reflectance measurements provided that some experimental conditions be fulfilled. The reflectance of a solid comprises the specular (or regular or mirror-like) reflectance and the diffuse reflectance, whose angular distribution is independent of the incidence angle, the latter being predominant on matte surfaces. We shall only recall here the main results of SKM theory: details are available elsewhere [3–7, 25]. Calling R and R0 the total reflectances of the sample and reference, respectively, collected by an integrating sphere, the detector of a double beam spectrometer gives the apparent absorbance:
 R0 A = log (1) R

which is generally not proportional to the concentration of the absorbing entity. According to the SKM approximation, the diffuse reflectance of a layer of infinite thickness R∞ is linked to the absorption coefficient K and the diffusion coefficient S by the expression F (R∞ ) =

(1 − R∞ )2 K(λ) = 2R∞ S

(2)

where F (R∞ ) is the remission or SKM function. The log form is sometimes useful: log F (R∞ ) = log K(λ) − log S

(3)

It is relevant to observe that double-beam spectrometers =R do not give R∞ but the ratio: R∞ ∞(sample) /R∞(ref .) ; hence R∞ and R∞(sample) are identical only if R∞(ref .) = 1, which is not fulfilled in the whole spectral range. The ‘‘infinite thickness’’ is generally obtained with a layer depth of 1–2 mm; however strongly diffusing powders (such as some silicas) may require up to 5 mm. Equations (2) and (3) show clearly that the remission function is proportional to K only if S is independent of λ. The variation of S with particle size is discussed in Ref. [25]. This phenomenon is significant only for λ < 300 nm (e.g. in some silicas). The occurrence of specular reflectance is another difficulty often met, particularly with samples showing high K values and large particle size. It results in the ‘‘compression’’ of

absorbance maxima. The specular contribution may be decreased significantly either by grinding the sample or diluting it in a white standard. In this case, it is recommended to take the white standard (diluent) as a reference [25]. b Experimental Considerations In diffuse reflectance measurements, the light re-emitted by the sample and the reference is collected by an integration sphere coated with a white standard showing a high diffuse reflectance (generally also used as a reference sample). Magnesium oxide and barium sulfate have been used for a long time but are now replaced by polytetrafluoroethylene (PTFE), which appears to be more stable. It is essential to control the quality of the standards as they are subjected to aging. It is also possible to use a diffuse reflectance attachment, which utilizes hemispherical mirrors to collect the light reflected from the sample [26]; although the light losses are important, this accessory offers significant advantages: it requires much smaller quantities of sample, covers a larger wavelength range and allows horizontal sample mounting. Most spectra are obtained at room temperature after pretreatment of the sample in controlled atmosphere (10−5 –1 atm) at various temperatures (up to 1400 K). However, it is possible to record them from about 100 up to 800 K with appropriate cells [3, 5, 26]. The use of optical fibers allows measurements in more severe conditions but incident light appears to suffer a strong attenuation. The windows of the cells are often made of optical quality OH-free silica as this material may be soldered on quartz tubes; a 2-mm thickness is appropriate. In some cases, CaF2 or NaCl windows may also be used. Very often special cells are designed for recording both DR and EPR spectra, as shown in the EPR section (Section 3.2.3.2.3). There are four important parameters attached to each UV–Vis–NIR spectrum, which records the molar extinction coefficient ε as a function of the wavenumber (cm−1 )/wavelength (nm) of photons sent onto the sample: (i) the number of bands (as shown, for instance, in Fig. 3 for dn configurations of TMIs); (ii) the position of the band indicating the energy of the corresponding electronic transition; (iii) the intensity of the band given by the value of ε; and (iv) the width of the band. Typical UV–Vis–NIR DR spectra are shown in Fig. 5 for the Ni(NH3 )2+ 6 complex after deposition on silica from its aqueous solution and further thermal treatments. D

Catalyst Preparation

a Control of Precursor Solutions Water is transparent in the 200–1300 nm range (and it is even possible to work up to 2500 nm when the sample thickness is

3.2.3 Valence States

40 30

20

15

10

n × 10−3 (cm−1)

5

SiO−H n2

n″′3 n″3 n′3 0.1

n1

0.1

SiO−H

e e NH3

d

OH

c b a 2

3

4

5

6

7

8

d c a b

NH3

H2O

9 11 13 15 17 19 21 23

l × 10−2 (nm) UV–Vis–NIR DR spectra of an Ni(NH3 )2+ 6 sample after deposition on silica from its aqueous solution and further thermal treatments: (a) exchange and drying at 353 K; (b) heating in oxygen at 773 K; (c) and (d) outgassing at 773 K for 1 and 15 h, respectively; (e) outgassing at 973 K for 15 h (after Ref. [30]).

Fig. 5

limited to 0.1 mm), which favors the examination of the predominant species of the precursor in aqueous solutions. This is of primary importance as e.g. TMIs may suffer hydrolysis, aquation and polymerization reactions [27], disturbing the preparation scheme and leading to the active phase, and the use of predominance diagrams [28] is highly recommended. b Deposition of the Precursor, Drying, Calcination Upon contact of an oxide support with a solution of the metal precursor, several reactions may occur that disturb the expected ionic exchange: (i) partial dissolution of the support in specific pH ranges, (ii) formation of neutral TMI complexes in the precursor solution through aquation processes, (iii) genesis of polynuclear species upon OH or O bonding and (iv) deposition–precipitation. These observations are of primary importance in selecting the TMI complexes and the conditions of interaction (pH, concentration). 2− 2+ Examples Ni(NH3 )2+ 6 , Pd(NH3 )4 and PdCl4 are well identified complexes in solution and in the solid state. They may be electrostatically adsorbed without exchange of ligand or with ligand exchange, the support acting as a mono- or polydentate ligand [29]; in fact, the species (MO)2 Ni(NH3 )4 , (MO)2 Pd(NH3 )2 and (MO)2 PdCl2 may appear, where M is the metal ion of the support. If this

1045

model applies, the ligand field strength may be estimated by using the average ligand field rule. As oxygen ions from the support are considered as weak π donor ligands, the ligation of the support to the transition metal ion (TMI) should lead to: (i) a red shift of the d–d bands in the ammonia complexes as NH3 is a strong σ donor ligand and (ii) a blue shift of the d–d bands in the chloro complexes as Cl− is a stronger π donor coordinate than oxygen ions. Unfortunately, aquation–anation equilibria may also be involved in the course of the ion-exchange procedure or during the washing or drying stage. Hence it is essential (i) to analyze the precursor solutions before, during and after immersion of the support and (ii) to examine the solid after filtration and after drying. Nickel Catalysts In octahedral symmetry, the spectrum of Ni2+ (d8 , the free ion has the fundamental term 3F and first excited term 3P ) shows three spin-allowed d–d transitions in the NIR and visible ranges, generally separated from CT transitions (Table 1). There is no Jahn–Teller effect. In tetrahedral symmetry, these d–d transitions are shifted towards the IR but are more intense because the Laporte rule is relaxed. In addition, spin-forbidden transitions appear, as spin–orbit coupling causes partial relaxation of the spin selection rule. The preparation of silica-supported nickel catalysts has been studied in from the complex Ni(NH3 )2+ 6 detail [30–32]. The latter shows three spin-allowed d–d transitions (Table 1). When the silica support enters the coordination sphere of Ni2+ , the d–d transitions are red shifted (Fig. 5, spectrum a) suggesting that the supermolecular SiO− surface ligand is weaker than H2 O and, a fortiori, than NH3 , i.e. with Ni2+ , the following spectrochemical series prevails: SiO− < H2 O < NH3 . NIR examination shows that NH3 ligands are still present after drying at 353 K. The spectral modifications observed upon various pretreatments (spectra b–e) reflect coordination and symmetry changes undergone by Ni2+ : coordination changes from 6 (Oh or pseudo-Oh ) to 4 Spin-allowed d–d transitions in some model Ni complexes [10]

Tab. 1

Complex

˜ν1 /cm−1

Ni(H2 O)2+ 6 Ni(NH3 )2+ 6 NiCl4− 6 NiCl2− 4 a Estimated

8500 10750 7700 (3500)a value.

References see page 1063

˜ν2 /cm−1

˜ν3 /cm−1

13800 17500 12700 6550

25300 28200 22600 14250

Symmetry Oh Oh Oh Td

1046

3.2 Chemical Properties

Tab. 2

Spin-allowed d–d transitions in some Ni catalysts [33]

Catalyst

Ni2+ (LaX) tetrahedral Ni2+ /ZnO tetrahedral Ni2+ (NaY) octahedral Ni2+ /MgO octahedral a Splitting

˜ν1 /cm−1

˜ν2 /cm−1

˜ν3 /cm−1

Crystal field strength, 10Dq/cm−1

Racah parameter, B/cm−1

3760 4600 6100–6400 8600

8200 8340 10400–10600 13500

16100; 17500a 15260; 16180a 19000–20000 24600

4460 4500–5300 6100–6400 8600

916 780–820 700–800 820–925

of the ν˜ 3 band by spin–orbit coupling.

(Td , D4h or C4v , see bands ν1 , ν2 , ν3 , ν3 and ν3 of spectrum b) and eventually 3 (D3h or C3v , spectrum e). Moreover, through photoreduction, Ni2+ is partially transformed into tetrahedral Ni2+ , characterized by a band at 838 nm (not shown). In the case of nickel–zeolites [33], the low value of the CF strength (Dq) in octahedral Ni2+ (zeol.) is ascribed to Ni–O distances larger than in Ni2+ /MgO (Table 2). On the other hand, the high B value measured in tetrahedral Ni2+ (zeol.) give evidence of the greater ionicity of bonds compared with Ni2+ /ZnO.

1E u

1 1A 2u′ Eu

CT

CT

1

B1g

n3 1E g

Palladium Catalysts Several studies have been devoted to Pd–silica [34–38], Pd–alumina [39–41], Pd–ceria [42] and Pd–zeolites [43–45]. Compared with nickel, palladium has a stronger preference for D4h symmetry, although lower symmetries cannot be excluded in zeolites [44, 45]; d–d transitions, located in the blue and UV ranges, are less distinct from CT. This is a softer cation than Ni2+ , which accounts for the low stability of Pd(H2 O)2+ 4 . The bands observed with commonly used complexes are given in Table 3 and energy levels of Pd2+ in a D4h environment are shown in Fig. 6. The d–d transitions ν1 and ν2 are generally not resolved and the resulting rather broad band is used to estimate the ligand field strength. Calcination Upon calcination, the ligands attached to the metal are replaced by oxygen ions and the DRS Spin-allowed d–d transitions in some model Pd complexes (D4h symmetry) [9, 10]

Tab. 3

Complex

Wavenumbers ˜ν /cm−1 and wavelengths λ/nm ˜ν1 (λ1 )

˜ν2 (λ2 )

PdCl4 2−

21300 (470)

Pd(H2 O)4 2+ Pd(NH3 )4 2+

26300 (380) 33700 (297)

˜ν3 (λ3 )

˜νCT (λCT )

30000 (335)

35700 (280) 44600 (224)

n2 1

A2g

3

B1g

n1

n′1 1A 1g

Energy levels of the terms for a Pd2+ (d8 ) ion in D4h symmetry (after Ref. [10]). Fig. 6

features depend on the nuclearity of metal–oxygen ensembles, which distinguishes: (i) isolated cations (e.g. Mn+ in zeolites), (ii) metal–oxygen ‘‘clusters’’ at the nanometer scale (d < 10 nm) and (iii) bulk oxide. The size of these ensembles is believed to govern the size of the metal particles obtained from a subsequent reduction. DRS gives a significant insight into the characterization of nanoparticles of semiconducting oxides, owing to the intervention of the quantum confinement effect or quantum size effect [46]. When the particle size is less than ∼ 6–8 nm, the bandgap width increases and the absorption threshold is shifted towards higher energies (shorter wavelengths). This size effect was first observed

3.2.3 Valence States

Active Phase Characterization

4 Intraframework

3

Extraframework

1

a Sulfur–Zeolite Systems Wark et al. [49] examined the encapsulation of zinc, cadmium and lead sulfides in an NaX faujasite. These materials were prepared first by cation exchange from acetates and then sulfurization by H2 S. Table 4 shows the blue shift (increase of the bandgap Eg ) occurring in the encapsulated sulfides compared with bulk sulfides. This is another example of the quantum size effect. b Titanium-Exchanged Zeolites DRS has recently been applied to titanium silicalite [50–54]. In the silicalite matrix, titanium appears as Ti4+ ions (3d0 ), which show only LMCT in the UV range. When isolated, tetrahedral Ti4+ ions lead to a transition at ∼48 000 cm−1 (210 nm) (Fig. 7), whereas octahedral ions are expected to lead to a transition at about 42 000 cm−1 (240 nm). On the other hand, an infinite array of Ti4+ O2− ions absorb near 24 000 cm1 (410 nm) in rutile or 27 000 cm−1 (370 nm) in anatase. For nanosized TiO2 particles, these absorption thresholds are shifted towards shorter wavelengths. Moreover, new LMCT transitions are detected upon interaction of adsorbates (e.g. NH3 ) with Ti4+ ions [54]. c Copper Catalysts Copper is found as Cu0 (3d10 4s1 ), Cu+ (3d10 4s0 ) or Cu2+ (3d9 4s0 ). In contrast to most transition metals in the zerovalent state, Cu0 absorbs in the visible range (3d–4s transition) near 18 200 cm−1 (550 nm) [55]. Whereas the spectrum of Cu+ species comprises only MLCT transitions in the UV range, Cu2+ entities also show d–d transitions appearing in either the visible (‘‘octahedral’’ Cu2+ ) or the NIR region (tetrahedral Cu2+ ). It must be emphasized that Cu2+ ‘‘octahedra’’ are

Td

TS1

Anatase

2

TS2 0 −1 10

15

20

25

30

(cm−1)

35

×

40

45

50

103

DR spectra of TiO2 (anatase) and titanium silicalite (TS-1 and TS-2) samples (after Ref. [50]).

Fig. 7

strongly distorted by the Jahn–Teller effect and belong to D4h symmetry. Copper–Alumina The spectral features depend on the copper content and changes upon aging in methane–air reaction mixtures [56]. Figure 8 illustrates the influence of the copper content. At low content, the six-coordinated Cu2+ ions are isolated, as shown by an asymmetric band (Jahn–Teller effect) near 13 000 cm−1 (750 nm); at higher content, the band is shifted to 15 400 cm−1 (650 nm) and loses its asymmetry; the spectrum is then like that of a mechanical mixture of CuO and Al2 O3 , the band near 40 000 cm−1 (250 nm) being ascribed to an MLCT.

a 1.5

d

F (R∞)

E

5

F (R∞)

by transmission on glasses and on colloidal sols [46, 47]. DRS observations on solid catalysts were reported in the 1990s [48–54]; for example, according to the particle size, narrow gap semiconductors, such as PdO, are black (large size) or colored (d < 6 nm). The same considerations apply to sulfides (see the next section).

1047

1

0.5

c

Increase of band gap width of faujasite-encapsulated sulfides [49]

Tab. 4

b Compound ZnS CdS PbS

Eg /eV Bulk Encapsulated Bulk Encapsulated Bulk Encapsulated

3.60 4.40 2.42 3.46 0.41 2.59

0

λg /nm 345 282 513 359 3025 480

200

1,000

2,000

Wavelength/nm DR spectra of fresh CuO/Al2 O3 catalysts: (a) mechanical mixture; (b), (c) and (d) catalysts containing 2.1, 4.8 and 9.2% CuO, respectively (after Ref. [56]).

Fig. 8

References see page 1063

1048

3.2 Chemical Properties

The influence of aging has also been investigated. When the reactant mixture is methane-rich, the bands in the visible range disappear, which suggests that Cu2+ ions are reduced to Cu+ , probably as CuAlO2 . If the methane proportion is lower, Cu2+ ions are still present but their environment is now tetrahedral, as shown by an intense band near 6700 cm−1 (1500 nm), also observed on Cu(II) aluminate, CuAl2 O4 [57].

a

F (R∞)

Copper–Silica Such catalysts were used by Tazekawa et al. [58] in the reforming of methanol. Through contact with the reaction mixture, the CuO-rich samples are reduced to Cu, as shown by the appearance of an absorption edge near 18 000 cm−1 (560 nm); on the other hand, when the copper content is low (0.5%), Cu(II) remains predominant.

b 3

2

d 1

Copper–Thoria These catalysts, active in the dehydrogenation of 2-propanol to acetone [59], have been studied by DRS, EPR and in catalysis; evidence was given that the activity is linked with the presence of both Cu2+ and Cu+ or Cu2+ and Cu0 .

c 0

200

400

600

800

l / nm DR spectra of Cr3+ ions in: (a) bulk perovskite, LaCrO3 ; (b) bulk Cr2 O3; (c) chromium-doped alumina; and (d) MgAl2 O4 (after Ref. [61]). Fig. 9

Copper–Zeolite Systems A detailed review concerning copper–zeolite systems has been presented [60]. The O → Cu charge-transfer transitions lie in the 39 000–44 000 cm−1 range; their wavenumber decreases when the Sanderson’s electronegativity increases and increases with the global softness (ZSM5, mordenite, chabasite, zeolite A and faujasite). The EPR (vide infra) and DRS fingerprints of Cu2+ in the different sites are distinguished. d Chromium Catalysts Perovskites are often used as total oxidation catalysts and are believed to be competitive with platinum metals. In LaCrO3 , the environment of Cr3+ is still octahedral; by comparison with Cr3+ /Al2 O3 and MgCrAlO4 , the three d–d transitions are red-shifted, which indicates greater chromium–oxygen distances (Fig. 9) [61]. e Cobalt Catalysts Cobalt appears generally as Co2+ (3d7 ) or Co3+ (3d6 ). In oxides and sulfides, Co2+ is high spin and 4 F and 4 P terms are involved; octahedral and pseudo-octahedral species show three spin-allowed d–d transitions. In tetrahedral complexes, three transitions are also observed at lower energy and their intensity is enhanced because they are allowed. Due to the occurrence of spin–orbit coupling, a detailed interpretation is difficult. Generally, it is only intended to identify the oxidation state and the symmetry of cobalt. The Co3+ complexes are mostly low-spin and the lowest energy

term in Oh symmetry is 1 A1g . The most frequently observed d–d transitions are given in Table 5 for some model complexes. The spectroscopic features of cobalt in CoAPO zeolites have been described [62]. Moreover, coupling UV–Vis DRS with magnetic measurements (SQUID) has given valuable information about the nature of apatite-supported cobalt species (isolated Co2+ ions, Cox Oy clusters or Co3 O4 ) [63]. f Molybdenum Catalysts Like vanadium, molybdenum(VI) has a propensity to form polynuclear complexes, which renders the interpretation of spectra difficult. For many years, the influence of the local symmetry (generally Td or Oh ) has been considered as predominant. However, subsequent studies [64, 65] have shown the prominent roles of (i) the distant environment, (ii) the degree of Spin allowed d–d transitions of some Co model complexes [10]

Tab. 5

Complex Co(H2 O)2+ 6 CoCl4− 6 CoCl2− 4 Co(H2 O)3+ 6

˜ν1 /cm−1

˜ν2 /cm−1

˜ν3 /cm−1

Spin

Symmetry

8100 6600 3000 16500

16000 13300 4780 24700

19400 17150 14300

High High High Low

Oh Oh Td Oh

3.2.3 Valence States

condensation, (iii) the size of the counter-cation and (iv) the degree of hydration. Raman spectroscopy and EXAFS [66] have thrown some light on the nature of the species present in Mo/SiO2 and Mo/Al2 O3 catalysts. g Interaction of CO with Supported Metal Ions CO is a molecular probe often used to examine (generally by IR or EPR spectroscopy) the electronic state of supported metals or metal ions. The IR wavenumber ν˜ (CO) in Mn+ –CO species may be considered as a fingerprint of the Mn+ ion; however, this is not direct evidence and, in some cases, the Mn+ target may be reduced by the probe. As electronic (UV–Vis–NIR) spectroscopy gives direct information about the presence of Mn+ , it is relevant to use it to check whether CO behaves as an ‘‘innocent’’ probe or a reducing agent. A recent study on supported Pd2+ [42] has shown the influence of the nature of the support: whereas CeO2 -supported Pd2+ is immediately reduced upon contact with CO, the reduction is very slow and only partial when alumina is the support [39]. F Characterization of Adsorbed Species In this section, attention is focused on the spectral modifications undergone by the reactants upon contact with the solid catalysts. a

Interaction of Various Reactants with Zeolites

Aromatics It is well known that aromatics give rise to protonated species upon contact with strong Brønsted acids. Thus benzene produces cyclohexadienyl entities; the basicity of substituted benzenes was shown to increase with the degree of substitution, p-xylene being taken as a reference. When contacted with various

1049

mordenites, this class of substrates shows a band A in the 210–275 nm range, ascribed to a neutral species, and a band B (330–360 nm), attributed to cyclohexadienyl cations [67]. In contrast with m-xylene, 1,2,4- and 1,3,5trimethylbenzenes (TMBs) are protonated by both HM and dealuminated HM but they behave differently (Fig. 10). When the Al content is lowered, the number of strong acid sites decreases: in the case of 1,3,5-TMB, the intensity of B (355 nm) decreases, whereas with 1,2,4TMB, B is slightly enhanced; this suggests that 1,2,4-TMB is protonated by acid sites of medium strength, whereas 1,3,5-TMB reacts with all acid centers. Further, neither xylenes nor TMBs are able to form cyclohexadienyl species when contacted with HNaY; this led to class the zeolites by increasing Brønsted acidity. Propene Interaction of propene with Y-zeolites and mordenites was studied by transmission (Fig. 11) in both the UV–Vis and IR ranges [68]. With NaHM and HM, propene gives rise first to band I (325 nm), then to bands II (375–380 nm) and III (450 nm), ascribed to monoenyl, dienyl and trienyl carbocations, respectively. Band I is red shifted for increasing contact time, because of substitution on the allylic system. Polyenyl carbocations are considered as coke precursors. Upon contact with HZSM-5 zeolite, allyl alcohol also leads to monoenyl cations giving polyenyls and, at higher temperature, to coke precursors. Similar results were reported with acrolein and neopentane [69, 70]. Sulfur Compounds A transmission study of the interaction of SO2 and H2 S (Claus reaction) on NaX zeolite References see page 1063

0.3 HM initial

0.2

355 1.2

354 355 HM D-2

0.1

HM initial

F (R∞)

F (R∞)

352

HM D-1 0.6

HM D-2 HM D-4

HM D-4

0 (a)

200

300

400

500

l (nm)

0

600 (b)

200

300

400

500

600

l (nm)

Fig. 10 DRS spectra (SKM function) of trimethylbenzenes (TMBs) adsorbed on H-mordenite: (a) 1,2,4-TMB and (b) 1,3,5-TMB. The Al content decreases in the order HM, HMD-1, HMD-2, HMD-4 (after Ref. [67]).

1050

3.2 Chemical Properties

1.2

1

NaHM (III) 450

(II)

440 295

HM

375

3 2

325

Absorbance

Absorbance

375

375

290

3 335 (II)

2

335

0 200 (a)

300

(I)

1

(I) 400

1 0 200

500

l (nm)

300

400

500

l (nm)

(b)

Fig. 11 DRS spectra obtained through interaction of propene with (a) NaH-mordenite; (b) H-mordenite upon admission of 2660 Pa of propene (1) immediately after exposure at 300 K, (2) after 1 h at 370 K and (3) after 13 h at 400 K (after Ref. [68]).

catalysts gave evidence on the nature of intermediates and 2− 2− products (HSO− 3 , S2 O4 , S2 , S8 ) [71]. b Interaction of Pyridine with Alkaline Earth Metal Oxides After outgassing at high temperature, MgO, CaO and SrO present low-coordinated surface oxygen ions, characterized by transitions in the 250–300 nm interval, sometimes called surface excitons [72]. Pyridine adsorption induces modifications in the spectrum of the solid and of the adsorbate itself: (i) pyridine N-coordinated to cus cations gives two bands in the 200–250 nm range ascribed to π → π ∗ transitions in weakly adsorbed pyridine and (ii) between 300 and 1000 nm several bands appear pertaining to strongly adsorbed anionic derivatives arising from heterolytic dissociative chemisorption (i.e. bipyridyl entities). These anions are generated without electron transfer from the solid but their formation is easier when the basicity of the oxide concerned increases. G Conclusion UV–Vis–NIR spectroscopy brings information on all the preparation stages of catalysts; it also permits one to follow changes of the active phase upon contact with reactants. Quality control and routine examination are readily performed. Moreover, new technical facilities now encourage to perform experiments on catalysts under their working state [73, 74] (Fig. 12). However, the transitions involved show considerable intensity differences (from 1 to 106 ) and bandwidths are usually large, which make the identification of species tedious. These difficulties may be overcome by comparison with pertinent model systems. In addition, the very intense CT transitions offer particular analytical interest when the concentration of (cationic) active species is low (e.g. 1/2) represents the electronic spin-spin interaction (fine structure tensor D), the third term is the References see page 1063

1052

3.2 Chemical Properties

electron spin–nuclear spin hyperfine interaction (tensor A). The last two terms account for nuclear Zeeman and, for I > 1/2, quadrupole interactions (tensor P). When two or more paramagnetic centers interact (which is not rare for catalytic materials), the EPR spectrum is described by a total spin Hamiltonian (Htotal ), which is the sum of the individual spin Hamiltonians Hi and the interaction Hamiltonian Hij , which accounts for the isotropic exchange (Jij Si ·Sj ), antisymmetric exchange (dij Si ×Sj ) and the anisotropic dipole–dipole coupling (Si ·Dij ·Sj ), can be written as

gII

g⊥

B q

gyy Ayy

f

g Axx xx (b)

Hij = Jij Si ·Sj + dij Si ×Sj + Si ·Dij ·Sj

A gzzzz

(a)

gzz gyy gxx

(5)

where, following the convention, Jij > 0 stands for ferromagnetic and Jij < 0 stands for antiferromagnetic couplings [100, 102]. The point symmetry at the paramagnetic center determines whether any of the principal values of g or A (and other tensors) have to be equal and whether their principal axes coincide [103, 105]. Therefore, for a paramagnet of unknown structure, providing the type of EPR spectrum can be correctly identified, then valuable structural information can be obtained [105]. b Determination and Molecular Interpretation of Parameters Catalyst samples are usually powders, composed of small crystallites randomly oriented in space. Fairly accurate principal values of the g and A tensors can be extracted from a polycrystalline spectrum only in the simplest cases, when interactions are not too complex, intrinsic linewidths are not too large and axes are coincident (Fig. 14a–c) [84, 87, 106]. For more complicated spectra (Fig. 14d), the number of lines may increase considerably depending on anisotropy, number of magnetic nuclei and overall symmetry, so that the analysis of EPR spectra becomes difficult. In such cases, computer simulations combined with advanced fitting appear to be practically the only way to analyze experimental powder EPR spectra reliably [102, 107, 108]. Analysis and interpretation of the magnetic parameters in terms of molecular structure may be further supported by quantum chemical calculations [109]. This computational spectroscopy approach is nowadays becoming well established for an advanced interpretation of complex EPR spectra, often encountered in heterogeneous systems [110–113]. This approach is well illustrated by Fig. 15, with the molecular analysis of the g tensor of VO2+ /ZSM-5. The principal couplings between the relevant orbitals, induced by the magnetic field, can be readily identified and their particular contributions to the g-shifts quantified [114].

Azz

Ayy Axx (c)

gxx

gzz Azz

gzz

Ayy

a

gyy gxx Axx

(d)

Fig. 14 Typical EPR powder spectra of paramagnetic species in (a) axial and (b) orthorhombic symmetry in the absence of hyperfine structure. The case of powder EPR spectra in orthorhombic symmetry illustrated by N− 2 radical trapped on an MgO surface with a hyperfine structure due to the presence of two atoms with nuclear spin I = 1 is given in (c), and the monoclinic spectrum with non-coincident g and A axes using a surface carbonyl complex of molybdenum (I = 5/2) as an example is shown in (d) (after Ref. [95]).

c Samples and Sample Holders – Practical Considerations The best-resolved EPR spectra are obtained for magnetically diluted samples for which intermolecular magnetic interactions practically disappear. For radicals with very narrow signals, it is recommended to remove oxygen, which being paramagnetic broadens the spectra and thus decreases the resolution. Samples are placed in quartz tubes, which may be connected to vacuum or gas handling lines for sample activation, gas adsorption and/or reaction. The quartz tubes (∼3 mm in diameter for usual X band measurements with 30 mm average height of powder) are placed in the EPR cavity. Figure 16 shows a cell for recording EPR spectra in a controlled atmosphere

3.2.3 Valence States

1053

gxx ≅ gyy = g⊥ 20a1

−67210

10b2

−19175

10b1

−19671

gzz = g

A = − 4/7P

SOMO

A⊥ = 2/7P 1754

2463

2547

3a2

8b1 gxx

5400

gyy

9b2

18a1

gzz 15a1

X-band EPR spectrum of VO2+ /ZSM-5 sample along with the orbital diagram for the most important paramagnetic contributions to the g shift in spin-restricted VWN LDA DFT calculations. MOs contributing more than 10% of the total giso due to magnetic field-induced coupling are indicated by arrows along with their contributions in ppm (after Ref. [114]).

Fig. 15

LN2

(a)

(b)

(c)

Fig. 16 (a) A typical cell to record EPR spectra and (b) a special cell allowing EPR and DR spectra to be taken on the same sample; (c) a quartz Dewar used for measurements at liquid nitrogen temperature.

in static (Fig. 16a) and flow (Fig. 16b) conditions. The latter is equipped with flat quartz cell for joint UV–Vis measurements. Spectra are often obtained at 77 K using a special Dewar fitting in the cavity (Fig. 16c). Variabletemperature measurements from 4.2 or 77 to 600 K can be carried out using commercial cryostats. There are

also available high-temperature EPR cavities designed to work at temperatures as high as 1000 K. They can be adapted for in situ monitoring of processes involved in catalyst calcination, redox treatment or phase transformations [115]. Reliable determination of the absolute concentration of spins in a sample is not a routine task. Theoretical and practical problems associated with quantitative EPR measurements have been summarized in numerous review articles [116–118]. The most common method of spin dosimetry consists in comparing the EPR signal intensity [double integral of the first derivative Sx = ## S(B)dB] of a sample with unknown spin number, Nx , with that of a reference with known spin number, Nref : ref G √P gref 2 S(S + 1)ref ηref Qref Hm ref ref Nx = √ xG gx 2 S(S + 1)x ηx Qx Hm P x x
 Sx × Nref (6) Sref where η = filling factor, Q = quality factor of the loaded cavity, Hm = modulation amplitude, G = receiver gain References see page 1063

1054

3.2 Chemical Properties

and P = microwave power. The reference materials commonly used as EPR standards are listed by Chang [116].

gxx,gyy X-band

d Characteristic Features of EPR Spectroscopy Applied to Catalysis and Surface Studies Catalytic materials are usually polycrystalline, exhibit high surface areas and are composed of small, randomly oriented crystallites or even of amorphous particles (e.g. silica-supported materials). When a crystalline specimen is finely divided, additional line broadening effects can appear in the EPR spectrum due not only to the effects of particlesize distribution (usually in the range 1–50 nm) but also to the distribution of dislocation strains, surfacerelated defects and local microheterogeneity in the nearest environment of paramagnetic centers. Because of the inherent heterogeneity of catalytic materials, several paramagnetic species very often coexist, in both real and model systems [84, 119, 120]. The corresponding component EPR signals then become difficult to identify and their parameters such as the g tensor and fine (D tensor) and hyperfine (A tensor) interactions appear dispersed, giving rise to extra line broadening [95]. The use of powders thus generates a number of difficulties connected with poor resolution, overlapping signals, low symmetry phenomena and g-tensor anisotropy too small to be observed at conventional frequencies [84, 85, 88]. Complex spectra can be analyzed by different physical and chemical methods, supported by computer simulation. Commonly they rely on the use of different microwave frequencies, the study of saturation behaviors of the centers depending on their spin–lattice relaxation times, and also on their thermal stability or reactivity [84, 87, 95]. A multi-frequency approach at frequencies of 0.5–90 GHz has proven to be particularly useful for such investigations [121–124]. It not only allows one to distinguish clearly g from hyperfine features, but also helps to determine the principal values of both tensors. Figure 17 illustrates how application of X-, Q- and W-band EPR allows to separate the g and A tensor components and to resolve the weak difference between gx and gy components [122]. High-frequency resolution enhancement is possible only when the contribution of g-strain to the linewidth is small [125]. Paramagnetic TMIs with large g-tensor anisotropy and sensitivity to local environment are more likely to exhibit pronounced g-strain broadening in heterogeneous systems. The g-strain is often correlated with A-strain, reflecting local microheterogeneity of the sites. Such a situation is typically encountered in the case of hydrated transition metal ions exchanged into zeolites, leading to a pronounced mI dependence of the widths of the hyperfine lines (Fig. 18). EPR studies on copper zeolites provide a typical example [126].

n = 9.3738 GHz

gxx Q-band

W-band

gzz

gyy

n = 33.96 GHz

gzz

gxx

×5 n = 93.972 GHz

gzz gyy −50

0

50

100

150

200

250

Sweep width, ∆B/mT Fig. 17 Experimental EPR spectra of NO adsorbed on Na-A at 10 K measured at different microwave frequencies (after Ref. [122]).

(a)

(b)

(c) 10 mT

Fig. 18 X-band EPR spectrum of hydrated Cu Beta sample recorded at 120 K, showing pronounced mI dependence of the hyperfine linewidths due to the microheterogeneity of the local environment of copper ions: (a) experimental spectrum; (b) simulated spectrum without inclusion of g and A strain; (c) with correlated g- and A-strain included (after Ref. [126]).

The resolution may also be improved by using higher derivative spectra and the third derivative is usually recommended [84]. In the context of catalytic systems, this method has been successfully applied to reveal unresolved features in the EPR spectra of 13 CO adsorbed on silicasupported NiI [127], MoV [128] and VIV [129] or more recently to unravel speciation of cobalt carbonyl complexes in ZSM-5 zeolites [130].

3.2.3 Valence States

O

O 6+

Mo

O

O 6+

5+

Mo

Mo

Mo

320 340 360 380

6+

Mo

O 6+

Mo

+N

−O Rearrangement of octahedra O O

O 6+

Mo

6+

Mo

O−

O 6+

Mo

Mo

N−

−O

O

6+

5+

Mo

320 340 360 380

+ e− O

O 6+

O 6+

6+

6+

Mo

Mo

Mo

−O 320 340 360 380 O 6+

Mo

5+

Mo

O

O 5+

Mo

6+

Mo

Rearrangement of octahedra

320 340 360 380

320 340 360 380

B/ mT

B/mT

The molecular structure of various paramagnetic centers produced during mild reduction of molybdena and the corresponding EPR spectra (after Ref. [131]). Fig. 19

A chemical approach can be also taken to identify overlapping signals. It consists in designing as many separate model experiments as there are components in the complex EPR spectrum. If properly designed, each model experiment should lead to essentially one type of paramagnetic center and thus one EPR component. This approach has been used to analyze the EPR spectra of various V4+ entities present in alkali metaldoped vanadium oxide supported on ZrO2 [120] and to investigate the various paramagnetic centers produced in the course of mild reduction of molybdena [131]. In the latter case, the shapes of individual component signals along with the proposed structures of the associated defects are shown in Fig. 19. The low site symmetry resulting in non-coincident g and A axes (Fig. 14c) [103, 132] and the vanishing of strongly angle-dependent lines restrains further the amount of information which can be extracted from experimental powder spectra. Although the inherent anisotropy at the interface suggests that low symmetry

1055

can often be expected for surface species, there are relatively few well documented cases [95, 133]. They include the {MoV = O(OD)CO}/SiO2 complex obtained upon adsorption of CO on 95 MoOx /SiO2 activated by UV irradiation in D2 at 77 K [133] and the {O− –MoVI = O}/SiO2 surface complex [134]. Further notable examples are provided by bent adducts of Cs symmetry, such as η1 {CoO2 }9 /Cox Mg1−x O [135] or η1 {NaNO}1 /A [136] and η1 {CuI NO}11 /ZSM-5 and η1 {CuI NO}11 /L [137, 138]. The power of EPR to study catalytic surfaces may be extended by the use of various probe molecules [85, 139]. In EPR work [87], the latter can be classified as spin probes used to study diamagnetic centers [85, 140], coordination and space filling probes used to obtain better insight into the stereochemistry of surface paramagnetic TMIs and to help assign EPR spectra [85, 87, 133] and spin traps applied to reveal elusive paramagnetic intermediates produced in the course of catalytic reactions [141, 142]. This approach is illustrated here for silica-supported NiI3c (3d9 ) ions, catalytically active in olefin dimerization, probed with natural and 13 C-enriched CO [127]. Depending on the 12 CO pressure, a series of EPR spectra can be obtained (Fig. 20). They correspond to complexes with the following structures: NiI (Os )3 (CO) (1), NiI (Os )2 (CO)2 (2), NiI (Os )2 (CO)3 (3) and NiI (Os )(CO)4 (4), where Os represents surface oxygen ions [127]. C Selected Applications For the sake of conciseness, a limited number of typical examples will be given to illustrate the applications of EPR to catalysis and surface studies. The interested reader is referred to numerous review articles [84–88, 95, 96, 141], given whenever available for each of the topics covered. a Catalyst Preparation To rationalize the preparation of oxide materials containing well-dispersed TMIs, an interfacial coordination chemistry (ICC) approach has been developed [84]. If the preparation parameters are controlled, it has become possible to identify the major steps involved in the deposition of transition metals on the surface, such as the formation of outer- and innersphere complexes or diffusion of surface TMI into the bulk of the support upon thermal activation. The usual strategy adopted is to select an appropriate complex of TMI and investigate how the coordination sphere is modified during the deposition. The examples below show how EPR can provide information on the way TMIs bind to the support and the way to control the deposition sites at gas–solid and solid–liquid interfaces. Highly dispersed silica-supported molybdenum can be obtained by grafting, which involves the reaction References see page 1063

1056

3.2 Chemical Properties

g1 g2

g3 g1

g2 g3

g||

g⊥

g⊥ > g|| ≅ ge Ground state dz 2

(a)

g|| > g⊥ > ge Ground state dx 2−y 2

(b)

g1 g2

g3 g⊥ > g|| ≅ ge Ground state dz 2

(c)

DPPH g⊥ > g|| = ge Ground state dz 2

100 G (d) g⊥

g||

EPR spectra (X band, 77 K) of NiI (CO)n (n ≤ 4) complexes obtained upon adsorption at 293 K of 12 CO under a pressure of (a) 1.33 kPa followed by outgassing for 15 min at 340 K, (b) 1.33 kPa, (c) 13.3 kPa and (d) 53 kPa. The structure of the predominant species is given in each case (after Ref. [127]).

Fig. 20

of surface Si–OH groups with MoCl5 , under air- and water-free conditions [143]. MoCl5 is particularly attractive here since it is paramagnetic (4d1 ), monomeric in the vapor phase and gives an isotropic signal at g = 1.952. The reactor used for grafting is equipped with an EPR tube to allow the in situ monitoring of the grafting process by EPR. When dehydrated silica is contacted with gas-phase MoCl5 at 200 ◦ C, an Mo5+ EPR signal with axial symmetry (g⊥ = 1.952, g|| = 1.968, A⊥ = 4.4 mT, A|| = 8.2 mT) appears and increases in intensity with time (Fig. 21). Its hyperfine pattern is due to the interaction of the unpaired electron with 95,97 Mo nuclei (I = 5/2) of ca. 25% natural abundance. The initial isotropic EPR signal of monomeric MoCl5 is modified upon grafting, suggesting that the Mo5+ coordination sphere has been affected. The parameters of the EPR signal and the UV–Vis spectrum obtained by DRS after grafting are found to be very similar

to those of the [MoOCl4 ]− ion, suggesting the following grafting reaction: MoCl5 + ≡SiOH −−−→ ≡SiOMoCl4 + HCl

(7)

After the grafting process, hexacoordinated Mo5+ species are formed [143] and upon further treatment in vacuum at increasing temperatures [144], penta- and tetracoordinated ions with a molybdenyl Mo=O bond are also obtained (Fig. 22). The preparation of Cu–zeolites by cation exchange illustrates processes occurring at liquid–solid interface in the course of catalyst preparation, which can be followed by EPR. Copper ions (3d9 , I = 3/2) have been much studied in zeolites since they give EPR spectra with resolved hyperfine structure and g anisotropy in the isolated and dimer states. EPR and DRS parameters (g and A tensors and d–d band maxima) are site

3.2.3 Valence States

Tab. 6

g|| = 1.968

DPPH

×1 (a)

× 10

g⊥ = 1.952

EPR signatures of Cu2+ on exchange sites [119]

Site

Zeolite

I II I

A X, Y Chabasite Y X Mordenite Mordenite

III I VI

1057

g||

A|| /mT

g⊥

A⊥ /mT

2.397 2.387 2.340 2.332 2.354 2.327 2.277

12.6 12.4 16.0 15.7 14.0 15.4 16.8

2.065 2.073 2.073 2.067 2.068 2.062 2.057

0.22 1.34 2.00 1.86 1.60 1.49 1.19

Table 6 gives signatures of Cu2+ on exchange sites for various types of zeolites, each site being characterized by a given number of oxygen ions binding the Cu2+ ion [119]. b

8.2 mT 4.4 mT (b)

EPR spectra (X band, 77 K) of an Mo/SiO2 catalyst obtained after reacting silica with MoCl5 vapor at 473 K at different amplifications: (a) ×1; (b) ×10 (after Ref. [143]).

Fig. 21

5+ g⊥Mo5c

hf1

5+ g⊥Mo6c

gIIMo5+ 6c hf2

gIIMo5+ 5c

(a) 5+ g⊥Mo4c

DPPH

(b) gIIMo5+ 4c (c) 5+ g⊥Mo5c

5+ g⊥Mo6c 5+ g⊥Mo4c

10 mT

(d)

Fig. 22 EPR spectra (X band, different temperatures) of hexa-, 5+ 5+ penta- and tetracoordinated Mo5+ 6c , Mo5c and Mo4c ions obtained on a reduced grafted Mo/SiO2 catalyst. EPR spectra recorded at (a) 300, (b) 77 and (c) 4.2 K. (d) A 3dr derivative spectrum of the perpendicular region (after Ref. [128]).

dependent, so they can be used as diagnostic features while preparing well-defined samples of Cu–zeolites.

Catalyst Characterization

Catalytic materials exist in various Surface Sites forms, of which supported systems, zeolites and layer materials and also simple or mixed oxides are probably the most important [84, 97]. For EPR purposes, the surface sites which determine the reactivity of such systems can roughly be classified as magnetically diluted single (mononuclear) or cluster (polynuclear) centers and magnetically undiluted extended centers, depending on the extent of interactions with the surrounding [95]. Their EPR spectra are determined by the relative magnitudes of the electronic Zeeman, hyperfine, electron–electron spin interactions and magnetic exchange. The effect of dipolar De and exchange J interactions on the shape of the signal depends on their relative magnitude as compared with hyperfine A and electronic Zeeman splitting. Following Mabbs [105], four generic cases may be distinguished (i) |De | < |A| but J ≈ 0, (ii) |J | < |A|, (iii) |A| < |J | < hν, (iv) |J | > hν. Understanding the intimate mechanism of catalytic reactions requires the knowledge of the structure of active centers before, after and, if possible, during the reaction. Single Centers Usually catalytic systems best suited for EPR studies are those that contain defect centers, such as Lix Mg1−x O [88], UV-irradiated MgO (F+ s centers) [145], partially reduced MoO3 [131] or isolated TMIs hosted in oxide matrixes of high crystallinity, because they give better resolution [85]. Typical examples of catalytic materials that exemplify the use of EPR, along with other techniques, to follow the changes in the coordination environment and oxidation state of the transition metal ion are mixed oxides, Mnx Mg1−x O [146], Cox Zn1−x O [147] and CeO2 –ZrO2 [148], and supported References see page 1063

1058

3.2 Chemical Properties

oxide materials, MoOx /SiO2 [128], CrOx /Al2 O3 [149] and WOx /ZrO2 [150]. Other common hosts, probably best suited for EPR studies, are zeolites with their crystalline microporous 3D structures and specific sites for hosting TMIs [119, 151–153]. Porous oxide materials also include layered compounds such as clays and their pillared variants and micro- and mesoporous molecular sieves. However, the multiplicity of hosting sites leads to speciation and complicates the analysis of the EPR spectra. In favorable cases, when Al atoms are fairly well dispersed, such as in zeolites with high Si/Al ratios, attachment of TMI guest ions is manifested by the appearance of a distinctly resolved 27 Al superhyperfine structure observed in the case of VIV (d1 ) in β-zeolite [154] and for chromium(V) in ZSM-5 zeolite [155]. Dipolar interactions with 27 Al nuclei usually remain unresolved, leading to mere broadening of the EPR spectra. This phenomenon has been exploited in practice to distinguish between the copper species located in the interlamellar spaces and those bound to the walls of the Al13 pillars in intercalated saponites [156]. EPR allowed the formation of surface inner-sphere complexes upon adsorption of copper, leading to [CuII (AlO)n (H2 O)4−n ]x+ and [CuII (AlO)n (H2 O)6−n ]x+ species, to be followed. Clustered Centers Clustered centers in catalytic materials have been less investigated with EPR techniques than more common mononuclear sites. The most spectacular results have been obtained with ionic clusters of (z < n) of alkali metals [157] various nuclearity, Mz+ n and silver [158, 159], formed within cavities of molecular sieves and clays or trapped on the oxide surfaces [160]. The well-resolved EPR spectra with clear hyperfine pattern provide a solid basis for their identification. Copper-doped CeO2 is probably one of the best examples of a simple bulk oxide that gives a clearly resolved EPR spectrum of centers with S = 1 formed by two nearly equivalent Cu2+ ions separated by an oxygen vacancy [161]. After calcination of the sample in a flow of dry air at 1173 K, two signals were detected: one was due to isolated Cu ions and the other (g|| = 2.2079, g⊥ = 2.0403, A|| = 8.5 mT, A⊥ = 1.35 mT, D = 0.066 cm−1 , Ddip = −0.081 cm−1 , Dex = 0.147 cm−1 and J = −52.5 cm−1 ) was assigned to copper dimers. The inter-ion separation of 0.34 nm determined from the relative intensity between the forbidden (ms = 2) and allowed (ms = 1) lines is considerably smaller than the Ce–Ce distance (0.541 nm) in the ceria matrix, indicating that the two adjacent copper ions are not in the substitutional positions. Copper dimeric centers have also been observed in the case of CuY zeolites [162] and Cu/MCM-41 mesoporous materials [163].

Extended Centers Transition metal oxides are rarely stoichiometric and there are two extreme types of structures (typified by ZrO2 and ReO3 ) that can accommodate significant non-stoichiometry. Whereas in the former abundant defects can be formed, in the latter faults known as shear planes are easily created. Both kinds of defect play an essential role in catalytic oxidation reactions [164]. Such systems (e.g. deeply reduced oxides such as V2 O5 and MoO3 ) often give rise to magnetically undiluted samples exhibiting EPR spectra subjected to inhomogeneous broadening, caused by electron–electron dipolar and exchange narrowing interactions. Both effects depend primarily on the distance between paramagnetic centers. The shape and intensity of those spectra depend on the strength of the spin–spin interactions between coupled ions, giving rise to valuable insight into temporal changes of structural and electronic properties brought about by the catalytic reaction. A moment ratio method ( B4 / B2 2 ) has been used for this purpose and its applicability can be demonstrated with, inter alia, V2 O5 –Fe2 O3 catalysts for fluorene oxidation [121] and (VO)2 P2 O7 catalyst for butane ammoxidation [165]. Extended magnetically coupled centers have also been observed in the case of MnAPO-44 molecular sieves [166]. With increasing manganese loading the hyperfine structure completely disappears due to superexchange interactions between Mn2+ ions incorporated in the framework via Mn–O–P–O–Mn bridges (Fig. 23). From the temperature dependence of EPR spectra, the value of |J | = 2.2 cm−1 has been determined. Other magnetically undiluted catalytic systems investigated by EPR include the V2 O5 –MoO3 [167, 168], deeply reduced MoO3 [169] g′ = 2 g′ = 4.3

100 mT

Fig. 23 X-band EPR spectrum of MnAPO-44 molecular sieves. With increasing manganese loading the hyperfine structure completely disappears due to superexchange interactions between Mn2+ ions incorporated in the framework via Mn−O−P−O−Mn bridges (after Ref. [166]).

3.2.3 Valence States

and mixed MoVNbTe oxide catalyst [170], all used for the oxidation reactions. c Migration and Bulk Diffusion of Surface Transition Metal Ions Surface transition metal ions exhibit distinct mobility in hydrated and dehydrated states with preference for some oxides over others. This was recently confirmed by EPR, used in tandem with DRS, on Cr-loaded silica and alumina. After mixing of CrOx /SiO2 with alumina, paramagnetic Cr5+ and Cr6+ ions are located on the Al2 O3 support [149]. The behavior is different in the case of zeolitic systems, for which it is generally observed that TMIs migrate from accessible sites to more hidden sites upon thermal activation and conversely migrate back to more accessible sites upon adsorption of gas-phase reactants. Such ion migration has been discussed elsewhere [171]. On the other hand, during thermal activation or catalytic reactions, oxide-supported TMIs may permeate from the support surface into the bulk and become catalytically inactive. It is therefore important to know the conditions under which the active phase is lost by such migration. When TMIs are paramagnetic, it is possible to monitor by EPR the permeation process and the Tammann temperature of the support can be determined in this way [123, 172]. d Characterization of Adsorbed Species and Reaction Adsorbed species are extensively studIntermediates ied because of their role as intermediates in catalytic reactions, and this aspect of EPR applications has been reviewed elsewhere [84–86, 98, 173]. Since, in contrast to surface spectator species, intermediates are often in low concentration, the high sensitivity of this method is of great advantage. When EPR and molecular probes are employed conjointly, important information such as location and mobility of adsorbed species, surface crystal fields, surface redox properties, active site identification, surface morphology and coordination of surface TMIs can readily be elucidated [84, 85, 87]. Paramagnetic reactant molecules amenable for EPR investigations can be stabilized by adsorption on the external surface and within the pores of the catalyst. Depending on the extent of their interaction with the surface, they are considered as bound or itinerant species [98]. Bound species are rather strongly attached to specific surface sites via chemical bonds or by electrostatic forces, as in the case of charged radicals trapped on highly ionic systems. Spectroscopic investigations allowed one to distinguish between (a) ligand-centered radicals ({NO–NaI }ZSM-5 III − VI [136], {O− 2 –Co }MgO [135], {O –Mo }SiO2 [134]), (b) metal-centered radicals ({CO–NiI }SiO2 [127], {(C2 H4 )2 –PdI }X [174]), (c) mixed metal–ligand radicals

1059

({cyclohexyl–Ru0 }SiO2 [175], intrazeolitic {NH3 –Ag} NaA [174]) and (d) radical ions stabilized electrostatically − by surface counter ions (e.g. O− 2 or N2 on MgO [177]). These radicals are usually sufficiently stable for straight observation and structural characterization by EPR. At elevated temperatures, bound radicals become mobile enough to spill over on the surface, which essentially influences the catalytic reactivity. The migratory disproportionation and reductive dissociation of superoxide radicals were observed on MoOx /SiO2 and Cox Mg1−x O catalysts, respectively [85, 97]. Those processes can be followed fairly easily by EPR spectroscopy, due to the high sensitivity of the g tensor of O− 2 to the environment. Typical itinerant radicals are neutral species such as • H, • CH3 and • OH that on surfaces such as those of silica or alumina are only weakly stabilized by the van der Waals interactions [98]. Such species are therefore fairly mobile and can easily tumble and diffuse on the surface (cf. • H on Pd/Al2 O3 [142]), even at relatively low temperatures. Their motion, however, may be considerably perturbed. For instance, in the case of methyl radicals trapped on silica at 77 K, the tumbling frequencies are found to be 2 × 107 s−1 , whereas the rate of the corresponding free reorientations is 103 times faster [178]. The recombination of neutral radicals is diffusion controlled and the chemical reactivity is usually only slightly modified. e Surface Elementary Processes One of the goals in fundamental approaches to catalysis is the exploration of pathways along which reactants can be activated. Elucidation of surface molecular events, ascertaining their stoichiometries and sequences of appearance, and identification of key intermediates undoubtedly belong to inherent steps of such studies. In this context, surface electron transfer processes provide a paramount illustration of the utility of EPR techniques. Electron Transfer Processes The formation of paramagnetic intermediates is often initiated by surface electron transfer (ET). Because the unpaired electrons are readily observed by EPR, in such studies spin can be regarded as a sensitive tracer of the dynamics of the processes occurring. Typical systems involve either closed-shell or non-Kramers TMIs (e.g. d10 Cu+ or d8 Ni2+ ) interacting with paramagnetic ligands (NO, NO2 ) [133, 137, 141, 151] or, more frequently open-shell TMIs (d1 Mo5+ , V4+ ; d3 Cr3+ ; d7 Co2+ ; d9 Ni+ , Cu2+ , etc.), contacted with diamagnetic ligands such as N2 O, CO, olefins, small alcohol molecules, benzene or with dioxygen [127–129, 134, 179]. References see page 1063

3.2 Chemical Properties

Surface ET processes observed on oxide systems can be classified as non-dissociative, dissociative and electroprotic (electron transfer coupled with proton transfer) [133]. Specifically, the electron transfer can occur from surface transition metal ions to coordinated ligands (MLET), from ligands to metal ions (LMET), between ligands (LLET) [85, 133] or between adsorbed molecules [98, 180]. The last process is sometimes termed surface intermolecular electron transfer (SIET) [98]. Mechanistically, it is important to distinguish between complementary and non-complementary ET, which depends whether or not the redox site can accommodate all electrons supplied by a reactant molecule [133]. Noncomplementary ET may give rise to surface migration of the radicals, leading to pronounced structure and dispersion sensitivity of catalytic reactions, observed mostly for catalysts with a heterodesmic structure such as V2 O5 and MoO3 , bulk or spread on inert supports [85]. Oxidation of propene over Bi2 O3 , MoO3 and Bi2 MoO6 catalysts illustrates this point well, revealing a complex fate of the surface-formed allyl radicals [88]. Non-dissociative ET is often observed in the case of dioxygen contacted with TMIs dispersed on oxides. It may be accompanied by spillover of the resulting superoxide species on to the support, driven by the increase in the configuration entropy. Examples are provided by variabletemperature interaction of dioxygen with the surface of diluted Cox Mg1−x O solid solutions [135] or MoOx /SiO2 catalyst [181]. In the former case at low temperatures (T < 100 K), reversible adducts with gz = 2.120–113, gx = 1.983, Az = 3.8, Ax = 1.75 mT, consisting with 3+ centers, are top-on O− 2 moiety stabilized on Co observed. On increasing the temperature to 290 K, the 2+ superoxide spillovers onto MgO, forming O− 2 /Mg side-on species (gz = 2.07, gy = 2.009, gx = 2.002). In the case of MoOx /SiO2 interacting with dioxygen, the molecular events initiated by ET are quite involved and depend on molybdenum dispersion. EPR investigations have revealed that the main reaction steps include the 5+ radicals (g = 2.018, formation of primary O− 1 2 /Mo g2 = 2.011, g3 = 2.005), their spillover on to silica to form O− 2 /SiO2 (g1 = 2.0272, g2 = 2.02, g3 = 2.005) − 2− ↑ and surface disproportionation (O− 2 + O2 → O2 + O2 ) − competing with the secondary reduction process (O2 + 6+ Mo5+ → O2− 2 + Mo ). Because linear N2 O and bent N2 O− molecules have different geometries, activation of N2 O by ET must be preceded by vibrational excitation to satisfy the Franck–Condon principle. Since the resulting N2 O− radical anion is unstable, the ET results in its immediate dissociation. Such processes have been investigated in detail by EPR with MoOx /SiO2 catalysts [134]. It has

been found that the entire reaction can be epitomized in three steps involving (i) bonding of N2 O to tetracoordi5+ nated molybdenum, N2 O + Mo5+ 4c → N2 O-Mo4c (g⊥ = 1.96, g|| = 1.87), (ii) temperature-induced vibronic preactivation (C∞v → Cs bending of the coordinated N2 O), (iii) dissociative metal to ligand electron transfer (N2 O6+ ‡ 6+ ↑ − − Mo5+ 4c → {N2 O − Mo4c } → O − Mo4c + N2 ), lead− ing to the formation of O species bound to molybdenum Mo A = complex, O− –Mo6+ ⊥ 4c (g⊥ = 2.020, g|| = 2.005, Mo 0.67, A|| = 0.78 mT) and N2 release. From the temperature dependence of the EPR signal and analysis of the g and 95 Mo superhyperfine tensors, the energetic barrier for ET and the molecular structure of O− − Mo6+ 4c have been determined [134]. Electroprotic processes appear in the reaction between CH3 OH and O− , both adsorbed on silica-grafted molybdenum catalyst [182]. In this study, the EPR technique was used to investigate the cycle of model partial reactions, which mimic the most important features of the catalytic methanol oxidation [85]. Analysis of the data shows that the reaction consists of two processes: an activation of the C−H(sp3 ) bond by O− to form a ·CH2 OH radical ({CH3 OH + O− }Mo6+ → {· CH2 OH + OH− }Mo6+ ) and subsequent conversion of the transient hydroxymethyl radical into final CH2 O with concomitant reduction of Mo6+ surface ions ({· CH2 OH + OH− }Mo6+ –O2− → CH2 Oads + OH− Mo5+ + OH− ). This electroprotic step involves a ligand-to-metal electron transfer (LMET) coupled with a ligand-to-ligand proton transfer (LLPT). The · CH2 OH radicals decay above 140 K, reducing the remaining Mo6+ ions through the process · CH2 OH + − 5+ Mo6+ + O2− surf → CH2 O + OHsurf + Mo . In the presence of molecular oxygen, it competes with the superoxide pathway (· CH2 OH + O2 + O2− surf → CH2 O + OH− surf + O− 2 ). The superoxide radicals disappear by surface disproportionation or secondary MLET reaction, discussed above (Fig. 24). The mechanistic subtlety of this latter process is that it exhibits dispersion sensitivity. On the isolated Mo centers, the hydroxymethyl radicals have

0.8

I(t )/Io

1060

0.4

O2−

CH2OH 0 0

200

400

600

t/s Kinetic curves of the interaction of · CH2 OH radicals (triangles) with dioxygen (squares) at 295 K (after Ref. [181]).

Fig. 24

3.2.3 Valence States

to spill over on to silica searching for the remaining Mo6+ sites for terminal reduction to occur. f EPR Techniques in Working Conditions (Operando) and Catalytic Reaction Mechanisms In many reactions, radicals, whether charged or neutral, play an important role either as catalytically active sites or as intermediates [84, 85, 88, 173]. Since Lunsford’s review on the application of EPR to study catalytic reactions [173], numerous papers have appeared on in-situ EPR studies not only at room temperature [174, 183] but also at temperatures up to 500 ◦ C [184–186]. Catalytic reaction studies by means of in situ EPR techniques are usually restrained by the temperature gap between the optimal conditions for spectra registration (T < 300 K) and for conducting surface reactions (usually T > 400–500 K). There are also important technical constraints implied by the use of resonant cavities. As a result, investigations of catalytic reactions in real conditions have been made possible only in favorable cases, where signal broadening is not a limiting factor [85, 86, 88, 184, 186]. They may consist in a mere post mortem analysis of the quenched samples, quasi-in situ and operando studies or investigations that use ex situ flow micro-reactors in tandem with matrix isolation devices, in the case when radicals generated during the reaction escape from the surface [173]. The most common experimental approaches developed for this purpose include various thermally shielded flow cells inserted into a high-temperature cavity [168, 187], specially designed micro-reactors [184] and radical freezing and matrix isolation devices [88, 95, 188] and photo-cells for in situ studies of photocatalytic processes [189, 190]. Recently, the construction of a fixed-bed flow reactor, heated with bifilar Pt winding and equipped with UV–Vis and Raman fiber-optic sensors, for application in coupled operando EPR, UV–Vis and Raman spectroscopy has been described [186]. It has been successfully used for monitoring VOx /TiO2 catalysts during oxidative dehydrogenation of propene in statu operandi. The mechanistic importance of the open-shell processes in catalysis derives from the high reactivity of radical intermediates that, even at low abundance, may initiate competitive processes leading to significant changes in the main reaction course (e.g. driving a partial oxidation into total combustion [164]). Radical species can also define the principal pathway of catalytic transformations. Decomposition of NOx [98], dimerization of light alkanes [88], oxidation of alcohols [182] and hydrocarbons [88, 191], including regioselective reactions [192], hydrogenation of unsaturated hydrocarbons [193] and photocatalytic reactions [194–196] may serve as relevant illustrations. On the other hand, free-radical intermediates are known to initiate undesired processes that may lead to catalyst

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deactivation, e.g. by contributing to coke [184] or green oil (oligomeric hydrocarbons ≥ C4 ) formation [142], which severely affect the catalyst performance. Other examples are discussed elsewhere in more detail [86, 88, 97, 98, 139]. Among them, in situ mechanistic studies of allyl alcohol oxidation over V2 O5 –MoO3 catalyst provides a noteworthy illustration of how EPR can be used to quantify kinetic rate constants of the component reduction and oxidation processes that contribute to the global kinetic equation of the oxidation reaction [168]. Radicals generated in the course of catalytic reactions can simply remain on the surface and be involved in further reaction, following either the Langmuir–Hinshelwood or Eley–Rideal type of mechanism. They may also escape from the surface into the gas phase at high temperatures, contributing to the formation of primary products or taking part in chain processes that are often non-selective [88]. The mixed heterogeneous–homogeneous route of partial oxidation of organic reactants, with surface-generated gas-phase radicals as key intermediates, belongs to those radical processes in catalysis for which the greatest EPR achievements have been made. Radical processes can be greatly modified in the presence of constrained media such as molecular sieves. Space confinement renders radicals not only persistent by supramolecular steric effects [196], but also the diffusing reactants gain easy access to the active sites only at the unrestrained positions. This may lead to a pronounced regioselectivity of the catalytic reaction, as has been observed in aluminophosphate materials such as CoAlPO-18 and MnALPO-18, containing isolated tetrahedral CoIII or MnIII substituted into the framework of the sieve. They are active in the free-radical autoxidation of n-octane and n-hexane with a high degree of regioselectivity [192]. D Conclusions EPR spectroscopy is a powerful, non-destructive and sensitive technique which can be successfully applied to resolve many of the problems encountered in surface chemistry and catalysis. In most catalytic and surface studies, polycrystalline samples, rather than single crystals, are used and information on catalysis phenomena is obtained from powder EPR spectra. We have outlined the importance of magnetic parameters extracted from powder EPR spectra to identify and characterize the paramagnetic centers together with their environment. The technique appears most efficient in characterizing active sites, reaction intermediates and in monitoring the elementary processes occurring within the coordination sphere of surface transition metal ions. The physical and chemical methods can often References see page 1063

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be combined to arrive at a better understanding and interpretation of complex EPR spectra. Typical examples have been given to illustrate the use of EPR in catalysis, ranging from catalyst preparation to catalytic reactions. For catalytic studies, however, the joint use of several techniques, especially those which can be correlated with EPR (essentially UV–Vis and photoluminescence), is more powerful than EPR alone. In recent years, an increasing number of papers have appeared on studies involving the coupling of UV–Vis and EPR with other techniques/spectrocopies (see Section 3.2.3.2.2.Bb), for example EPR–UV–Vis [86], EPR–UV–Vis–LRS [197] and EPR–Raman–IR [198].

In UV–Vis–NIR, the usually more intense chargetransfer bands often overlap with the weak d–d bands, making the latter invisible; in such cases, they can be revealed in EPR, using Eqs. (8) and (9) as seen above.

g⊥ = ge−

2λ (Eyz,zx − Exy )

(8)

g|| = ge−

8λ (Ex 2 −y 2 – Exy )

(9)

B Differences EPR, unlike UV–Vis–NIR, is confined exclusively to the detection of paramagnetic species. There is generally a low-temperature limitation for the observation of EPR spectra, but in UV–Vis–NIR temperature is less of a restriction. The samples in EPR cannot possess high dielectric constant or conductivity, whereas in DRS optical properties are important (in order to reflect they cannot be black). Charge-transfer transitions have no analogs in EPR, and vice-versa, hfs (shfs) interactions (providing information about spin density distribution in the ground state) have no counterpart in UV–Vis–NIR. In UV–Vis–NIR, electrons are promoted from the ground term to excited terms, without changing spin, following the S = 0 rule, whereas in EPR, electrons change spin, obeying the ms = 1 rule. Although photons are required for both spectroscopies, the energies of the photons are distinctly different, much higher for UV–Vis–NIR than for EPR (electronic transitions are of the order of 10 000 cm−1 versus 1 cm−1 for spin transitions). In UV–Vis–NIR, energy levels are determined only by the nature and oxidation state of the TMI and by the nature and number of ligands of the complex (Fig. 14c) whereas in EPR, the energy levels depend on the strength of the external magnetic field as indicated by Eq. (6), via the Zeeman effect (Fig. 14d).

where λ is the metal spin–orbit coupling constant and Eyz,zx − Exy and Ex 2 −y 2 − Exy are two of the three electronic transitions indicated in Fig. 14c. From the experimental g-tensor components and knowing the effective value of λ, the values of the two electronic transitions Eyz,zx − Exy and Ex 2 −y 2 − Exy can be derived. Alternatively, when the electronic transitions are available from UV–Vis–NIR (DRS) spectra, these expressions can lead to the value of λ. In some more sophisticated treatments [199], theoretical g, hf and shf tensor expressions can enable to derive covalency factors of the chemical bonds between TMIs and their surrounding ligands. It has been seen (see Section 3.2.3.2.2.Bb) that the nephelauxetic effect corresponds to d orbital expansion towards the ligands upon complexation. Whereas in UV–Vis–NIR this effect is measured globally for the complex by the value of the Racah B parameter, in EPR it is measured selectively for atoms which have I = 0 nuclei. In such cases, shf structures and constants can be obtained from which the distribution of the spin density on atoms surrounding the paramagnetic center can be deduced.

Conclusions As with most spectroscopies, UV–Vis–NIR and EPR techniques have their own advantages and disadvantages, which have been outlined. There is, however, one major asset shared by both techniques, namely the possibility of studying ET processes. The latter are of the utmost importance since they lie at the core of chemistry which deals with reactivity. As an example, ET processes are known to play an important role in the activation of many molecules by TMIs. Although many studies have been conducted at room temperature, there appears to be a trend for in situ studies not only at room temperature but also at higher temperatures. Despite their interest, any results obtained by UV–Vis–NIR and EPR should always be compared, at both the qualitative and quantitative levels, with those from other spectroscopic and non-spectroscopic (thermal, structural, etc.) techniques. It is always important to keep in mind that the joint use of several complementary techniques (UV–Vis–NIR and EPR are indeed such techniques) is the best approach to obtain a more general understanding of catalysis phenomena.

3.2.3.2.4 Comparison between UV–Vis–NIR (DRS) and EPR Spectroscopies A Analogies and Complementarities Although UV– Vis–NIR and EPR spectroscopies rest on different principles, they can provide similar chemical information on oxidation states, coordination and local symmetry and energy levels of TMIs. They are complementary: whereas UV–Vis–NIR provides direct information on energy levels such as those given in Fig. 14c, in EPR such energy levels can only be derived from theoretical expressions of g-tensor components. For instance, for a d1 ion with the electron in a dxy ground state, the latter are calculated to be

3.2.3.2.5

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3.2.3 Valence States 139. M. Che, E. Giamello, in Catalyst Characterization: Physical Techniques for Solid Materials, B. Imelik, J. C. V´edrine (Eds.), Plenum Press, New York, 1994, Chapter 6. 140. O. O. Parenago, Yu. N. Pushkar, A. O. Turakulova, G. P. Muraveva, E. V. Lunina, Kinet. Katal. 1998, 39, 268. 141. R. F. Howe, Adv. Colloid. Interface Sci. 1982, 18, 1. 142. A. F. Carley, H. A. Edwards, B. Mile, M. W. Roberts, C. Rowlands, F. E. Hancock, S. D. Jackson, J. Chem. Soc., Faraday Trans. 1994, 90, 3341. 143. C. Louis, M. Che, J. Catal. 1992, 135, 156. 144. C. Louis, M. Che, M. Anpo, J. Catal. 1993, 141, 453. 145. E. Giamello, M. C. Paganini, D. M. Murphy, A. M. Ferrari, G. Pacchioni, J. Phys. Chem. B 1997, 101, 971. 146. M. Labanowska, Colloids Surf. 1993, 72, 177. 147. M. Che, G. Djega-Mariadassou, K. Dyrek, Z. Olech, Z. Phys. Chem. N. F. 1987, 152, 131. 148. A. Martinez-Arias, M. Fernandez-Garcia, C. Belver, J. C. Conesa, J. Soria, Catal. Lett. 2000, 65, 197. 149. B. M. Weckhuysen, I. E. Wachs, R. A. Schoonheydt, Chem. Rev. 1996, 96, 3327. 150. S. Kuba, P. Lukinskas, R. G. Grasselli, B. C. Gates, H. Kn¨ozinger, J. Catal. 2003, 216, 353. 151. D. Bigliano, H. Li, R. Erickson, A. Lund, H. Yahiro, M. Shiotani, Phys. Chem. Chem. Phys. 1999, 1, 2887. 152. L. Kevan, Electron Spin Reson. 1991, 12B, 99. 153. C. Nozaki, C. G. Lugmair, A. T. Bell, T. D. Tilley, J. Am. Chem Soc. 2002, 124, 13194. 154. E. El Malki, B. Manohar, A. M. Davidson, P. Massiani, S. Sivasanker, M. Che, J. Chim. Phys. 1997, 94, 1848. 155. A. V. Kucherov, A. A. Slinkin, Zeolites 1987, 7, 38. 156. L. Bergaoui, J. F. Lambert, H. Suquet, M. Che, J. Phys. Chem. 1995, 99, 2155. 157. M. R. Harrison, J. Edwards, J. Klinowski, J. M. Thomas, D. C. Johnson, C. J. Page, J. Solid State Chem. 1984, 54, 330. 158. J. Michalik, L. Kevan, J. Am. Chem. Soc. 1986, 108, 4247. 159. J. Michalik, Appl. Magn. Reson. 1996, 10, 507. 160. D. M. Murphy, E. Giamello, J. Phys. Chem. 1994, 98, 7929. 161. A. Aboukais, A. Bennani, C. F. Aissi, G. Wrobel, M. Guelton, J. C. Vedrine, J. Chem. Soc., Faraday Trans. 1 1992, 88, 615. 162. H. J. Chen, M. Matsuoka, J. L. Zhang, M. Anpo, J. Catal. 2004, 228, 75. 163. A. P¨oppl, M. Newhouse, L. Kevan, J. Phys. Chem. 1995, 99, 10019. 164. A. Bielanski, J. Haber, Oxygen in Catalysis, Marcel Dekker, New York, 1991, 467 pp. 165. A. Br¨uckner, B. Kubias, B. L¨uecke, Catal. Today 1996, 32, 215. 166. U. Lohse, A. Br¨uckner, E. Schreier, R. Bertram, J. J¨ahnchen, R. Fricke, Microporous Mater. 1996, 7, 139. 167. K. Dyrek, M. Labanowska, J. Catal. 1985, 96, 32. 168. N. Then Tai, B. V. Rozentuller, O. V. Krylov, Kinet. Katal. 1998, 39, 392. 169. K. Dyrek, M. Labanowska, J. Chem. Soc., Faraday Trans. 1991, 87, 1003. 170. P. Botella, B. Solsona, A. Martinez-Arias, J. M. L´opez Nieto, Catal Lett. 2001, 74, 149. 171. C. Louis, C. Lepetit, M. Che, Mol. Eng. 1994, 4, 3. 172. A. Davidson, M. Che, J. Phys. Chem. 1992, 96, 9909. 173. J. H. Lunsford, Catal. Sci. Technol. 1987, 8, 227. 174. A. K. Ghosh, L. Kevan, J. Phys. Chem. 1988, 91, 4439. 175. M. Kobayashi, T. Shirasaki, J. Catal. 1975, 38, 394. 176. H. Yahiro, K. Manabe, Y. Itagaki, M. Shiotani, J. Chem.Soc., Faraday Trans. 1 1998, 94, 805.

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177. Z. Sojka, M. Chiesa, E. Giamello, M. C. Paganini, Stud. Surf. Sci. Catal. 2001, 140, 443. 178. B. D. Flockart, in Surface and Defect Properties of Solids, M. W. Robert, J. M. Thomo (Eds.), Spec. Period. Rep., Vol. 2, 1973, p. 69. 179. B. M. Weckhuysen, R. A. Schoonheydt, F. E. Mabbs, D. Collison, J. Chem. Soc., Faraday Trans. 1996, 92, 2431. 180. O. Diwald, E. Kn¨ozinger, J. Phys. Chem. B 2002, 106, 3495. 181. Z. Sojka, M. Che, Top. Catal. 2001, 15, 211. 182. Z. Sojka, M. Che, J. Phys. Chem. 1995, 99, 5418. 183. F. X. Cai, C. Lepetit, M. Kermarec, D. Oliveer, J. Mol. Catal. 1987, 43, 93. 184. H. G. Karge, J. P. Lange, A. Gutsze, M. Laniecki, J. Catal. 1988, 114, 144. 185. A. V. Kucherov, J. L. Gerlock, H.-W. Jen, M. Shelef, J. Catal. 1995, 152, 63. 186. A. Br¨uckner, Chem. Commun. 2005, 176. 187. O. B. Lapina, B. S. Balzhinimaev, S. Boghosian, K. M. Eriksen, R. Fehrmann, Catal. Today 1999, 51, 469. 188. T. A. Garibyan, L. Yu. Margolis, Catal. Rev. Sci. Eng. 1989–90, 31, 355. 189. Y. Inoue, T. Kubokawa, K. Sato, J. Phys. Chem. 1991, 95, 4059. 190. A. M. Volodin, V. I. Sobolev, G. M. Zhidomirov, Kinet. Katal. 1998, 39, 775. 191. S. Kuba, P. C. Heydorn, R. K. Grasselli, B. C. Gates, M. Che, H. Kn¨ozinger, Phys. Chem. Chem. Phys. 2001, 3, 146. 192. J. M. Thomas, Angew. Chem. Int. Ed. 1999, 38, 3588. 193. G. C. Chinchen, M. S. Spencer, Catal. Today 1991, 10, 293. 194. K. I. Zamaraev, Stud. Surf. Sci. Catal. 1996, 101, 35. 195. M. Anpo, M. Matsuoka, H. Yamashita, Catal. Today 1997, 35, 177. 196. X. G. Lei, S. Jockusch, M. F. Ottaviani, N. J. Turro, Photochem. Photobiol. Sci. 2003, 2, 1095. 197. A. Br¨uckner, E. Kondratenko, Catal. Today 2006, 113, 16. 198. M. Brandhorst, S. Cristol, M. Capron, C. Dujardin, H. Vezin, G. Le Bourdon, E. Payen, Catal. Today 2006, 113, 34. 199. M. Che, M. Fournier, J. P. Launay, J. Chem. Phys. 1979, 71, 1954.

Photoluminescence Spectroscopy and Its Application to the Characterization of Active Sites and Reaction Dynamics in Catalysis

3.2.3.3

Masaya Matsuoka, Takashi Kamegawa, and Masakazu Anpo∗

3.2.3.3.1 Introduction Recent advances in various spectroscopic technologies have allowed the acquisition of spectra with high resolution and sensitivity on a shorter time-scale, leading to a better understanding of the chemical nature of catalytically active sites and of the mechanisms of catalytic reactions on a molecular level. Time-resolved spectroscopy has been especially successful in investigations of the reaction dynamics at a time resolution of the orders of 100 fs, based on recent progress References see page 1072 ∗ Corresponding author.

1066

3.2 Chemical Properties

0→4 0→3 0→2 0→1

Wavelength

S1 5 4 3 2

S0

1

0→0

0

Intensity 5 4 3 2 1 0

Wavelength

Absorption spectrum (Excitation)

0→4 0→3 0→2 0→1 0→0

Intensity

Intensity

2nd derivative Photoluminescence photoluminescence spectrum spectrum

Schematic energy diagram showing the origin of absorption, photoluminescence and second-derivative photoluminescence spectra.

Fig. 1

in laser technology. Among the new and advanced spectroscopic techniques, photoluminescence spectroscopy is one of the most powerful and sensitive methods to investigate coordinatively unsaturated surface sites of solid catalysts and the local structure of highly dispersed oxide catalysts, both of which exhibit unique photoluminescence under UV irradiation [1–14]. Moreover, time-resolved photoluminescence spectroscopy is also applied to elucidate the dynamics that initiate photocatalytic reactions involving a recombination process of the photo-formed electrons and holes [13–16]. In this chapter, the applications of photoluminescence spectroscopy, especially time-resolved photoluminescence analysis of the characterization and reactivity of solid catalysts, are presented in relation to the reaction dynamics of photocatalysis. 3.2.3.3.2 Characterization of Catalysts by Photoluminescence and Time-Resolved Photoluminescence Spectroscopy The sequence of steps involved in the photoluminescence process is described in Fig. 1. The initial absorption excites the molecule from the ground electronic singlet state (S0 ) to the electronically excited singlet state (S1 ). The photoexcited molecule relaxes down the ladder of

vibrational levels, releasing its energy into its surroundings as heat and finally ends up in the lowest vibrational level of S1 . If the thermal environment cannot accommodate the electronic excitation energy (transition from S1 to S0 ), the molecule emits this energy as radiation, which can be observed as photoluminescence (fluorescence) [8]. When intersystem crossing occurs between S1 and the triplet state (T1 ), phosphorescence can be observed as the radiative deactivation process from the lowest vibrational state of T1 to S0 [8]. If the excitation energy is localized on a certain bond of the molecule, the photoluminescence spectrum exhibits a well-resolved vibrational fine structure which reflects the vibrational energy levels of the S0 state. The second derivative photoluminescence spectrum can be utilized to determine the energy value between the vibrational energy levels of the S0 state, since the local minimum of the second-derivative photoluminescence spectrum corresponds exactly to the local maximum of the vibrational fine structure of the photoluminescence spectrum [2–4]. hν

hν

[Ti4+ − O2− ] −−−→ [Ti3+ − O− ]∗ −−−→ [Ti4+ − O2− ] Normally, the photoluminescence or steady-state photoluminescence spectra are measured under continuous irradiation by a light source such as an Xe or high-pressure

1067

3.2.3 Valence States

Delay time

Sampling time

Lamp pulse

Lamp pulse

t=0

Time

Sampling

t=0

Time

Typical sequence of data acquisition for time-resolved photoluminescence measurements.

Fig. 2

Hg lamp. In contrast, time-resolved photoluminescence spectra are measured by pulsed irradiation of the samples using light sources such as a nitrogen laser or pulse Xe lamp. Time-resolved photoluminescence spectra are obtained by collecting the photoluminescence radiation within a certain time interval (sampling time) after a certain time period (delay time) following the pulse excitation, as shown in Fig. 2. As an example, threedimensional plots of time-resolved photoluminescence of Ti/MCM-41 measured at various delay times are shown in Fig. 3a. A broad photoluminescence due to the surface OH groups (Si−OH) of MCM-41 can be observed at around 400 nm in the spectrum measured at a short delay time (0.5 ms). However, the intensity of the photoluminescence at around 400 nm decreases with increasing delay time (∼1.0 ms) and another photoluminescence process was observed at around 460 nm. This emission can be attributed to the following charge-transfer processes on the Ti−O moieties of the isolated tetrahedrally coordinated Ti4+ oxide species [Ti(IV)O4− 4 ], involving an electron transfer from the O2− to Ti4+ ions and a reverse radiative decay from the charge-transfer excited triplet state [2–4]. The lifetime of the excited state of two different photoluminescence moieties can be determined by decay analysis at certain wavelength regions where the overlap of the two photoluminescence moieties are negligible [Si−OH, 0.11 ms at 370 nm; Ti(IV)O4− 4 , 11.8 ms at 480 nm]. It should be noted that, under continuous irradiation, the photoluminescence of the Ti4+ oxide species can be observed only as a small shoulder on the photoluminescence of MCM-41, as shown in Fig. 3b. Time-resolved photoluminescence techniques are therefore useful in distinguishing different photoluminescent moieties having different photoluminescence lifetimes. Time-resolved photoluminescence of the tetrahedral Ti4+ oxide species at 460 nm revealed them to be efficiently quenched by the addition of NO or propane, accompanied by a decrease in the photoluminescence lifetime. This indicates that the photoexcited Ti4+ oxide species interacts dynamically with the added gas-phase molecules. In fact, UV irradiation of Ti/MCM-41 in the presence of NO and propane led to the efficient photocatalytic reduction of

Delay time 0.5 ms

(a)

400

350

0.6

600

(b)

450

550

Wavelength / nm 200

Intensity / a.u.

Delay time

Time

NO by propane to produce N2 and acetone [1–5]. These results clearly demonstrate the advantage of time-resolved photoluminescence measurements for the investigation of the reaction dynamics of the photoexcited active site with reactant molecules as compared with steady-state photoluminescence measurements. Figure 4 shows the time-resolved photoluminescence spectra of V-containing β-zeolite catalysts [V/Siβ (0.05 wt.%)] and its corresponding second-derivative spectra [9, 10]. Here, the second-derivative photoluminescence spectra are shown upside down by multiplying the original spectra by −1 in order to emphasize the good correspondence between vibrational fine structure and its second-derivative spectrum. The photoluminescence spectrum observed under steady-state irradiation did not show a well-defined vibrational fine structure (Fig. 4a), whereas its second-derivative spectrum exhibited welldefined peaks (Fig. 4a ), suggesting that the steady-state photoluminescence spectrum may consist of several overlapping photoluminescence moieties having different vibrational fine structure. In fact, optimizations of the delay and sampling times allow the discrimination of photoluminescent moieties having different lifetimes. Figure 4b and c show the time-resolved photoluminescence spectra measured at different delay times and sample windows, the former and the latter corresponding to the photoluminescence moieties having short and long lifetimes, respectively. It can be clearly seen that both timeresolved photoluminescence spectra exhibit well-defined vibrational fine structures. Comparisons of the time-resolved spectra with the steady-state spectrum made it possible to conclude that Intensity / a. u.

To computer / software

0.7 1.0

0.8 0.9

360

400

0

5 10 15 440

480

520

560

Wavelength / nm Time-resolved photoluminescence spectra (a) and steady-state photoluminescence spectrum (b) of Ti/MCM-41 (Si : Ti = 200) measured at 298 K. (a) Sampling time, 0.1 ms; excitation wavelength, 260 nm. (b) Excitation wavelength, 260 nm. Fig. 3

References see page 1072

1068

3.2 Chemical Properties

6

Intensity / a. u.

4

α

α β

β

α β

α β

2

α γ

γ

γ

β (a)

γ

0 −2

(a′) (× − 1)

−4

Intensity / a. u.

0.8

α

α

α

α α

(b)

0.4 0 (b′) (×−1)

−0.4

Intensity / a. u.

0.3 0.2

β

β β

β

(c)

0 (c′) (× −1)

−0.1 500

550

Investigations of the Dynamics of Photocatalysis by Time-Resolved Photoluminescence Spectroscopy

3.2.3.3.3

β

0.1

450

V=O bond length decreases in the order α < γ < β. The distortion of the coordination sphere of V oxide species from an ideal Td (tetrahedral) symmetry increases in the reverse sequence: α < γ < β [9, 10]. It was also found that the photoluminescence lifetime increases with increasing distortion of the coordination symmetry of the vanadium oxide species, i.e. 28 ms (α) OFF > FAU. A maximum of strong acid sites (corresponding to more than 88% H2 SO4 ) was expected for the ratio Si/Al ≈ 9.4 in H-MOR or 6.8 in H-FAU. A number of reactions are mentioned in Ref. [66] which indeed exhibited a maximum of conversion over, e.g., H-MOR at Si/Al ≈ 10 or H-FAU at Si/Al ≈ 7. In another interesting contribution to the elucidation of the phenomenon of acid site strength, Barthomeuf [56, 64] interpreted the increase in acid strength of H-forms of zeolites with decreasing Al content in analogy with the behavior of inorganic oxy acids of the general formula XOn (OH)m . The strength of such acids increases with increasing n or decreasing m. An example is the sequence of oxy acids of chlorine: Cl(OH) < ClO(OH) < ClO2 (OH) < ClO3 (OH). The corresponding formula for zeolites is TOn (OH)m with T = Al or Si and n + m = 2. In the case of Y-type zeolite with Si/Al = 2.42, one obtained n = 1.71 and m = 0.29, whereas for a dealuminated Y-zeolite with Si/Al = 4.12 the result was n = 1.81 and m = 0.19. From experiments on the temperatureprogrammed desorption of pyridine (see below), it was inferred that a material with a higher value of n exhibited a higher strength of acidity. B Utilization of Indicators and Titration in Aprotic Solvents The utilization of indicators and titration of acidic and basic sites in aprotic solvents such as benzene or hexane was reported relatively early and extensively for surfaces of oxides [68, 69], especially for the classical acidic silica–alumina solids [25, 69, 70], the predecessors of acid zeolites as cracking catalysts. This method and its applications were excellently reviewed by Tanabe [71]. From the color of the appropriate indicators, which in aprotic suspension are simply contacted with the (white) powder of the oxides or zeolites, one may qualitatively estimate the (maximum) acid or basic strength of the surface sites. Titration with a suitable titrator (in the case of acidity usually butylamine) enables one to determine the concentration of the respective sites. The titration cannot be conducted in aqueous suspension, because this would eliminate the differences in acid strength of the various sites. The experimental procedures are described in the literature (see, e.g., Refs. [25, 70–73]). The use of indicators and titration of the acidic and basic sites [74] are, in fact, the most chemistry-related techniques for zeolite characterization. References see page 1118

1104

3.2 Chemical Properties

A useful basis for a comparison of various strengths of acids is the capability of proton transfer to a neutral base. This may be illustrated by the following equilibrium: + −−  B + H+  −− − −BH

(6)

with B = neutral base and BH+ = conjugated form of the base. Quantitatively, the proton transfer may be described by the acidity function H0 , which was introduced by Hammet and Deyrup [75]:
 aH+ fB (7) H0 ≡ − log fBH+ where aH+ = proton activity and fB and fBH+ = activity coefficients of B and BH+ , respectively. Similarly, the acid strength of a solid surface may be visualized according to Walling [76] as the capability of an acidic solid surface to transform an adsorbed neutral indicator base, I, into its conjugated form, IH+ :  

aH+ fI c + H0 ≡ − log = pKIH+ − log IH (8) fIH+ cI where pKIH+ is the negative Briggs’ logarithm of the thermodynamic equilibrium of the indicator (frequently simply termed pKa values) and cIH+ and cI are the concentrations of its acidic and basic form. When cIH+ and cI become equal (around the color change in the case of titration), it follows that H0 = pKIH+ or pKa . Relationships analogous to Eq. (8) can be derived for pure Lewis acids of solid surfaces and also for the use of Hirschler indicators of the arylmethanol type (see Ref. [77], where also preparations of a large number of appropriate Hirschler indicators are described). However, it turned out that none of the usually used indicators (i.e. neither Hammett [74] nor Hirschler [25, 77] indicators; cf. Table 2) are really selective for either Brønsted or Lewis acid centers. The titration method essentially fails if both types of acid sites are simultaneously present. Moreover, as mentioned above, the titrations unavoidably have to be conducted in aprotic media (e.g. benzene, hexane), whereas the relevant pKa values of the indicators, the color change of which indicates the strength of a certain fraction of the sites (corresponding to the respective pKa s), are defined for aqueous solutions. Thus, the basis of the applicability of the titration technique remains debatable. On the other hand, the argument frequently advanced against the titration method, namely that the bulky indicator molecules are hindered from diffusing into the interior of the zeolite crystallites, is not relevant, since the visual observation of the indicator molecules and their change upon protonation by acid sites is, at any rate, restricted to the outer surface, except that optical

spectroscopy is employed in transmission for determining the end-point of titration. However, in many cases the competition between the adsorption of the molecules of the aprotic solvent and that of the titrator base (e.g. butylamine) impedes the approach to a true chemical equilibrium between the indicator base on the external zeolite surface and the acid sites inside the zeolite structure within a reasonably extended time of the experiment [78]. Neglecting the above-mentioned fundamental limitations of the titration method and taking a phenomenological point of view, one may state, however, that the titration of zeolite acidity has led in several cases to reasonable results, in particular with more open zeolite structures such as faujasite-type materials [25, 77, 79–83]. Thus, the total number of Brønsted acid sites as determined via titration (corresponding to the amount of butylamine consumed with 4-benzeneazodiphenylamine as an indicator, pKa = +1.5) was found to agree very well with the number calculated from the chemical composition of the zeolite [72, 83]. In a few publications, successful titration of zeolite acidity was reported with potentiometric or thermometric determination of the end-point of titration [84]. C IR Spectroscopy of Brønsted Acidity with and without Probe Molecules IR spectroscopy, as one of the most powerful tools for identifying the nature of and quantifying the acidity of solids, was already employed before the advent of acid zeolites as catalysts, viz. in the investigation of amorphous silica/alumina cracking catalysts [85]. The same technique, i.e. the study by diffuse transmission IR spectroscopy of thin wafers, which were pressed from zeolite powders, was then used for characterization of these materials [86, 87]. In the pioneering work by Uytterhoeven et al. [88], the formation of acid OH groups through deammoniation of the ammonium form of faujasite-type zeolite Y was, for the first time, monitored by transmission IR spectroscopy and, also, the subsequent dehydroxylation at higher temperatures (cf. Scheme 3, Fig. 4). The OH groups were easily identified by the typical OH stretching bands at 3640 and 3550 cm−1 (Fig. 4). Similarly, OH groups as Brønsted acid sites were detected in a large number of hydrogen forms of other zeolite types (for the abbreviations, i.e. the three-letter codes, see Ref. [17]), for example in H-MOR (3610 cm−1 [89, 90]), H-HEU (heulandite and clinoptilolite structure, 3600 cm−1 [90]), H-MFI or H-ZSM-5 (3605 cm−1 [91]), H-OFF (3610 and 3550 cm−1 [92]), H-BEA (3612 cm−1 [93]) and H-FER (3640 and 3600 cm−1 [94]), to mention just the most important examples. A more extended list is provided in Table 6 in Ref. [95], in which review the application of IR spectroscopy to zeolites is dealt with in much more detail.

3.2.4 Acidity and Basicity

1105

Hammett indicators and Hirschler (arylcarbinol) indicators suitable for the visible indication of the end-point in the titration of colorless solid acidsa

Tab. 2

Indicator Hammett indicators Natural Red Phenylazonaphthylamine Butter Yellow 4-Benzeneazodiphenylamine Dicinnamalacetone Benzalacetophenone Anthraquinone Hirschler (arylcarbinol) indicators 4,4 ,4 -Trimethoxytriphenylmethanol 4,4 ,4 -Trimethyltriphenylmethanol Triphenylmethanol 3,3 -Trichlorotriphenylmethanol Diphenylmethanol 4,4 ,4 -Trinitrotriphenylmethanol 2,4,6-Trimethylbenzyl alcohol a After

Basic color

Acid color

pKa

Acid strength (wt.% H2 SO4 )

Yellow Yellow Yellow Yellow Yellow Colorless Colorless

Red Red Red Purple Red Yellow Yellow

+3.3 +4.0 +3.3 +1.5 −3.0 −5.6 −8.2

8 × 10−8 5 × 10−5 3 × 10−4 2 × 10−2 48 71 90

Colorless Colorless Colorless Colorless Colorless Colorless Colorless

Yellow Yellow Yellow Yellow Yellow Yellow Yellow

+0.82 −4.02 −6.63 −11.03 −13.3 −16.27 −17.38

1.2 36 50 68 77 88 92.5

Benesi and Winquist [7].

The acidic character of the above-indicated OH groups was proven also by IR spectroscopy: Contact of the hydrogen form of, e.g., faujasite-type zeolite (H-Y) with basic probe molecules such as NH3 or pyridine resulted in the elimination of the corresponding OH groups. Simultaneously, the IR bands of ammonium and pyridinium ions developed at 1640 and 1540 cm−1 , respectively. Such experiments were also first carried out in Hall’s group [88] (cf. also, e.g., Refs. [96, 97]). Ammonia and pyridine still belong to the most popular probes for acidity employed in spectroscopic and non-spectroscopic methods (for the latter, see below, e.g. temperature-programmed desorption and microcalorimetry, Sections 3.2.4.1.3H and I, respectively). Whereas the spectra of ammonia adsorbed on zeolites usually exhibit relative broad bands, the bands of pyridine are rather sharp (cf. Fig. 5) and well-suited for quantitative evaluation of the exact wavenumbers and the band areas. The concentration of sites can be determined from the depths or areas of the IR bands if the intensity (absorbance A) is plotted against wavenumber. For this purpose, use is made of the Lambert–Beer–Bouguer law: A = εν˜ cd

(9)

where A is the absorbance [see Eq. (10)], εν˜ is the extinction coefficient at wavenumber ν˜ of a sample with concentration c of centers absorbing at ν˜ and a thickness d. If in transmission spectroscopy (see below) the spectrometer does not provide directly the absorbance of the sample but the so-called transmittance, T , then one has to convert T into A according to

Eq. (10): A = − log T = log(I0 /I )ν

(10)

with I0 being the incident and I the transmitted radiation energy. In principle, Eq. (9) is valid only for dilute systems. For instance, in the case of zeolites populated with OH groups or adsorbate–zeolite systems, one has to be aware of the possibility that εν˜ may be dependent on the OH group concentration and adsorbate coverage. Frequently, it is sufficient to use as a first approximation the maximum absorbance defined by Amax = log T ∗ (˜νmin )/ log T (˜νmin )

(11)

where T ∗ (˜νmin ) is the background (baseline) transmittance and T (˜νmin ) the actual transmittance, both measured at the wavenumber ν˜ min of minimal transmittance (cf. Fig. 6 and Ref. [98]). The integral absorbance, as a more precise measure of the concentration of the absorbing species, is expressed as in Eq. (12):   ν˜e   ν˜ e T (˜ν ) Aint = cd log εν˜ d˜ν (12) d˜ ν = cd T (˜ν ) ν˜ b ν˜ b where ν˜ b and ν˜ e are the wavenumbers of the beginning and the end of the band, respectively, and T (˜ν ) and T (˜ν ) are the transmittances along the baseline and the band contour, respectively. Modern instrumentation usually allows routine baseline determination and band integration to evaluate Aint . In IR and Raman References see page 1118

1106

3.2 Chemical Properties

100

Transmission

80

25°C 60

555°C 3580

40

110°C 475°C 230°C

290°C

20

385°C

415°C

3670 Wavenumber / cm−1

Development of the stretching bands of acid OH groups of an NH4 ,Na-Y zeolite upon deammoniation in high vacuum at increasing temperatures and subsequent dehydroxylation [88].

Fig. 4

Transmittance

H-Y

H - Mordenite

L 1452

L 1452

+

B 1490 1542

B 1542

1700

Na - Mordenite

1300 1700

C (Na ) 1440

1300 1700

1300

Wavenumber/cm−1

IR bands of pyridine adsorbed on Brønsted acid sites (B-sites), ‘‘true’’ Lewis acid sites (L-sites) and Lewis acid cations (C-sites, here Na+ ) in H-faujasite (H-Y), H-MOR and Na-MOR [31].

Fig. 5

spectroscopy of zeolites or oxides and adsorbate-zeolite or adsorbate–oxide samples, the extinction coefficient, εν˜ or εν˜ (c), is usually unknown. Therefore, if knowledge of the absolute concentrations is required, εν˜ has to be determined in separate experiments (cf., e.g., Refs. [99–104]). In such experiments, the absorbance has to be measured of zeolite samples covered with a known number of functional surface groups or loaded with well-defined amounts of adsorbate in order to obtain calibration curves (Aint vs. concentration, c). So far, such determinations of εν˜ have been carried out for a limited number of systems (cf., e.g., Refs. [105–110]).

Examples of results are collected in Table 2 in Ref. [95]. In some cases, the agreement between the results of different authors is satisfactory. However, frequently significant deviations occur. Therefore, it is advisable not simply to adopt literature data, but to check them or rely on ones own data obtained with ones own range of samples, instrumentation and techniques (for experimental details, compare Refs. [95, 98]). In any event, the (absolute) density of sites is defined as their number per unit cell, N (u.c.)−1 or millimoles per gram of dehydrated material (mmol g−1 ). However, since the distribution of sites over the zeolite crystals is not necessarily homogeneous,

3.2.4 Acidity and Basicity

∼ ∼ ∼ A int = c ⋅d ⋅ ve e∼ dv∼ = ve log T ′(V∼) dv∼ T ′′(V ) v∼b v v∼b

1107

A max = e∼ c ⋅d ⋅ = log T* v T

min

Transmittance/%

100

T ′ (v∼)

T* (v∼min) T ′′ (v∼ )

Background

Tmin

0

∼ v∼b v∼min v b

∼ Wavenumber, v /cm−1

Fig. 6

Quantitative evaluation of IR spectra – integrated and maximum absorbance (see text) [98].

densities given in N (u.c.)−1 or mmol g−1 may provide an average value. In the above-described way, the density of sites may be evaluated from the intensities of the bands of the acid OH groups and also from those of the corresponding ammonium or pyridinium ions (see above). However, both ammonia and pyridine are strong bases and, therefore, rather weak acid sites may be indicated, which in some cases is a disadvantage. Moreover, pyridine molecules are bulky and do not fit into smaller channels or voids of zeolite structures. In fact, if one wishes to determine the so-called external Brønsted (or Lewis) acidity on the outer surfaces of zeolite crystallites, bulky probe molecules such as substituted pyridines (2,6-di-tert-butylpyridine, etc.), which cannot penetrate into the zeolitic channels and cavities, may even be very helpful (cf. examples in Refs. [111, 112]). However, in the majority of cases one is interested in the characterization of the internal acidity, i.e. inside the zeolite structures. Therefore, many researchers looked, because of the above-mentioned reasons, for other appropriate probe molecules. Examples of adsorbates tested for probing acidity include water [113, 114], amines [115], light hydrocarbons [116], hydrogen [117], carbon monoxide [118, 119], nitrogen monoxide [120, 121] and acetonitrile [122, 123], and carbon dioxide [124] and pyrrole [63] for basicity, to list just a selection of probes and related references (cf., e.g., Ref. [125] and Table 8.6 in Ref. [126]). In many instances, such probes can be employed in attempts to evaluate both the density (via the absorbance of the respective bands resulting from the interaction of the probes with the acid sites) and the strength of the respective sites (see below).

Kazansky (sometimes transcribed as Kazanskii) and coworkers [127–129] and Beck and Pfeifer [130] undertook pioneering work on the application of light paraffins and dihydrogen as probe molecules. In their experiments, diffuse reflectance IR spectroscopy was employed, which has several advantages: high sensitivity, extension into the region of overtones and combination modes (cf. Table 4.3 in Ref. [131]) and use of zeolite powder instead of pressed wafers. In contrast to diffuse reflectance spectroscopy in the UV-visible range, however, a drawback of this technique lies in the difficulties in obtaining quantitative results. As an example, Fig. 7 (adopted from Ref. [132]) displays the spectrum of H2 adsorbed on OH groups of H-ZSM-5 and H-Y. The utilization of two other frequently used probes, benzene and carbon monoxide, for the characterization of the Brønsted acid strength by IR spectroscopy is illustrated in Fig. 8, Table 3 and Fig. 9. In both cases the wavenumber shift, ˜νOH , of the original OH band, which indicated the Brønsted acid sites, was adopted as an approximate measure of the acid strength. More recently, Kazansky and coworkers [133, 134] suggested characterizing the strength of acid (active) sites not through the shift upon interaction with probe molecules but by the intensities of the correlated bands. Note, however, that the use of probes can hardly provide more than a ranking of members of a homologous series of acid solids, since the interaction of the probe molecules with the acid sites is a dynamic phenomenon which probably modifies the original acid strength and its distribution. Thus, when probes are involved one cannot expect to determine the ‘‘intrinsic’’ acid (or basic) strength. References see page 1118

1108 Tab. 3

3.2 Chemical Properties Characterization of the strength of hydroxyls in zeolites and SAPOs by CO adsorption [118]

Parameter

Type of OH Si(OH)

ν˜ OH /cm−1 ˜νOH /cm−1 Origin

Al(OH)

3748 90 Terminal or defects

Diffuse remission / A.U.

4060

2.6

4105

4125 d c b a

2.5

2.4

2.3

2.2

Wavelength/µm

Diffuse reflectance spectra of hydrogen adsorbed on H-ZSM-5 at 77 K: a, b and c, after heat-pretreatment in high vacuum at 770, 970 and 1220 K, respectively; d, on H-Y after ‘‘deep-bed treatment’’ at 770 K. Abscissa linear in wavelength; important bands indicated in wavenumbers (cm−1 ) [132].

Transmittance

Fig. 7

∆v∼ 4000

3500

Si(OH)Al

3677 175–195 Terminal or defects

3650–3615 237–332 Framework; bridging

angle spinning (MAS) NMR spectroscopy (cf. the work of Brunner [135] and Sauer and coworkers [136]).

Figures indicate wavenumbers / cm−1

4035 4010

3700 3660 140 195 Non-framework; silica–alumina

P(OH)

3000

Wavenumber/cm−1

Wavenumber shift, ˜ν , of the high-frequency OH band (at 3640 cm−1 ) of H-Y zeolite on adsorption of benzene: dotted curve, spectrum before adsorption; solid curve, spectrum after adsorption of benzene.

Fig. 8

A more reliable measure of the strength of Brønsted acid sites is provided by the deprotonation energy, Ep (see above), which may be calculated via quantum chemical methods (cf. the work of Sauer and coworkers, e.g. Refs. [50, 51]) or experimentally determined by magic

D NMR Spectroscopic Characterization of Brønsted Acidity with and without Probe Molecules Similarly to the IR method, Brønsted acid sites may be easily identified as such, i.e. without probe molecules, through 1 H MAS NMR spectroscopy. In their classical papers, Pfeifer and coworkers [57, 137, 138] (see also Ref. [139]) detected that acidic OH groups give rise to characteristic chemical shifts, e.g. of δ = 1.8 (line a), 3.9–4.6 (line b), 4.8–5.6 (line c) and 7.0–7.5 (line d) ppm [referenced to tetramethylsilane, TMS, therefore usually indicated as δ(TMS)] for the nonacidic OHs (silanol groups, corresponding to IR bands at about 3740 cm−1 ), so-called high-frequency (HF) OH groups in the large cavities of H-Y (corresponding to the IR band at 3640 cm−1 ), low-frequency (LF) OH groups in the smaller β-cages (corresponding to the IR band at 3550 cm−1 ) and to residual NH+ 4 cations (cf. also Ref. [99] and Fig. 10). The OH groups in H-ZSM-5 give rise to chemical shifts of δ = 2.0, 4.2 and 7.0 ppm (cf. Fig. 11), corresponding to IR bands of silanol groups, Si−OH, at 3740 cm−1 , bridging OH groups, [≡Si−OH−Al≡], at ca. 3614 cm−1 and bridging OH groups with an additional electrostatic interaction to the framework at 3250 cm−1 , respectively [139, 141]. The weakly acidic or non-acidic silanol groups (corresponding to IR bands around 3740–3720 cm−1 ) generally produce a signal at about δ = 2.0 ppm [135, 139–141]. A more detailed list with lines and corresponding IR bands of various species is given in Ref. [141] (see also below, Fig. 13). With respect to the quantitative spectroscopic characterization of zeolitic systems, MAS NMR [98] is usually in a better situation than IR, because the former technique provides more directly quantitative results in that only a comparison with a standard is required. Unlike quantitative IR determinations, no extinction coefficients (see above) need to be determined in separate experiments. However, when reliable IR extinction coefficients were available and the pretreatment of the samples was identical, good agreement between the results obtained by IR and 1 H MAS NMR spectroscopy was observed (Fig. 12).

Absorbance

1 2

1109

3383

3413

3650

3.2.4 Acidity and Basicity

6′ 5′

0.1

4′ 3 0.1

0.05

3′ 2′

4 5 6

3800

2 1 3600

3400

3200

Wavenumber/cm−1

IR spectra of OH groups of H,Na-Y zeolite before (1) and, after (2–6) adsorption of increasing amounts of CO and, at lower wavenumbers, the bands shifted according to the strength of the OH groups involved, viz. to 3383 → 3413 cm−1 , which indicates a decrease from 2 to 6 [118].

Fig. 9

(b) (c)

(d)

(a)

(a)

10 (b)

5

0

dH/ppm

(a) 1 H MAS NMR spectrum of H-Y. (b) Decomposition into the four lines a, b, c and d, corresponding to non-acid OHs (silanol groups, IR bands at 3740 cm−1 ), acid OH groups in the supercages (HF IR band at 3640 cm−1 ), acid OH groups in the small cages (β-cages, LF IR band at 3550 cm−1 ) and OH groups with additional electrostatic interaction with oxygen atoms of the framework (IR band around 3250 cm−1 ), respectively (see text and Ref. [135]).

Fig. 10

Moreover, in the hydrogen form of a classical zeolite, i.e. in an aluminosilicate, a 27 Al MAS NMR signal at δ = 60 ppm [referenced to Al(H2 O)3+ 6 , i.e. aqueous solution of Al(NO3 )3 ] indicates the presence of Brønsted acid sites (cf., e.g., Ref. [139]). This signal is due to tetrahedrally coordinated framework aluminum and, therefore, corresponds to structural acid OH groups (see top of Scheme 3). Thus, from the area of the signal at δ = 60 ppm the number of Brønsted sites can also be determined. Pfeifer and coworkers suggested using the chemical shifts of the OH group directly as a measure of Brønsted acid strength [57, 137, 138]. Basically, the shifts should increase with increasing strength. Exceptions are the shifts of OH groups, which are affected by additional electrostatic interactions with adjacent oxygen atoms of the framework (see e.g., the above-mentioned lines c and d of H-Y or the OH groups located in channels of ferrierite (FER) formed by eight-membered rings compared with those in the wider channels built by 10-membered rings [94]; see also below). The accuracy of the method could be significantly improved by increasing the differences in the chemical shifts by adsorption of CO or halogenated hydrocarbons such as C2 Cl4 , which induced additional chemical shifts (δOH ) and thus increased the resolution (see e.g. Table 4). Both the total chemical shifts (δOH + δOH ) and the induced chemical shifts (δOH ) were suggested as an appropriate measure of acidity strength, since they correlate well with the induced wavenumber shifts, References see page 1118

1110

3.2 Chemical Properties Tab. 4 1 H MAS NMR original chemical shifts of OH groups and shifts induced on CO adsorption [141, 150]

4.2 ppm

Zeolite

Hydroxyl group structure

δH / ppm

δH···CO / ppm

δ/ ppm

Si−OH (silanol groups) Al−OH (on nonframework Al)

2.0

2.0

0.0

2.9

3.9

1.0

Si−(OH)−Al (free) Si−(OH)−Al (in the large cavities)

4.2

6.2

2.0

3.9

4.8

0.9

2.0 ppm 7.0 ppm

H-ZSM-5 H-ZSM-5 hydrothermally treated

20

10

0

−10

0.3 H,Na-Y

−20

dH /ppm 1 H MAS NMR spectrum of H-ZSM-5 measured at 123 K with a sample spinning rate of 6 kHz. The asterisks denote spinning sidebands. The lines at 2.0, 4.2 and 7 ppm correspond to Si−OH groups at the external surface of the crystallites or framework defect sites (IR band at 3740 cm−1 ), bridging hydroxyl groups (Brønsted centers, IR band at 3614 cm−1 ) and OH groups affected by additional interaction with the zeolite framework (broad IR band at about 3250 cm−1 ) [141].

Fig. 11

˜νOH (see above; cf., e.g., Refs. [118, 119] and, for δOH , Ref. [141]). Simultaneously, it turned out that upon loading with the adsorbates the OH groups indicated by the line at about δOH = 7 ppm behaved like free bridging OH groups, i.e. the action of, for instance, CO molecules on those OH groups annihilated their electrostatic interaction with adjacent framework oxygen atoms [141]. Fleischer et al. [136] have shown that under restricted conditions [only oxygen atoms surrounding the T atom (T = Al, Si, P, etc.) bearing the bridging OH group], a linear relationship exists between the deprotonation

energy (EDP ) as the measure of acid strength (see above) and the chemical shifts related to the acid OH groups, δOH . Later, Brunner and coworkers [141, 142] provided evidence for a corresponding linear relationship between the wavenumbers of Brønsted acid OH groups, ν˜ OH , and the chemical shifts, δOH (TMS) (cf. Fig. 13 and Refs. [141, 142]). From the proportionalities EDP ∝ δOH and δOH ∝ ν˜ OH , it was several times inferred that a relationship also exists between the Brønsted acid strength and the wavenumbers ν˜ OH of the respective hydroxyl groups (cf., e.g. Refs. [143, 144]); however, more experimental and theoretical work seems to be required to establish this firmly. For MAS NMR investigations a number of probe molecules have also been suggested as being suitable for the investigation of acid sites (cf., e.g., Refs. [146–148]). Examples of probes frequently employed in NMR investigations of acid sites are water, alcohols, benzene, acetone, acetylacetone and phosphines. In several cases 13 CO and

3670 3580 cm−1

n (acidic OH's) = 48.8 OH/u.c.

B

Bridging acidic OH's c b +

Transmittance

NH4

OH (non acidic)

d

505 K 565 K 660 K IR-Spectroscopy

n (acidic OH's) = 49.2 OH/u.c.

a

7 4.5 2 x = 90% (90H, Na-Y) dH / ppm Tactivation: 565 K 575 K 1H

MAS NMR Spectroscopy

Comparison of the determination of the density of acid OHs in H-Y (degree of exchange: x = 90%) by IR spectroscopy and 1 H MAS NMR spectroscopy under almost identical conditions [88, 137].

Fig. 12

3.2.4 Acidity and Basicity

15

H-Y

SiOH SiOHAI (HF) SiOHAI (LF) H-ZSM-5 SiOH SiOHAI H-ZSM-5 SiOH (deal.) AIOH (non-framework Al) SiOHAI H-SABO SiOH SiOHB SiOHAI SAPO-5 SiOHAI (I) SiOHAI (II) SAPO-11 SiOHAI (I) SiOHAI (II) Mg-Y Mg2+OH− Ca-Y Ca2+OH−

3

5

5

8

2

dH / ppm

4

13 11 14

3

10

12

9 1

7

2 4

6

1 16

17

0

3500

3600

∼

VOH

3700

1111

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17

3800

/cm−1

Dependence of the 1 H chemical shift, δ H , of surface hydroxyls in zeolites on the wavenumber of OH stretching vibrations, [142].

Fig. 13

ν˜ OH

15 N-containing

Si

Si OH

Al

O

C

OH

C O

Al (I)

(II)

complex (II) is the stable one [150]. Fraissard and coworkers [113, 114] suggested the use of H2 O as a probe molecule for zeolitic acidity in order to evaluate the hydroxonium ion concentration (i.e. their number, N ) related to the original bridging OH groups (Brønsted acid centers), i.e. N(H3 O)+/Ninit. (Si−OH−Al), in analogy with the pH scale (negative logarithm of the proton activity in aqueous solutions). The ratio N(H3 O)+/Ninit. (Si−OH−Al) was evaluated from broad line NMR spectra as a measure of the acid strength. These spectra were obtained at 4 K. Their analysis was based on the assumption that only the following entities coexist: (i) isolated OHs; (ii) isolated H2 O molecules; (iii) H2 O molecules attached to OH groups via hydrogen bonds; and (iv) hydroxonium ions, H3 O+ . Evidence was offered that, up to room temperature, the temperature effect on the equilibrium is negligible. In the case of non-dealuminated H-Y, it was claimed that an equilibrium, i.e. N(H3 O+ )/Ninit. (Si−OH−Al) = constant, was reached for one H2 O molecule per initial (Si−OH−Al). With dealuminated H-Y, a second

0.5 N(H3O+)/N(Initial OHb)

ammonia, pyridine, amines or acetonitrile were used and 13 C MAS NMR [139, 147–150] and 15 N MAS NMR [145, 151], respectively, were employed. It turned out, perhaps surprisingly, that from two possible complexes of Brønsted acid sites with CO, viz.

0.4 0.3

Dealuminated H-Y H-Y

0.2 0.1 0.0

0

1

2

3

4

5

6

7

N(H2ADS.)/N(Initial OHb)

Formation of hydroxonium ions, H3 O+ , as a function of loading of non-dealuminated H-Y zeolite () and dealuminated H-Y (
) with H2 O [113, 114].

Fig. 14

increase in N(H3 O)+ /Ninit. (Si−OH−Al) was observed for N(H2 O)/Ninit. (Si−OH−Al) > 2 (Fig. 14). The latter finding was explained by a reaction of H2 O with defect sites. Note, however, that according to quantum chemical calculations by Sauer and coworkers [152, 153], an isolated, i.e. non-hydrated, hydroxonium ion as the cation participating in an ion pair such as [≡Si−O −Al − ≡(O3 )−Si]− · · · H3 O+ is in aqueous solutions not a stable but a transition structure. Its energy corresponds to a saddle point separating two stable hydrogen-bonded neutral structures, [≡Si−OH−Al≡(O3 )−Si] · · · H2 O, by only 10 kJ mol−1 (see Fig. 15). References see page 1118

1112

3.2 Chemical Properties

Neutral complex O O Si

Al O

O

H

Si

Si

O

Al −

O

O

H + H O

H

H

H

Neutral complex O O

Ion pair O O Si Si

Al O

O

H

Si

H O H

∆Eads ∆EPT Fig. 15 Sketch of the potential energy surface (PES) for the proton transfer from a Brønsted acid site on to adsorbed H2 O; Eads ≈ 50 kJ mol−1 , energy change upon adsorption; EPT = 10 kJ mol−1 , proton transfer energy, i.e. energy of the saddle point [153].

E IR Spectroscopic Characterization of Lewis Acidity with Unlike Brønsted acid centers (see Probe Molecules above), ‘‘true’’ Lewis sites (L-sites) and Lewis acid cations (C-sites) are not directly observable by IR spectroscopy but only with the help of probe molecules. Most of the probes used for the determination of Brønsted acid sites (e.g. ammonia, pyridine, dihydrogen, carbon monoxide, nitriles, nitric oxide) may also be employed for the investigation of Lewis acidity. Assignments are made through the characteristic shifts of the wavenumbers. The quantitative evaluation of the respective adsorbate bands is analogous to that of bands of OH groups or probes associated with Brønsted acid centers (see cf. above). An application of ammonia in the case of Lewis sites produces bands at around 1640 cm−1 (cf., e.g., Refs. [154, 155]). Pyridine attached to ‘‘true’’ Lewis sites (L-sites) and cations (C-sites) gives rise to very sharp bands of ring deformation vibrations around, e.g., 1452 and 1440–1451 cm−1 , respectively (see, for instance, the IR spectrum in Fig. 5 and Refs. [89, 90, 97]). The position of the latter bands indicating C-sites depends on the Coulomb field of the cation involved [156]. Dihydrogen was used by Kazansky [132], who claimed that also Lewis sites in the form of only threefold-coordinated ≡Al and ≡Si+ were indicated by bands at 4010 and 4035 cm−1 , respectively. Carbon monoxide was studied by Angell and Schaffer [157] and Manoilova et al. [158]. Nitric oxide (NO) was especially utilized in cases of transition metal cations, for instance by Che et al. [159] and Ziolek et al. [160]; the latter authors studied Cu2+ -containing ZSM-5 and also mesoporous materials (Cu−[Al]MCM-41). On the other hand, MAS NMR is able to detect and measure directly Lewis acid sites in the form of

extra-framework Al-containing species (AlEF ). In partially dehydroxylated or dealuminated zeolites, new lines appear besides the signal at about 60 ppm [referenced against Al(H2 O)3+ 6 , i.e. an aqueous solution of Al(NO3 )3 ], which is ascribed to tetrahedrally coordinated framework aluminum, AlF (see above). Octahedrally coordinated AlEF species (Lewis sites) manifest themselves by a signal around 0 ppm, when referenced against Al(H2 O)3+ 6 . In addition, sometimes a line at 30 ppm was observed and ascribed to pentacoordinated AlEF species (cf., e.g., Refs. [161–163]). Even the occurrence of fourfoldcoordinated extra-framework aluminum was claimed. In any event, for quantitative measurements care must be taken to ensure that the whole amount of aluminum present in the sample is ‘‘MAS NMR visible’’ [163, 164]. MAS NMR may also utilize probes to investigate Lewis acidity. Thus, 13 CO was employed in 13 C MAS NMR studies [146, 148, 165]. Acetonitrile adsorbed as a probe on cation-containing faujasite-type zeolites was studied by 15 N MAS NMR (cf., e.g., Refs. [145, 166]). Another probe molecule sometimes used for characterizing Lewis acid centers is trimethyl phosphine, where 31 P MAS NMR is applied [167, 168]. Typical measured chemical shifts lie in the range 2–22 ppm referenced against liquid trimethylphosphine. F ESR Spectroscopic Characterization of Lewis Acidity with NO as a Probe Karge and coworkers developed a method which applies selectively to the identification and characterization of Lewis acid sites, using ESR spectroscopy with NO as a probe [169, 170]. The NO molecule possesses an unpaired electron but usually does not give rise to an ESR spectrum because its orbital momentum just compensates the electron spin. However, the NO molecule does produce a very typical ESR spectrum with a characteristic hyperfine splitting when adsorbed on ‘‘true’’ Lewis sites (see Fig. 16), but a different one upon interaction with cations. The parameters of the experimantal spectrum are gxx ≈ gyy = 1.997, gzz = 1.950, with hyperfine splitting constants Axx = Ayy = 45 MHz, Azz unresolved. NO adsorption on, for instance, Na+ causes a different ESR signal with gII = 1.840 and Axx = Ayy = Azz = 0 MHz [171]. No ESR signal similar to that in Fig. 16 is observed upon interaction of NO with Brønsted acid sites (H-ZSM-5). Its appearance in the case of Lewis centers is explained by the quenching of the orbital momentum, equal in value but opposite in sign to the spin, when NO interacts with a charged Lewis site [171]. G Spectroscopic Characterization of Basicity As in the case of Lewis acid sites, the basic sites are usually identified with the help of probe molecules.

1113

NO/H-ZSM-5

Tact p (NO) Tads τads Tmeas

= = = = =

∆H = 10 mT

1075 K 50 Pa 298 K 1h 77 K

Transmittance

3.2.4 Acidity and Basicity

v∼ [NH] v∼ [CH] ?

Base line Tact. = 775 K (15 h)

?

Comb. bands : Pyrrole Sorbate Tad. Tdes. : R.T Evacuation : 1 h

∆v∼

3500

3000

2500

2000

Wavenumber/cm−1 Fig. 17

NO/H-ZSM-5 computer - smulated spectrum

Fig. 16 ESR spectrum of NO adsorbed on H-ZSM-5 containing ‘‘true’’ Lewis sites and comparison with the computer-simulated spectrum [170].

Unfortunately, to date there are only a few probes which perhaps can be successfully used [63, 172–175]. Carbon dioxide has been employed several times, but it may suffer from the drawback that it can form various carbonates when interacting with zeolites. Other probes suggested for the characterization are acetic acid, boric acid trimethyl ester, acetylene derivatives, halogenated paraffins (e.g. deuterochloroform, Cl3 CD [173]) benzene and pyrrole [172]. In some instances, benzene was used for probing basic oxygen sites in the 12-membered ring windows of faujasite-type zeolites, where a benzene molecule can weakly interact with basic oxygen atoms through its hydrogen atoms. It was suggested that the shift of the CH deformation band provides a measure of the basic strength [173]. Similarly, the shift of the IR band due to the NH vibration of the amphoteric probe molecule pyrrole was used for characterizing the basicity of, e.g., zeolites containing cations of alkali metals. Figure 17 shows as an example the spectrum and the shift ˜ν after adsorption of pyrrole on KLTL [63, 174, 175]. This approach, advanced by de Mallman and Barthomeuf, was supported by the finding that the thus determined basicity parallels the value of the partial negative charges on the basic oxygen sites.

Adsorption of pyrrole on basic K-LTL zeolite [63].

H Temperature-programmed Desorption (TPD) of Probe Molecules from Acidic (or Basic) Sites Temperatureprogrammed desorption of bases such as ammonia, amines and pyridine is a popular method for characterizing the acid strength of the sites from which the probe molecules were desorbed. Analogously, desorption of weak acids may be tested for characterizing the strength of basic sites. The desorption may be monitored by gas chromatography (GC) or mass spectrometry (MS). In particular, when GC is employed one has a tool to measure simultaneously via the amount of the desorbed probes the density of sites (to the best of our knowledge, examples have not yet been reported illustrating the characterization of the strength of basic sites by appropriate TPD experiments). However, there are some possible pitfalls with the TPD technique. First, TPD is not selective, i.e. one cannot decide whether the probes are desorbing from Brønsted or Lewis acid sites when both types are present. Therefore, in order to characterize the strength of acidic sites, it is advisable to combine TPD with an independent technique, e.g. simultaneous insitu IR spectroscopy. Observation of the developing TPD peaks and the simultaneous decrease in the IR bands indicating the Brønsted and/or Lewis acid sites permits an assignment of the TPD peaks to the respective types of sites [154, 155]. Figure 18 shows the desorption of NH3 from mordenite monitored simultaneously by MS and IR spectroscopy. Second, the TPD results may be corrupted by readsorption of the species desorbed from acid sites. Here, the use of thin samples or sample layers may be helpful [176, 177]. Models have been developed that take into account the kinetics of the desorption and the possibility of a distribution of acid strength. They enable one to derive the number, n, of types of sites and evaluate References see page 1118

1114

3.2 Chemical Properties

p (NH3)

results of other, independent techniques; an example is displayed in Table 5. ESR-monitored TPD of NO adsorbed on ‘‘true’’ Lewis sites provided for the first time the possibility of measuring selectively the acid strength of true Lewis sites [170].

TPD / MS Tact. = 673 K

568 363

1.00 Relative absorbance

TPD / IR 0.75

a

IR bands

353

a : 1450 cm−1 (Bronsted)

553

0.50

b

b : 1620 cm−1 (cations)

0.25

0.00 300

400

500

600

700

T/K

Temperature of desorption

Temperature-programmed desorption of NH3 from BeMOR simultaneously measured by MS [increase in p(NH3 )] and IR spectroscopy (decrease in the absorbances at 1450 and 1620 cm−1 ), due to desorption from Brønsted acid sites and Be cations, respectively. The maxima in the upper part of the figure correspond to the steepest decrease in the respective absorbance [154, 155].

Fig. 18

the activation energies of desorption, En . The latter may provide a quantitative measure of the acid strength and at least lead to a reliable ranking of the members of a sample series with respect to their acidity strength [177, 178] (see Table 5). Even though TPD is a method the application of which must often be treated with caution, there are many instances where it led to a ranking of acid zeolites with respect to acid strength in good agreement with the

I Microcalorimetric Measurement of the Strength of Acidic In microcalorimetric investigations, Sites in Zeolites the differential heat of adsorption, Qdiff , of basic probe molecules is used as a measure of the strength of the acid sites. Again, the method is non-selective in that one has to employ independent techniques in order to recognize whether the interaction occurs on, e.g., ‘‘true’’ Lewis sites, Brønsted acid sites or cations if various types of sites exist in the zeolitic adsorbent. However, once this problem has been solved, microcalorimetric measurements, when properly conducted, seem to be a very reliable tool for characterizing the strength and homogeneity or heterogeneity of the acidic internal and external surface of zeolites (cf., e.g., Refs. [179–182]). Moreover, the amount of basic probe molecules consumed for neutralization of the acid sites measures directly their density in the adsorbent sample. Figure 19 compares the results of microcalorimetric measurements of NH3 adsorption on a heteropoly acid, which was used as a standard and exhibited a high degree of homogeneity (constant Qdiff until complete neutralization), with an almost homogeneous sample of H-ZSM-5 and an La,Na-Y zeolite with a rather broad distribution of acid strength. J Catalysis and Test Reactions Since the advent of zeolites in research on and application of heterogeneous [183, 184] catalysis by about 1960, a huge body

Ranking of the acid strength of Brønsted sites of three H-ZSM samples through various techniquesa

Tab. 5

TPD/MS (NH3 )

H-ZSM-5 Tact = 673 K

Sample A Sample B Sample C

C6 H6 adsorption IR spectroscopy

Microcalorimetry NH3 adsorption

Tdes (max.)/K

¯ des /kJ mol−1 E

˜νOH /cm−1

Qdiff /kJ mol−1

655 619 620

110 104 103

359 337 340

157 147 145

activation temperature; Tdes (max .), maximum temperature of desorption; E¯ des , most frequently occurring energy of desorption; ˜νOH , wavenumber shift upon adsorption of benzene; Qdiff , differential heat of NH3 adsorption (see text). Sample A: H4.0 Na0.4 [Al4.4 Si91.6 O192 ]; Si/Al (total) = 21; AlF = 2.5; AlNF = 1.9. Sample B: H2.4 Na0.4 [Al2.8 Si93.2 O192 ]; Si/Al (total) = 33; AlF = 2.1; AlNF = 0.7. Sample C: H2.6 Na0.2 [Al2.8 Si93.2 O192 ]; Si/Al (total) = 33; AlF = 2.8; AlNF = none. AlF = framework Al; AlNF = non-framework Al. aT , act

3.2.4 Acidity and Basicity

Differencial heat of adsorption Q (diff) / kJ mol−1

200

H3PW12O40,Tact. = 423 K H-ZSM-5, Tact. = 675 K LaNa-Y,

Tact. = 675 K

150

100

50 0

1

2

3

Adsorbed amount / mmol g−1

Differential heat of adsorption of NH3 on H3 [PW12 O40 ] (for calibration of the microcalorimetric equipment), H-ZSM-5 and La,Na-Y as a function of the amount adsorbed measured via microcalorimetry [182].

Fig. 19

of literature has been accumulated in this field. Very informative reviews were provided by, inter alia, Venuto and Landis [185], Venuto [186], Poutsma [187], Csicsery [188], H¨olderich [189], Haag [190] and Weitkamp and Puppe [191]. With respect to catalysis of inorganic reactions, there was and still is remarkable research activity on zeolites in environmental catalysis (for example, in decomposition of NO, selective reduction of NO by NH3 , etc.) where, however, the catalytically relevant centers are formed by exchange cations such as Cu2+ , Cu+ and Co2+ (cf., e.g., Refs. [192, 193]). For the Claus reaction (reaction of SO2 with H2 S to form water and elemental sulfur [194]) aluminas or bauxites are still the dominant catalysts [195] and in this catalysis basic Lewis sites are strongly involved, as has been shown by gas-phase titration with BF3 and monitored by IR spectroscopy [196]. The IR spectroscopic observations in the latter study [196] are in agreement with the model proposed by Kn¨ozinger and Ratnasamy [48]. Zeolites also catalyze the Claus reaction but are usually rapidly deactivated through deposition and pore blocking by sulfur [197]. Thus, the main domain of zeolite catalysis lies in the area of hydrocarbon reactions. Similarly to catalysis by other solids, e.g. oxides and metals, the question soon arose as to which type of sites are the catalytically relevant centers. In the case of oxides, it was found that Lewis acid sites, i.e. coordinatively unsaturated cations, play the most important role. As an example, the dehydration of alcohols over alumina should be mentioned (cf., e.g., Ref. [198]). Regarding metal catalysts, frequently certain surface positions of the metal atoms, for example in so-called kinks (HalbkristallLagen), on steps or corners were claimed to be active centers. With respect to zeolites, three types of sites

1115

were controversially discussed as loci of acid-catalyzed organic reactions, namely (i) extra-framework cations, which should exhibit a high electrostatic field with a strong carboniogenic potential; in particular, cations introduced by ion exchange and balancing the negative charge of the zeolite framework were assumed to be catalytically active centers (cf., e.g., Ref. [199]); (ii) Lewis acid sites in the form of only threefold-coordinated framework Al or Si (see above; cf., e.g., Refs. [200, 201]) or extra-framework Alx On+ y complexes, so-called ‘‘true’’ Lewis sites (see above; cf., e.g., Ref. [35]); and (iii) Brønsted acid sites in the form of proton-donating OH groups (see Section 3.2.4.1.2B). Regarding the extra-framework exchangeable cations, it soon became clear that in the exclusive presence of monovalent cations, i.e. when even traces of multivalent cations were carefully avoided, the respective zeolites did not catalyze hydrocarbon reactions such as cracking, alkylation, dealkylation and isomerization. In cases where multivalent cations were not excluded, acid OH groups form according to the Hirschler–Plank mechanism [see above; cf. Refs. [25, 26] and Eqs. (1) and (2)]. Correspondingly, when Karge and coworkers prepared lanthanum-containing zeolite catalysts via solid-state ion exchange and also carefully excluded even traces of water during the pretreatment and admission of the reactant (ethylbenzene), no catalytic activity was observed. Only when small doses of water vapor were intermittently admitted (1 min contact with 102 Pa water vapor), the reaction started to occur, because now the Hirschler–Plank mechanism became operative; simultaneously, on the zeolite in the IR transmittant flow-reactor cell, the typical OH bands at 3616 and 3518 cm−1 were observed [31]. Furthermore, the alternative ‘‘Lewis acid sites vs. Brønsted acid sites’’ was clarified in favor of the latter, since a close correlation was found between the catalytic activity and the density of Brønsted centers in, for instance, the alkylation of benzene by ethene or propene (cf. Ref. [89] and Fig. 20). Here, the decrease in the activity (measured via the conversion of benzene) upon dehydroxylation (see above; see, for instance, Scheme 3) followed exactly the decrease in the density of the acid OH groups as determined by IR spectroscopy with and without the probe pyridine. At higher dehydroxylation temperatures, the density of Lewis sites as determined via IR spectroscopy with pyridine as a probe is still high but the activity is zero. It is not yet entirely clear why the density of Lewis sites decreases at all when the pretreatment temperature is above 450 ◦ C. Most likely this is due to an agglomeration of AlO+ to bulkier Alx On+ y complexes. Similar relations between the concentration of Brønsted centers and catalytic activity have been observed in many other cases of acid-catalyzed hydrocarbon reactions. References see page 1118

3.2 Chemical Properties

1.0

d (Py.-Lewis sites)

:

X and d

0.8 0.6

X:

0.4 0.2

300

d:

Conversion of benzene upon alkylation with ethylene; normalized to Tact. = 450 °C

d (Py.-Bronsted sites) : : d (OH groups)

Density of sites; measured via A(OH), A(PyB) and A(PyL); normalized to Tact. = 450 °C

Rel. catalytic activity

1116

100

10

1

0.1 10

400 500 600 700 Pretreatment temperature,Tact. / °C

100

1000

10000

Al-content / ppm

Linear relationship between the n-hexane cracking activity over and the Al content (measuring the density of acid OH groups) in H-ZSM-5 [190].

Comparison of activity in benzene alkylation by ethylene and acid site density of H-mordenite as a function of pretreatment temperature [89].

Fig. 21

It is not definitively clarified whether or not the Lewis sites play any role in heterogeneous catalysis on zeolites. Several authors claim that they are relevant in coke formation (cf., e.g., Ref. [202]) and it seems very likely, however, that Lewis sites are able to increase the acid strength of adjacent Brønsted OH centers through a quasi-inductive effect, which lowers the deprotonation energy of the OH groups. Moreover, it was shown that Lewis acid sites are able to form carbenium ions such as triphenylmethyl cations (CPh+ 3 ) [123] and catalyze the Meerwein–Ponndorf–Verley reaction and its reverse, the Oppenauer oxidation [203]. Several times it has been reported that such correlations between the density of acid sites of a zeolite sample and the conversion or rate of a reaction catalyzed by these sites may be used as tests and means of determination of the site density. A very impressive example was provided by the work of Haag and coworkers [39, 190], who used n-hexane cracking (alpha test) over H-ZSM-5 as a test reaction (cf. Fig. 21). However, ‘‘mild steaming’’ (up to 13.3 kPa for 2.5 h at 813 K) increased the activity to values much above the linear correlation curve [190]. The authors proposed that mild steaming creates sites of enhanced activity and that these sites are formed only from Al atoms in close proximity, such as paired Al centers, which explains the strong dependence of the effect on the Al content [190]. A scheme for the formation of this type of site was proposed [39, 190] (see also Ref. [204] and Scheme 4, where state IV represents a site of enhanced activity). On more severe treatment this site is transformed to a conventional ‘‘true’’ Lewis site (see above; cf. Refs. [33, 34]). Similarly, a linear relationship was measured between the conversion in the disproportionation of ethylbenzene and the density of Brønsted acid sites in a series of

mordenites or Y-type samples as measured through the absorbance of the respective IR bands of the acid OH groups or pyridinium ion bands (cf. Fig. 22 and Refs. [72, 143, 205]). This reaction, when properly conducted, does not suffer from deactivation through coke formation [72]. The fact that the straight line in Fig. 22 does not go through the origin indicates that a minimum strength of the Brønsted acid site is required to operate as catalytically active centers, which is not fulfilled in the case of the barium-modified mordenites. Sigl et al. [206] employed the same reaction to check the acid strength of [Al]-, [Ga]and [Fe]-H-ZSM-5 samples with almost exactly identical Si/Al ratios (about 25) under standardized conditions. In agreement with the results of various other tests, it turned out that the level of quasi-stationary conversion 10

Conversion, X / %

Fig. 20

HM-D

8

BeM-D

6

MgM-D

4 CaM-D

2

SrM-D BaM-D

0.1

0.2

0.3

0.4

0.5

0.6

Absorbance A(OH) ∝ C(Bronstedt) Fig. 22 Steady-state conversion of ethylbenzene disproportionation over dealuminated mordenites as a function of the density of acidic OH groups as measured by the absorbance of the OH stretching band around 3600 cm−1 . Conditions: Tactiv. = 673 K, pactiv. = 10−4 Pa, Treact. = 448 K, p(EB) = 1.33 × 103 Pa, mass of catalyst, m = 0.25 g, absorbances A(OH) normalized to equal sample thickness of 10 mg cm−2 [143].

3.2.4 Acidity and Basicity

H O Si

H

O Al

O Si

I

Al O

O

Si

Si

d− O Al O δ+ O O Si Al Si

− H2O

O Si

H

IV

H

O

O

+ 2H2O

Si

Si

H

O Al

O Si

II

H

H

H

O

O

O Si

Si

− H2O

O Si

O

O

H H

H H

O

O

Si

Si

Si

H O

O

+

Si

H

H

H

H

O

O

O

O

Si

Si

Si

Si

III

O Al

H

O

Al d− d+ O O O Si Al Si

1117

Si

[AlO]+ O Si

V

Scheme 4

−

Al

O

O Si

O Si

H

Si

H

O

O

Si

Si

Formation of acid sites of enhanced activity in H-ZSM-5 after mild steaming.

reflected the acid strength decreasing in the sequence [Al]-H-ZSM-5 > [Ga]-H-ZSM-5 > [Fe]-H-ZSM-5. Several other test reactions have been proposed, such as conversion of toluene, xylene isomerization [207] and dehydration of cyclohexanol. The last reaction was studied by Karge et al. [111] inter alia over mordenites and clinoptilolites as catalysts with different acid strength, which was systematically varied by appropriate cation exchange. Using 2,6-di-tert-butylpyridine as a probe for IR spectroscopic investigation, it was demonstrated that the reaction occurred in the case of the clinoptilolites only on the external surface. Similarly to what was observed with the disproportionation of ethylbenzene, the reaction rate depended linearly on the density of the Brønsted acid sites. Also, series of test reactions (cf. Table 6) requiring different activation energies or reaction temperatures under otherwise standardized conditions were suggested by Guisnet and coworkers [208] in order to measure the different acid strengths of the zeolites employed to catalyze these reactions (Table 6).

As test reactions for the basicity of zeolites, inter alia the already mentioned Knoevenagel condensation of benzaldehyde with malononitrile [173], side-chain alkylation of alkylbenzenes, conversion of 2-propanol with formation of acetone and of methanol to carbon monoxide have been suggested and to some extent successfully employed. 3.2.4.1.4 Conclusion The Brønsted acidity of oxides and, in particular, of zeolites and its role in heterogeneous catalysis has been extensively studied and seems to be relatively well understood. Its qualitative identification is largely routine work and in most cases reliable. Also, a large number of techniques for the quantitative determination of the density of Brønsted and Lewis acid sites and their strength is available. However, the methods for characterization of Brønsted and Lewis acid sites usually allow only a ranking of the strengths within a series of similar acid materials under standardized References see page 1118

1118

3.2 Chemical Properties

Characterization of Brønsted acid strength of zeolite catalysts through various reactions [208]

Tab. 6

Reactant 3,3-Dimethyl1-butene Cyclohexene

(3,3-DMB-1) (c Hexe)

2,2,4Trimethylpentane 2,4Dimethylpentane 2-Methylpentane

(2,2,4-TMP)

(2-MP)

n-Hexane

(n-Hx)

o-Xylene

(o-X)

1,2,4Trimethylbenzene

(1,2,4-TMB)

(2,4-TMP)

Reaction

TR a

Skeletal isomerization Skeletal isomerization, hydrogen transferb Cracking

200

Isomerization, cracking Isomerization, cracking Isomerization, cracking Isomerization, disproportionationb Isomerization, disproportionationb

350

200

350

400 400 350 350

a T : reaction temperature/◦ C. R b Bimolecular reaction.

conditions. Here, further refinement of the techniques and the development of additional methods are certainly desirable. Advances towards a deeper understanding of Lewis acidity, especially of its specific nature and possible role in heterogeneous catalysis, remain a challenge for further research. References 1. H. M. Hey, Trans. Br. Ceram. Soc. 1937, 36, 84. 2. D. W. Breck, Zeolite Molecular Sieves – Structure, Chemistry and Use, Wiley, New York, 1974, Chapter II, pp. 29–185; reprinted by Krieger, Malabar, FL, 1984. 3. E. M. Flanigen, in Introduction to Zeolite Science and Practice, 2nd Ed., H. van Bekkum, E. M. Flanigen, P. A. Jacobs, J. C. Jansen (Eds.), Studies in Surface Science and Catalysis, Vol. 137, Elsevier, Amsterdam, 2001, Chapter 2, pp. 11–35. 4. J. C. Vartuli, W. J. Roth, J. S. Beck, S. B. McCullen, C. T. Kresge, in Molecular Sieves – Science and Technology, Vol. 1, H. G. Karge, J. Weitkamp (Eds.), Springer-Verlag, Berlin, 1998, Chapter 4, pp. 97–119. 5. D. Barthomeuf, in Molecular Sieves – II; Proceedings of the 4th International Zeolite Conference Chicago, IL, 1977, J. Katzer (Ed.), ACS Symposium Series, Vol. 40, American Chemical Society, Washington, DC, 1977, pp. 453–472. 6. P. A. Jacobs, Carboniogenic Activity of Zeolites, Elsevier, Amsterdam, 1977, Chapters III and IV, pp. 33–181. 7. H. A. Benesi, B. H. C. Winquist, in Advances in Catalysis, D. D. Eley, H. Pines, P. B. Weisz (Eds.), Vol. 27, Academic Press, New York, 1978, pp. 97–181. 8. W. E. Farneth, R. J. Gorte, Chem. Rev. 1995, 95, 615. 9. R. G. Pearson, J. Am. Chem. Soc. 1988, 110, 7684. 10. G. Klopman, Chemical Reactivity and Reaction Path, MIR, Moscow, 1977.

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170. F. Witzel, H. G. Karge, A. Gutsze, in Proceedings of the 9th International Zeolite Conference, R. von Ballmoos, J. B. Higgins, M. M. J. Treacy (Eds.), Montreal, Canada, 1992, Vol II, Butterworth-Heinemann, Boston, 1993, p. 283. 171. A. Gutsze, M. Plato, H. G. Karge, F. Witzel, J. Chem. Soc., Faraday Trans. 1 1996, 92, 2495. 172. H. G. Karge, E. Geidel, in Molecular Sieves – Science and Technology, H. G. Karge, J. Weitkamp (Eds.), Vol. 4, SpringerVerlag, Berlin, 2003, p. 1, especially pp. 147–149. 173. U. Rymsa, M. Hunger, H. Knoezinger, J. Weitkamp, in Porous Materials in Environmentally Friendly Processes; Proceedings of the 1st International FEZA Conference, I. Kiricsi, G. P´al-Borb´ely, J. B. Nagy, H. G. Karge (Eds.), Eger, Hungary, 1999, Studies in Surface Science and Catalysis, Vol. 125, Elsevier, Amsterdam, 1999, p. 197. 174. A. de Mallman, D. Barthomeuf, Zeolites 1988, 8, 292. 175. D. Barthomeuf, in Catalysis and Adsorption by Zeolites; Proceedings of ZEOCAT ’90, G. Oehlmann, H. Pfeifer, R. Fricke (Eds.), Leipzig, Germany, 1990, Studies in Surface Science and Catalysis, Vol. 65, Elsevier, Amsterdam, 1991, p. 157. 176. R. J. Cvetanovi´c, Y. Amenomiya, Catal. Rev. 1972, 6, 21. 177. H. G. Karge, V. Dondur, J. Phys. Chem. 1990, 94, 765. 178. B. Hunger, M. von Szombathely, in Zeolites and Related Microporous Materials: State of the Art 1994; Proceedings of the 10th International Zeolite Conference, GarmischPartenkirchen, Germany, 1994, J. Weitkamp, H. G. Karge, H. Pfeifer, W. Hoelderich (Eds.), Studies in Surface Science and Catalysis, Vol. 84, Elsevier, Amsterdam, 1994, p. 669. 179. A. Auroux, V. Bolis, P. Wierzchowski, P. C. Gravelle, J. C. Vedrine, J. Chem. Soc., Faraday Trans. 1 1979, 75, 2544. 180. A. Auroux, in Innovation in Zeolite Materials Science; Proceedings of the International Symposium, Nieuwpoort, Belgium, 1987, P. Grobet, W. J. Mortier, E. F. Vansant, G. SchulzEkloff (Eds.), Studies in Surface Science and Catalysis, Vol. 37, Elsevier, Amsterdam, 1988, p. 385. 181. A. Auroux, in Molecular Sieves – Science and Technology, Vol. 6, H. G. Karge, J. Weitkamp (Eds.), Springer-Verlag, Berlin, 2008, Chapter 4, and references therein, in press. 182. H. G. Karge, L. Josefowicz, in Zeolites and Related Microporous Materials: State of the Art 1994; Proceedings of the 10th International Zeolite Conference, Garmisch-Partenkirchen, Germany, 1994, J. Weitkamp, H. G. Karge, H. Pfeifer, W. Hoelderich (Eds.), Studies in Surface Science and Catalysis, Vol. 84, Elsevier, Amsterdam, 1994, p. 685, and Refs. 13, 14 therein. 183. J. A. Rabo, P. E. Pickert, D. N. Stamires, J. E. Boyle, in Actes du Deuxi`eme Congr`es International de Catalyse, Technip, Paris, 1960, p. 2055. 184. P. B. Weisz, V. J. Frilette, J. Phys. Chem. 1960, 64, 342. 185. P. B. Venuto, P. S. Landis, Adv. Catal. 1968, 18, 259, and references therein. 186. P. B. Venuto, Microporous Mater. 1994, 2, 297, and references therein. 187. M. L. Poutsma, in Zeolite Chemistry and Catalysis, J. A. Rabo (Ed.), ACS Monograph, Vol. 171, American Chemical Society, Washington, DC, 1976, Chapters 8 and 9, pp. 437. 188. S. M. Csicsery, in Zeolite Chemistry and Catalysis, J. A. Rabo (Ed.), ACS Monograph, Vol. 171, American Chemical Society, Washington, DC, 1976, Chapter 12, pp. 680. 189. W. H¨olderich, in Proceedings of the 10th International Congress on Catalysis, Budapest, Hungary, 1992, L. Guczi, F. Solymosi and P. T´et´enyi (Eds.), Elsevier, Amsterdam, 1993, p. 127.

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190. W. O. Haag in Zeolites and Related Microporous Materials: State of the Art 1994; Proceedings of the 10th International Zeolite Conference, Garmisch-Partenkirchen, Germany, 1994, J. Weitkamp, H. G. Karge, H. Pfeifer, W. Hoelderich (Eds.), Studies in Surface Science and Catalysis, Vol. 84, Elsevier, Amsterdam, 1994, p. 1375. 191. J. Weitkamp, L. Puppe (Eds.), Catalysis and Zeolites – Fundamentals and Applications, Springer-Verlag, Berlin, 1999, especially Chapters 5, 6 and 7, pp. 327. 192. M. Iwamoto, in Zeolites and Related Microporous Materials: State of the Art 1994; Proceedings of the 10th International Zeolite Conference, Garmisch-Partenkirchen, Germany, 1994, J. Weitkamp, H. G. Karge, H. Pfeifer, W. Hoelderich (Eds.), Studies in Surface Science and Catalysis, Vol. 84, Elsevier, Amsterdam, 1994, pp. 1395. 193. I. Kiricsi, G. P´al-Borb´ely, J. B. Nagy, H. G. Karge (Eds.), Porous Materials in Environmentally Friendly Processes, Studies in Surface Science and Catalysis, Vol. 125, Elsevier, Amsterdam, 1999, pp. 815. 194. A. V. Deo, I. G. Dalla Lana, H. W. Habgood, J. Catal. 1971, 21, 270. 195. Z. Dudzik, M. Bilska, J. Czeremuzinska, J. Bull. Acad. Pol. Sci., Ser. Sci. Chim. 1974, 22, 307. 196. H. G. Karge, I. G. Dalla Lana, J. Phys. Chem. 1984, 88, 1538. 197. H. G. Karge, J. Ladebeck, in Proceedings of the of the 5th International Conference on Zeolites, Naples, Italy, 1980; Recent Progress Reports and Discussion, R. Sersale, C. Colella and R. Aiello (Eds.), Gianni, Naples, 1980, p. 180. 198. H. Kn¨ozinger, Angew. Chem. 1968, 80, 778. 199. P. E. Pickert, J. A. Rabo, E. Dempsy, V. Schomaker, in Proceedings of the 3rd International Congress on Catalysis, W. M. Sachtler, G. C. A. Schuit, P. Zwietering (Eds.), Amsterdam, Netherlands, 1964, Wiley, New York, 1965, p. 714. 200. B. V. Liengme and W. K. Hall, Trans. Faraday Soc. 1966, 62, 3229. 201. J. Turkevich, F. Nozaki, D. N. Stamires, in Proceedings of the 3rd International Congress on Catalysis, W. M. Sachtler, G. C. A. Schuit, P. Zwietering (Eds.), Amsterdam, Netherlands, 1964, Wiley, New York, 1965, p. 586. 202. T. J. Weeks Jr., A. P. Bolton, in Proceedings of the 3rd International Conference on Molecular Sieves, Zurich, Switzerland, 1973; Recent Progress Reports, J. B. Uytterhoeven (Ed.), University of Leuven Press, Leuven 1973, p. 426; see also H. G. Karge, in Introduction to Zeolite Science and Practice, 2nd Ed., H. van Bekkum, E. M. Flanigen, P. A. Jacobs, J. C. Jansen (Eds.), Studies in Surface Science and Catalysis, Vol. 137, Elsevier, Amsterdam, 2001, Chapter 16, pp. 707–746, especially p. 734. 203. W. F. H¨olderich, H. van Bekkum, in Introduction to Zeolite Science and Practice, 2nd Ed., H. van Bekkum, E. M. Flanigen, P. A. Jacobs, J. C. Jansen (Eds.), Studies in Surface Science and Catalysis, Vol. 137, Elsevier, Amsterdam, 2001, Chapter 18, p. 821, especially p. 890. 204. N. Arsenova, B. Bludau, W. O. Haag, H. G. Karge, Microporous Mesoporous Mater. 2000, 35–36, 113. 205. H. G. Karge, S. Ernst, M. Weihe, U. Weiss, J. Weitkamp, in Zeolites and Related Microporous Materials: State of the Art 1994; Proceedings of the 10th International Zeolite Conference, Garmisch-Partenkirchen, Germany, 1994, J. Weitkamp, H. G. Karge, H. Pfeifer, W. Hoelderich (Eds.), Studies in Surface Science and Catalysis, Vol. 84, Elsevier, Amsterdam, 1994, p. 1805. 206. M. Sigl, S. Ernst, J. Weitkamp, H. Kn¨ozinger, Catal. Lett. 1997, 45, 23.

207. M. R. Guisnet, Acc. Chem. Res. 1990, 23, 392. 208. G. Bourdillon, C. Gueguen, M. Guisnet, Appl. Catal. 1990, 61, 123.

3.2.4.2

Thermochemical Characterization

Nelson Cardona Mart´ınez and James A. Dumesic∗

3.2.4.2.1 Introduction Heterogeneous catalysis involves specific chemical interactions between the surface of a solid and molecules from the reacting gas or liquid phase. Catalytic cycles on solids are generally composed of adsorption steps, surface reaction processes and desorption steps. The energetics of these surface chemical events play an important role in determining the catalytic properties of the surface. The acid–base character of solid catalysts is a determining factor for their application in many industrial reactions. This chapter addresses studies of the characterization of solid acidity and basicity using temperature-programmed desorption (TPD) and adsorption microcalorimetry. Temperature-programmed desorption and microcalorimetric studies employ probe molecules to examine interactions of surfaces with gas- or liquid-phase molecules. The probe molecules are chosen with respect to the nature of the adsorbed species believed to be important in the catalytic reaction under study or chosen to provide information on specific types of surface sites. High-valent coordinatively unsaturated metal cations or anionic vacancies are considered to act as Lewis acid centers and OH groups may act as Brønsted acid sites or basic sites, whereas oxygen ions (O2− ) account for the basic character of most metal oxide catalysts. Surface spectroscopic studies and well-chosen catalytic tests are typically conducted in conjunction with thermochemical techniques to provide information about the nature of the adsorbed species and to determine the conditions under which the thermochemical studies should be carried out to produce the desired adsorbed species. Basic molecules, such as ammonia (e.g. [1–4]), pyridine [5–7], methylamine [4, 8], dimethylamine [4, 8], trimethylamine [2, 4, 8], triethylamine [2, 4], ethylamine [4], isopropylamine [4] and n-butylamine [4, 9, 10], are generally used to titrate acidic sites. Ammonia and pyridine are widely used probe molecules since there is well-established literature concerning the use of infrared spectroscopy to distinguish between Lewis and Brønsted acid sites [11–15]. Acidic probe molecules, such as carbon dioxide [16–19], sulfur dioxide [20] and pyrrole [21–23], have been used to titrate basic sites. Carbon dioxide is a ∗

Corresponding author.

3.2.4 Acidity and Basicity

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particularly useful probe molecule for basic sites because it can be used to titrate basic oxygen anions and hydroxyl species on the surface [17, 24–27]. We note here that the catalysis literature contains a long history of research in which results from TPD and microcalorimetric measurements have been used in attempts to interpret the performance of catalysts containing acidic and/or basic sites. Clearly, the results from TPD and microcalorimetric measurements are of direct relevance for determining the numbers of acidic and basic sites; however, these studies have been less successful in determining how the strengths of these sites are related to catalyst performance. One of the reasons for this mixed success is that TPD and microcalorimetric measurements probe the energetics of adsorbed species, whereas catalyst performance is controlled by the energetics of transition states. Accordingly, when the adsorbed species probed by TPD and microcalorimetric measurements resemble the transition states that control chemical reaction kinetics, then good correlations of TPD and microcalorimetric results with catalyst performance may be achieved. However, such correlations may fail when the interactions probed by TPD and microcalorimetric measurements are different from the interactions that determine the energetics of the kinetically controlling transition states. Fortunately, results from density functional theory (DFT) calculations can be used to bridge this gap. In particular, DFT calculations can be used to probe the interactions that determine the energetics of both stable adsorbed species and reactive transition states. Thus, DFT calculations can be used to determine first which types of surface sites give heats of adsorption that are consistent with the results from TPD and microcalorimetric measurements and subsequent DFT calculations can be carried out to determine the energetics of transition states on these surface sites. In this respect, we recommend that results from TPD and microcalorimetric measurements be combined with studies of catalyst performance for well-defined reactions that probe the acidic and/or basic sites (e.g. isobutane cracking [28], 2-propanol dehydration and dehydrogenation [29–32]) and the findings from these experimental studies can then be combined with results from DFT calculations on surface sites consisting of clusters of atoms or periodic arrangements of atoms on various surfaces.

Falconer and Schwarz [33] have reviewed the applications of temperature-programmed desorption and reaction. In a typical TPD experiment, the catalyst is contained in a reactor that can be heated at a linear rate. After pretreatment, the catalyst is saturated with a probe molecule under well-defined adsorption conditions. After the excess gas has been flushed out of the reactor, the sample is heated in a flowing inert gas stream. (Temperature-programmed reaction studies are conducted by replacing the inert gas with a reactive gas feed to the reactor.) A thermocouple inserted in the catalyst measures the temperature and a detector downstream measures the effluent gas composition. The concentration of the desorbing gas in the effluent gas may be monitored by absorption/titration, thermal conductivity, flame ionization or mass spectrometry. Alternatively, TPD experiments may be carried out by using a microbalance to measure the changes in sample mass during heating. Inert gas flow-rate, catalyst particle size, catalyst pore size and catalyst bed depth affect the TPD spectra. TPD studies of porous catalysts are generally carried out using reactors designed to minimize concentration gradients in the reactor. Demmin and Gorte [34] have developed dimensionless groups to test for readsorption and transport effects in desorption from packed catalyst beds. The size of the probe molecule may affect the accessibility to specific sites and also influence the rate of diffusion in TPD experiments. In particular, large probe molecules influence the concentration gradients within the catalyst particles of zeolites during TPD experiments. For example, Sharma et al. [35] showed that pyridine desorption from MFI (ZSM-5) is limited by molecular diffusion in the zeolite crystals. For TPD studies of porous catalysts carried out in well-mixed reactors designed to minimize concentration gradients, differential equations [36] may be derived that describe the desorption of the probe molecule, the nonactivated readsorption of the probe molecule on the catalyst surface and the transport of the probe molecule from the TPD cell. Specifically, for first-order desorption these equations are as follows:
 d a −Ed = νd exp a − ka Pa (1 − a ) rd = − dt RT

3.2.4.2.2 Principles of Temperature-Programmed Desorption TPD measurements are generally carried out in two experimental regimes: studies of materials having low surface areas at ultra-high vacuum conditions and studies of porous catalysts at ambient pressures. The primary difference in the analysis of these experiments is that readsorption during desorption must be addressed in the latter experiments.

(adsorption–desorption) (1) −Pa Fs0 + P 0 rd dPa = dT Ns β dT =β dt References see page 1133

(reactor mass balance) (2) (heating rate) (3)

1124

3.2 Chemical Properties

where rd is the net desorption rate in turnover-frequency units (s−1 ); a is the fractional surface coverage of species a; νd is the frequency factor for desorption (s−1 ); Ed is the desorption activation energy (kJ mol−1 ); ka is the adsorption rate constant (bar−1 s−1 ); Pa is the gas-phase pressure of species a (bar); Fs0 is the inlet molecular site velocity of carrier gas (molecules per site); P is the gas phase pressure in the cell (bar); Ns is the number of gaseous molecules in the cell per site (molecules per site); and β is the heating rate (K s−1 ). These analyses make no assumptions about the relative rates of adsorption and desorption during the TPD experiment, but the values of νd and ka are constrained by the relation   0 Sads ka (4) = exp vd R 0 is the standard entropy change of adsorption. where Sads In the limiting case where readsorption is fast, the above analysis indicates that the temperature corresponding to the maximum rate of desorption Tm may be given by the following expression: 
Ed vd Fs0 −Ed (5) = exp RTm2 β RTm [ka P (A − Am )2 ]

where Am is the surface coverage at the peak maximum, approximately equal to half the initial surface coverage. The relationship in Eq. (5) can be rearranged and the enthalpy and entropy of adsorption may be determined from TPD data if spectra are collected over varying samples sizes or carrier-gas flow-rates [37, 38]. For this case where adsorption and desorption are equilibrated, the slope from a plot of 2 ln Tm − ln(W/F ) versus 1/Tm is equal to −Hads /R, where W is the catalyst mass (g) and F is the carrier-gas flow-rate (cm3 s−1 ) measured at room temperature. The ordinate intercept of this plot is equal to   Sads −A0 T0 βHads P (1 − m )2 ln + (6) P0 R where A0 is the acid site density (mol g−1 ), T0 is the room temperature, P0 is the standard state pressure (1 bar) and m is the surface coverage at temperature Tm . Sharma et al. used this procedure to determine the enthalpy and entropy of adsorption of ammonia on H-MOR [35] by varying the sample size for two sets of carrier-gas flow-rates. Good agreement was obtained between the two series of runs with an average heat of ammonia adsorption on H-MOR equal to 153 kJ mol−1 . This heat was consistent with the value determined using adsorption microcalorimetry. The value of the standard entropy change extracted from the intercept of the plot was −156 J mol−1 K−1 .

Applications of Temperature-Programmed Desorption Temperature-programmed desorption or reaction of simple bases or acids is a widely used method to assess the total number and strength of acid sites (e.g. [35, 39–44]) or basic sites [30, 45, 46]. The next two sections provide a concise summary of selected TPD studies performed to characterize the acidity and basicity of zeolites and metal oxides. 3.2.4.2.3

A Acidity of Oxide Catalysts Ammonia TPD has been employed to characterize the acidity of a variety of zeolites including MFI (ZSM-5) [35, 47–50], MOR (mordenite) [35], FAU (faujasite) [48, 49, 51], BEA (beta) [49, 52], MAZ (mazzite) [53] and MWW (MCM-22) [54]. Pyridine TPD has also been used to study the acid properties of zeolites (e.g. [35, 48, 55]). However, pyridine desorption is limited by molecular diffusion in the zeolite particles, and for this reason it is a poor choice as a probe molecule for assessment of acid-site strength on these materials by TPD [35]. Gorte and co-workers [56–58] suggested that the use of temperature-programmed reaction (TPR) studies of reactive amines is a more appropriate method to investigate Brønsted acid sites than TPD of ammonia. In this case, the alkylamines used as probe molecules undergo a decomposition reaction only at Brønsted sites. In contrast, ammonia adsorbs on both Lewis and Brønsted sites, and this may lead to conflicting results [59]. For example, Juskelis et al. [60] have shown using ammonia TPD that CaO has stronger acid sites than USY, while no reaction of isopropylamine was observed in desorption from CaO. Gorte and co-workers used their TPD–TGA method to study the Brønsted acidity of MFI, MTW (H-ZSM-12), MOR and FAU [41–44, 59]. They showed that the sites responsible for the decomposition of isopropylamine also appear to be responsible for hydrocarbon cracking on steamed H-FAU catalysts [56, 57]. Dumitriu et al. [47] observed that the amount of desorbed ammonia correlated well with the increase in aluminum content in MFI zeolites, but not with the iron content in Fe-MFI. No difference in the amount of acid sites was found for Fe-MFI with Si/Fe ratios of 26 and 12. The presence of extra-framework species was given as an explanation of the discrepancy. Borges et al. [50] used the deconvolution of ammonia TPD to obtain the acid strength distribution of a series of sodium-exchanged MFI zeolites. Molecular modeling/quantum calculations (DFT calculations) were used to establish that the activation energy for the desorption of ammonia was a reasonable measure of acid strength and that the activation energy for n-hexane cracking may be related to the activation energy for ammonia desorption using either the Polanyi or Marcus

3.2.4 Acidity and Basicity

models. They were able to show that experimental activities for the cracking of n-hexane can be correlated satisfactorily with theoretical values evaluated using the Polanyi or the Marcus model. Intermittent temperature-programmed desorption (ITPD) or stepwise temperature-programmed desorption (STPD) is an interesting application of TPD [45, 61]. The ITPD technique uses a differential approach with a sawtoothed heating program to generate a sequence of interrupted desorptions. The lower part of these TPDs occurs at quasi-constant surface coverage and may consequently be interpreted by a standard Arrhenius plot. Gaillard et al. [53] studied the acidity of H-MFI and H-MAZ using ammonia ITPD. Using this technique, the researchers were able to determine activation energies of desorption and frequency factors as a function of ammonia coverage and their results were consistent with literature values obtained using adsorption microcalorimetry. Ammonia [62] and pyridine [63] TPD have been used to study the effect of replacing part of the Si atoms in MCM-41 with Al or Ti. Addition of Al or Ti enhances both the number and strength of the acid sites in the modified MCM-41. Deutsch et al. [52] used ammonia TPD to study the acid properties of H-BEA and sulfated zirconia. The zeolite had a higher number of acid sites, especially of those with intermediate and weak strength. The acid properties of alkaline earth metal salts of 12-tungstophosphoric acid were also studied using ammonia TPD [64]. In this instance the addition of the alkaline earth metal cations caused a significant decrease in both the total number of acid sites and the number of the strongest acid sites. B Basicity of Oxide Catalysts Li and Davis used CO2 STPD and adsorption microcalorimetry to evaluate the total number and strength of basic sites of Cs and K oxideloaded X zeolites [45, 46]. The values of CO2 adsorption capacity from STPD were slightly higher than those from adsorption microcalorimetry. The small difference was attributed to the fact that they neglected the amount of CO2 adsorbed with heats of adsorption less than 60 kJ mol−1 . The results from CO2 STPD indicated that about four CO2 molecules adsorbed per unit cell at 373 K. Most of the CO2 desorbed between 373 and 573 K and little desorbed above 673 K. Results from CO2 poisoning of the catalysts indicated that about 80% of the activity of CsOX /CsX and CsOX /KX in 1-butene isomerization arose from the basic sites the CO2 desorption temperature of which was between 673 and 773 K. However, the amounts of the basic sites were only about 5% or less of the basic site inventory on these catalysts. Fishel and Davis [30] used TPD of CO2 and 2-aminopropane and TPR of 2-propanol to study the basic

1125

properties of Mg−Al mixed oxides derived from hydrotalcites. The TPR method is similar to that of Gorte and co-workers [56–58] described above. In this case, 2-propanol is used to probe the active sites, since it generally dehydrogenates to form propanone over basic catalysts and dehydrates to form propene over acidic catalysts. Their results indicated that the mixed oxides had surface adsorption capacities and reactivities similar to those of MgO and that these characteristics were significantly influenced by the hydrotalcite synthesis method. D´ıez et al. [32] studied the same reactions at 533 K and characterized the acid–base properties using TPD of NH3 and CO2 coupled with IR spectroscopy of adsorbed CO2 on similar materials. The chemical nature and acid–base properties of the active sites and also the catalyst bulk structure and product formation rates for the elimination reactions depended strongly on the Mg:Al ratio. In Mg-rich catalysts, the presence of increasing concentrations of more electronegative Al3+ cations decreased the solid average basicity and the overall and dehydrogenation catalytic activity. The base site density determined from the evolved CO2 decreased with increasing catalyst Al content whereas the NH3 desorbed increased almost linearly. The Al-rich catalysts (Mg:Al < 1) were more active than the Mg-rich ones, converting 2-propanol mainly to propene via an E2 mechanism. Al-rich samples were structurally heterogeneous oxides that contained a separate quasi-amorphous Al2 O3 -like phase where dehydration took place at high turnover rates. In some instances acid–base pairs are needed for adequate catalytic activity (e.g. [65]). Tago et al. proposed a TPD method for simultaneously characterizing the acidic and basic properties of solid catalysts by utilizing the co-adsorption of NH3 and CO2 [66]. This method was employed to study the acid–base properties of a series of catalysts. For catalysts with low concentration of basic sites, there was little difference in the NH3 TPD spectra between single and co-adsorption experiments. In contrast, for amphoteric catalysts such as alumina, a remarkable difference in the NH3 TPD spectra was observed between single adsorption and co-adsorption experiments. The difference observed was ascribed to a strong induction effect appearing on the acidic and basic sites. In another study, the surface acidity and basicity of La2 O3 , LaOCl and LaCl3 catalysts were characterized with various probe molecules using infrared spectroscopy, TPD and DFT calculations [67]. The acidic sites were probed with CO, pyridine and 2,6-dimethylpyridine (DMP) and the basic sites with CO2 . The results from the three techniques were consistent. References see page 1133

1126

3.2 Chemical Properties

To end this section, we provide an example where results from CO2 TPD measurements were combined with studies of catalyst performance for n-hexane reforming, and the findings from these experimental studies were then combined with results from DFT calculations on cation-exchanged Pt−M−ETS-10 [where M = Li, Na, K, Rb, Cs, Mg(OH), Ca(OH), Sr(OH) and Ba(OH) ions] [68]. The results of this study demonstrate that the catalytic activity and selectivity of n-hexane reforming to benzene of Pt−M−ETS-10 is altered by the activation of Pt (electron density on Pt) by the support, the basicity of which is influenced by the exchanged metal ion (M) and the Pt cluster location. There was good agreement between the theoretical and experimental observations, demonstrating the potential of combining such characterization techniques. 3.2.4.2.4 Principles of Adsorption Microcalorimetry A microcalorimeter is a tool that provides a direct measurement of the heat evolved during adsorption [3]. Microcalorimeters are typically connected to volumetric systems with sensitive pressure-measurement devices for accurately measuring adsorbed amounts. Small mounts of adsorbate (ca. 1–10 µmol) are sequentially admitted to the adsorbent to produce differential heats of adsorption. Various microcalorimeters have been developed, which may be classified into three basic categories: adiabatic, isothermal and heat-flow (or quasi-isothermal) calorimeters (e.g. [69]). Heat-flow calorimeters, in particular Tian–Calvet microcalorimeters, are well suited for the study of slow adsorption processes over a wide range of temperatures (77–873 K) [3, 69–72]. Heat-flow calorimeters consist of three basic parts:

(i) the calorimeter cell, where the adsorption process occurs (ii) the calorimetric or thermal block, used as a heat sink to keep the process at a specific temperature (iii) the heat conductor or thermopile, which connects parts (i) and (ii) [70]. When heat is liberated in the calorimeter cell, a thermal flux (dQ/dt) is created in the heat conductor that flows to the calorimetric block until thermal equilibrium of the system is restored. The thermal block has a high heat capacity relative to the calorimetric cell and thus the block temperature should not change with the influx of heat. The heat flux through the heat conductor is measured by the difference in temperature (T ) between the calorimetric cell and the thermal block, following the Tian equation: dT dQ = hT + c dt dt

(7)

where h is the heat transfer coefficient of the heat conductor and c is the heat capacity of the calorimetric cell and its contents [73]. The resulting heat response, or thermogram, appears as a peak that returns to the baseline (the calorimetric response in the absence of heat production) after typically 20–40 min [72]. Integration of the Tian equation over the entire time interval of the thermal response (t1 , t2 ), gives the integral heat:  t2 Q=s T dt (8) t1

Note that the integral heat is proportional to the area under the thermogram. The proportionality constant (s) is known as the sensitivity, which is discussed later. A more complete derivation of the Tian equation is given elsewhere [70]. The differential heat (qd ) is defined as the amount of heat, dQ, liberated when dnS mol of gas adsorb on a solid surface at a constant temperature:
 dQ (9) qd = dnS VG ,VS ,AS where the volumes of gas (VG ) and solid (VS ) phases and the surface area of the adsorbent (AS ) remain constant. Calorimetry is often employed at temperatures for which the adsorption process is not reversible and true differential heats cannot be defined. The heats measured calorimetrically are, therefore, average integral heats (Q) evolved during the adsorption of a small dose of gas (nS ):
 Q (10) qd = nS VG ,VS ,AS These differential heats are commonly plotted against the quantity of adsorbed gas. The most commonly used heat-flow calorimeter for catalytic adsorption and reaction studies is the Tian–Calvet microcalorimeter [69, 70, 72]. It consists of a thermal block with two identical calorimetric cells, one for the adsorbent and the other serving as a reference. Two identical thermopiles surround the calorimetric cells and these conduct heat to the thermal block and measure the resulting heat flux. Each thermopile consists of a series (ca. 500) of thermoelectric junctions (or thermocouples) arranged normal to the calorimetric cells. The thermoelectric junctions, connected in series, record voltage changes associated with the heat liberated within the calorimetric cell. The two thermopiles are connected in opposition to one another such that the resulting signal is attributed solely to the process under study. This differential arrangement improves the stability of the baseline by compensating for electrical noise and temperature fluctuations within the calorimetric block.

3.2.4 Acidity and Basicity

The thermopiles are capable of measuring millijoule quantities of heat (hence the name microcalorimeter) [70]. Although heats of adsorption are typically large (ca. 10–500 kJ mol−1 of adsorbed gas [72]), differential heats of adsorption are obtained by dosing micromolar quantities of gas on the sample, resulting in temperature changes as low as 10−4 K [73]. In a typical experiment, 0.5–5 g of catalyst is loaded into the sample calorimetric cell followed by catalyst treatment. The calorimetric cells are then evacuated (ca. 10−4 –10−3 Pa) at the treatment temperature and subsequently immersed in the calorimetric block (usually by raising the calorimeter around the calorimetric cells). The calorimeter is allowed to equilibrate at the adsorption temperature (ca. 4–6 h), at which point a stable differential heat response (baseline) is achieved. The experiment begins by gradually introducing small quantities (ca. 1–10 µmol) of adsorbate to the sample. The resulting heat response, recorded by the thermopiles, is amplified (typically by a factor of 103 ) and plotted on a chart recorder or sent to a computer. The amount of gas adsorbed is determined volumetrically from the dose and equilibrium pressures and the system volumes and temperatures [74]. The differential heat is then calculated by integrating the heat response according to Eq. (8) and normalizing this heat by the amount adsorbed, as in Eq. (10). Sequential doses are admitted to the sample and the experiment continues until the catalyst becomes saturated. External leaks into the volumetric system are typically kept below 10−3 Pa min−1 (ca. 10−5 µmol min−1 ), and the system temperature is kept to ±0.1 K to ensure accurate measurement of differential heats and adsorbate coverages. The choice of adsorption temperature has a significant effect on the adsorption selectivity and the apparent heat of adsorption. At moderately high temperatures (e.g. 400–500 K), equilibrated adsorption occurs leading to preferential adsorption on the stronger sites. Thus, surface heterogeneities and effects of lateral interactions between adsorbed species can be measured. If the temperature is too high, the adsorbed molecules may undergo further surface reactions, and higher pressures are required to populate the adsorption sites. Conversely, if the temperature is too low such that equilibration of the adsorbate on the adsorbent is not achieved, then only an average heat of adsorption is measured. Therefore, the choice of adsorption temperature represents a compromise between these effects. Evidence for these conclusions has been reported [3, 75–77] for ammonia and pyridine adsorption on amorphous silica-alumina and zeolite catalysts. Some of the heat produced in the calorimetric cells is lost to thermal leaks, by conduction, convection and radiation. As a result, only a fraction of the heat

1127

produced in the calorimetric cells is recorded by the thermopiles. Therefore, the sensitivity of the calorimeter (usually expressed in µW mV−1 ) must be calibrated to determine the measurable fraction of heat. This calibration can be determined in two ways: chemically, by using a catalyst–adsorbate system with a known heat of adsorption, and physically, using Joule-effect devices. Chemical calibrations are easy to apply and the test conditions can be nearly identical to the experimental conditions. Oxygen adsorption on NiO (200) [78], carbon monoxide adsorption on a Pt/SiO2 catalyst [79] and ammonia adsorption on H-mordenite [80] have been used for this purpose; however, no chemical calibration standard has been defined. The major disadvantage with this approach is that heats of adsorption may depend on the preparation and pretreatment of the sample. The Joule-effect method is most commonly used for calibration of the calorimeter [72]. In this approach, a resistor is placed within the calorimetric cell. A voltage pulse is applied across the resistor and the heat response is recorded as a function of time. The sensitivity can be determined by comparing the heat measured calorimetrically with the heat generated by the electrical impulse. A DC voltage source (0.1–10 mW), capable of generating millijoule quantities of heat, is required [79]. Applications of Adsorption Microcalorimetry Heat-flow microcalorimetry is a powerful tool for understanding the acidity and basicity of solid catalysts. It has also been used extensively for the study of acidic and basic sites in zeolites and metal oxides. The following is a brief summary of selected calorimetric studies performed mostly over the last decade. 3.2.4.2.5

A Acidity of Zeolites and Mesoporous Materials The use of adsorption microcalorimetry for the characterization of the acidity of zeolites has been reviewed over the last decade [40, 42–44, 48, 77]. The number and strength of acid sites were discussed as functions of basicity of the probe molecules, framework structure, concentration and nature of framework cation, chemical modifications and influence of aging and coke deposition [40, 77]. The heat of adsorption was observed to increase with an increase in the gas-phase proton affinity of the probe molecule. For H-MFI, a good correlation between the average differential heats of adsorption and the proton affinities of the bases was found for a series of amines [42], substituted pyridines [43] and nitriles [44]. The differential heats of adsorption of ammonia and pyridine adsorption were found to vary with the zeolite structure. A compilation from various sources shows the following order of average References see page 1133

3.2 Chemical Properties

acid strength for some common zeolites: USY (ultra-stable FAU) > H-MOR > H-MFI ≈ H-ZSM-12 (H-MTW) > H-MWW > H-BEA > H-FAU [7, 40, 43, 48, 81, 82]. For high-silica zeolites without extra-framework Al and particularly for H-[Al]MFI, the differential heat of ammonia and pyridine adsorption was nearly constant up to the coverage of one adsorbed molecule per framework Al, independent of the Si:Al ratio [42–44]. Similar heat plateaus have been observed when little or no extraframework Al is present in Y zeolites (e.g. [83, 84]). Isomorphous substitution of Al in MFI for Ga or Fe also generated surfaces that displayed constant heats of pyridine adsorption up to stoichiometric adsorption and with the same acid strength as H-[Al]MFI independent of the cation substituting Al [85]. On the other hand, samples of H-[Al]MFI, H-[Fe, Al]MFI, H-[Fe]MFI, H-[Ga]MFI and H-[B]MFI that had extra-framework cations and both Lewis and Brønsted acidity displayed significantly different acid strength distributions [47]. Similar results were found for the isomorphous substitution of Al by Ga in both the FAU [86] and BEA structures [87]. Using ammonia adsorption microcalorimetry and pyridine FTIR spectroscopy, it was shown that partial replacement of Al with Ga in the FAU framework increased the Brønsted acid site density but decreased the acid strength [86]. For H-[Ga]FAU, both Brønsted site density and strength were found to decrease [87]. In all instances where zeolites displayed homogeneous acid strength distributions, only Brønsted acid sites were observed. Conversely, the acid strength distribution was heterogeneous for zeolites with both Brønsted and Lewis acid sites (e.g. [82, 84, 88–94]). Commercial FCC catalysts are complex materials containing at least USY, ZSM-5 and an acidic binder [48]. These catalysts display heterogeneous acid strength distributions that change with sample pretreatment and possess both Brønsted and Lewis sites. For example, the acidity of USY-based FCC catalysts steamed at various severities was studied using adsorption microcalorimetry of ammonia and IR spectroscopy of adsorbed pyridine [89–92]. Figure 1 shows that increasing the steaming time and/or temperature shifted the acid strength distribution to weaker sites. Steaming resulted in the preferential loss of Brønsted acid sites corresponding to NH3 adsorption with a differential heat of 130 kJ mol−1 . These results were consistent with earlier work using adsorption microcalorimetry and IR of pyridine that indicated that the sites of intermediate strength in Y-zeolites are predominantly Brønsted acids and their number and strength are reduced by steaming [93]. The Brønsted site population decreased non-linearly with decrease in the number of framework Al and the relationship was not stoichiometric, with less Brønsted acid sites observed than framework Al present [92]. Consequently, some acid sites in Y zeolites

160 Differential heat / kJ mol−1

1128

140 120 100 USY-1

80 USY-6 60

0

100

USY-3 200

300

Coverage / µmol

400

500

600

g−1

Differential heats of ammonia adsorption on USY-1 (steaming for 2 h at 840 K), USY-3 (2 h at 1030 K) and USY-6 (3 h at 1090 K) at 423 K [92].

Fig. 1

appear to be inaccessible to pyridine and hence to most hydrocarbons during reaction. Enhanced catalytic cracking activity was observed with an increase in the strength of the Brønsted acid sites [89–93]. Detailed quantitative kinetic models for analyzing the catalytic cracking of isobutane [90, 91, 95] and 2-methylhexane [89, 92, 96] over USY-zeolite-based catalysts were formulated that incorporated the catalyst acid properties, number and strength of the Brønsted acid sites. The number of acid sites was that obtained from the adsorption microcalorimetry experiments and the Brønsted acid strength of the catalyst was given in terms of an adjustable parameter, H+ . This parameter may be defined as the enthalpy of a carbenium ion transition state relative to the enthalpy of stabilization of a surface proton. The models accurately predicted the experimentally observed changes in activity and selectivity with conversion and steaming for both reactions by adjusting only the parameter that reflects acid strength. The model predictions were in agreement with the microcalorimetric results. Microcalorimetric measurements have been used to probe the changes in the acid properties caused by deactivation processes for commercial FCC catalysts during gas oil cracking under MAT conditions [97] and during the isobutane/butene alkylation reaction over various zeolites [82]. Coke formation decreased the acid site density and strength on both occasions. The combination of adsorption microcalorimetry with N2 porosimetry and AFM showed that during gas oil cracking under MAT conditions, coke was deposited both inside and outside the catalyst micropores [97]. The coke uniformly coated the micropore walls, decreasing their pore size and causing a moderate decline in acid site strength and density. Coke formation during alkylation caused larger changes in the

3.2.4 Acidity and Basicity

widely studied of these catalysts is sulfated zirconia because it has high activity for butane isomerization at low temperatures. Because of its unexpected catalytic properties, Hino and Arata suggested that these materials might be superacidic [105]. However, calorimetric measurements indicate [106–110] that sulfated zirconia catalysts exhibit strong acidity, but not superacidity, with acid strengths similar to those of strong solid acids, for example H-mordenite, as shown in Fig. 2. Microcalorimetric and infrared spectroscopic studies of ammonia adsorption showed that there are ca. 70 µmol g−1 of acid sites with adsorption heat between 125 and 165 kJ mol−1 and these acid sites are mainly of the Brønsted type with a smaller number of Lewis acid sites [106]. Kinetic studies using selective poisoning of acid sites over sulfated zirconia with adsorbed ammonia showed that these acid sites are responsible for the high catalytic activity [107]. In addition to catalyst acidity, the hydration state of the catalysts also affects the catalytic activity significantly [108]. The initial catalytic activity changed with the drying temperature for sulfated zirconia and the activity was a function of the amount of water added to the dehydrated catalyst. An amount of 75 µmol g−1 of water was needed to promote the dehydrated catalyst to maximum initial activity [111]. Dehydration of the catalyst at either 588 or 773 K did not alter the heat or the extent of ammonia adsorption, or the type of acid sites, that is, Brønsted acid sites [109]. The promotion effect of water may have resulted from hydroxyl groups generated during drying treatments or dissociative water adsorption. 200

30 µmol g−1 47 µmol g−1

180 Differential heat / kJ mol−1

acid strength distribution for USY, REY and Beta zeolites [82]. The Y-based zeolites lost a significant fraction of the strongest sites whereas Beta zeolite lost a larger fraction of the weaker sites. Although the REY had lower acid site strength and density than the USY, it displayed a higher activity than the USY for the alkylation reaction. It was suggested that Brønsted sites with intermediate acid strength might be appropriate for maintaining good catalytic performance in alkylation reactions [82]. In addition to reviewing the use of adsorption microcalorimetry for the acidity characterization of zeolites, Auroux has reviewed some studies for mesoporous materials [40]. These materials have attracted increased interest recently because of their potential as supports or catalysts to process high molecular weight compounds of the type used in certain pharmaceutical and homogeneous reactions or the cracking of gas oils. It has been shown that isomorphous substitution of Si atoms in mesoporous silica with the MCM-41 structure by Al [63], Ti or Zr [98] increases both the concentration of acid sites and the acid strength. In addition, the acidity of hexagonal and cubic thermally stable mesoporous tin(IV) fluorophosphates [99] and also aluminophosphates (AlPOs) and silicoaluminophosphates (SAPOs) has been characterized calorimetrically [100]. The tin phosphates had higher acid strength than MCM-41 and had sites that were as strong as on some H-MFI zeolites [99]. The hexagonal phase had a higher concentration of intermediate strength acid sites than the cubic phase. These catalysts were active for NO reduction and C2 H4 oxidation. The mesoporous AlPOs and SAPOs had higher acid strength than the microporous analogues and the amount of strong sites increased with the Si content [100]. Pillared clays may have a bimodal distribution of pores in the 1.0–20 nm range [101]. This pore-size distribution can be useful for the formulation of new catalysts for processing large molecules. Occelli et al. [101] prepared synthetic saponite expanded with stable SiO2 · TiO2 colloidal particles. The inclusion of SiO2 · TiO2 clusters significantly increased the total concentration of acid sites (Brønsted and Lewis) and acid strength of the clay [101]. The pillared clay dried under supercritical conditions had good catalytic cracking activity and selectivity. The acidity of montmorillonite pillared clays with microporosity [102] and mesoporosity [103] has been characterized with ammonia microcalorimetry.

160 140 120 8 µmol g−1

100

69 µmol g−1

80 60 40 0

50

100

150

200

250

NH3 coverage / µmol g−1

Differential heat of NH3 adsorption vs. adsorbate coverage at 423 K for a sulfated zirconia sample calcined at 848 K in oxygen. The amounts of NH3 used in the selective poisoning of acid sites over sulfated zirconia were based on this result, as indicated in the plot [110].

Fig. 2

B Acidity of Binary Metal Oxides and Supported Metal Sulfated metal oxide catalysts have attracted Oxides much interest as a potential replacement for liquid catalysts in processes such as isomerization, hydrocracking, alkylation and oligomerization [104]. The most

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References see page 1133

1130

3.2 Chemical Properties

The absence of a correlation between the heat of adsorption of ammonia or pyridine and catalytic activity for the isomerization of n-butane over sulfated zirconia [108, 112, 113] and over Pt-promoted sulfated zirconia [113] and for the isomerization of n-pentane over Cu-promoted sulfated zirconia [114] suggests that these reactions are not strictly acid catalyzed. The lack of correlation between the heat of adsorption of ammonia and catalytic activity for the isomerization of n-pentane was also observed for a series of tungsten oxide–zirconia catalysts [115]. The presence of acid sites appears to be a necessary but not a sufficient condition for high catalytic activity for isomerization of n-butane or n-pentane. It has been suggested that butane isomerization at 423 K over sulfated zirconia can be viewed as being a surface chain reaction comprising initiation, propagation and termination steps [116]. The surface chain reaction can be initiated by adsorption of alkenes present in the feed stream on acid sites to form protonated species or the alkenes can be generated on sulfated zirconia by either dehydrogenation of butanes over redox sites or by protolysis of butanes over strong acid sites. Experimental evidence for the generation of alkenes by sulfated zirconia catalysts was provided by detecting the production of dihydrogen during butane isomerization [116]. Although the activation of isobutane over sulfated zirconia may proceed via protolysis on strong acid sites, dihydrogen evolution during isobutene isomerization was still observed over sulfated zirconia catalysts that had been selectively poisoned by preadsorbed ammonia. Hence protolysis of isobutene by strong acid sites did not appear to be the primary mode for initiation for butane isomerization over sulfated zirconia. Instead, it appears that the dehydrogenation of butane over sulfated zirconia takes place via a redox process. Auroux and co-workers [52, 117] used a modification of the adsorption microcalorimetric technique to correlate between data from the calorimetric experiments and the Friedel–Crafts catalytic acylation of various aromatic compounds over sulfated zirconia. The results for gaseous ammonia adsorption did not correlate with the catalytic results and the results were inconclusive for gaseous pyridine. Adsorption of pyridine diluted in anisole resulted in the best correlation with catalytic activity. Pyridine in anisole seemed to be closer to the reaction conditions of the liquid-phase acylation and was the most reliable method to correlate the acidity with the activity under these experimental conditions. Another family of solid acids that has received attention as promising candidates for green catalysts are heteropoly acids (HPAs) (e.g. [118]). These materials are unique catalysts since they may display surface acidity and pseudo-liquid acidity. Pseudo-liquid behavior is observed for some heteropoly acids in the solid state because

of their flexible lattice (variable secondary structure) that allows molecules to be absorbed into the threedimensional solid bulk [118]. The number of water molecules of hydration found in the solid determines the secondary structure [119]. The water content has a strong influence on the acidity and catalytic properties of HPAs. For example, Bardin et al. [120] found that increasing the pretreatment temperature from 473 to 573 K caused a decrease in acid strength as measured by the heat of ammonia sorption of between 15 and 30% for bulk H3 PW12 O40 (HPW) and from 20 to 75% for bulk H3 PMo12 O40 (HPMo). In the case of HPMo, the acidity also decreased to less than 20% of the original ammonia uptake. Results from TGA measurements showed that within that temperature range HPW lost water of hydration whereas HPMo underwent decomposition. Auroux and co-workers [64, 121] and Shikata et al. [122] observed similar initial heats of ammonia sorption (ca. 200 kJ mol−1 ) for hydrated HPW. The acid strength decreased in the order HPW > H4 SiW12 O40 > HPMo [120]. Partly exchanging the protons of HPW with an alkali metal (Cs) [120] or with a series of alkaline earth metals [64] considerably decreased both the number and strength of the acid sites of the samples. DFT calculations were used to compare the acid strengths of anhydrous HPW and HPMo by computing the adsorption energy of ammonia on model clusters of each heteropoly acid [123]. The adsorption of ammonia on a phosphotungstic acid cluster was stronger than the adsorption on a phosphomolybdic acid cluster. The predicted adsorption heats were similar to the experimental heats of ammonia sorption determined from microcalorimetry. Comparable results were obtained for the adsorption of a series of alkenes on HPW [124]. The DFT-calculated chemisorption energies were consistent with the experimentally measured heats of adsorption on anhydrous HPW. Supporting HPW on silica [122, 125] or on activated carbon [121] caused distinct modifications to the surface acidity. When silica was used as a support, both the acid strength and density declined noticeably compared with the bulk HPA and became heterogeneous. The number of acid sites increased approximately proportionally to the HPA loading [125]. Supporting HPW on activated carbon reduced the acid strength of the material [121]. In all the studies mentioned above, only Brønsted acidity was observed for all heteropoly acids. These materials were active for double-bond isomerization of 1-butene and cis-2-butene [120, 125], butane and pentane skeletal isomerization [125], synthesis of methyl tert-butyl ether [122] and methanol dehydration [121]. In all cases the results from ammonia sorption microcalorimetry correlated with those from catalytic reactions.

3.2.4 Acidity and Basicity

C Basicity of Zeolites The use of basic zeolites has recently attracted increased attention due to their environmentally benign character and their potential use in fine chemical synthesis [126–128]. The basicity of zeolites is a result of the framework negative charge. Thus, the basicity (both number and strength) increases with increase in the aluminum content (a decrease in the framework Si:Al ratio) of the zeolite. The basic strength also increases with increase in the electropositive nature of the counterion [23]. The framework oxygens are considered the basic sites in zeolites as they function as effective electron donors. Auroux and co-workers have studied the acidic and basic properties of alkali metal-exchanged X and Y zeolites [23, 129, 130] and of Na-A zeolite [130]. Using adsorption microcalorimetry of adsorbed pyrrole, IR spectroscopy and XPS, they identified the coexistence of basic sites with two different strengths that were associated with two different cations. The base strength distribution of the alkali metal-exchanged zeolites displayed two maxima. The lower value (117 kJ mol−1 for the X zeolite samples) coincided with the sole peak observed for the parent Na zeolite and was therefore assigned to adsorption on the remaining Na cations after exchange. The higher value peak was assigned to adsorption on the basic sites associated with the exchange cations. A good correlation between these differential heats and the negative charge on the oxygens, as determined using Sanderson’s electronegativity equalization method [131], was observed. These results supported their conclusion that the Lewis basicity in alkali metal-exchanged zeolites is a local property strongly influenced by the adjacent alkali metal cation. An equivalent correlation was observed between the average integral heat of sulfur dioxide adsorption and the partial negative oxygen charge for alkali metalexchanged X and Y zeolites [129]. Similarly, the average integral heat of ammonia adsorption correlated with the average Sanderson electronegativity of the zeolites. The decrease in acidity with increase in electropositivity of the alkali metal cation was compensated for by a corresponding increase in basicity. The alkali metalexchanged X zeolites (Si:Al = 1.4) were more basic than the corresponding Y zeolites (Si:Al = 2.5), in agreement with the difference in Si:Al ratio. The same trend was observed for Na-A (Si:Al = 1.0) > Na-X (Si:Al = 1.2) > Na-Y (Si:Al = 2.5) [130]. The Lewis acid and base properties also depend on the crystallinity of the zeolite. In a study where Na zeolites were crushed using progressive high-energy ball-milling, it was found that the collapse of the zeolite crystal results in a decrease in the strength and number of both Lewis acid and Lewis base sites [130]. The dependence of the Lewis

1131

acid strength on the level of crystallinity was stronger than the dependence of the Lewis base strength. The base strength and catalytic activity of alkali metalexchanged zeolites can be increased significantly by occlusion of alkali metal oxide clusters via impregnation and decomposition of alkali metal compounds [128, 129, 132–134]. These occluded moieties are apparently more strongly basic than the framework oxygen atoms of the zeolite. Using adsorption microcalorimetry of carbon dioxide, it was found that zeolites containing occluded CsOX clusters had higher CO2 adsorption capacities and heats of adsorption than the alkali metalexchanged zeolites [129]. The number of basic adsorption sites in zeolites containing occluded CsOX was directly proportional to the amount of occluded cesium with one CO2 molecule being adsorbed for every four occluded Cs atoms [45, 133, 134]. A linear relationship was also observed between the rate of 1-butene isomerization and the number of occluded Cs atoms [45, 133]. The rate of ethylene carbonate formation during the cycloaddition of carbon dioxide to epoxides also increased with increasing basicity as determined by CO2 adsorption microcalorimetry [132]. A comparison of the differential heats of carbon dioxide adsorption suggested that bulk cesium oxide is a much stronger base than zeolite-supported cesium oxide [134]. A computational study of CO2 adsorption on model cesium oxide surfaces suggested that the cesium oxide occluded in the zeolite pores was not a stoichiometric oxide, but instead might have been a peroxide [128, 135]. A further study using Raman spectroscopy found no evidence for cesium peroxide or superoxide and it was proposed that the occluded cesium species is an oxycarbonate, which is a metastable intermediate between cesium carbonate and cesium oxide [46]. Alkali metal-containing zeolites prepared by decomposing alkali metal azides impregnated on the zeolites were found to have a higher base strength than alkali metal oxide-containing zeolites [128, 134]. These strong bases are capable of catalyzing toluene alkylation with ethylene, but are prone to deactivation by poisoning and migration of the metal to the external surface [128, 133]. D Basicity of Binary Metal Oxides and Supported Metal Oxides Ammonia and carbon dioxide adsorption microcalorimetry were used to study the surface acid–base properties of a series of γ -Al2 O3 -supported basic metal oxides [18, 136]. Impregnation of various amounts of potassium, magnesium, lanthanum [18] and europium [136] oxides on alumina neutralized acid sites and generated basic sites. On silica, europium oxide generated acidic sites, References see page 1133

1132

3.2 Chemical Properties

whereas the formation of basic sites was inhibited. The effectiveness of basic metal oxides to generate basic sites on alumina was related to the Sanderson electronegativity of the oxide [131]. Sanderson’s method has also been used to develop semiempirical correlations between the ionic character of a series of metal oxides and their basicity and activity in the reaction of CO2 and CS2 [137]. Among the oxides studied, the highly ionic lanthana exhibited the highest basicity and catalytic activity. The acid–base characteristics of a series of materials prepared by impregnating salts of rubidium or strontium on γ -alumina, titania, carbon and silica were also studied using adsorption microcalorimetry of ammonia and carbon dioxide [19]. Alumina and titania adsorbed CO2 and NH3 with appreciable initial heats, indicating the amphoteric nature of their surfaces. Addition of alkali and alkaline earth elements partially neutralized the acid sites and created new basic sites on the supports. The extent of these modifications depends on the compositions of both the basic additive and the support. Rubidium was more effective than strontium at altering the acid–base character of all the supports studied. Carbon and silica are generally considered to be neutral supports and do not adsorb CO2 substantially. After incorporation of Rb and Sr, new basic sites were formed on these carriers. Similar results were obtained by Auroux and coworkers using ammonia and sulfur dioxide adsorption microcalorimetry to study the effect of adding small amounts of ions on γ -Al2 O3 , SiO2 or MgO [138, 139]. Li and Ni on alumina generated the highest number of acid sites, and Zr on alumina the highest average acid strength. Ca and Nd on alumina generated the highest number of basic sites and Li on magnesia displayed the highest average base strength. Increasing the tin oxide loading on γ -Al2 O3 decreased both the strength and the number of basic sites until close to a monolayer of tin oxide was formed [140]. Bulk SnO2 exhibited heats and total coverages similar to those of bulk alumina. An acidity–basicity scale defined by the specific effect of ions (ISE) was proposed to correlate the number of acidic or basic sites generated on alumina or silica [138, 139]. The specific effect of ions was defined as follows: I SE =

nion−supp − nsupp nion

(11)

where nion−supp is the irreversible amount adsorbed of the respective probe molecule for the ion-modified support, nsupp is the irreversible amount adsorbed of the respective probe molecule for the reference support and nion is the amount of the metal ion deposited on the host oxide. The average acid strength of the catalysts, as measured by the average integral heat of ammonia adsorption,

correlated with the charge-to-radius ratio and with the electronegativity of the doping ions. The average base strength of the catalysts, as measured by the average integral heat of SO2 adsorption, correlated with the partial oxygen charge of the corresponding oxides. Hydrotalcite (magnesium aluminum hydroxycarbonate) is an anionic clay that decomposes upon hightemperature calcination and can be used to prepare Mg−Al−O mixed oxides with high surface area and moderate acidity and basicity [141–143]. Microcalorimetric measurements of ammonia and carbon dioxide adsorption showed that calcined hydrotalcites with Mg:Al ratios from 3 to 12 exhibited similar surface acid–base properties [141]. The number and strength of the acidic sites were considerably lower than for γ -Al2 O3 but higher than for MgO. In contrast, the number and strength of the basic sites for calcined hydrotalcites were lower than for MgO but higher than for γ -Al2 O3 . Infrared spectroscopic measurements of NH3 and CO2 adsorption indicated that the acidic sites are mainly Lewis acid sites whereas the basic sites are surface oxygen anions. Exchanging the Cl− anions with CO2− increases 3 the initial differential heat of CO2 adsorption of the Mg−Al−O mixed oxides to values similar to those observed for MgO, but the adsorption capacity is still lower than for MgO [142]. On the carbonated samples, two types of Lewis acid sites related to a linear coordination of CO2 with Mg2+ and Al3+ cations were assigned. Most of the basic sites were identified as OH groups with a small number of strongly basic sites, adsorbing CO2 with 140 kJ mol−1 , probably corresponding to O2− centers. Using the condensation of benzaldehyde and acetone as a test reaction, the authors concluded that the base strength of the high carbonate content Mg−Al−O mixed oxides was comparable to that of piperidine in homogeneous catalysis. It was suggested that the catalytically active sites are OH groups of moderate base strength. In some instances, the proximity of surface Lewis acid sites to surface base sites is needed for a particular reaction to occur. Davis [128] and Tu and Davis [132] studied the cycloaddition of carbon dioxide to various epoxides over a series of solid basic catalysts and found that the high catalytic activity of Cs/Al2 O3 for all of the epoxides studied suggested that coexistence of Lewis acid sites and basic sites was required for this reaction. The acidic sites are believed to stabilize the adsorbed epoxide whereas contiguous basic sites adsorb carbon dioxide. Bolis and co-workers [144, 145] studied the basicity of pure and sulfated zirconia and found that the sulfating process causes the basic O2− /Zr4+ pair sites present on the parent ZrO2 to be virtually consumed. The basic sites were not recovered by the partial selective elimination of sulfates through the calcination step.

References

Conclusions We have shown that TPD and adsorption microcalorimetry are powerful tools for the characterization of solid acidity and basicity. For some catalytic systems, TPD or microcalorimetric measurements combined with other characterization techniques can provide adequate interpretation of the catalyst performance. When direct correlations are not found, such as when the interactions probed by TPD and microcalorimetric measurements are different from the interactions that determine the energetics of the kinetically controlling transition states, results from DFT calculations can be used to bridge the gap. 3.2.4.2.6

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3.2.4 Acidity and Basicity 126. 127. 128. 129. 130. 131. 132. 133. 134. 135. 136. 137. 138. 139. 140. 141. 142. 143. 144. 145.

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Infrared Spectroscopy for the Characterization of Surface Acidity and Basicity

3.2.4.3

.. Helmut Knozinger∗

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Acid–base concepts which apply to solution chemistry have also been introduced into the surface chemistry of solids, although acid–base equilibria on surfaces and in solution frequently cannot be compared because solvation phenomena, which play an important role in solution equilibria, will not occur on surfaces when probe molecules (or reactants) are adsorbed from the gas phase. When classical techniques, including titration methods, determination of adsorption isotherms and isobars and calorimetric measurements, are being used for the characterization of the acid and base properties of solids, a discrimination between sites of different quality (e.g. sites of Brønsted or Lewis type) is often difficult or even impossible. Typical solid acids or bases are binary or ternary amorphous or crystalline oxides which frequently expose both hydroxy groups (Brønsted acids or basic sites) and coordinatively unsaturated (cus) cations (Lewis acid sites) and anions (e.g. basic O2− ions) on their surfaces simultaneously, thus creating acidobasic properties of the surface under consideration. The relative densities of these sites (or functional groups) depend on the state of hydroxylation or hydration of the particular material under the conditions of the investigation. For discrimination between these various types of surface centers and for the characterization of their individual properties, the nature of the interaction between suitably selected probe molecules and the oxide surfaces must be analyzed at a molecular level. Infrared spectroscopy has a very high potential for investigations of intermolecular interactions and has proved to be the most powerful method for investigations into acid–base interactions on solid surfaces, namely those of oxide materials [9, 23–38]. Infrared spectroscopy permits discrimination between various types of surface sites based on the effects induced by the acid–base interaction in the spectral features of the adsorbed probe molecule and of surface groups such as O−H oscillators (see Chapter 3.1.3.8). It also has the potential for the determination of the density of surface centers of either quality when the integrated intensities of characteristic vibrational modes can be measured and it permits the estimation of quantities that are related to the interaction energy between an acidic or basic surface center and the probe molecule (and, hence, to the acid or base strength, respectively). Infrared spectroscopy of adsorbed probe molecules should therefore provide information on the intensive and extensive properties of solid acids and bases when appropriate probe molecules are available.

Introduction Acid- and base-catalyzed reactions belong to the technologically most important classes of catalytic conversions, the catalysts preferably being solids which allow the application of fixed- or fluidized-bed reactors [1–3]. Easy separation of products and reactants is possible when batch reactors are to be used and corrosion problems are minimized. A detailed understanding of solid acid- or base-catalyzed reactions requires knowledge on the surface acidity or basicity, in terms of quality (e.g. protonic versus non-protonic in the case of acidic solids), acid and base strength and site densities. Extensive attempts to characterize acid and base properties of solid catalysts have therefore been made, and a large number of review articles and monographs on this subject have been published [4–22] (see also Chapter 3.2.4.1).

3.2.4.3.2 Criteria for the Selection of Probe Molecules It is useful to define the properties a molecule should

∗

References see page 1158

3.2.4.3.1

Corresponding author.

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3.2 Chemical Properties

possess when it is to be used as a probe for surface acidic and/or basic centers of a given catalytic material by means of infrared spectroscopy. Criteria for the selection of probe molecules have been formulated by Paukshtis and Yurchenko [24], Kn¨ozinger [31] and later Lercher et al. [32]. The criteria [31] summarized below are considered to be important guidelines for the choice of one or more suitable probe molecules for a given problem, although occasionally a particular probe may not fulfill all criteria simultaneously. Criterion I: The detectable spectral response induced by an acid–base interaction between a basic probe and a surface acid site or an acidic probe and a surface basic site must permit an unequivocal analysis of this intermolecular interaction. In the case of acidic properties, the measurable spectral response must in particular discriminate between protic (Brønsted) and aprotic (Lewis) sites. This analysis may be based on changes of intramolecular vibrational characteristics of the probe molecule (frequency shifts and/or activation of IR-silent normal modes by symmetry reduction) and/or on perturbations of the surface acid groups, such as O−H oscillators. Criterion II: The probe molecules should preferentially interact selectively with either acidic or basic sites. This means that a probe molecule should possess either dominant electron-pair donor or electron-pair acceptor properties, since otherwise it may form several types of acid–base surface complexes. As a consequence, the superposition of the respective vibrational spectra may lead to complications in the spectral analysis and band assignments. Criterion III: Quantities, such as frequency shifts, must be measurable with sufficient accuracy. Hence, frequency shifts induced by the acid–base interaction should be larger than the bandwidth of the normal vibration under consideration. The measured spectral quantity should correlate with the strength of interaction between the probe molecule of a given base or acid strength and a surface acid or base site, respectively. The strength (energy) of interaction may be considered as a measure of the acid or base strength of the centers on the catalyst surface. Even though the acid or base strength of centers on solid surfaces cannot be expressed in terms of thermodynamically well-defined quantities such as pK a values by this approach, one can at least establish a relative acidity/basicity scale for a series of materials. Criterion IV: Extinction coefficients of characteristic vibrational modes should be high so that high detection sensitivity can be achieved even with materials of low specific surface areas. Criterion V: Extinction coefficients for characteristic vibrational modes of the adsorbed probe molecules should be experimentally available (for example, by simultaneous volumetric or gravimetric and spectroscopic measurements)

so that the densities of surface centers and perhaps the relative abundance of centers of different quality or strength can be determined. Criterion VI: The probe molecule should have the appropriate acid or base strength so as to induce an optimal acid–base interaction. Probe molecules can be classified with respect to the strength of interaction with the surface sites of a solid material on the basis of the principle of hard and soft acids and bases (HSAB) [39, 40]. This approach was first proposed by Burwell and coworkers [41, 42] and later advocated by Kn¨ozinger [6] as a basis for a qualitative discussion of the adsorption properties of solids. In general, ‘‘hard’’ indicates low polarizability and small size, whereas ‘‘soft’’ indicates high polarizability and large size. Hardness will increase and softness correspondingly decrease with increasing positive oxidation state and degree of coordinative unsaturation and with decreasing size of an ion on a surface. Thus, the most typical members of the group of hard acid centers on solid (oxide) surfaces are the proton (Brønsted site) and small cus cations in high oxidation states (aprotic Lewis sites). As regards potential probe molecules, ammonia (NH3 ) is to be classified as a hard base, whereas carbon monoxide (CO) is a soft base. The HSAB concept has been formulated as follows: hard acids prefer to associate with hard bases and soft acids preferably bind to soft bases. This rule successfully predicts the relative strength of adsorption (acid–base) interactions. Ammonia as a hard base is known to chemisorb strongly on oxide surfaces; in contrast, the soft carbon monoxide molecule interacts only weakly with the hard acid centers on oxide surfaces. Criterion VII: The probe molecule should provide high specificity so as to discriminate between sites of the same quality having (perhaps only small) differences in acid or base strength. Hard bases such as ammonia will typically allow determination of the quality of acid centers but will not necessarily discriminate between centers of different strength. Paukshtis et al. [43] reported that the frequency shifts induced in hard bases such as pyridine and nitriles, when coordinated to hard cationic (Lewis acid) centers, were characteristic for the particular element and largely insensitive to its local coordinative environment. In contrast, as will be discussed below, soft bases such as carbon monoxide respond sensitively to the state of a given metal in different coordinations and oxidation states. As a consequence, it may be recommended to use a series of probe molecules of varying hardness for the full characterization of the surface centers of a catalytic material. Criterion VIII: The molecular size of the probe molecule should be as small as possible, for several reasons. First, the surface center may be located in narrow pores or in cavities which are accessible only through channels or windows having small diameters (zeolites). Second, an

3.2.4 Acidity and Basicity

acid site, namely aprotic Lewis acid sites, may be exposed in an oxygen vacancy and it may be sterically shielded by neighboring oxide ions if the probe (or the basic functional group of the probe molecule) is too large. Third, a large probe molecule may adsorb on one surface site and sterically block another site located in close proximity. It is therefore recommended to use probe molecules of varying size so as to test the accessibility of surface sites not only for the probe molecules but also for potential reactant molecules. Benesi [44] and Jacobs and Heylen [45] suggested that the steric hindrance in 2,6-disubstituted pyridines may be utilized to detect specifically protic sites in the presence of aprotic sites. Spectroscopic evidence, however, indicated that this approach must be treated with caution [46]. Criterion IX: The reactivity of the probe molecule at the applied temperatures and pressures should be low so as to avoid any surface chemical transformations other than simple acid–base associative interactions. The formation of surface compounds via chemical transformations of the probe would modify the intrinsic acidobasic properties of the original catalytic material. However, surface chemical transformations induced by acid–base sites of a catalytic material may provide valuable information on the reactivity of the surface, as outlined in Section 3.2.4.3.4D. An example is the transformation of carbon dioxide into various surface carbonate structures on basic surfaces [6, 9, 29, 30, 47]. Characterization of Acid Sites on Solid Catalysts An acidic catalyst surface may provide protic (Brønsted) sites and aprotic (Lewis) sites. The former are typically surface hydroxy groups OHs on oxide surfaces, the latter cus surface cations L. A basic probe molecule B will interact with hydroxy groups via hydrogen bonding: 3.2.4.3.3

−  OHs + B −  −− − − OHs · · · B

(1)

Eventually, proton transfer may occur, provided that the hydroxy group is sufficiently acidic and the proton affinity of the base is sufficiently high: − + −  OHs · · · B −  −− − − Os · · · H B

(2)

In the case of aprotic sites L, the base B will form a Lewis acid–base adduct: −−  L + B −− − − L ←−−− B

(3)

Infrared spectroscopy is known to be extremely useful in probing H-bonded systems since it responds sensitively to the perturbations induced in the O−H oscillator and in the H-bonded base B [48–50]. The protonated base H+ B can easily be discriminated from the parent base when proton transfer occurs and the H-bonding of the H+ B species to the conjugated surface base O− s in the

1137

+ ion-pair complex O− s · · · H B can be studied. When Lewis acid–base adducts are formed, infrared spectroscopy can only provide information on the perturbation of the base B. These perturbations may lead to modifications of the bond energies and, hence, force constants and normal mode frequencies without any symmetry reduction. Depending on the molecular structure of the basic probe, adduct formation may also lead to symmetry reduction so that a symmetry analysis of the vibrational spectra will provide information on the nature and structure of the adsorption complex.

A Basic Probe Molecules and Their Properties The following classes of probe molecules have been used in conjunction with infrared spectroscopy for the study of acid properties of solid surfaces (the probes are listed approximately according to decreasing hardness): aliphatic amines, ammonia, pyridine and substituted pyridines, nitriles, ethers, ketones and aldehydes, aromatics, alkenes, carbon monoxide, alkanes, dinitrogen and dihydrogen. Within each class of organic probe molecules, the individual molecular properties can be modified by substitution in a controlled and predictable fashion. Several of the mentioned classes of probe molecules obviously contain potential reactants in acid-catalyzed reactions. Their use may in fact be the best choice as they would provide the type of information that is relevant for the catalytic reaction under consideration. However, surface modifications may occur when the probe molecule undergoes catalytic transformations under the conditions of the test experiments. Careful control of the catalytic chemistry that might occur in such cases is definitely required. A brief evaluation of the properties of each of the abovementioned groups of basic probe molecules will follow below based on the above criteria I–IX. a Ammonia Ammonia is probably one of the most frequently used probe molecules for the characterization of acid properties of solid catalysts. It is a hard Lewis base and it is small in size. The various forms of expected chemisorbed NH3 species, namely coordinated, H-bonded and protonated NH3 , can be detected by infrared spectroscopy and discriminated by their characteristic normal vibrations [28, 51, 52], although the N−H stretching and deformation regions are typically rather complex because of H-bonding effects in which N−H and O−H groups (surface hydroxy groups and spuriously present water) are involved and because of band overlaps of fundamental stretching modes with combination modes and overtones of N−H deformation References see page 1158

1138

3.2 Chemical Properties

vibrations. The protonated form NH+ 4 gives rise to bands near 1450 and 3130 cm−1 , whereas NH3 coordinated to aprotic sites shows bands near 1250, 1630 and 3330 cm−1 . Several complications may arise with the use of ammonia as a probe for surface acidity. Due to its strong basicity (hardness), ammonia can be very strongly bonded to a wide variety of sites. Therefore, it cannot be considered as a very specific probe molecule and it may not always provide ideal properties for studies of surface acidity. Moreover, dissociative adsorption of NH3 with formation of NH2 and NH surface groups has been observed at elevated temperatures (≥500 K) on partially dehydroxylated oxide surfaces [53, 54]. Finally, it should be noted that quantum chemical calculations have shown [55] that protonation of NH3 occurs if the resulting NH+ 4 forms two or three H-bonds with oxygen atoms in the vicinity of bridging OH groups in zeolites. The surface in this case can be considered to function as a ‘‘pseudo-solvent’’ and the formation of the NH+ 4 species · · · H+ NH3 cannot be described as a simple ion pair O− s [see Eq. (2)]. It should be noted that exposure of oxide materials including aluminum orthophosphate [56], zirconium phosphate [57], SAPO-11 [58], SBA-15 [59] and several others to ammonia at elevated temperatures leads to surface and bulk incorporation of nitrogen with formation of NH2 and NH groups. Thereby, basic properties were induced in these materials. b Alkylamines Amines, particularly trimethylamine and n-butylamine, have been widely used for surface acidity determinations [4, 5, 8]. Nevertheless, comparatively little spectroscopic information on adsorbed alkylamines is available. As with ammonia, coordinatively adsorbed amines bonded to aprotic sites have been detected on pure oxides [60–62] (–NH2 bending mode of nbutylamine near 1605 cm−1 ) and protonated species (symmetric and antisymmetric deformation modes of the −1 −NH+ 3 group, near 1590 and 1510 cm , respectively) were additionally identified on silica–alumina and zeolites [61]. The protonated species was even reported to be formed on n-butylamine adsorption on transition aluminas [61, 62] because of the high basicity of alkylamines. In addition, there is evidence for dissociative adsorption of n-butylamine [61] and chemical transformations at higher temperatures [63, 64]. Catalytic elimination of ammonia may particularly occur when protic acidity is present [65, 66]. The interaction of alkylamines with oxide surfaces is certainly less specific than that of ammonia because of their higher basicity and their greater molecular size. In particular, trialkylamines will give rise to significant steric hindrance and their large kinetic diameter may

lead to inaccessibility of sites in narrow pores. The influence of steric effects on alkylamine adsorption has been discussed [64, 67] and Medema et al. [67] concluded that the adsorbed amount of an alkylamine on γ -Al2 O3 primarily depended on its molecular cross-sectional area rather than on its basicity. It must therefore be concluded that alkylamines cannot be considered as ideal probes for the characterization of surface acidity. c Pyridine and Substituted Pyridines Pyridine is classified as a weaker base than ammonia, this being consistent with its pKa value as determined in the liquid phase. However, gas-phase basicities indicate that pyridine should be the stronger base relative to ammonia, and Parillo et al. [68] were able to demonstrate that pyridine is more easily protonated than ammonia and that the pyridinium ion is thermally more stable than the ammonium ion. In any case, pyridine is to be classified as a relatively hard base. The molecular size of the molecule may give rise to steric hindrances in intermolecular interactions. The pyridine molecule can undergo coordination to aprotic sites, it can be protonated to form the pyridinium ion PyH+ on acidic OH groups and it can undergo H-bonding with less acidic OH groups. The infrared spectra of pyridine coordination compounds [69] are clearly distinct from those of PyH+ [70, 71] and of H-bonded pyridine, so that the corresponding surface species can be fairly easily distinguished. The ring vibration modes 19b and 8a – according to the nomenclature introduced by Kline and Turkevich [72] – respond most sensitively to the nature of intermolecular interactions via the nitrogen lone pair electrons. These two modes are observed at 1440–1447 and 1580–1600 cm−1 , respectively, for H-bonded pyridine, at ca. 1540 and around 1640 cm−1 for PyH+ and at 1447–1460 and 1600–1633 cm−1 for coordination compounds (see Table 1). In addition, H-bonded species lead to characteristic shifts of the O−H stretching modes of surface hydroxy groups to lower frequencies, the shift νOH being a measure of the H-bond energy. The 8a mode of coordination compounds shifts to increasingly higher frequencies as the coordination bond energy (hardness of the coordination site) increases. It has also been suggested [73–75] that this vibrational mode of coordinated pyridine would be sensitive to the coordination number of the metal center and a possible discrimination of Al3+ in octahedral and tetrahedral sites on alumina surfaces was inferred. Alkyl substitution in the 4-position increases the basicity of pyridines via inductive effects. The same is true for substitution in the 2- and 6-positions; in this case, however, the nitrogen lone pair is

3.2.4 Acidity and Basicity Infrared bands (cm−1 ) of pyridine adsorbed on solid acid catalysts

Tab. 1

H-bonded pyridine

Coordinatively bonded pyridine

1400–1447 1485–1490

1447–1460 1488–1503

1580–1600

1580 1600–1633

1139

In conclusion, pyridine and substituted pyridines, although rather hard bases, provide interesting properties for studies of surface acidity.

Pyridinium ion

1485–1500 1540

1640

sterically shielded. The influence of these effects on the adsorption interaction with alumina surfaces was first investigated by Kn¨ozinger and Stolz [46]. Benesi [44] has suggested that 2,6-dimethyl-substituted pyridine may be used as a specific probe for protonic sites. However, the work of Kn¨ozinger and Stolz [46] had previously shown that 2,4,6-trimethylpyridine can still undergo coordination to Al3+ sites on alumina surfaces, although the interaction is sterically hindered and hence weaker than with unsubstituted pyridine. Following Benesi’s paper [44], several workers have used substituted pyridines as probe molecules for acid sites on various oxide materials including zeolites [45, 62, 76–82]. 2,6Di-tert-butylpyridine is sterically so strongly hindered that coordination on to Lewis acid sites is definitely impossible [83]. It has been inferred from infrared studies that protonation of this probe molecule by protic sites should be possible [84]. However, 2,6-di-tert-butylpyridine does not form N· · ·H−O hydrogen bonds with silanol groups on silica surfaces but rather interacts with silanol groups via the π-electron system [85], as shown by a frequency shift of the O−H stretching vibration of 100 cm−1 which is typical for OH· · · π interactions. The same type of interaction was also observed on alumina [83], so that protonation turns out to be unlikely. However, 2,6-di-tert-butylpyridine is being used for the characterization of Brønsted acid sites [86, 87]. Corma et al. [88] used this probe molecule for the investigation of the external surface acidity of MFI and MOR zeolites. However, 2,6-di-tert-butylpyridine penetrated easily into the pore system of three-dimensional BEA zeolite. Kotrla and coworkers [89, 90] suggested the use of the N−H stretching mode of protonated pyridine to probe the strength of the conjugated basic center after proton transfer has occurred. The authors demonstrated that the N−H stretching frequency shifts of pyridinium ions formed in zeolites increased with increasingly negative charge on the corresponding lattice oxygen atom, as calculated according to Sanderson’s electronegativity, thus suggesting decreasing acid strength of the corresponding original O−H group.

d Nitriles Nitriles are soft bases and they are sterically less crowded than pyridines in the vicinity of the lone pair electrons on the nitrogen atom. Nitriles may therefore be expected to provide attractive properties as specific probe molecules. The C−N stretching mode is sensitive to intermolecular interactions and typically shifts to higher frequencies by 30–60 cm−1 when coordination bonds are formed to aprotic sites via the nitrogen lone-pair electrons [91]. Upon H-bonding, C−N stretching frequency shifts of 10–30 cm−1 have been reported [92–94]. In addition, H-bonding also induces characteristic shifts of the O−H stretching mode of surface hydroxy groups. Acetonitrile (CH3 CN), the smallest nitrile molecule, has been used as a probe for surface acid sites on various materials [62, 91–98]. However, the assignment and interpretation of the C−N stretching mode of CH3 CN (ν2 mode) in coordination compounds and H-bonded species is complex because the ν2 mode is in Fermi resonance with the (ν3 + ν4 ) combination band [91]. An experimental solution to this problem is to use perdeuterated acetonitrile (CD3 CN). Pelmenshikov and coworkers [99, 100] and Kubelkov´a and coworkers [90, 93, 94] demonstrated that acetonitrile forms H-bonded complexes with acidic hydroxy groups in zeolites. This conclusion, drawn from the experimental infrared spectra, was supported by quantum chemical calculations. The H-bond interaction of CD3 CN with acidic hydroxy groups in HZSM-5 gave three bands at ca. 2800, 2400 and 1700 cm−1 (so-called A, B, C bands well-known for strong H-bond complexes in vapors, liquids and solids) [101, 102]. The difference spectrum obtained after adsorption of CD3 CN on H-ZSM-5 is shown in Fig. 1. The negative band at 3610 cm−1 indicates the perturbation of the bridging OH groups, which leads to the broad features in the frequency range 1500–3200 cm−1 . The three bands seen at 2770, 2400 and ca.1700 cm−1 are caused by Evans transmission windows at 2600 and 1900 cm−1 in the broad ν(OH) ± kν(OH· · ·B) superposition band (ν1/2 ≥ 800 cm−1 ) of the H-bonded complexes. These windows have been shown [101] to result from Fermi resonances of the δ(OH) in-plane overtone at ca. 2600 cm−1 and the γ (OH) out-of-plane overtone at ca. 1900 cm−1 with the ν(OH) + kν(OH· · ·B) modes in these frequency regions. The range around the 1700 cm−1 band may be complicated by combination modes δ(OH) + ν(TO), the latter being a lattice vibration References see page 1158

1140

3.2 Chemical Properties

of the OH group oxygen atom. The possible occurrence of Evans windows must be taken into account when OH stretching frequency shifts are determined from experimental spectra and interpreted in terms of acid strength. It should be noted that similar phenomena have also been observed for several other H-bonded adsorption systems [101, 102]. The simultaneous existence of Lewis and Brønsted acid sites in metal-substituted AlPOs has been probed with acetonitrile [103, 104]. Acetonitrile is relatively reactive, particularly on basic oxide surfaces which expose nucleophilic hydroxy groups [91]. Its applicability as a probe molecule must therefore be carefully tested for each individual material. The surface reactivity of acetonitrile can be suppressed by substitution of the methyl group for stronger electronreleasing groups. tert-Butyl cyanide has been shown [91] to adsorb on alumina by coordination to cus Al3+

2770

2400

Absorbance

0.4

0.2

0.0 1500

2000

2500

3000

3500

4000

Wavenumbers/cm−1

Spectral changes in the 1500–4000 cm−1 range of H-ZSM-5 induced by adsorption of CD3 CN (0.05 mbar, 295 K). (Adapted from Ref. [102].)

Fig. 1

cations, in addition to weak H-bonding with no chemical transformation occurring below 500 K. Successful applications of tert-butyl cyanide were reported by Scokart and Rouxhet [97] and Lavalley and Caillod [105]. Odenbrand et al. [106], in a study of silica–titania mixed oxides, also used tert-butyl cyanide as a probe molecule. Alkanenitriles with a range of alkyl groups of different structures and sizes have been used to test the accessibility of acid sites in confined spaces in zeolites and to discriminate intrazeolitic acid sites from those located on the external surface [107–110]. e Ethers Ethers can coordinate to Lewis acid sites or undergo H-bonding interactions via the oxygen lone-pair electrons. The C−H stretching modes in dimethyl ether respond sensitively to the oxygen environment and shift toward higher frequencies when the lone pair electrons become engaged in coordination bonds. Lavalley and Caillod [105] used deuterated dimethyl ether (CD3 OCD2 H) as a probe molecule, because complications due to Fermi resonance occur with the undeuterated compound. C−H stretching frequency shifts to higher values of 37 and 70 cm−1 have been reported when CD3 OCD2 H was adsorbed on ZnO [105] and on Al2 O3 [105, 111], respectively, whereas H-bonding on silanol groups of SiO2 led to a shift of only 28 cm−1 [105]. The reported C−H stretching frequencies of CD3 OCD2 H adsorbed on several oxides are compared with corresponding frequency shifts of coordination sensitive normal modes of other probe molecules, namely pyridine, ammonia and tert-butyl cyanide in Table 2. The relatively soft base CD3 OCD2 H may find further application as a probe molecule for surface acid sites, although the steric situation may not be favorable. Di-tert-butylphenyl ether has been used to study the acidity of the outer surface of MOR and BEA zeolites [112].

Characteristic wavenumbers (cm−1 ) of different probe molecules adsorbed on silica, zinc oxide and alumina (adapted from Ref. [105])

Tab. 2

Oxide

Pyridine

Ammonia

CD3 OCD2 Ha

(CH3 )3 CCN

ν8a : 1582

ν

δs : 988

ν

νC≡N : 2238

ν

νCH : 2886

ν

SiO2 ZnO

1596b 1605

+14 +23

+13 +37

2914b 2923c

+28 +37

1614d

+32

(+62) +219 +194 +237

2251b 2275

Al2 O3

1050 1207 1182 1225

2299

+61

2956c

+70

a The

wavenumbers reported for the ether are relative to the most intense νCH band. They correspond to following amounts adsorbed: 5 µmol g−1 (ZnO); 40 µmol g−1 (Al2 O3 ); 270 µmol g−1 (SiO2 ). b After desorption at room temperature at 10−2 Pa. c Band shifts to lower wavenumbers on further addition of ether. d Accompanied by a shoulder towards higher wavenumbers.

f Ketones and Aldehydes These molecules are less sterically crowded around the carbonyl oxygen than are ethers around the ether oxygen atom and they are soft bases. Their reactivity, however, may create major problems [6]. The carbonyl stretching mode responds sensitively to intermolecular interactions, namely coordination to Lewis acid sites and H-bonding to Brønsted sites. Shifts to lower frequency of up to 150 cm−1 are induced when acetone is coordinated to cationic centers while H-bonding leads to shifts of 20–40 cm−1 . Even strong solid Brønsted acids do not protonate ketones, although the ion-pair formation could not entirely be ruled out for the adsorption of acetone on H-ZSM-5 [93]. Examples of the use of acetone have been reported [6, 113–115]. As in the case of nitriles, substitution at the carbonyl group by electron-releasing substituents leads to reduced reactivity, as has been shown by Schulz and Kn¨ozinger [116] for the adsorption of diisopropyl ketone and di-tert-butyl ketone on γ -Al2 O3 . Aldehydes have less frequently been used as probe molecules [117, 118]. g Benzene and Substituted Benzenes Benzene and substituted benzenes are soft molecules which form π-bonds to Lewis and Brønsted acid sites. These interactions can be monitored by the perturbations of C−H stretching and ring modes and by the shift of surface O−H stretching vibrations [119, 120], the νOH values being considered as a measure of the OH–π interaction [9, 50]. A ranking of proton acidities of various surface OH groups on modified silica surfaces established in this way is as follows [121]:

BOH < SiOH < GeOH < POH

(4)

Amorphous silica–aluminas show typically only one O−H stretching band of unperturbed surface OH groups at 3750 cm−1 characteristic of silanol groups [9], although their catalytic properties indicate the existence of protic sites having much higher acid strength than SiOH groups. The adsorption of benzene resulted in two bands of perturbed OH groups (Fig. 2), which were shifted relative to the position of unperturbed groups at 3750 cm−1 by 110 and 310 cm−1 , respectively [122]. The former shift is characteristic of silanol groups and the latter is typical of OH groups bridging an Al3+ and a Si4+ atom: Si(OH)Al. Hence, two Brønsted acid centers of different acid strengths can be discriminated by their interaction with benzene, although only one O−H stretching band is detected in the adsorbate-free state. Similarly, benzene has been used to characterize the heterogeneity of acid sites in faujasites and ZSM-5 zeolites [120, 123]. There is evidence, however, that the νOH frequency shifts induced by benzene adsorption in different families of zeolites and molecular sieves may not reflect their relative

1141

Absorbance

3.2.4 Acidity and Basicity

3800

3400

3000

2600

Wavenumbers/cm−1

Infrared spectra of silica–alumina (75% SiO2 ) outgassed overnight at 873 K: dashed line, before adsorption; solid line, after adsorption of benzene; dot-dashed line, deconvoluted bands. (Adapted from Ref. [96].)

Fig. 2

acid strength in an unequivocal fashion, since larger shifts have been observed for phosphate-based molecular sieves such as MeAPO-5 than for H-ZSM-5 which is the stronger acid material [124]. The interaction of benzene with Lewis acid sites is typically stronger than the OH· · · π interactions with surface hydroxy groups [125]. The interaction with cation sites induced significant frequency shifts of C−H stretching modes when benzene and toluene were adsorbed on alkali metal-exchanged faujasites, and the frequency shifts were proportional to the size of the cation [126]. This trend is unusual since the Lewis acid strength is expected to decrease with increasing cation radius. It has been argued [32] that this is due to the fact that the interaction between benzene and a Lewis acid site is governed by the electron pair acceptor strength of the cation and by the geometric match between the π-electron system of the probe molecule and the cation. An additional interaction of C−H bonds in alkyl substituents of alkylsubstituted benzenes was proposed when the zeolite matrix contained basic lattice oxygens [127]. The available experimental evidence suggests that benzene may in fact be considered as a useful, highly specific probe molecule for the detection of protic and aprotic sites. In addition, alkyl-substituted benzenes may give some insight into the basic properties of lattice oxygen atoms in the vicinity of cation sites and they permit testing of the accessibility of acid sites. h Alkenes Alkenes have not been used as probe molecules very frequently [128–133], although they are References see page 1158

1142

3.2 Chemical Properties

to be classified as soft bases. They are able to undergo OH· · · π interactions with acidic OH groups and to form π-complexes with cus cation sites. The major drawback of alkenes is their high reactivity (isomerization, polymerization) on acidic surfaces. A few successful applications of ethene to the characterization of protic acidity in zeolitic materials have been reported [131, 133] and a correlation between O−H stretching frequency shifts induced by adsorption of ethene and CO has been observed. With this in mind, and taking the possible reactivity into account, it is preferable to probe the acid properties of catalysts with other probe molecules, such as CO (see below), and to use an alkene in addition only to test the accessibility of acid sites if the alkene is to be used as a reactant in a catalytic conversion. i Carbon Monoxide A molecule that combines the softness of benzene or alkenes with small molecular size is carbon monoxide. As a consequence of the softness of CO, heats of adsorption are typically low, and low-temperature measurements are therefore required, particularly when site saturation for quantitative determination of site densities are to be achieved. Under these low-temperature conditions, the CO probe is entirely unreactive. CO forms donor bonds with prevalent σ -character to metal centers via its 5σ orbital. Considerable contributions to the stabilization of the metal–carbon bond may come from the participation of metal d orbitals and the antibonding π ∗ orbitals of the CO ligand (π back-donation). It is the increased electron density in the π ∗ orbital originating from the π back-donation which leads to a decrease in the C−O bond order and stretching force constant and, hence, stretching frequency; this effect may overcompensate an increase of the force constant by formation of the σ bond. The σ bond, however, becomes dominant when metal centers with zero or negligible d electron densities are involved. Comparably weak bond energies and carbonyl stretching frequencies higher than the gas-phase frequency are the consequence. The relevant literature up to 2002 was exhaustively reviewed by Hadjiivanov and Vayssilov [134]. The carbonyl stretching frequency responds very sensitively to coordination on to cationic sites, leading to positive frequency shifts of up to 90 cm−1 relative to the gas-phase frequency of 2143 cm−1 when CO is coordinated to the hardest d0 cation sites such as lowcoordinated Al3+ . Negative shifts may occur according to the bonding mechanism when CO is coordinated to low-valent transition metal centers possessing finite d-electron density. The carbonyl stretching mode is also sensitive relative to H-bonding to protonic sites. A high extinction coefficient of the C−O stretching mode provides good spectral sensitivity. Hence, carbon

monoxide is an almost ideal probe molecule that can discriminate between aprotic [23, 24, 31, 135, 136] and protic [31, 135, 137] acid sites with high specificity for sites with different acid strength. It must be noted, however, that CO is sensitive to through-space dipole–dipole interactions and to throughlattice interactions between adjacent CO molecules. These interactions lead to frequency shifts to higher and lower carbonyl frequencies, respectively, with increasing CO coverage [138–140]. These shifts are negligible at sufficiently low CO coverages. Alternatively, carbonyl stretching frequencies free of contributions from dipole–dipole interactions can be measured using CO isotope mixtures in which the two components are systematically varied. Frequencies determined in this way at high coverage would, however, still be influenced by chemical through-lattice interactions. When CO coordinates to cus cation centers, it forms typically terminal monocarbonyl surface complexes at least at low coverages. However, when the cation is large enough and when it is spatially not shielded like, for example, the alkali metal ions in dehydrated Y-zeolites, diand tricarbonyl species may be formed [134]. These structures can be characterized by adsorption of 13 CO−12 CO isotopic mixtures. Several examples have been reported recently [134, 141–143] which provide information on the steric situation in the local vicinity around a cation site. Spoto et al. [144] reviewed the application of CO as a probe molecule for the characterization of the surface properties of magnesium oxides. j Alkanes Alkane adsorption for the characterization of acid properties has been reported very infrequently. Adsorption of alkanes at low temperatures leads to perturbations of acidic OH groups in zeolites producing O−H stretching frequency shifts of 100–140 cm−1 [31, 131, 145–148]. Lercher et al. [32] recommended the use of alkanes such as n-pentane and n-hexane for probing the accessibility of protic sites in zeolites. Alkanes are sorbed at protic acid sites via dipole-induced H-bonding [149]. The adsorption of methane is a special case because of the Td symmetry of the molecule. Alkanes, and methane in particular, can undergo interactions with cus cation sites via polarization of C−H bonds at low temperatures, which lead to symmetry reductions and consequently to band splittings and activation of vibrational modes which are silent in Td symmetry. Table 3 shows the correlation table for XY4 molecules in Td , C3v , C2v and C1 point groups. Cohen de Lara and coworkers [150–154], in a series of papers, studied the effect of the electric field in alkali and alkaline earth metal-exchanged A zeolites, and Huber and Kn¨ozinger [155] reported on methane adsorption on Na-Y and Cs-Y zeolites. As an example,

3.2.4 Acidity and Basicity

Fig. 3 shows the FTIR spectrum of CH4 adsorbed on CsNa-Y at 88 K. The appearance of the strong band at 2889 cm−1 [ν1 (a1 )] and of the weak band at 1526 cm−1 [ν2 (e)] and the splittings of the ν3 (f) and ν4 (f) modes are clear evidence for the symmetry reduction of the methane molecule from Td to C3v which is induced by the adsorption interaction. The most likely adsorption structure should therefore be a Cs+ −H3 C−H complex. In Na-Y, an additional splitting of the ν1 (a1 ) mode was observed, which could only be explained by the adsorption of methane on Na+ ions in different crystallographic sites [155]. Methane is thus an excellent example for the use of highly symmetric extremely soft probe molecules, for which the symmetry analysis provides direct information on the structure of the adsorption complexes. Frequencies, band splittings and intensities of normal modes which are infrared silent in the free molecule can be considered as measures of the strength of interaction which should relate to the polarizing power of the cation sites and, hence, to their Lewis acid strength. In a series of papers [156–159], Kazanski and coworkers carefully measured the intensities of C−H infrared Tab. 3

Correlation table for XY4 molecules

Point group

ν1

ν2

A1 (R) E(R) A1 (IR/R) E(IR/R)

C2v

A1 (IR/R) A1 (IR/R) + A2 (R)

C1

A(IR/R)

Absorbance

Td C3v

2A(IR/R)

ν3

ν4

F2 (IR/R) A1 (IR/R) + E(IR) A1 (IR/R) + B1 (IR/R) + B2 (IR/R) 3A(IR/R)

F2 (IR/R) A1 (IR/R) + E(IR) A1 (IR/R) + B1 (IR/R) + B2 (IR/R) 3A(IR/R)

0.2

stretching bands of short-chain alkanes adsorbed by various zeolites and oxides. They claimed that these intensities may be used as a new spectral criterion of the chemical activation of alkanes via polarization by stretching of chemical bonds. k Dinitrogen The N2 molecule is isoelectronic with CO. As a homonuclear diatomic molecule it is infrared inactive. Symmetry reduction can, however, lead to infrared activity of the N−N stretching mode when the molecule is brought into an anisotropic environment, for example by adsorption. Even though specific interactions with OH groups or cationic centers may occur, the interaction energies are very low and experiments with dinitrogen as a probe must therefore be carried out at low temperatures. The N2 molecule has been used increasingly as a probe for surface acidity because it is entirely unreactive and highly specific as a very weak base. When N2 is adsorbed on alkali and alkaline earth metal-exchanged zeolites A, the N−N stretching mode becomes IR active and is shifted to higher frequencies by 10–20 cm−1 relative to the Raman band of the free molecule at 2331 cm−1 [160]. This blue shift suggests an end-on configuration for the cation–dinitrogen interaction. Interestingly, when N2 was adsorbed on Cu-ZSM-5, a red-shifted band was observed at 2294 cm−1 [161]. This band must be due to an interaction of N2 with Cu+ centers, since the formation of this species was quantitatively blocked by coadsorption of CO which gave rise to a Cu+ −CO complex that was characterized by a carbonyl band at 2158 cm−1 . Further studies using N2 as a probe have been carried out by several research groups [131, 162–169]. The weak H-bonding interaction between N2 and acidic OH groups deserves particular interest. When N2 was adsorbed on H-ZSM-5 at 80 K, a frequency shift of the O−H stretching mode of the bridging OH groups to lower wavenumbers of νOH = 109 cm−1 was induced and the N−N stretching mode appeared at 2332 cm−1 almost at the gas-phase Raman shift [133]. Density functional calculations by Neyman et al. [166] suggested an end-on configuration 1 for the H-bonded molecule. N 1.097

N 2.057

H 0.983

H H 3200 3100 3000 2900 2800

Si

O

H 1

Wavenumbers /cm−1 Fig. 3

References see page 1158

H

Al H

H

1700 1600 1500 1400 1300

FTIR spectrum of methane (0.5 hPa) adsorbed on CsNa-Y-68 at 88 K. (Adapted from Ref. [155].)

1143

1144

3.2 Chemical Properties

The specificity of the N2 probe was demonstrated by Marchese et al. [168] when studying the Brønsted acidity of CoAPO-18 catalysts. This material contains two types of OH groups, namely POH (νOH = 3681 cm−1 ) and bridging Co(OH)P (νOH = 3573 cm−1 ) groups. In addition, the Co2+ centers may act as Lewis acid sites when the Co−O bond becomes weakened or opens under the action of a sufficiently strong base. The experiments showed that N2 did not interact with Co2+ centers, but it did specifically discriminate between the properties of the two types of OH groups, as shown in Fig. 4. The POH band was shifted by ca. 70 cm−1 and the bridging OH band by ca. 105 cm−1 , indicating that the latter was the more strongly acidic group. Simultaneously, the N−N stretching mode was activated and blue shifted. In conclusion, the dinitrogen molecule is strongly recommended as a highly specific probe molecule for the characterization of acid sites. l Dioxygen The O2 molecule may also be used as a probe for surface acidity, although as yet it has hardly been applied for this purpose. Makarova et al. [131] recorded characteristic red shifts of the O−H stretching modes in zeolites and aluminophosphates when O2 was adsorbed. Field-induced infrared absorptions of dioxygen sorbed in cation-exchanged A zeolites have been reported by Jousse and coworkers [170, 171].

3468

m Nitric Oxide The electronic structures of NO, N2 and CO are similar, and it was therefore predicted that

3573 3610

0.1

3681

Absorbance

0.2

a

b

0.0 3700

3600

3500

3400

3300

Wavenumbers /cm−1 IR spectra of the OH stretching mode region of CoAPO-18 catalyst [Co/(Co + Al + P) = 0.02]: (a) background; (b) in the presence of N2 (13.3 hPa) at 77 K. (Adapted from Ref. [168].)

Fig. 4

the NO molecule should respond to electric field perturbations as known for N2 and CO. NO was in fact recommended as a probe molecule of the environment of cationic sites [169], and a successful attempt at a comparative application of NO, N2 and CO adsorption on H-, Li-, Na- and K-exchanged ferrierite was reported [169]. Complementary experiments using CO and NO including coadsorption experiments demonstrated that cation oxidation states of transition and noble metal ions on supported catalysts could be convincingly determined [172–175]. The infrared spectra of adsorbed NO have been reviewed by Davydov [28] and Hadjiivanov [176]. n Dihydrogen Molecular H2 (and D2 ) was first used for the characterization of oxide and zeolite surfaces by Kazansky and coworkers [177–182] using a highly sensitive infrared diffuse reflectance technique. F¨orster and Frede [183] reported infrared spectra of dihydrogen adsorbed on CaNa-A zeolites. Three H−H stretching bands were detected for dihydrogen adsorbed at 77 K on η-Al2 O3 that was thermally activated at 973 K, namely at 3975, 4020 and 4105 cm−1 [180, 181]. These bands were red shifted relative to the gas-phase Raman band of H2 at 4162 cm−1 . The band at 4105 cm−1 was associated with H2 interacting with surface OH groups, and the two low-frequency bands were attributed to H2 adsorbed on surface defect sites on which the H2 molecule is thought to interact with Al−O acid–base pairs. Bands in the range 4060–4010 cm−1 were generated on mild steaming of various zeolitic materials and were assigned to framework Lewis sites [179, 180]. It was suggested by Kazansky and coworkers that a band at 4035 cm−1 could be associated with tricoordinated framework Si+ sites. Surprisingly, however, the same band position was also observed for dihydrogen adsorbed on SiO2 and Al2 O3 but not on silica–alumina. Frequency shifts of surface O−H stretching modes to lower frequencies were induced when H2 (or D2 ) was adsorbed on SiO2 [131, 184, 185] and on H-forms of various zeolites [179, 180, 185]. Quantum chemical calculations [184] showed that a T-configuration with a side-on adsorption of the H2 molecule on an SiOH group of SiO2 is energetically less favorable than the so-called F-configuration in which dihydrogen acts as a proton donor to the oxygen atom of the SiOH group. The spectra shown in Fig. 5 suggest that adsorption of D2 at 88 K on H-ZSM-5 occurs selectively on bridging OH groups (νOH = 3618 cm−1 ), resulting in a red shift of 52 cm−1 towards 3566 cm−1 [185]. Only at higher D2 pressures are the SiOH groups perturbed and suffer a frequency shift from 3748 to 3715 cm−1 . The corresponding D−D stretching vibrations were

3.2.4 Acidity and Basicity Vibrational characteristics for H2 (D2 ) adsorption on SiO2 and zeolites

Tab. 4

O−H vibrations/cm−1

Sample

νO – H (H2 )

νO – H

a νH –H

3748 3748 3620 3609 3609 3618 3618 3622 3622

3742 3743 3580 3546 3546 3566 3573 3581 3587

−6 −5 −40 −63 −63 −52 −45 −41 −35

4129 4129 4108

H-MOR

H-[Al]ZSM-5 H-[Ga]ZSM-5

3748

3715

Absorbance

3700

3566

2971 2952

0.2

3800

νDb – D

2947 4103 2952 4108 2955 4112

Ref.

ν −33 −33 −54 −47 −59 −44 −54 −39 −50

[185] [184] [184] [185] [185] [185] [185] [185] [185]

of free H2 4162 cm−1 . of free D2 2994 cm−1 .

3618

H−H bν D−D

H−H (D−D) vibrations/cm−1

νO – H (free) SiO2

aν

1145

3000

3600

3500

2950

3400

0.05

2900

3300

Wavenumbers/cm−1

IR spectra of the O−H stretching and D−D stretching regions after adsorption of D2 on H-ZSM-5 at 88K and at 0 (background), 20, 50, 200 and 500 hPa D2 . (Adapted from Ref. [185].)

Fig. 5

observed at 2952 and 2971 cm−1 , respectively. The H2 adsorption resulted in almost identical effects on the O−H stretching mode and gave H−H stretching frequencies at 4108 and 4157 cm−1 , respectively. Table 4 summarizes spectral data that have been reported for the interaction of dihydrogen with acidic OH groups in several zeolites. Recently, Zecchina et al. [186] discussed the use of H2 for probing acid sites in confined spaces of microporous materials such as zeolites, heteropoly acids and sulfonated membranes (Nafion). o Water and Methanol Water and alcohols, CH3 OH in particular, possess multifunctional properties in that they act as H-bond donors and acceptors and they can coordinate to Lewis acid sites. Moreover, these molecules

may be protonated by strong Brønsted acid sites and the formation of the hydronium ion, H3 O+ , would strongly influence the surface acidity in the presence of trace amounts of H2 O. The discussion of infrared spectra of H2 O adsorbed in zeolites continues to be controversial. Several research groups [187–189] observed doublet bands at ca. 2900 and 2450 cm−1 on adsorption of small doses of water, which were interpreted as antisymmetric and symmetric vibrations of surface H3 O+ species. However, Pelmenschikov and van Santen [190] strongly emphasized that these two bands are detected on adsorption of many other basic molecules and belong to the so-called (A, B, C) trio of OH bands, which are generated as pseudobands due to Fermi resonance phenomena (see Section 3.2.4.3.3Ad on nitriles). From their ab initio calculations, Pelmenschikov and coworkers [190, 191] concluded that the water molecule was H-bonded, forming one acceptor bond to a bridging OH group (band at 3390 cm−1 ) and one donor bond to the nearest AlOSi bridging oxygen (3695 cm−1 ). Wakabayashi et al. [192, 193] reached the same conclusion based on an infrared spectroscopic study of the water adsorption on H-ZSM-5 zeolite. Hence the preferred adsorption state of H2 O seems to be represented by structure 2a rather than 2b. H H Si

O

O

Al

H H O

2a

References see page 1158

H Si

Si

O

O +

H

−

O

Al 2b

Si

1146

3.2 Chemical Properties

Buzzoni et al. [194] investigated the interaction of the H2 O molecules with the ‘‘super-acidic’’ perfluorosulfonic membrane Nafion and succeeded in recording the full set of normal vibrations of the adsorbed H3 O+ . They were also able to demonstrate the formation of dimeric and oligomeric species at increasing water loading. Diffuse reflectance infrared spectroscopy, combined with thermogravimetry of progressively dehydrated SAPO-34, revealed that there is a stoichiometric proton transfer at room temperature from the acidic bridging hydroxy site to the physisorbed water molecule [195]. All six predicted normal modes for the H3 O+ ion in Cs or C1 symmetry could be identified. Similarly, the vibrational spectra obtained for water adsorbed on CoAPO-18 catalysts were taken as evidence for the simultaneous presence of H-bonded and protonated water molecules [196, 197], as shown in structures 3a and 3b. Marchese et al. [196] stated: ‘‘In view of the small energetic difference between structures 3a and 3b, the hypotheses of the simultaneous presence of both H-bonded and protic species appears very reasonable [198]’’.

H

H

H

O

O

H

H

+

H

O O O O Co P P O O O O O O

O O − O O Co P P O O O O O O

3a

3b

Although the interpretation of the infrared spectra of adsorbed water on strong solid acids has not yet been established unequivocally, it is believed that the spectra – although complex in nature – will provide highly relevant information on the properties of solid acid catalysts. B Characterization of Aprotic Acid Sites The characterization of aprotic sites is based on the adduct formation [Eq. (3)] between a basic (electron pair donor) molecule and a cus surface cation. a Adsorption of Pyridines on Catalytic Aluminas An excellent overview of the published spectroscopic data on this system was given by Morterra and Magnacca [34]. Partially dehydroxylated surfaces of transition aluminas expose Al3+ ions in various coordination environments [9, 199]. Pyridine coordinates to these sites, as shown by infrared spectroscopy [6, 9, 24, 25, 71, 200, 201]. An ‘‘outer’’ complex with a characteristic 8a mode at 1616 cm−1 was formed at about room temperature, which was then transformed into an ‘‘inner’’ complex via an activated process

upon heating to above 520 K, the latter being characterized by an 8a mode at 1624 cm−1 [200, 201]. The activation barrier between the two forms of coordinated pyridine was attributed to a steric restriction of the approach of the Al3+ cation to the pyridine nitrogen. In contrast, Morterra and coworkers [74, 75] did not find any evidence for an activated chemisorption of pyridine on alumina. They reported the simultaneous formation at room temperature of three chemisorbed pyridine species of different bond energies, which were characterized by 8a modes at 1598–1600, 1613–1617 and 1623–1627 cm−1 . Only the low-frequency species was observed on α-Al2 O3 . It was therefore assigned to pyridine bonded to a doubly uncoordinated octahedral Al3+ cation or to octahedral Al3+ cations being exposed in a bridging vacancy. The two highfrequency species were assigned, in order of increasing polarizing power of the adsorption sites, to coordinated pyridine bridging an octahedral and a tetrahedral Al3+ cation and to pyridine bonded to a tetrahedral Al3+ cation. The studies of Kn¨ozinger and Stolz [46] indicated that 2,4,6-trimethylpyridine can still coordinate to cus Al3+ sites despite the steric crowding around the nitrogen atom. However, the steric restrictions made a close approach between the Al3+ and N atoms impossible so that the coordination interaction of this molecule was significantly weaker than that of unsubstituted pyridine on the same surface. In fact, preadsorbed 2,4,6-trimethylpyridine was almost quantitatively displaced by the weaker base pyridine [46]. These observations practically eliminate 2,6dimethyl-substituted pyridines as selective probes for the detection of protonic sites as originally suggested by Benesi [44]. However, it should be noted that a possible protonation of 2,6-dimethylpyridine has been reported for several zeolites [202–204] and also for alumina [204]. Corma et al. [205] demonstrated that even 2,6-di-tertbutylpyridine could be used to probe the proton acid sites on the external surface of MFI and MOR zeolites and that this probe molecule could easily penetrate into the pore system of three-dimensional BEA zeolite. Paukshtis et al. [206] proposed a quantitative correlation between the heat of adsorption Hads (kJ mol−1 ) of pyridine and the frequency of the 8a vibrational mode ν8a (cm−1 ): Hads = A(ν8a − 1580) − B

(5)

where A and B are constants which are characteristic for the oxide under consideration. b Adsorption of Carbon Monoxide on Oxides Carbonyl frequency shifts varying between the gas-phase frequency (2143 cm−1 ) and values as high as 2240 cm−1 have been reported, depending on the type of oxide and its pretreatment. Among the systems studied

3.2.4 Acidity and Basicity

Stretching frequencies of CO adsorbed on various Na+ zeolites and corresponding frequency shifts relative to the gas-phase value (adapted from Ref. [232]) Tab. 5

Zeolite Na-A Na-X Na-Y Na-MOR Na-ZSM-5

Si:Al ratio

ν (C−O)/cm−1

ν (C−O)/cm−1

1.0 1.25 2.35 5.0 35

2163 2164 2172 2177 2178

+20 +21 +29 +34 +35

*

Absorbance

are the binary oxides Al2 O3 [135–137, 140, 207–210], MgO [209–214], TiO2 [140, 209–217] and Cr2 O3 [136, 214, 218, 219] and more complex multicomponent systems such as SiO2 −Al2 O3 [135, 209], solid solutions (CoO−MgO, NiO−MgO [214, 220] and CoO−ZnO [214, 221]), vanadium oxide [222] and tungsten oxide [223] supported on titania, sulfated zirconia [224, 225], SiO2 -TiO2 mixed oxides [226] and cobalt and copper spinels [227]. An exhaustive review on infrared spectra of CO adsorbed on oxide surfaces was published by Hadjiivanov and Vayssilov [134]. The carbonyl infrared spectra of CO adsorbed at liquid nitrogen temperature on several types of metal ionexchanged zeolites showed that the CO probe molecule specifically interacts with the extra-framework metal ions [228–238]. Na+ · · ·CO complexes in Y-zeolites give rise to carbonyl stretching frequencies near 2175 cm−1 . However, the exact position of this band seems to be influenced by the charge distribution in the vicinity of the extra-framework Na+ cation, as evidenced by its dependence on the zeolites structure and Si : Al ratios [232]. The data summarized in Table 5 highlight this dependence for a series of different zeolite types. It has also been found that the C−O stretching frequency is sensitive to the local environment of the Na+ cation within a given zeolite matrix [209, 233] and the carbonyl bands of Na+ · · · CO in different crystallographic sites in NaY could be resolved [233]. As expected, the carbonyl stretching frequency depends strongly on the size of the alkali metal cation [230, 231, 233]. Figure 6 shows infrared spectra of CO adsorbed on alkali metal-exchanged ZSM-5 zeolite. Clearly, the strongest band in the spectra, which is characteristic for the M+ ← CO complex, shifts from 2178 cm−1 for NaZSM-5 to 2157 cm−1 for Cs-ZSM-5. This dependence was interpreted on the basis of an electrostatic model, and the carbonyl frequencies were correlated with the electric field strength along the CO dipole axis, as can be seen from data for ZSM-5 and mordenite summarized in Table 6. The observed trends are well supported by

1147

Cs+

3

* Rb+ 2

K+ 1

Na+ 0 2250

2200

2150

2100

2050

Wavenumbers/cm−1 Infrared spectra of CO adsorbed on alkali metal-exchanged ZSM-5 at 77 K at an equilibrium pressure of ∼102 Pa. The asterisk indicates contributions from Na+ · · · CO complexes. (Adapted from Ref. [229].)

Fig. 6

C−O stretching frequencies, νCO , relative frequency shifts, νCO (with respect to the free molecule, 2143 cm−1 ), and relative experimental electric field strength, E, along the CO dipole axis (adapted from Ref. [231])

Tab. 6

Zeolite Na-ZSM-5 K-ZSM-5 Rb-ZSM-5 Cs-ZSM-5 Na-MOR K-MOR Rb-MOR Cs-MOR

νCO /cm−1

νCO /cm−1

Rm /nm

E/V nm−1

2178 2166 2162 2152 2177 2163 2159 2155

+35 +23 +19 +14 +34 +20 +16 +12

0.099 0.137 0.152 0.162 0.099 0.137 0.152 0.162

6.26 4.01 3.28 2.39 6.07 3.46 2.47 2.05

frequency–electric field strength correlations deduced from quantum chemical calculations [239]. It must be noted that in addition to the M+ ← CO complex, a less stable isomeric M ← OC complex is formed, characterized by a carbonyl frequency below the gas-phase frequency of 2143 cm−1 [240, 241]. References see page 1158

3.2 Chemical Properties

The potential of CO to discriminate between adsorption sites of different coordinations was demonstrated for transition aluminas which expose cus Al3+ in octahedral 3+ and tetrahedral sites [Al3+ 5c and Al3c , respectively (the subscripts 5c and 3c indicate five- and three-coordination, respectively)]. Carbonyl bands near 2190 cm−1 and below were associated with CO coordinated to Al3+ 5c , whereas carbonyl bands beyond 2200 cm−1 were related to CO coordinated to Al3+ 3c sites [34, 135, 136, 207, 242]. Only the high-frequency species were observed on SiO2 −Al2 O3 [135, 136, 209] which exclusively exposes Al3+ cations. Studies of multicomponent systems, such as tungsten oxide on TiO2 , are of particular interest, since here the infrared spectra of coordinated CO permit the discrimination between different elements, such as W6+ and Ti4+ [223]. The relative acid strength (hardness) of the two sites, being characterized by carbonyl stretching frequencies of 2196 and 2180 cm−1 , respectively, can be qualitatively detected by the temperature and pressure dependence of the C−O stretching band intensities and the relative abundance of the two types of sites is correlated with their intensity ratios at saturation coverage. It has been shown by theoretical considerations [243] that the carbonyl stretching frequency should increase linearly with the electric field strength at the cation (d0 electron configuration) center. Several examples of such correlations have in fact been reported [135, 136, 209, 244–246]. Zaki and Kn¨ozinger [135, 136] proposed that oxidation states and coordination numbers of cation (Lewis acid) centers can be estimated from experimental carbonyl stretching frequencies by means of a frequency–field strength correlation shown in Fig. 7, provided that the carbonyl stretching frequencies are determined free of shifts due to coverage-dependent dynamic and static lateral interactions. The electric field strength Fm in this correlation is defined as Fm = sRm −2

(2)

2220

Al3+ 3c (1) Al3+ 3c 5+ Cr5c

2000 Ti3+ 4c

v (CO)/cm−1

1148

Ti4+ 5c

3+ Cr3c

3+ Cr3c

4+ Cr5c

2180 Al3+ 5c

2160

Mg2+ 5c

3+ Cr5c

Ni2+ 5c

COgas

2140

0

0.1

0.2

0.3

Fm / a.u.

Correlation between ν(CO) and electric field strength Fm 3+ − (arbitrary units): Al3+ 3c (1) refers to Al2 O3 , Al3c (2) to SiO2 Al2 O3 ; 0 cations (•) having a d electron configuration; cations (◦, ∇) having a dn (n = 0) electron configuration. (Adapted from Ref. [136].) Fig. 7

Paukshtis et al. [249] proposed a linear correlation between the heat of adsorption Hads (kJ mol−1 ) and the carbonyl stretching frequency shift νCO relative to the gas-phase frequency (2143 cm−1 ): Hads = 10.5 + 0.5νCO

(7)

In contrast, Morterra et al. [250] reported a non-linear, parabolic correlation between carbonyl frequency shifts and enthalpies of adsorption of CO. They strongly emphasized that such correlations are real, although there might be no unique correlation curve but, rather, individual correlations for families of similar and comparable adsorbents.

(6)

where Rm is equal to the sum of the effective ionic radius of a cation in the considered coordination (data compiled by Shannon [247]) and the van der Waals radius of the C atom (0.15 nm). The parameter s is Pauling’s strength of the electrostatic bond [248], which is defined as the ratio Z : N of the cation charge Z and its coordination number N . Carbonyl frequency–electric field strength correlations have also been proposed for alkali metal-exchanged zeolites and an almost linear correlation between the carbonyl band position and (RM + RCO )−2 [230, 231] was found, where RM and RCO are the cation radius and the van der Waals radius of CO, respectively.

C Characterization of Protic Acid Sites Binary and multicomponent oxides typically contain surface hydroxy groups [9] and zeolites may contain protons as chargecompensating ions [119]. The O−H stretching frequency of unperturbed hydroxy groups depends on the electronegativity of the coordination center of the group, the number of these centers and the overall bond character of the particular material (see Chapter 3.1.3.8). The proton donor strength of a hydroxy group should increase with increasing polarizing power and number of coordination centers and, hence, with decreasing O−H stretching frequency. However, the absolute frequency cannot be taken as a direct measure of the proton donor strength or acid strength of OH groups when comparing different types of

3.2.4 Acidity and Basicity

materials, although the frequencies of different OH configurations on one oxide material may with some caution be taken as a relative measure of an acid strength sequence, as e.g. on aluminas [9, 199]. Kazansky et al. [251] suggested that the extinction coefficients of O−H stretching vibrations of bridging OH groups in zeolites may be used as a measure of the intrinsic acid strength. The proton donor or acid strength of an OH group determines the ability of the group to either protonate an adsorbed basic probe molecule to form an ion pair [Eq. (2)] or to undergo H-bonding with softer probe molecules as in Eq. (1). Whether or not proton transfer occurs, depends on the proton donor strength of the OH group and on the hardness (basicity) of the basic acceptor molecule. Paukshtis and Yurchenko [24] strongly emphasized the proton affinity as a characteristic molecular property that controls proton transfer. The proton affinity of a base is defined as the enthalpy of the gas-phase reaction H+ B −−−→ H+ + B

(8)

The respective intermolecular interaction will lead to a characteristic double minimum potential well at fixed O· · ·B distance, and the shape of the potential separating the two minima (anharmonicity) will determine the probability of proton transfer and the lifetime of the protonated species will depend on the relative depths of minima at the O atom and the B molecule, respectively. Kazansky and coworkers [252–255] first identified these aspects and suggested how to determine experimentally the anharmonicities x and dissociation energies D0 of O−H oscillators interacting with probe molecules by measuring the overtones of the H-bonded system. The data summarized in Table 7 clearly demonstrate that the anharmonicity of an H-bonded O−H oscillator (silanol group) increases with the hardness of the base B while simultaneously its dissociation energy D0 decreases. For a given material, the formation of protonated species is obviously most likely with the hardest probe molecules such as aliphatic amines, ammonia and pyridine, which can be easily detected by means of the characteristic normal modes of their protonated forms.

Thus, silica–aluminas and H-zeolites will typically form + species, whereas only NH+ can be NH+ 4 and PyH 4 detected on aluminas and neither of these protonated species is formed on silica [9]. Experimental results of this kind provide information on the ability of surface OH groups to stabilize a protonated species of the applied base molecule for a period of time which must at least correspond to one vibrational period. Softer probe molecules (or reactants in catalytic reactions) may be protonated for shorter periods of time and will thus escape detection by infrared spectroscopy. Nevertheless, such species may be relevant in proton-catalyzed reactions. It must be emphasized, however, that the stabilization of a protonated species will depend strongly on H-bonding interactions of this species with the solid ‘‘pseudosolvent’’, as discussed above for the case of the NH+ 4 ion. In certain cases, the individual protonated species may escape detection via their molecular normal modes. Instead, a diffuse continuous absorption over the entire spectral range may be observed [85]. This phenomenon has been explained by the formation of extremely highly polarizable H-bonds which are presumably produced in dimeric species of the type H3 N · · · H+ · · · NH3

or

Py · · · H+ · · · Py

References see page 1158

Vibrational frequencies, anharmonicities and dissociation energies of SiOH groups (adapted from Ref. [239])

Adsorbed molecules

ν0 – 1 /cm−1

ν0 – 2 /cm−1

νe /cm−1

x

D0 /kJ mol−1

SiOHa SiOH + C6 H12 SiOH + CH3 COCH3 SiOH + NH3

3749 3720 3460 2960

7325 7220 6400

3921 3943 3979 3670

2.3 × 10−2 3.0 × 10−2 7.6 × 10−2 12 × 10−2

502 397 146 84

surface SiOH group.

(9)

These bridges are characterized by a double minimum potential well, which can be easily deformed due to the anisotropic and inhomogeneous environment of the surface species and thus leads to a quasi-continuum of vibrational states [50, 85]. Moreover, the proton can fluctuate between the two equilibrium positions within the H-bond fairly rapidly and thus lead to a dynamic polarization. The unequivocal detection of protonated probe molecules on solid acid catalysts has been largely limited to hard probe molecules such as ammonia, amines and pyridines. Although these molecules possess very useful properties from the spectroscopic point of view, they are very unselective as regards the proton donor strength of surface acid groups and their proton affinity is typically

Tab. 7

a Isolated

1149

3.2 Chemical Properties

800 19 600

18

17 13

14

16

15 400 7 5 200

1 2

6

8

12 11 9 10

4 3

0

40

80

120

160

Donicity / kJ mol−1

Correlation between wavenumber shift (νOH ) and donicity (see text for definition) of adsorptives: (1) acetyl chloride; (2) benzoyl chloride; (3) nitrobenzene; (4) acetic anhydride; (5) benzonitrile; (6) acetonitrile; (7) isobutyronitrile; (8) propionitrile; (9) methyl acetate; (10) ethyl acetate; (11) n-butanenitrile; (12) acetone; (13) diethyl ether; (14) tetrahydrofuran; (15) dimethylformamide; (16) diethylformamide; (17) dimethylacetamide; (18) diethylacetamide; (19) pyridine. (Adapted from Ref. [260].)

Fig. 8

1 4

800

∆VOH/cm−1

far beyond that of most reactants. Softer probe molecules that only undergo H-bonding may therefore provide information that is catalytically more relevant. H-bonding will induce characteristic shifts of internal vibrations of the probe (as described above) but, more importantly, typically rather significant shifts νCO to lower wavenumbers occur in the O−H stretching region accompanied by the characteristic broadening of the O−H band and increase in integrated band intensity [48, 49]. The analysis of these spectral characteristics of H-bonded systems provides significant insight into the acid properties of protic sites on solid surfaces. A few examples are discussed below. H-bonding interactions have been studied most extensively on silica surfaces, since the single sharp band of free silanols near 3750 cm−1 permits an accurate determination of νOH relative to this band. Values of νOH fall into the range ca. 100–200 cm−1 for π-electron donor molecules and may reach values as high as 1000 cm−1 for hard molecules containing lone-pair electrons [50]. Evidently the energy of interaction increases in this order. The wavenumber shifts of SiOH groups were correlated with various molecular properties of the adsorbate, such as the pKa values in aqueous solutions [256], the ionization potentials [257, 258], Hammett’s inductive parameters [259] and Taft’s inductive parameters [258] and with the proton affinities [24]. Linear correlations were also reported between νOH and the specific H-bond interaction energy [259]. Horill and Noller [260] found a general linear correlation for various classes of adsorbate molecules between νOH and the donicity (see Fig. 8) as defined by Gutman [261] by the negative enthalpy of formation of a donor–acceptor complex between the electron-pair donor molecule (H-bond acceptor) and SbCl5 in dilute 1,2-dichloroethane. Frequency shifts νOH between 515 and 855 cm−1 were observed when substituted pyridines were adsorbed on silica [85]. These shifts were correlated with Hammett’s polar substituent parameters σ that resulted in a straight line (see Fig. 9) for a large number of substituted pyridines, which is also fitted by those probes that carry one or two methyl or ethyl substituents in the ortho position. Even the point for 2,6-diisopropylpyridine falls on the straight line within the limits of accuracy. This result seems to indicate that the so-called ortho effects do not influence the H-bonding of ortho-substituted pyridines to surface silanol groups. Steric restriction should therefore not remarkably influence the H-bond strength for ortho substituents with volumes smaller than that of the isopropyl group. There are deviations apparent for three other ortho-substituted adsorbates in Fig. 9, particularly 2-methoxy-, 2-chloro- and 2-fluoropyridine. Various reasons may account for these deviations, which most probably work together. All these substituents possess p-electron pairs and the substituent effect

∆vOH/cm−1

1150

3

7 14 15 12 6 9 11 16 13 8 17 5

700 20 10

600

21

19

500

22

18

−0.4

+0.4

0 s

Correlation of vOH with Hammett’s polar substituent parameter: (◦) para; ( ) meta; (•) ortho substitution; () 2,6-disubstitution; (∇) other disubstituted pyridines. (1) 2,4,6-Trimethyl; (2) 2,6-tert-butyl; (3) 2,4-dimethyl; (4) 2,6-diisopropyl; (5) 2,6-dimethyl; (6) 2,5-dimethyl; (7) 4-tert-butyl; (8) 2-tert-butyl; (9) 4-methyl; (10) 2-methoxy; (11) 2-methyl; (12) 4-ethyl; (13) 2-ethyl; (14) 4-isopropyl; (15) 3,5-dimethyl; (16) 3-ethyl; (17) pyridine; (18) 2-fluoro; (19) 2-chloro; (20) 3-chloro; (21) 4-cyano; (22) 3,5-dichloro. (Adapted from Ref. [85].) Fig. 9

therefore may not be purely polar but also mesomeric in nature. However, a second observation probably accounts

3.2.4 Acidity and Basicity

predominantly for the deviations of these compounds. Besides the broad and intense absorption band of the silanol groups interacting with the nitrogen lone-pair electrons, weaker and less displaced absorptions are observed in the infrared spectra on adsorption of the three pyridines, indicating dual interaction with silanol groups. These interactions will mutually affect each other. In solution, the wavenumber shifts of two H-bond donors can be correlated by plotting the relative wavenumber shifts ν/ν0 of one donor against those of another [262]. Such graphs give linear relationships [Bellamy, Hallam and Williams (BHW) plots] including chemically very different H-bond acceptors, provided that the proton-bearing atoms of the donors are identical [263]. When using p-fluorophenol, phenol and methanol as reference compounds, Rouxhet and Sempels [264] succeeded in relating the wavenumber shifts of surface silanol groups analogously. The slopes of the BHW plots could be correlated with pKa values of H-bond donors in aqueous solutions (Fig. 10) and the value for silanol groups of SiO2 surfaces also fitted the same correlation when a pKa value of 6.8 as determined by Schindler and Kamber [265] was adopted. Using the same procedure, Hertl and Hair [266] determined pKa values of 15.5 for hydroxy groups on MgO, 8.8 for BOH on silica, 7.1 for silica SiOH and 0.4 for POH on silica. Lercher and Noller [113] have shown that the wavenumber shift of the carbonyl stretching band of acetone observed upon adsorption on pure and mixed oxides correlated with Sanderson’s intermediate electronegativity [267] of the materials. Similarly, a correlation between the O−H stretching frequency shift νOH induced by

2.0

BHW slope

1.5

1.0

0.5

0 −5

0

5

10

15

20

pKa Correlation between BHW slope and the pKa of various proton donors with OH as the active group: (•) phenol and p-fluorophenol (references for the BHW slope); (◦) molecules with OH attached to a carbon atom (alcohols, phenols, organic acids); () water, acetoxime, trimethylsilanol, silica gel. (Adapted from Ref. [264].)

Fig. 10

1151

adsorption of CD3 CN on several zeolites having different Si : Al ratios and structures and Sanderson’s intermediate electronegativity was reported [90]. This correlation, however, may not be entirely unambiguous since the possible influence of Fermi resonance phenomena in the O−H stretching region (see above) was not taken into account. Carbon monoxide undergoes H-bonding interactions with surface OH groups [135] and, as discussed above, is expected to be a very specific probe at low temperatures. CO adsorption on silica has been investigated in detail by Ghiotti et al. [268] and Beebe et al. [269]. Undoubtedly, an H-bonded complex 4 is formed at 77 K, as indicated by the decrease in intensity of the band of unperturbed OH groups and the appearance of an intense broad band shifted to lower frequency by 93 cm−1 [269]. Simultaneously, the CO stretching mode was observed at 2158 cm−1 , i.e. shifted to higher frequency by 15 cm−1 as compared with the gas-phase frequency of CO. CO H O Si 4

Aluminas are amphoteric and their surfaces contain five different OH configurations, which give rise to five individual O−H stretching bands [9, 199] (see also Chapter 3.1.3.8). Figure 11 shows a low temperature spectrum of γ -Al2 O3 (calcined at 773 K) in which these five bands can be clearly recognized [137]. After admission of CO, the two bands at 3785 and 3775 cm−1 remain almost entirely unaffected, whereas the three bands at lower frequency, namely at 3725, 3715 and 3695 cm−1 , are completely eroded. Instead, a partially resolved triplet appears at lower frequencies with band positions at 3655, 3635 and 3600 cm−1 . The relative intensities of individual bands in the two triplets permit a correlation to be made between the unperturbed and associated OH groups, which is summarized in Table 8 together with the corresponding frequency shifts νOH . The results indicate – in agreement with the sequence of acid strengths predicted on the basis of the estimated net charges of the free hydroxy groups – that the lowfrequency bands represent the stronger H-bond donor groups. Moreover, the weak interaction with CO permits the distinction between AlOH· · ·CO surface complexes of only slightly different H-bond energies. The increasing frequency shift for the individual associated complexes as the O−H stretching frequency of the unperturbed groups References see page 1158

1152

3.2 Chemical Properties

3600

3655

Transmission

3695

3635

3725 3715

3785

3775

g-Al2O3

3640

MgO

3746

10%

3800

3500

Wavenumbers/cm−1 IR hydroxy spectra of γ -Al2 O3 and MgO taken before (solid line) and after (dashed line) adsorption of CO (53 hPa) at 80 K. (Adapted from Ref. [137].)

Fig. 11

Hydroxy band positions (cm−1 ) of free and associated OH groups on γ -Al2 O3 before and after CO adsorption at 77 K

Tab. 8

Isolated OHs Type I

Type II

Associated OHs

νOH

3785 3775 3655 3635 3600

70 80 95

Type III

3785 (Ib) 3775 (la) 3725 (IIb) 3715 (IIa) 3695

decreases, provides strong support for the increasing H-bond donor capacity with increasing coordination number of the OH groups. The value of 95 cm−1 for the lowfrequency type III hydroxy groups is almost identical with the frequency shift for silanol groups (see above [269]) and, hence, suggests that the corresponding OH groups

should possess an H-bond donor strength (acid strength) comparable to that of SiOH groups. It is interesting that the two high-frequency (type Ia and Ib) hydroxy bands remain unperturbed by CO. This behavior resembles that of the OH groups of MgO (see Fig. 11) and CeO2 and clearly indicates their strongly basic nature [135, 136]. Such detailed information of the individual hydroxy groups coexisting on oxide surfaces can hardly ever be obtained when stronger bases are used as probe molecules. Silica–aluminas are known to develop stronger protonic acidity than pure silicas [4]. One would therefore expect a larger value for νOH . However, low-temperature infrared spectra of CO adsorbed on SiO2 −Al2 O3 (‘‘High-alumina’’ from Akzo Research) gave a broad band characteristic of H-bonding with a maximum at 3658 cm−1 , shifted relative to the position of unperturbed OH groups at 3738 cm−1 by only 70 cm−1 [137]. In agreement with earlier reports [270], bands characteristic of free OH groups associated with Al3+ were not detected. The small frequency shift seems to be in contrast with expectation. A closer comparison of the hydroxy spectra of SiO2 −Al2 O3 [268] and of SiO2 [269] after CO adsorption shows, however, significant differences. The band, typical for associated OH groups on SiO2 [268, 269], is relatively symmetric and the absorbance has declined to the baseline at wavenumbers below 3550 cm−1 . In contrast, the corresponding band for SiO2 −Al2 O3 [137] shows some structure on its low-frequency flank and a finite absorption extends to wavenumbers even below 3400 cm−1 . This suggests that the OH groups on SiO2 −Al2 O3 are heterogeneous in nature, although individual O−H stretching bands cannot be resolved in the absence of the H-bonded adsorbate. The broad absorption extending from 3600 cm−1 towards lower frequencies particularly, suggests the existence of OH groups with high H-bond donor strength. These groups might well be considered tentatively as those developing protonic acidity. This example nicely documents that the fundamental O−H stretching frequency of unperturbed OH groups cannot be taken as a measure of acid strength because the 0 → 1 transition does not reflect the anharmonicity of the potential at higher energy which is relevant for the intermolecular interactions. Hydrogen bonding also occurs when CO is adsorbed at 77 K on zeolites and leads to frequency shifts νOH of the O−H stretching mode of bridging Si(OH)Al groups of 250–340 cm−1 , depending on the structure and the composition of the zeolite. Simultaneously, the CO stretching mode shifts to higher frequency relative to the gas-phase frequency by as much as 30–35 cm−1 [135, 233, 270–276]. Density functional calculations have shown that the OH· · ·CO is energetically more favorable than OH· · ·OC [166].

3.2.4 Acidity and Basicity

1153

OH stretching frequencies of zeolites prior to and after CO adsorption at 77 K and the corresponding frequency shifts, νOH

Tab. 9

Zeolite

νOH /cm−1

SiO2 :Al2 O3 ratio

H-[Al]ZSM-5

24 13.6 26.8 41.5 –

OH· · ·CO

3620 3617 3619 3617 3623 3652 3643 3647 3650 3640

3305 3305 3290 3295 3330 3377 3347 3369 3385 3355

3324 3623

Absorbance 3800

3600

3400

Ref.

315 312 339 322 293 275 296 278 265 285

[270] [271] [275] [275] [275] [270] [271] [271] [271] [272]

The examples presented clearly demonstrate the potential of CO as a highly specific probe molecule for the characterization of protic acidity on surfaces, which permits a ranking of acid strength to be established on the basis of νOH shifts.

3747

H-[Ga]ZSM-5 H70 Na30 -Y H95 Na5 -X H70 Na30 -Y H20 Na80 -Y H90 Na10 -Y

Free OH

νOH /cm−1

3200

3000

Wavenumbers /cm−1 Hydroxy infrared spectra of H-[Ga] ZSM-5 (SiO2 : Ga2 O3 ratio 26.8) at 80 K (below) prior to and (above) after adsorption of CO (1 mbar). (Adapted from Ref. [275].) Fig. 12

Figure 12 shows the effect of CO adsorption on the bridging OH stretching mode in H-[Ga]ZSM-5 (SiO2 : Ga2 O3 ratio 26.8). Clearly, the band at 3623 cm−1 corresponding to Si(OH)Ga groups is shifted by H-bonding to CO to 3324 cm−1 , while the band of SiOH at 3747 cm−1 remains almost unperturbed at the low CO pressure of the experiment. Table 9 summarizes reported data on CO adsorption on HNa-Y, H-[Al]ZSM-5 and H-[Ga]ZSM-5. The data demonstrate that the νOH value measured upon adsorption of CO for the Ga form of ZSM-5 is significantly lower than that obtained for the Al form, consistent with the known relative acid strengths of these two materials. The data in Table 9 also show that the acid strength of HNa-Y zeolites increases with decreasing Na+ content and that that of H-[Ga]ZSM-5 falls into the range of the highly H+ -exchanged Y-zeolites.

D Quantitative Determination of Acid Site Densities Site densities can be determined when the extinctions of characteristic vibrational modes can be measured with sufficient accuracy. The sensitivity of the infrared transmission–absorption technique is dependent on the extinction coefficient of the chosen vibrational mode of a surface group (e.g. O−H group) or adsorbed probe molecule. Extinction coefficients may vary from about 5 × 10−18 cm2 per molecule for the carbonyl stretching mode in CO ligands to between 10−20 and 10−19 cm2 per molecule for C−H stretching modes in saturated hydrocarbon chains. A high surface-to-volume ratio of the catalyst is certainly desirable for high sensitivity, the more so as the possible increase in sample thickness is limited by concomitant increasing energy losses by absorption and scattering. Thus, the sensitivity of transmission infrared spectroscopy depends strongly on the nature of the surface group or probe molecule and on the physical properties of the solid sample. Assuming typical values of 100 m2 g−1 for the catalyst surface area, of 20 mg cm−2 for the weight of solid material within the irradiated geometric area, of 10−19 cm2 per molecule for the extinction coefficient and of 5% for a desirable absorption in order to obtain good quality spectra, the estimated lower detection limit for the coverage is θ = 0.02. This shows that transmission infrared spectra can be obtained at coverages below 10%. Quantitative measurements of site densities are thus undoubtedly possible, provided that References see page 1158

1154

3.2 Chemical Properties

the extinction coefficients of surface groups or adsorbed molecules are known and the Lambert–Beer law is applicable. It must be kept in mind, however, that this law is strictly valid only for optically homogeneous materials and deviations may occur for disperse substances. The extinction coefficients for OH groups are unfortunately available for only a few materials with relatively poor accuracy. Values of 12.2 and 19.5 cm µmol−1 have been reported for the O−H stretching bands at 3640 and 3550 cm−1 in HNa-Y zeolite, respectively [277], whereas Makarova et al. [274] measured values of 3.2 and 8.5 cm µmol−1 , respectively. For silanol groups of silica surfaces, values between 9.8 and 35 cm µmol−1 [278, 279] were found. The relative numbers of protic and aprotic acid sites can be estimated if the ratio of the extinction coefficients ε is known for two bands that are characteristic for the protonated and coordinated forms of the probe molecule. Attempts have been made to evaluate this ratio for pyridine normal modes, using the 19b mode, which occurs near 1450 cm−1 for coordinated pyridine species and near 1545 cm−1 for the pyridinium ion [76, 277, 280–282]. The most reliable value is probably that given by Hughes and White [277]: ε(1450) = 1.08 ± 0.09 ε(1545)

(10)

which was later verified by Matulewicz et al. [76], who found a ratio of 0.9 ± 0.1. Using the integrated absorption intensities reported by Hughes and White [277], Ward and Hansford [283] estimated the limits of detectability of Brønsted acidity to be of the order of magnitude of 10−2 meq g−1 for silica–aluminas with BET surface areas between 350 and 500 m2 g−1 . Bielanski and Datka [284] reported an integral extinction coefficient of (1.06 ± 0.07) cm µmol−1 for the 1545 cm−1 band of the pyridinium ion adsorbed on decationized Y zeolites. The integral intensity of this band and that of the O−H stretching band of bridging OH groups in H-mordenites and H-ZSM-5 was shown to correlate linearly with H+ concentrations as determined by conductometric titration [285, 286]. Thibault-Starzyk et al. [287] used a combined IR/TG method for the determination of extinction coefficients of adsorbed pyridine. Jacobs and Heylen [45] and Matulewicz et al. [76] also estimated the relative extinction coefficients for the corresponding normal modes of coordinated (1579 cm−1 ) and protonated (1642 cm−1 ) 2,6-dimethylpyridine, which are also very close to unity. Onfroy et al. [288] reported extinction coefficients for 2,6-dimethylpyridine adsorbed on several binary oxides and supported oxides. Extinction coefficients for CO coordinated to cation sites on a series of transition metal oxides have been

reported [289]. It must be noted, however, that the extinction coefficient for the carbonyl stretching frequency of coordinated CO ligands depends on the strength of the σ -donor bond in the donor–acceptor complex and hence on the C−O stretching frequency. Brown and Darensbourg [290] pointed out that adduct formation with CO via 5σ donation from the carbon end should shift the electron distribution in CO towards that characteristic of a homopolar molecule so that the dynamic dipole moment of the CO ligand would be reduced. As a consequence, the dynamic dipole moment may pass through zero and reverse sign at increasing frequency, i.e. increasing adduct bond strength. These predictions have in fact been supported by experimental results for CO adsorption on ZnO and CuO/SiO2 [291], TiO2 [250], NiO and MgO [292] and for alkali metal halides [292, 293], which showed that the dynamic dipole moment passed through zero at a C−O stretching frequency near 2175 cm−1 . However, as pointed out by Morterra et al. [74], extinction coefficients of adsorbed species cannot be transferred from one system to another even if chemical and spectroscopic characteristics are similar. In contrast to the examples mentioned above, it has been shown that the extinction coefficients of CO coordinated to alkali metal ions [294] or to Lewis sites [276] in ZSM-5 zeolites decrease with increasing carbonyl frequency and pass through zero at approximately 2250 cm−1 . This trend is consistent with the theoretical calculations of Hush and Williams [243] based on a purely electrostatic model, while being in contrast to the results of Gruver and Fripiat [295], who concluded, using first-order perturbation theory, that the CO extinction coefficient should increase linearly with increasing frequency. Neyman and R¨osch [296] confirmed that purely electrostatic interactions should result in a decrease in the extinction coefficient of CO coordinated to Mg2+ ions at a MgO surface. However, when the effect of nearest oxygen anions was taken into account in LCGTO–LDF cluster calculations [297, 298], intensity enhancements of around 30–40% were predicted. A very informative compilation of extinctions of surface carbonyl species can be found in the review by Hadjiivanov and Vaysslov [134]. In the case of dinitrogen, the extinction coefficient is clearly enhanced when N2 is coordinated to cationic sites or to H+ centers [276, 299]. Wakabayashi et al. [276] reported relative extinction coefficients for N2 interacting with acidic OH groups and CO interacting with Lewis sites and with acidic OH groups in ZSM-5, assuming that the extinction coefficient of N2 coordinated to Lewis acid sites was unity. From the above discussion, it becomes clear that there is no simple procedure to determine acid site densities quantitatively with high accuracy. In principle, extinction coefficients have to be determined for each particular

3.2.4 Acidity and Basicity

system, for example by combining volumetric or gravimetric chemisorption measurements with determinations of infrared band intensities. Characterization of Basic Sites on Solid Catalysts Basic sites on oxide surfaces may be constituted by surface 2− OH− s and Os anions. The characterization methods of basic surface sites are much less advanced than those for acidic sites. This is largely the case because suitable acidic probe molecules which fulfill the criteria mentioned in Section 3.2.4.3.2 are difficult to find. In principle, soft acids that contain X–H functions may undergo H-bonding with basic surface sites: 3.2.4.3.4

2− −−  O2− −− − −Os · · · H − X s + H − X

(11)

and the frequency shift of the X−H stretching mode in the H-bonded complex relative to that of the free acid may be taken as a measure of the H-bond acceptor or basic strength of the surface sites. Eventually, proton + − transfer with formation of an ion pair O2− s −H · · ·X may occur. Among the molecules that can be considered for this purpose are C−H and N−H acids. A major problem in the characterization of basic surface properties is the fact that many potential probe molecules possess high reactivity in basic environments and would thus lead to surface chemical transformations and in turn to modifications of the surface properties. Nevertheless, the detailed investigation of such surface reactions may also provide valuable information on the basic and nucleophilic character of catalyst surfaces. Therefore, the following discussion focuses on the potential of the H-bond method and is followed by an account of the surface chemical reactivity as related to basic character of the surface. Reviews on the characterization of basic sites using infrared spectroscopy of probe molecules were published by Paukshtis and Yurchenko [24] and more recently by Lavalley [29, 33]. A Hydrogen Bonding Method The most promising soft probe molecules are probably C−H acids. a Trihalomethanes Trichloromethane (chloroform) (CHCl3 or CDCl3 ) was first proposed by Paukshtis and coworkers [300, 301] and later advocated by Berteau et al. [302] as a suitable probe molecule for various oxides and chemically modified aluminas [300–303]. Depending on whether or not hydroxy groups were present on the oxide surfaces, three different adsorption structures 5 were suggested. In a study of alkali metal-exchanged Y-zeolites [304, 305], a strong competitive interaction of trichloromethane with the alkali metal ions was detected in addition

Cl

Cl C H

C Cl

Os

Cl

Cl H

Cl Cl

Cl C

Cl H

H

H Os

Os 5a

1155

5b

Os 5c

to the formation of the hydrogen-bonded structure 5c. The latter gave C−D stretching frequencies of 2225, 2220, 2205 and 2173 cm−1 for Na−Y, K−Y, Rb−Y and Cs−Y, respectively, which were clearly shifted to lower frequencies relative to the gas-phase frequency of CDCl3 at 2253 cm−1 . The decreasing C−D frequency with increasing cation radius may be taken as evidence for the increasing base strength of the framework oxygen atoms. Further examples of the use of trichloromethane as an acidic probe are investigations of the basic properties of magnesia [306] and of activated hydrotalcites [307]. Trifluoromethane has also been tested as a C−H acid probe molecule [233, 306, 308]. Blue shifts of the C−H stretching mode were consistently observed for the Na+ forms of several different zeolite frameworks and for the entire series of alkali metal-exchanged Y-zeolites [233]. This is unequivocal evidence for cation–fluorine interactions, suggesting that CF3 H is not selective for basic sites when cus cations are present on the surface. It must be emphasized that trihalomethanes can undergo basic hydrolysis reactions [309] and Gordymova and Davydov [310] observed decomposition products when CHCl3 and CDCl3 were adsorbed on γ -Al2 O3 at room temperature. Therefore, such reactivity of trichloromethane may restrict its general applicability as an acidic probe molecule. b Methane Adsorption of methane at low temperatures has been studied on MgO [311–313], and a series of characteristic spectra in the C−H stretching regime is shown in Fig. 13. The appearance of the symmetric CH stretching mode at 2894 cm−1 and the broadening of the antisymmetric CH stretching mode suggest a symmetry reduction for the adsorbed CH4 molecule, although the spectra do not yet permit a complete symmetry analysis. Coadsorption experiments with CO, which coordinates exclusively to Mg2+ cations, indicated that the CH4 adsorption remained almost unaffected by the presence of CO [311, 313]. It was therefore concluded that CH4 interacted with 2− basic O2− s ions probably via a C−H· · ·Os bond, forming a structure which has C3v symmetry. Additional adsorption References see page 1158

1156

3.2 Chemical Properties

3004

3260

Na−Y

VCH/cm−1

3250

K−Y Rb−Y

3240 3230

Cs−Y

3220 0.10

2894

Absorbance

0.03

2 hPa 1 hPa 0.7 hPa 0.3 hPa 0.2 hPa 0.1 hPa

3200

3100

3000

2900

2800

2700

Wavenumbers /cm−1 IR spectra in the C−H stretching region of methane adsorbed on MgO at 80 K (MgO pretreated in O2 at 773 K followed by evacuation at 773 K).

Fig. 13

structures in which methane bridges an Mg2+ s and an O2− site have also been proposed [312, 313]. s c Acetylene Acetylene and acetylene derivatives, RC ≡ CH [R = Me, (EtO)2 CH, Ph], have been mentioned as possible C−H acids for probing basic sites on zeolites and various modified oxides [314]. The applicability of methylacetylene (propyne) was tested by its adsorption on the series of alkali metal-exchanged Y-zeolites, which was mentioned above for the adsorption of trichloromethane [306]. As shown in Fig. 14, the ν1 (a1 ) mode (C−H stretching) which occurs at 3334 cm−1 in the free molecule is increasingly shifted to lower frequencies in the series Na−Y < K−Y < Rb−Y < Cs−Y and thus supports the increasing base strength in this sequence, as suggested already from the results of trichloromethane adsorption. The use of unsubstituted acetylene may be attractive because of the high symmetry of the molecule. However, the tendency of this molecule to undergo base-catalyzed transformations may limit its applicability. d Pyrrole Pyrrole has found relatively frequent application as an N−H acid, which can undergo H-bonding to

0.12 0.14 Ionic radius/nm

0.16

0.18

Correlation between C−H stretching frequency of methylacetylene adsorbed on alkali metal-exchanged Y-zeolites and cation radius.

Fig. 14

basic sites or form the pyrrolate anion via H+ transfer. It must be taken into account, however, that pyrrole can relatively easily polymerize or decompose, reactions that may occur on the catalyst surface or even in the liquid, thus leading to impurity adsorptions unless the pyrrole has been freshly distilled prior to use. Pyrrole adsorption on various oxides was first studied by Scokart and Rouxhet [97, 315] using infrared spectroscopy. The N−H stretching mode of pyrrole in dilute solutions is found at 3497 cm−1 and was shifted to lower frequencies when adsorbed on oxide surfaces such as Al2 O3 , MgO and ThO2 [97, 315] and CeO2 [316], suggesting that H-bonding to basic (H-bond acceptor) sites had occurred. On reduced ceria, the formation of adsorbed pyrrolate anions (C4 H4 N− ) via dissociative chemisorption was also detected [316]. Pyrrole has been strongly recommended for the characterization of basic sites in zeolites [10, 16, 317]. It could be demonstrated that the N−H stretching frequency shift of pyrrole adsorbed in alkali metal-exchanged X- and Y-zeolites correlated with the charge on framework oxygen ions as calculated with the Sanderson electronegativity equalization principle [33, 317, 318]. The detected changes in basic properties with the Si : Al ratio were also consistent with the calculated oxygen charge, except for mordenite [317]. The oxygen acceptor sites were thought to be framework atoms adjacent to the cations [318]. Huber and Kn¨ozinger [319] also observed a consistent increase in N−H stretching frequency shifts with increasing alkali metal cation radius in X- and Y-zeolites and noted that additional interactions occurred between the cations and the double bonds of the pyrrole ring. Binet et al. [316] also reported that for the more strongly basic materials such as K-X- and Cs-X-zeolites, H+ transfer with formation of H-bonded pyrrolate anions C4 H4 N− · · ·HOs occurred. These trends were later confirmed for alkali metal cation-exchanged EMT zeolites [320] and for alkaline earth metal-exchanged X-zeolites [321].

3.2.4 Acidity and Basicity

e Pyridinium Ions The pyridinium ion (PyH+ ) is typically formed on H-forms of zeolites and is thought to be H-bonded to the bridging framework oxygen atom. Kubelkov´a et al. [94] were able to correlate the N−H stretching frequency of the PyH+ · · ·O(zeol) complex with the charge on the − O(zeol) site, which is to be considered as the conjugate base of the original acidic OH group (see Section 3.2.4.3.4B). f 2,6-Di-tert-Butyl-4-methylphenol 2,6-Di-tert-butyl-4methylphenol might be tested as a probe molecule for strongly basic sites. However, very little is known at present about the behavior of this molecule on basic surfaces [322]. B Lewis Acids as Probe Molecules The chemisorption of Lewis acids such as BF3 , BCl3 and B2 H6 , especially on silica surfaces, has been studied by infrared spectroscopy [9, 323]. These molecules readily react with surface hydroxy groups and have been used to determine the OH density on silica surfaces. Because of their high reactivity, they probably cannot be used for unequivocal characterization of basic sites on catalyst surfaces. Li et al. [324] proposed boric acid trimethyl ester [B(OCH3 )3 ] as a probe molecule for the detection of Lewis basic sites. The reported infrared spectra of B(OCH3 )3 adsorbed on various oxides indicated a splitting of the degenerate vibration involving the B–O bond at 1360 cm−1 and a shift towards higher frequencies of the ν(C−O) band at 1036 cm−1 . It was concluded that the B(OCH3 )3 formed a Lewis acid–base adduct with basic oxygen anions and that the base strength of the tested oxides should decrease in the sequence

MgO, CaO > ZrO2 , TiO2 > SnO2 , Sb2 O5

(12)

Further experience needs to be accumulated before a final conclusion can be drawn as to whether or not B(OCH3 )3 can be recommended as a suitable probe molecule. C Surface Chemical Transformations Among the reactive molecules that have been used to study basic properties of solid catalyst surfaces are CO and CO2 [47], SO2 , alcohols, ketones, carboxylic acids and acetonitrile. The surface chemical transformations of these molecules have been extensively reviewed [6, 9, 29, 33, 47]. These transformations lead to new surface groups, the nature of which depends strongly on the degree of hydroxylation of the oxide surface. For example, chemisorption of CO2 on hydroxylated surfaces will typically produce surface bicarbonate HCO− 3 , whereas variously coordinated carbonates CO2− will be formed on dehydroxylated surfaces. 3 In many of the reported surface reactions, adjacent acidic

1157

and basic sites (acid–base pair sites) act in a concerted manner. This means that the information obtained from the reactive interaction of probe molecules on oxide surfaces is typically more complex than simple acid–base adduct formation can provide. However, understanding of the surface chemical reactivity may perhaps be even more catalytically relevant. Surface chemical transformations of various reactive probe molecules on alumina have been studied in detail by Kn¨ozinger et al. [325]. It was observed that typically the basic high-frequency OH groups (see Chapter 3.1.3.8) of the aluminas were consumed during these surface reactions. The mechanisms of these reactions were therefore described as a nucleophilic attack of the surface OH group on a molecule being coordinated to an adjacent Lewis acid site. The coordination of the chemisorbed molecules was thought to activate it for nucleophilic attack and it was suggested that the reaction occurred in a cooperative manner, assuming that the Lewis acid–base adduct and the basic OH group would undergo throughlattice inductive communication. Examples of this type of surface chemical transformations are shown in Scheme 1. The gas-phase products of these reactions have been detected and the infrared spectra of the adsorbates were consistent with the proposed surface ligands. 3.2.4.3.5 Conclusion Infrared spectroscopy of probe molecules is a valuable tool for the characterization of acid–base properties, at least in a qualitative manner. The choice of suitable probe molecules is much broader for the characterization of acid sites than of basic sites. The quantitative determination of site densities is possible, although it may be difficult to measure extinction coefficients with high accuracy in many cases. Spectroscopic data can often be correlated with quantities CO + OHs−

HCOOs−

CO2 + OHs−

− HCO3s

CH3(CO)CH3 + CH3CHO +

OHs−

[326] CH3COO s−

CH3COOs−

OH−s

CH3CH2OH + OHs−

+ CH4

[326, 327]

+ H2

CH3COOs−

CH3CN + OHs−

[326]

+ 2H2

[328, 329]

CH3(CO)NHs−

N Al

[326]

O

[326, 327]

OH

N

Al

Al

Scheme 1 References see page 1158

+

O O

Al

H2

1158

3.2 Chemical Properties

such as heats of adsorption and proton affinities, which should be related to the acid or base strength of surface sites. Although quantitative determination of acid or base strength appears to be extremely difficult, relative sequences of these quantities for comparable groups of solid catalyst materials can be established. New technical developments and the application of new methods will provide important insight into the acid–base properties of solid catalysts that were not hitherto available. For example, FTIR microscopy permitted the analysis of acid site distributions across large ZSM-5 crystals [330]. Time-resolved vibrational spectroscopic methods may provide information on the dynamics of H-bonded systems that could be directly related with mechanisms of acid-catalyzed reactions. Thus, O−H vibrational population lifetimes in Y-zeolites could be measured [331]. The results suggested that the lowfrequency OH groups undergo H-bonding interactions with adjacent framework oxygen atoms, whereas the high-frequency OH groups do not. Transient infrared hole-burning spectroscopy showed large variations in the homogeneous linewidth of the O–D stretching vibration of differently H-bonded hydroxy groups in zeolite catalysts [332]. The linewidth has been shown to be very sensitive to the frequency of the O−H· · ·B H-bond vibration. It is expected that particularly important information on the detailed properties (including dynamic properties) of acid catalysts will emerge from timeresolved spectroscopic methods as they become more frequently applied in catalysis research in the future. References 1. B. C. Gates, J. R. Katzer, G. C. A. Schuit, Chemistry of Catalytic Processes, McGraw Hill, New York, 1979, 464 pp. 2. C. L. Thomas, Catalytic Processes and Proven Catalysts, Academic Press, London, 1970, 284 pp. 3. B. C. Gates, in Ecyclopedia of Catalysis, I. T. Horv´ath (Ed.), Wiley, Hoboken, NJ, Vol. 2, 2003, p. 104. 4. K. Tanabe, Solid Acids and Bases, Kodansha, Tokyo, 1970, 175 pp. 5. L. Forni, Catal. Rev. Sci. Eng. 1974, 8, 65. 6. H. Kn¨ozinger, Adv. Catal. 1976, 25, 184. 7. H. A. Benesi, B. H. C. Winquist, Adv. Catal. 1978, 27, 97. 8. K. Tanabe, in Catalysis – Science and Technology, J. R. Anderson, M. Boudart (Eds.), Springer-Verlag, Berlin, Vol. 2, 1981, p. 231. 9. H.-P. Boehm, H. Kn¨ozinger, in Catalysis – Science and Technology, J. R. Anderson, M. Boudart (Eds.), Springer-Verlag, Berlin, Vol. 4, 1983, p. 39. 10. D. Barthomeuf, G. Coudurier, J. C. Vedrine, Mater. Chem. Phys. 1988, 18, 553. 11. S. Malinowski, M. Marczewski, in Catalysis, G. C. Bond, G. Webb (Eds.), Royal Society of Chemistry, Cambridge, Vol. 8, 1989, p. 107. 12. K. Tanabe, in Proceedings of the 9th International Congress on Catalysis, Calgary 1988, M. J. Phillips, M. Ternan (Eds.), Chemical Institute of Canada, Ottawa, Vol. 5, 1988, p. 85.

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3.2.4 Acidity and Basicity 308. V. I. Zakharov, A. M. Litkevich, A. A. Tsyganenko, A. Travert, A. Daudin, F. Maug´e, Poster P1-449 presented at the 13th International Congress on Catalysis, Paris, 2004. 309. J. Hine, A. M. Dowell Jr., J. E. Singley Jr., J. Am. Chem. Soc. 1956, 78, 479. 310. D. A. Gordymova, A. A. Davydov, React. Kinet. Catal. Lett. 1983, 23, 233. 311. P. Thomasson, O. S. Tyagi, H. Kn¨ozinger, Appl. Catal. A 1999, 181, 181. 312. A. A. Davydov, A. A. Budneva, S. M. Aliev, V. D. Sokolovskii, React. Kinet. Catal. Lett. 1988, 36, 491. 313. C. Li, G. Li, Q. Xin, J. Phys. Chem. 1994, 98, 1933. 314. L. M. Kustov, E. B. Uvarova, V. B. Kazansky, F. Figueras, D. Tichit, Abstracts, 2nd EUROPACAT Congress, Maastricht, 1995, Poster p-25. 315. P. O. Scokart, P. G. Rouxhet, J. Chem. Soc., Faraday Trans. 1 1980, 76, 1476. 316. C. Binet, A. Jadi, J. Lamotte, J. C. Lavalley, J. Chem. Soc., Faraday Trans. 1996, 92, 123. 317. D. Barthomeuf, J. Phys. Chem. 1984, 48, 42. 318. M. Huang, S. Kaliaguine, J. Chem. Soc., Faraday Trans. 1992, 88, 751. 319. S. Huber, H. Kn¨ozinger, unpublished results. 320. D. Murphy, P. Massiani, R. Franck, D. Barthomeuf, J. Phys. Chem. 1996, 100, 6731. 321. J. Xie, M. Huang, S. Kaliaguine, Catal. Lett. 1994, 29, 281. 322. J. Barrios, R. Rojas, A. R. Alcanrara, J. K. Sinisterra, J. Catal. 1988, 112, 528. 323. V. M. Bermudez, J. Phys. Chem. 1971, 75, 3249. 324. C. Li, S. Fu, H. Zhang, Q. Xin, J. Chem. Soc., Chem. Commun. 1994, 17. 325. H. Kn¨ozinger, H. Krietenbrink, H. D. M¨uller, W. Schulz, in Proceedings of the 6th International Congress on Catalysis, London, 1976, G. C. Bond, P. B. Wells, F. C. Tompkins (Eds.), Chemical Society, London 1977, p. 183. 326. P. Fink, Rev. Roum. Chim. 1969, 14, 811. 327. A. V. Deo, T. T. Chuang, J. G. Dalla Lana, J. Phys. Chem. 1971, 75, 234. 328. R. O. Kagel, J. Phys. Chem. 1967, 71, 844. 329. A. V. Deo, J. G. Dalla Lana, J. Phys. Chem. 1969, 73, 716. 330. F. Sch¨uth, R. Althoff, J. Catal. 1993, 143, 388. 331. M. J. P. Brugmans, A. W. Kleyn, A. Lagendijk, W. P. J. H. Jacobs, R. A. van Santen, Chem. Phys. Lett. 1994, 217, 117. 332. M. Bonn, M. J. P. Brugmans, A. W. Kleyn, R. A. van Santen, H. J. Bakker, Phys. Rev. Lett. 1996, 76, 2440.

1163

3.2.4.4.1 Introduction In recent decades, solid-state NMR spectroscopy has been developed into a very useful tool for the investigation of surface sites on activated solid catalysts [1–14]. This chapter reviews techniques suitable for the NMR characterization of Brønsted acid sites, Lewis acid sites and base sites, either directly or by application of probe molecules. Generally, solid-state NMR spectroscopy

allows the determination of the type, strength, accessibility and concentration of surface sites [1–6]. By application of sophisticated one- and two-dimensional solid-state NMR techniques, structural parameters are available [7–14]. Hydroxyl groups acting as catalytically active Brønsted acid sites can be directly investigated by 1 H magic angle spinning (MAS) NMR spectroscopy. The resolution of modern solid-state NMR spectrometers allows the determination of the types of various hydroxyl groups via their isotropic chemical shifts. As in the case of Lewis acid sites and base sites, however, the strength and accessibility of Brønsted acid sites are preferentially characterized by application of probe molecules. Often, these probe molecules are investigated by solid-state 1 H, 2 H, 13 C, 15 N and 31 P NMR spectroscopy. An important advantage of NMR spectroscopy is the quantitative behavior of the intensities of signals due to spin I = 1/2 systems, such as of 1 H, 13 C, 15 N and 31 P nuclei. This behavior is utilized to determine the concentration of surface sites using an internal or external intensity standard. The most widely applied technique for reaching the necessary resolution of solid-state NMR spectra is rotation around the magic angle. The principle of MAS is a rapid rotation of the sample around an axis in an angle with respect to the external magnetic field B0 (Fig. 1) [15–17]. Using turbines with a gas bearing system, sample spinning frequencies νrot of up to 30 kHz can be reached. Maximum line narrowing occurs for the magic angle of = 54.7◦ . Hamiltonians of solid-state interactions containing the geometric term (3 cos2 − 1) become zero or are strongly decreased under these conditions. However, narrowing by MAS strongly depends on the thermal mobility of the nuclei under study. Narrowing by MAS is possible if the correlation time of thermal mobility, τc , is large in comparison with the period of the macroscopic sample spinning, that is, τc  τrot = 1/νrot [18]. In the case of strong homonuclear dipolar interactions between more than two nuclear spins I , improved narrowing of the NMR signals can be reached by a combination of MAS with the application of a multiple pulse sequence [19]. This technique is denoted CRAMPS (combined rotational and multiplepulse spectroscopy). In Chapter 3.1.3.7, Section 3.1.3.7.3, a survey of the techniques of high-resolution solid-state NMR spectroscopy is given. To pave the way for solid-state NMR studies of surface sites on activated solid catalysts, techniques for the preparation of dehydrated samples sealed in glass ampoules or gas-tight MAS NMR rotors were developed. Several commercial glass inserts for 4-, 5- and 7-mm MAS NMR rotors are available, and techniques for

∗

References see page 1176

NMR Spectroscopy for the Characterization of Surface Acidity and Basicity

3.2.4.4

Michael Hunger∗

Corresponding author.

1164

3.2 Chemical Properties

the manufacture of symmetric rotor inserts have been proposed [5, 6]. The glass insert is filled with the catalyst material under study and, subsequently, connected with a vacuum line. The catalyst material is calcined under vacuum and is sealed immediately after the calcination or after the loading of probe molecules. Problems associated with heating of the loaded catalyst during the sealing of the glass insert can be overcome by a simultaneous cooling of the sample volume in liquid nitrogen. Another approach is based on the preparation of the catalyst under study inside a vacuum system or directly inside an MAS NMR rotor. Different types of such equipments have been described [6, 20–22]. As an example, Fig. 2 shows a vacuum system allowing the calcination of powder material in a horizontal tube inside an oven [22]. This calcination can be performed either under vacuum or in flowing gas. Immediately after the calcination or after subsequent loading of the catalyst with probe molecules, the powder material is filled into an MAS NMR rotor, sealed with a rotor cap from a plug rack and transferred to the NMR spectrometer without exposure to air. 3.2.4.4.2 Investigation of Brønsted Acidic Hydroxyl Groups by 1 H MAS NMR Spectroscopy 1 H MAS NMR signals of hydroxyl groups in calcined solid catalysts cover a range of isotropic chemical shifts from δ1 H = 0 to ca. 15 ppm (see Table 1). The lowest chemical shifts have been observed for unperturbed metal OH groups such as AlOH groups at the outer surface of γ -Al2 O3 [24] and Mg(OH)+ groups in the supercages of zeolite Y [23]. SiOH groups at the outer surface of silicate or aluminosilicate particles or at framework defects in zeolites are responsible for 1 H MAS NMR signals at δ1 H = 1.2–2.2 ppm [7, 18, 26–29]. Hydroxyl groups located in small structural units, such as in narrow pores or in the small sodalite cages of

z, B0 Rotation axis

Rotor with sample in glass insert Angle of Θ = 54.7°

Scheme of the MAS technique which is based on a rapid rotation of the sample around an axis in the angle = 54.7◦ with respect to the flux density B0 of the external magnetic field.

Fig. 1

zeolite Y, are often involved in hydrogen bonding or electrostatic interactions with neighboring oxygen atoms. In the subsequent text, these hydroxyl protons are marked with a prime ( ). Yesinowski et al. [42] introduced an empirical equation describing the dependence of the lowfield resonance shift δ1 H of hydroxyl protons involved in hydrogen bonding on the oxygen-oxygen distance dOH – O : δ1 H /ppm = 79.05 − 0.255dOH – O /ppm

(1)

Brunner and Sternberg [43] found correlations for the dependence of the 1 H NMR shift of hydroxyl protons involved in hydrogen bonds to basic atoms B on the H−B distance rHB . Calcium hydroxyl protons (CaOH ) located in the sodalite cages of zeolite Ca,Na-Y, for example, are lowfield shifted by δ1 H = 2.8 ppm with respect to the position of the unperturbed metal OH groups occurring at

Tab. 1 1 H NMR shifts and assignments of hydroxyl groups in solid catalysts. Hydroxyl protons involved in hydrogen bonding or electrostatic interactions with neighboring oxygen atoms are marked with a prime ( ) 1H

NMR shift, δ1 H /ppm

Abbreviation

−0.5 to 0.5

MeOH

1.2–2.2 2.4–3.6

SiOH AlOH

2.8–6.2

CaOH , AlOH , LaOH SiOHAl SiOH Al SiOH Al

3.6–4.3 4.6–5.2 5.2–8.0 ca. 15

SiOH

Type of hydroxyl group

Refs.

Metal or cation OH groups in large cavities or at the outer surface of particles Silanol groups at the external surface or at lattice defects OH groups bonded to extra-framework aluminum species located in cavities or channels involved in hydrogen bonding Cation OH groups located in sodalite cages of zeolite Y and in channels of ZSM-5 involved in hydrogen bonding Bridging OH groups in large cavities or channels of zeolites Bridging OH groups in small channels and cages of zeolites Perturbed bridging OH groups in zeolites H-ZSM-5, H-Beta and MCM-22 Internal SiOH groups involved in strong hydrogen bonding

[23–25] [7, 26–31, 34] [27, 28–34] [23, 24, 35–37] [5, 7, 10, 26–34, 38] [4, 5, 7, 10, 33, 38] [27, 29–31, 34, 39] [40, 41]

3.2.4 Acidity and Basicity

Bellows

Valve

Heater Valves

Plug rack

Sealing rod Adsorptive Rotor rack

Scheme of equipment for the evacuation, loading and sealing of solid catalyst samples in a gas-tight MAS NMR rotor [22].

Fig. 2

ca. 0 ppm [20, 23, 37]. According to Eq. (1), this lowfield shift corresponds to an oxygen–oxygen distance of dOH – O = 0.299 nm. Using X-ray diffraction (XRD) of dehydrated zeolite Ca,Na-Y, extra-framework oxygen atoms with an oxygen–oxygen distance to the next nearest framework oxygen atom of 0.297 nm were localized [44, 45]. In the 1 H MAS NMR spectra of dealuminated zeolites, signals of hydroxyl protons bound to extra-framework aluminum species occur at δ1 H = 2.4–3.6 ppm [25, 27, 32–34]. This assignment was supported by 1 H/27Al TRAPDOR (transfer of population in double resonance) NMR experiments [46–49], which are suitable for the detection of a dipolar coupling of 1 H and 27Al nuclei (see Chapter 3.1.3.7 of this Handbook, Section 3.1.3.7.3E). The low-field shift of these AlOH groups by ca. 2.4–3.6 ppm in comparison with the AlOH groups occurring at δ1 H ≈ 0 ppm indicates hydrogen bonding in the former case. In the as-synthesized form, high-silica zeolites of the framework type NON, DDR, AFI (SSZ-24) and MFI synthesized with organic quaternary ammonium cations show an 1 H MAS NMR signal at 10.2 ppm, which does not originate from the organic structure-directing agents [40]. This signal was assigned to silanol groups involved in internal hydrogen bonding (SiOH ) between defect sites and neighboring framework oxygen atoms. In agreement with Eq. (1), an oxygen–oxygen distance of 0.27 nm was found [40]. A similar effect was observed for the layered material RUB-18 containing strongly hydrogenbonded silanol groups occurring at the resonance position of 15.9 ppm [41] corresponding to an oxygen–oxygen distance of 0.25 nm.

Catalytically active Brønsted acid sites in zeolite catalysts are caused by bridging OH groups (SiOHAl) formed at Si−O−Al bridges in the local structure of negatively charged framework aluminum atoms. Adsorption of probe molecules [5, 26, 32] and application of the 1 H/27Al TRAPDOR NMR technique [46–49] evidenced that 1 H MAS NMR signals occurring at δ1 H = 3.6–8 ppm are due to bridging OH groups in dehydrated zeolites in the H-form. Depending on the cation exchange degree of zeolites H,Na-Y, the 1 H MAS NMR spectra consist of two signals at δ1 H = 3.6–4.0 and 4.8–5.2 ppm due to bridging OH groups located in the supercages (SiOHAl) and in the small sodalite cages (SiOH Al), respectively [5]. The larger chemical shift of the SiOH Al groups is caused by weak hydrogen bonding or electrostatic interactions with neighboring framework oxygen atoms. The 1 H NMR shift of unperturbed SiOHAl groups in large structural units, such as in the supercages of zeolite Y or 10- and 12-ring pores of zeolite H-ZSM-5 and H-mordenite, respectively, depends on the framework nSi /nAl ratio [5, 32], which affects the mean electronegativity of the zeolite framework. In Fig. 3, the 1 H NMR shifts δ1 H of unperturbed SiOHAl groups [4, 5, 7] are plotted as a function of the mean Sanderson electronegativity S m of the zeolite frameworks [50, 51]. S m is defined as the geometric average of the electronegativities Si of the atoms i. For zeolites of the composition HAlO2 (SiO2 )x , S m is calculated by means of the equation [50, 51]  1/(3x+4) x S m = SH SAl SO 2x+2 SSi

(2)

50H,Na-X 52H,Na-Y 90H,Na-Y

4.3 4.2

d1H / ppm

O-ring

1165

37H,K-ERI

4.1

H-MOR

4.0

H-ZSM-5/1

3.9

H-ZSM-5/2

3.8 3.7 4.00

4.05

4.10

4.15

4.20

4.25

4.30

Sm 1 H MAS NMR shift δ1 of bridging OH groups in large H cages and pores of dehydrated zeolites plotted as a function of the mean Sanderson electronegativity Sm . Zeolites X, Y and ERI have degrees of cation exchange of 50, 52, 90 and 37%. Zeolites ZSM-5/1 and ZSM-5/2 are characterized by different framework nSi /nAl ratios [5].

Fig. 3

References see page 1176

1166

3.2 Chemical Properties

83Mg,Na-Y

3.9

83Ca,Na-Y

72La,Na-Y 3.9 5.6

0.5 −0.5

Dealuminated H-Y

3.9 2.8 2.6 1.8

4.2

6

0.5 0.0

1.8

4.8

4

2

0

6

4

2

1.8

0

8

d1H / ppm (a)

(b)

6

4

2

6

4

2

0

d1H / ppm (c)

(d)

Fig. 4 1 H MAS NMR spectra of magnesium-exchanged (exchange degree 83%) zeolite Y (83Mg,Na-Y) (a), calcium-exchanged (exchange degree 83%) zeolite Y (83Ca,Na-Y) (b) and lanthanum-exchanged (exchange degree 73%) zeolite Y (73La,Na-Y) (c) calcined at 433 K and of dealuminated zeolite H-Y calcined at 673 K recorded at resonance frequencies of 300.13–400.13 MHz [4].

with SH = 3.55, SAl = 2.22, SO = 5.21, SSi = 2.84 and the nSi /nAl ratio x. The curve shown in Fig. 3 indicates an increasing chemical shift δ1 H with decreasing framework aluminum content or increasing mean electronegativity S m of the framework. This correlation is valid for unperturbed bridging OH groups with similar local structures, which are not involved in hydrogen bonding or electrostatic interactions with the zeolite framework. An increase in the mean electronegativity S m of the zeolite framework leads to a decrease in the partial charge at the hydroxyl protons of the bridging OH groups [50]. Therefore, an increase in the 1 H NMR shift of bridging OH groups as a function of the mean electronegativity S m can be correlated with an increase in the acid strength of these surface sites [1]. In many cases, however, the above-mentioned correlation is not valid due to steric effects causing low-field shifts of the 1 H MAS NMR signals, such as hydrogen bonding and electrostatic interactions in small cages or pores. To obtain 1 H NMR shift values with high accuracy allowing the study of the above-mentioned correlation, an internal chemical shift standard, such as adsorbed methane, has to be used. Typical 1 H MAS NMR spectra of dehydrated zeolites Y are shown in Fig. 4 [4]. The spectrum of zeolite 83Mg,Na-Y consists of signals of MgOH groups at −0.5 and 0.5 ppm, silanol groups at 1.8 ppm and bridging OH groups in the supercages and in the sodalite cages at 3.9 and 4.8 ppm, respectively (Fig. 4a). In the spectrum of zeolite 83Ca,NaY, a signal due to CaOH groups in the sodalite cages occurs at 2.8 ppm (Fig. 4b). Lanthanum hydroxyl groups of lanthanum cations and oxide complexes located in the

sodalite cages cause a signal at 5.6 ppm in the spectrum of zeolite 72La,Na-Y (Fig. 4c). The dealumination of zeolite H-Y is accompanied by the formation of hydroxyl groups at extra-framework aluminum complexes leading to a signal at ca. 2.6 ppm in Fig. 4d. The decrease of the framework aluminum content due to dealumination causes an increase of the mean framework electronegativity and, therefore, a shift of the signal of unperturbed bridging OH groups in the supercages from 3.9 ppm for the parent zeolite H-Y to 4.2 ppm for the dealuminated material (Fig. 4d). In addition to the 1 H MAS NMR signals of unperturbed bridging OH groups, the spectra of a number of dehydrated zeolites contain signals of bridging OH groups, which are low-field shifted for different reasons. In the 1 H MAS NMR spectra of dehydrated zeolites H-Y, for example, the signal at 4.8 ppm is due to bridging OH groups in sodalite cages pointing into the centers of sixmembered oxygen rings of hexagonal prisms [5, 7, 32, 38]. The low-field shift of this signal in comparison with the signal of unperturbed SiOHAl groups in the supercages occurring at 3.9 ppm is caused by the electrostatic interactions with neighboring framework oxygen atoms. Different types of bridging OH groups were also found for zeolites H-ZSM-5, H-Beta and H-MCM-22 [27, 29–31, 33, 34, 39]. Applying diffuse reflectance IR spectroscopy, Zholobenko et al. [52] observed a broad band at νOH = 3250 cm−1 in the spectrum of zeolites H-ZSM-5 (nSi /nAl = 21 and 35). As in the case of SiOHAl groups in sodalite cages in zeolite H-Y, this band was explained by perturbed bridging OH groups, which are involved in electrostatic interactions with neighboring

3.2.4 Acidity and Basicity

oxygen atoms. The same assignment was suggested for 1 H MAS NMR signals at 5–7 ppm occurring in lowtemperature and high-field 1 H MAS NMR spectra of zeolites H-ZSM-5 and H-Beta (Fig. 5) [27, 39, 49, 53]. The assignment of the broad signal occurring at ca. 5 ppm in the 1 H MAS NMR spectra of zeolite H-Beta to bridging OH groups was supported by the study of cesiumexchanged materials [29]. After exchanging zeolite H-Beta in a 0.1 M cesium chloride solution, the narrow signal at 4 ppm and the broad signals at ca. 5 ppm were absent in the 1 H MAS NMR spectrum of the dehydrated material. On the other hand, the signals of SiOH groups at 1.2–2.2 ppm were not affected by the cation exchange. Hence both signals at 4 and ca. 5 ppm are due to bridging OH groups being replaceable by cesium cations. Based on a comparison of experimental data of 1 H MAS NMR and FTIR spectroscopy of solid catalysts, the following correlation between the chemical shift δ1 H and the wavenumber νOH of the IR stretching vibrations was found [54]: δ1 H /ppm = 57.1 – 0.0147νOH /cm−1

(3)

for unperturbed hydroxyl groups. In the case of hydrogenbonded hydroxyl groups, the correlation is [54] δ1 H /ppm = 37.9 – 0.0092νOH /cm−1

(4)

According to Eqs. (3) and (4), the FTIR band observed by Zholobenko et al. [52] in the spectra of dehydrated zeolites H-ZSM-5 at νOH = 3250 cm−1 corresponds to an 1 H MAS NMR signal at δ1 H = 8.0–9.3 ppm. In the low-temperature 1 H MAS NMR spectra of zeolite H-ZSM-5, a corresponding broad signal was found at ca. 7 ppm [53]. The resolution of 1 H MAS NMR spectra is influenced by the full width at half-maximum of the central MAS . The line, which is denoted the residual linewidth ν1/2 MAS NMR central line and the spinning sidebands occurring at multiples of the sample spinning frequency may be significantly broadened in the presence of homogeneous interactions such as homonuclear magnetic dipole–dipole interactions between more than two nuclei with spin I = 1/2 [55–57]. The influence of homonuclear MAS can be magnetic dipole–dipole interactions upon ν1/2 reduced either by increasing the sample spinning frequency νrot or by the application of a multiple-pulse MAS is affected by inhomosequence [19]. Furthermore, ν1/2 geneities of the external magnetic field, misadjustment of the magic angle, thermal motions and exchange processes [18], heteronuclear magnetic dipole–dipole interaction with quadrupole nuclei [58, 59], anisotropy of the magnetic susceptibility [60] and the distribution of the isotropic chemical shift. If the signals are influenced by inhomogeneous line broadening effects, the resolution

1167

can be improved by an increase in the flux density B0 of the external magnetic field. Hence NMR experiments in high magnetic fields and application of high sample spinning frequencies are suitable tools for reaching a good resolution of 1 H MAS NMR spectra of hydroxyl groups in solid catalysts. As an example, Fig. 5 shows 1 H MAS NMR spectra of dehydrated zeolites H-Beta (a) and H-ZSM-5 (b), recorded at a resonance frequency of 800.13 MHz and with a spinning rate of 12 kHz [27]. The simulation of these spectra leads to the separation of up to seven signals of SiOH, AlOH , SiOHAl and perturbed SiOH Al groups (see Table 1). The residual linewidth of these signals is mainly caused by a distribution of the chemical shift due to small differences in the local structures of the corresponding hydroxyl groups. A suitable way to quantify the concentration of hydroxyl groups in solid catalysts by 1 H MAS NMR spectroscopy is the comparison of the signal intensities of the sample under study with the intensity of an external intensity standard. In the presence of a background signal due to proton-containing materials of the NMR probe, a spectral subtraction or a spin-echo experiment should be performed. For quantitative studies, the repetition time of the pulse experiments has to be large in comparison with the spin–lattice relaxation times T1 of the different OH species, which are of the order of 1–10 s [30]. A suitable external intensity standard should be characterized by spectral parameters similar to those of the samples under study. This allows the application of identical spectrometer parameters for the measurement of the samples and of the intensity standard. Kennedy et al. [30] proposed octakis(trimethylsiloxy)silsesquioxane, commonly known as Q8 M8 , as a suitable intensity standard for 1 H MAS NMR investigations of hydroxyl groups in solid catalysts. Often, a well-characterized and stable catalyst material, such as a dehydrated zeolite H,Na-Y with a degree of cation exchange of 35% is used [36, 61–63]. The total concentration ci of the hydroxyl groups in a catalyst material i can be calculated by [64] ci =

cst mst Ai mi Ast

(5)

with concentration cst , weight mst and total integral Ast of the 1 H MAS NMR signal of the standard and the weight mi and the total integral Ai of the 1 H MAS NMR signal of the catalyst i. The determination of the concentration of the different OH types occurring in the catalyst under study requires a separation of the 1 H MAS NMR spectrum. This signal separation gives the relative intensities and, together with the total concentration ci , the concentration of the different OH types. In the case of a References see page 1176

1168

3.2 Chemical Properties

H-Beta (nSi /nAl = 30)

AlOH′ SiOH

SiOHAl Perturbed SiOH′Al

8

6

4

2

0

d1H / ppm

(a) H-ZSM-5 (nSi /nAl = 20)

SiOHAl

AlOH′ SiOH

Perturbed SiOH′Al

8 (b)

6

4

2

0

d1H / ppm

High-field 1 H MAS NMR spectra of dehydrated zeolites H-Beta (a) and H-ZSM-5 (b) recorded at a resonance frequency of 800.13 MHz [27].

Fig. 5

sample spinning frequency νrot that is large in comparison st (ν st with the static NMR linewidth ν1/2 rot  ν1/2 ), no spinning sidebands occur and only the central line must be quantitatively evaluated. If spinning sidebands do occur st ), these sidebands also have in the spectrum (νrot ≤ ν1/2 to be included in the quantitative evaluation of the 1 H MAS NMR spectrum. 3.2.4.4.3 Solid-State NMR Characterization of the Accessibility and Strength of Brønsted Acid Sites by Probe Molecules The most frequently used method for investigating the accessibility and strength of Brønsted acid sites is the application of probe molecules with different sizes and base strengths. Complexes formed by adsorption of probe molecules at surface sites can be the first step of a heterogeneously catalyzed reaction. In some cases, therefore, it is hard to distinguish between probe molecules and adsorbates formed by a conversion of the probe molecules at active surface sites. On the other hand, also reactants of heterogeneously catalyzed reactions are used as probe

molecules at low temperatures. A survey of the most popular probe molecules for solid-state NMR investigations of Brønsted acids sites is given in Table 2. The ability to protonate strongly basic probe molecules, such as pyridine or trimethylphosphine, or to form hydrogen bonds to these molecules is utilized to distinguish between acidic and non-acidic sites. A more quantitative comparison of the acid strength of Brønsted acid sites is possible with weakly basic probe molecules, such as acetonitrile, acetone, perchloroethylene and trimethylphosphine oxide, which generally interact via hydrogen bonding. The adsorbate-induced low-field shifts δ of the NMR signals caused by the interacting surface OH groups (1 H MAS NMR) or due to interacting functional groups of the probe molecules (13 C, 15 N or 31 P MAS NMR) depend on the strength or accessibility of the Brønsted acid sites. However, a rapid exchange between probe molecules in the gas phase and hydrogen-bonded probe molecules may occur, which requires measurements at low temperatures (footnote a in Table 2).

3.2.4 Acidity and Basicity

1169

Probe molecules applied for the characterization of Brønsted acid sites. 1 H and 13 C NMR shifts are referenced to tetramethylsilane (δ1 H = 0 ppm, δ13 C = 0 ppm), and the 15 N and 31 P NMR signals are related to liquid 15 NH3 at δ15 N = 0 ppm, corresponding to liquid pyridine at δ15 N = 317 ppm and to 85% H3 PO4 at δ31 P = 0 ppm, respectively

Tab. 2

Probe molecule

Refs.

ammonium ions at δ1 H = 6.5–7.0 ppm hydrogen-bonded pyridine at δ1 H ≈ 10 ppm (SiOH) and pyridinium ions at δ1 H = 12–20 ppm (SiOHAl) 1 H: adsorbate-induced low-field shift by δ1 = 4.3 (H-Y) to 7.1 ppm (H-ZSM-5) H 1 H: adsorbate-induced low-field shift by δ1 = 3.0 (H-Y) to 4.9 ppm (H-ZSM-5) H 1 H: adsorbate-induced low-field shift by δ1 = 0.75 (SiOH) to 1.9 ppm (SiOHAl) H 1 H: adsorbate-induced low-field shift of accessible OH groups by δ1 = 0.23 H (SiOH) to 0.47 ppm (AlOH) 1 H: activation energy of the H−D exchange 13 C: hydrogen-bonded acetone at δ13 = 216.8 (H-SAPO-5) to 225.4 ppm C (H-ZSM-22) 15 N: hydrogen-bonded pyridine at δ15 = 295 ppm and pyridinium ions at N 198 ppm 31 P: δ31 = −2 to −3 ppm for TMP protonated by strong acid sites P 31 P: hydrogen-bonded TMPO at δ31 = 53 (H-Y) to 63 ppm (USY) P 31 P: δ31 = 11.1 to 14.8 ppm for PPh at accessible acid sites 3 P

[1, 4, 31] [1, 4, 5, 26, 36, 65, 66] [67, 68] [67] [69] [22, 70, 71]

1 H:

Ammonia Pyridine-d5

1 H:

Acetonitrile-d3 Trichloroacetonitrile Perchloroethenea Perfluorotributylamine Deuterated alkanes and aromatics [2−13 C]Acetone [15 N]Pyridinea Trimethylphosphine (TMP) Trimethylphosphine oxide (TMPO) Triphenylphosphine (PPh3 ) a

Resonance/effect

[72–75] [76–78] [79–81] [31, 82–88] [13, 89–91] [92]

Rapid exchange requires NMR measurements at low temperatures (T ≈ 120 K).

H-ZSM-5

Mo/H-ZSM-5 1.7 2.4 3.9 3.9 4.9

4.9 5.8

5.8

2.4 1.8

∆ d1H = 0.23 ppm

∆ d1H = 0.25 ppm

∆ d1H = 0.33 ppm

∆ d1H = 0.47 ppm

+ Perfluorotributylamine

12 (a)

9

6

3

0

d1H / ppm

12 (b)

9

6

3

0

d1H / ppm

1 H MAS NMR spectra of calcined zeolite H-ZSM-5 (a) and molybdenum-impregnated (6 wt.-%) H-ZSM-5 (b) before (top) and after adsorption (bottom) of perfluorotributylamine [22].

Fig. 6

Most of the probe molecules mentioned in Table 2 can be applied for studying the location of acid sites in the pores of solid catalysts. A probe molecule suitable to distinguish hydroxyl groups at the outer surface of zeolite particles and hydroxyl groups located inside the pores and cages is perfluorotributylamine. Perfluorotributylamine

(diameter ≈0.94 nm [22]) is too large to enter the micropores of zeolites (H-ZSM-5, 0.53 × 0.55 nm; H-Y, 0.74 nm [93]). Figure 6 shows 1 H MAS NMR spectra of calcined zeolite H-ZSM-5 before (a) and after References see page 1176

1170

3.2 Chemical Properties

(b) modification with molybdenum. The spectra were recorded before (top) and after (bottom) adsorption of perfluorotributylamine [22]. Low-field resonance shifts of the signals due to SiOH and AlOH groups by ca. 0.25 and 0.47 ppm, respectively, indicate a location of these hydroxyl groups at the outer surface or in secondary mesopores of the H-ZSM-5 particles. In contrast, no lowfield resonance shift was found for bridging OH groups occurring at 3.9–5.8 ppm, which indicates their location inside the pore system of the zeolite catalysts under study. Wang et al. [92] utilized triphenylphosphine (PPh3 ) to study Brønsted acid sites occurring at the outer surface of zeolite H-MCM-22. Only 6% of the Brønsted acid sites were detected by this probe molecule, indicating a corresponding content of these hydroxyl groups at the outer surface. The existence of two different Brønstedbound PPh3 molecules causing 31 P MAS NMR signals at 11.1 and 14.8 ppm was supported by theoretical studies [92]. It was suggested that PPh3 is able to adsorb at Brønsted acid sites in the local structure of aluminum atoms incorporated at T1 and T4 positions into the framework of zeolite H-MCM-22. The effect of adsorption of strongly basic pyridined5 on dehydrated zeolites 88H,Na-Y and H-ZSM-5 is demonstrated in Fig. 7 [4, 94]. If pyridine interacts with

acidic hydroxyl protons, such as SiOHAl groups, pyridinium ions (PyrH+ ) are formed. 1 H NMR spectroscopy of organic acids dissolved in pyridine yields signals of protonated pyridine at chemical shifts of 14–20 ppm [95]. The interaction of pyridine with SiOHAl groups in zeolites 88H,Na-Y and H-ZSM-5 causes signals of pyridinium ions at 1 H NMR shifts of 16.5 and 15.5–19.0 ppm, respectively. Simultaneously, the 1 H NMR intensity of the acidic OH groups, which are involved in the proton transfer from the zeolitic framework oxygen atoms to the basic probe molecules, is decreased. In comparison, hydrogen bonding of pyridine at non-acidic SiOH groups causes a resonance shift of the 1 H MAS NMR signal from 2 to ca. 10 ppm only [4, 94]. Therefore, adsorption of pyridine is useful for distinguishing acidic from non-acidic OH groups. In the case of zeolite 88H,Na-Y (Fig. 7b), only SiOHAl groups in the supercages are accessible for pyridine and the 1 H MAS NMR signal of SiOHAl groups in the small cages (sodalite cages), occurring at 4.8 ppm, are not influenced by this probe molecule. Upon loading of zeolite H-ZSM-5 with pyridine (Fig. 7d), two signals of pyridinium ions appear, which indicates an interaction of the probe molecules with two different types of accessible SiOHAl groups. This effect is due to the

3.8 88H,Na-Y

H-ZSM-5

4.3

4.8

2.0

20 (a)

10

0

−10

d1H / ppm

88H,Na-Y + 1C5ND5/SiOHAl 16.5

20

10

0

−10

d1H / ppm

(c)

H-ZSM-5 + 1C5ND5/SiOHAl

4.8 6.9

19.0 15.5 2.0

20 (b)

10

0

d1H / ppm

−10

20 (d)

10

0

−10

d1H / ppm

1 H MAS NMR spectra of dehydrated zeolite H,Na-Y with a degree of cation exchange of 88% (a, b) and zeolite H-ZSM-5 (c, d), in the unloaded state (a, c) and loaded with one molecule of pyridine-d5 per bridging OH group [4].

Fig. 7

3.2.4 Acidity and Basicity

presence of unperturbed and perturbed bridging OH groups in zeolite H-ZSM-5 (see Fig. 5, bottom). In a number of studies on zeolite acidity, the strong base trimethylphosphine (TMP) was used as a probe molecule [82–86]. Lunsford et al. [83] performed 31 P MAS NMR studies of TMP adsorbed on zeolite H,Na-Y calcined at different temperatures. After adsorption of TMP at room temperature, the samples were degassed at 353 K for 30 min to avoid signals of physisorbed probe molecules. The 31 P MAS NMR spectrum of TMP adsorbed on zeolite H,Na-Y, which was calcined at 673 K, is dominated by a signal at −2.5 ppm (referenced to 85% H3 PO4 ). This signal is due to (CH3 )3 PH+ complexes (TMPH+ ) arising from chemisorption of TMP at Brønsted acid sites (Fig. 8a) [83]. Zeolite H,Na-Y, which was calcined at 773 K, shows additional 31 P MAS NMR signals in the region from about −32 to −67 ppm due to coordination of TMP at Lewis acid sites (see below) [83]. On increasing the calcination temperature up to 973 K (Fig. 8c and d), the 31 P MAS NMR signal of TMPH+ complexes decreases, which indicates strong dealumination and dehydroxylation of the zeolite. When using TMP, it must be considered that this probe molecule is rather bulky (kinetic diameter 0.55 nm [83]).

1171

Therefore, the characterization of acid sites in zeolites by TMP is restricted by the maximum adsorption capacity of the pore system and by the pore diameter of the catalyst under study [84]. The local structure of an ion pair TMPH+ –zeolite complex in zeolite H-Y was investigated by Kao et al. [12]. Applying 27Al/31 P and 27Al/1 H rotational echo doubleresonance NMR experiments (see Section 3.1.3.7.3E), Al−P and Al−HB distances (HB = hydroxyl proton of the Brønsted acid site) for the acid site–TMP complex of 0.395 ± 0.005 and 0.28–0.31 nm, respectively, were determined (Fig. 9). By fitting the spinning sidebands in the 1 H MAS NMR spectrum, a P−HB distance of 0.140 ± 0.002 nm was obtained. These internuclear distances are within the range of data calculated by ab initio methods for the ion pair TMPH+ –zeolite complex, which is formed by transferring a Brønsted acidic hydroxyl proton of a bridging OH group to the adsorbed TMP molecule [12]. A frequently applied probe molecule for characterizing the strength of Brønsted acid sites in solid catalysts in a more quantitative manner is [2−13 C]acetone [76–78]. References see page 1176

−54.5

−2.5

−4.3

−5

(a)

(c)

−58 −4.2

−32 −67 ∗

100 (b)

50

0

d31P / ppm

−50

−100

100 (d)

50

0

−50

−100

d31P / ppm

Fig. 8 31 P MAS NMR spectra of trimethylphosphine (TMP) adsorbed on zeolite H,Na-Y calcined at temperatures of 673 (a), 773 (b), 873 (c) and 973 K (d) [83]. Asterisks denote spinning sidebands.

1172

3.2 Chemical Properties

Me Me H P 0.14 nm

C

HB

H

0.395 nm H 0.28 – 0.31 nm

O

O

O Al

O

Experimentally obtained distances in an ion pair TMPH+ –zeolite complex [12].

Fig. 9

Using this molecular probe, Biaglow et al. [76] studied the acid strength of bridging OH groups in various acidic zeolite catalysts and observed 13 C NMR shifts of the carbonyl atom in [2−13 C]acetone of δ13 C = 216.8 (H-SAPO–5), 219.6 (H-Y), 221.8 (H-MOR), 222.8 (H-[Ga]ZSM-5), 223.4 (H-ZSM-12), 223.6 (H-ZSM-5) and 225.4 ppm (H-ZSM-22). In comparison, the chemical shifts of the carbonyl atom in [2−13 C]acetone dissolved in CDCl3 and in 100% sulfuric acid are 205 and 245 ppm, respectively [74]. Based on the experimentally determined dependence of the resonance positions of carbonyl atoms in [2−13 C]acetone dissolved in aqueous sulfuric acid of varying concentrations, a scale of the Brønsted acid strength could be introduced [3]. It appears from the results that bridging OH groups in acidic zeolites, such as in zeolite H-ZSM-5 (δ13 C = 223.6 ppm [76]), are not stronger than 80% H2 SO4 in water [3]. A limitation of acetone as a probe molecule is its reactivity in the case of an adsorption at strongly acidic sites, which can be overcome by experiments at low temperatures. Therefore, mesityl oxide, which is a product of the conversion of acetone on acidic catalysts, is another interesting candidate for the study of the strength of Brønsted acid sites by solid-state NMR spectroscopy [3]. Jaenchen et al. [67] utilized the weak basesacetonitrile and trichloroacetonitrile as probe molecules for the characterization of the strength of Brønsted acid sites by 1 H MAS NMR spectroscopy. They studied the influence of the framework nSi /nAl ratio and the structure type of zeolites on the adsorbate-induced low-field shifts δ1 H caused by hydrogen bonding with deuterated acetonitrile and trichloroacetonitrile. As an example, Fig. 10 shows 1 H MAS NMR spectra of zeolites H,Na-Y and H-ZSM-5

recorded before and after adsorption of deuterated acetonitrile [67]. The spectra of the unloaded materials are similar to those in Fig. 7a and c, respectively. The 1 H MAS NMR signals occurring in Fig. 10a at 4.1 and 4.7 ppm were assigned to bridging OH groups in the supercages and small cages, respectively, of zeolite H,Na-Y. The signal at 4.1 ppm in Fig. 10c is due to bridging OH groups located in 10-ring pores of zeolite H-ZSM-5. Upon adsorption of acetonitrile, the signal of bridging OH groups in the supercages of zeolite H,Na-Y is shifted to a resonance position of 10.5 ppm (Fig. 10b), whereas in the spectrum of the loaded zeolite H-ZSM-5 a low-field shifted signal occurs at 11.2 ppm (Fig. 10d). The latter signal has a broad low-field shoulder, which may be explained by the presence of a second type of acidic hydroxyl protons, such as perturbed bridging OH groups. The observation that there is no influence of acetonitrile on the 1 H MAS NMR signal of bridging OH groups in the small cages of zeolite H,Na-Y (signals at ca. 4.7 ppm) is caused by the steric hindrance of these hydroxyl protons. Table 3 gives a survey on the experimentally observed low-field shifts δ1 H of the 1 H MAS NMR signals of unperturbed bridging OH groups in various zeolite catalysts upon adsorption of deuterated acetonitrile and trichloroacetonitrile [67]. Generally, the adsorbateinduced low-field shifts δ1 H are larger upon adsorption of acetonitrile in comparison with trichloroacetonitrile. The qualitative behavior of the low-field shifts, however, is similar for samples loaded with both types of probe molecules. A larger low-field shift δ1 H was found for zeolites with a higher nSi /nAl ratio, which corresponds to a higher acid strength of the SiOHAl groups, in comparison with zeolites having the same structure type and a lower nSi /nAl ratio. In addition, bridging OH groups in zeolites with different structure types but similar nSi /nAl ratios show similar adsorbate-induced low-field shifts δ1 H . This finding agrees with the data shown in Fig. 3 and their discussion: a higher mean electronegativity S m (or lower aluminum content) of the

Low-field shifts δ1 H of the 1 H MAS NMR signals of unperturbed bridging OH groups in various zeolite catalysts upon loading of deuterated acetonitrile and trichloroacetonitrile [67]

Tab. 3

Sample

H,Na-Y H,Na-Y H-mordenite H-mordenite H-ZSM-5

nSi /nAl

δ1 H upon adsorption of CD3 CN/ppm

δ1 H upon adsorption of CCl3 CN/ppm

5.0 18.0 6.7 10.0 52.0

6.4 7.0 6.2 6.7 7.1

4.4 4.9 4.4 4.7 4.9

3.2.4 Acidity and Basicity

H,Na-Y

1173

H-ZSM-5

4.1 4.7

4.1

2.8 2.0 2.0

14

12

10

8

6

4

2

0

d1H / ppm

(a)

H,Na-Y + 0.01 mbar CD3CN

14

12

10

4.7

8

6

4

2

0

d1H / ppm

(c)

H-ZSM-5 + 0.01 mbar CD3CN

11.2 10.5

14 (b)

12

10

2.0

8

6

4

2

2.0

14

0

d1H / ppm

(d)

12

10

8

6

4

2

0

d1H / ppm

1H

MAS NMR spectra of zeolite H,Na-Y with an nSi /nAl ratio of 5.0 (a) and of zeolite H-ZSM-5 with an nSi /nAl ratio of 52.0 (c) recorded before (a, c) and after (b, d) loading with 0.01 mbar (1bar = 105 Pa) deuterated acetonitrile [67].

Fig. 10

zeolite framework correlates with a higher acid strength of the unperturbed bridging OH groups in these materials. Low-temperature 1 H MAS NMR studies of zeolite HZSM-5 loaded with weak bases interacting with zeolitic hydroxyl groups via hydrogen bondings have been reported by White et al. [73], Haw et al. [96] and Brunner et al. [39]. Upon adsorption of about one probe molecule per bridging OH group in zeolite H-ZSM-5, adsorbateinduced 1 H MAS NMR low-field shifts for ethene of 2.7 ppm, for carbon monoxide of 1.8 ppm and for ethane of 0.6 ppm were found at 123 K [96]. Hence the adsorbateinduced resonance shift is extremely sensitive to the type of probe molecule used. Koch et al. [97] studied the lowfield shift of the 1 H MAS NMR signal of bridging OH groups after loading of carbon monoxide as a function of the temperature. For zeolite H-ZSM-5 loaded with one molecule carbon monoxide per bridging OH group, a low-field shift of 2.0 ppm was observed at 123 K. In this case, the adsorbate-induced wavenumber shift of the IR stretching vibration of bridging OH groups amounts to νOH = 300 cm−1 [97]. From the linear correlation between the IR wavenumbers and 1 H NMR shifts of hydrogen bonded hydroxyl protons [Eq. (4)], a corresponding low-field NMR shift of δ1 H = 2.8 ppm follows. This indicates that for adsorption of CO on zeolite

H-ZSM-5, even at 123 K, the influence of thermal motions on the 1 H resonance position of the adsorbate complex is not completely eliminated. The interaction of hydroxyl groups in zeolites with perchloroethylene has been investigated by Sachsenroeder et al. [69]. At the beginning of the 1980s, Paukshtis and Yurchenko [98, 99] developed a method, which allows the calculation of the deprotonation energy EDP of hydroxyl groups in solids by their adsorbate-induced IR wavenumber shifts, νOH . Because of the linear correlation between δ1 H and νOH , also the adsorbate-induced 1 H MAS NMR low-field shift δ1 can be used to deterH mine the deprotonation energy EDP of SiOHAl groups contributing to weakly hydrogen bonded complexes [69]: −1

EDP /kJ mol

!
! ! δ1 H,SiOHAl ! 1 ! ! = − log ! A δ1 H,SiOH !

(6)

where δ1 H,SiOH and δ1 H,SiOHAl are the 1 H MAS NMR low-field shifts induced by the adsorbate molecule for the resonance positions of silanol and bridging OH groups, respectively, and A is a constant, 0.00226 [98]. With References see page 1176

1174

3.2 Chemical Properties

δ1 H,SiOH = 0.75 ppm and δ1 H,SiOHAl = 1.9 ppm, observed after adsorption of 12 Cl2 C=CCl2 /u.c. on dehydrated zeolite H-ZSM-5, a deprotonation energy of the SiOHAl groups in zeolite H-ZSM-5 of EDP = −179 kJ mol−1 was calculated. With the same method, a deprotonation energy of EDP = −146 kJ mol−1 was found for bridging OH groups in the supercages of zeolite 30H,Na-Y [69]. An interesting approach for the determination of the strength of Brønsted acid sites under reaction conditions is the study of the H−D exchange kinetics between surface hydroxyl groups and adsorbed probe molecules or reactants [72–75]. In these experiments, for example, deuterated benzene is loaded on the activated catalyst and the growth of the 1 H MAS NMR signal of the benzene protons formed via an H−D exchange with surface hydroxyl protons is quantitatively evaluated as a function of time. In this way, the H−D exchange rates are determined for at least three different temperatures. The slope of the Arrhenius plot of the H−D exchange rates gives the activation energy EA of this reaction. Upon loading of benzene-d6 on zeolites H,Na-Y, USY and H-ZSM-5, Beck et al. [75] determined activation energies of the H−D exchange of EA = 107.0, 84.9 and 60.2 kJ mol−1 , respectively. 3.2.4.4.4 Solid-State NMR Investigations of Lewis Acid A direct study of Lewis Sites by Probe Molecules acid sites, such as extra-framework aluminum species in activated zeolites by solid-state NMR spectroscopy, is limited by the large linewidth of the corresponding 27Al NMR signals for samples in the dehydrated state. For the study of hydrated aluminum-containing catalysts by solid-state 27Al NMR spectroscopy, a number of sophisticated techniques, such as echo sequences, DOR (double-oriented rotation) and MQMAS (multiple quantum magic-angle spinning) are available [100, 101] (see also Chapter 3.1.3.7). However, the hydration of activated solid catalysts is accompanied by strong modification of the surface sites. The 27Al NMR signals of octahedrally coordinated extra-framework aluminum species in hydrated catalysts cannot be correlated in a direct manner with the occurrence of Lewis sites. Therefore, the investigation of Lewis acid sites in activated solid catalysts requires the application of probe molecules. Table 4 gives a survey on the most important probe molecules utilized for the study of Lewis acid sites in solid catalysts. Most of them are also applied for the characterization of Brønsted acid sites. In the presence of Lewis acid sites, various complexes are formed by a coordination of probe molecules at these sites and are available for solid-state NMR spectroscopy. So far, however, no correlation between the NMR shift of probe

Probe molecules applied for the characterization of Lewis acid sites. 13 C NMR shifts are referenced to tetramethylsilane (δ13 C = 0 ppm) and the 15 N and 31 P NMR signals are referenced to liquid 15 NH3 at δ15 N = 0 ppm, corresponding to liquid pyridine at δ15 N = 317 ppm and to 85% H3 PO4 at δ31 P = 0 ppm, respectively Tab. 4

Probe molecule

Resonance/effect

Refs.

[2−13 C]Acetone

δ13 C = 233 ppm for acetone at zeolite USY 13 C: δ13 ≈ 770 ppm for C CO at dealuminated zeolite H-ZSM-5 15 N: δ15 ≈ 265 ppm for N pyridine at dealuminated zeolites 31 P: δ31 = −32 to P −67 ppm for TMP at dealuminated zeolites 31 P: δ31 = 37 ppm for P TMPO at dealuminated zeolites Y and γ -Al2 O3

[77]

[13 C]Carbon monoxidea [15 N]Pyridinea

Trimethylphosphine (TMP) Trimethylphosphine oxide (TMPO)

13 C:

[102, 103]

[79–81]

[82–84, 86]

[89, 90]

a Rapid exchange requires NMR measurements at low temperatures

(T ≈ 120 K).

molecules coordinated at Lewis acid sites and the strength of these surface sites has been found. A typical example of an NMR study of Lewis acid sites by probe molecules is shown in Fig. 8 [83]. Upon calcination of zeolite H,Na-Y at 673–973 K, trimethylphosphine (TMP) was adsorbed. The 31 P MAS NMR spectra of the strongly calcined materials consist of signals at −32 to −67 ppm, hinting at the formation of different Lewis acid sites in the activated materials. After calcination at temperatures of up to 873 K, a significant increase in the 31 P MAS NMR signals of TMP coordinated at Lewis acid sites occurs. In contrast, the calcination at 973 K (Fig. 8d) leads to a decrease in the 31 P MAS NMR signals at −32 to −67 ppm in comparison with the sample calcined at 873 K (Fig. 8c). This finding indicates that the preparation of zeolite catalysts with a high concentration of Lewis acid sites requires an optimum calcination temperature, such as 873 K in the present case [83]. The local structure of adsorbate complexes consisting of TMP coordinated at Lewis acid sites in dehydroxylated zeolite H,Na-Y was investigated by Kao and Grey [14]. By 31 P/27Al TRAPDOR (transfer of population in double resonance) NMR experiments (see Chapter 3.1.3.7, Section 3.1.3.7.3E), a 31 P MAS NMR signal occurring at −47 ppm could be assigned to TMP directly bound to an aluminum Lewis acid site. The 27Al/31 P INEPT (insensitive nuclei enhanced by polarization transfer) NMR technique was utilized to determined an 27Al/31 P J -coupling of 270 ± 10 Hz between the above-mentioned nuclei [14]. Upon adsorption of TMP on mixtures of

3.2.4 Acidity and Basicity

AlCl3 and zeolites H-Y and H,Na-Y, in which the AlCl3 molecules coordinated to one (four-fold coordinated aluminum) and two (five-fold coordinated aluminum) TMP molecules, respectively, two different TMP–AlCl3 complexes were studied [104]. 31 P/27Al J -coupling constants and internuclear distances of 299.5 Hz and 0.258 nm and of 260 Hz and 0.296 nm, respectively, were determined for the two above-mentioned adducts. The 27Al/31 P J -coupling constant of 270 Hz obtained for TMP in dehydroxylated zeolite H,Na-Y (see above) is intermediate between the values for the four- and five-fold coordinated TMP–AlCl3 complexes. Considering the high electronegativity of oxygen in comparison with chlorine, however, the aluminum in the TMP–Lewis acid adduct was assigned to a five-fold coordinated aluminum site [14]. This suggests that the pure Lewis acid site comprises a tetrahedrally coordinated aluminum site. Solid-State NMR Investigations of Base Sites in Solid Catalysts Due to the improvement of the NMR spectroscopic technique during the past decade, solidstate 17 O NMR spectroscopy became a suitable method for the direct investigation of basic oxygen atoms in zeolite catalysts. Superconducting magnets with flux densities of up to B0 = 18.8 T and the application of the DOR (double-oriented rotation) and MQMAS (multiplequantum magic angle spinning) techniques overcame the problems of strong signal broadening by the second-order quadrupolar interaction of 17 O nuclei having a nuclear spin of I = 5/2 [101, 105–108]. For low-silica faujasite with a framework nSi /nAl ratio of 1 (LSX), which contains Si−O−Al bridges only, 17 O DOR and MQMAS NMR spectra show four lines due to oxygen atoms at the four different crystallographic positions. For oxygen atoms at O-1 sites in zeolite LSX, a correlation was found between the 17 O chemical shift, δ17 O , and the cation radius, r, in the hydrated zeolites Li-LSX, Na-LSX and Cs-LSX [107]: 3.2.4.4.5

δ17 O /ppm = 9.2r/A˚ + 41.47

(7)

This indicates that the framework oxygen atoms in zeolites LSX are affected by the introduction of different alkali metal cations. The investigation of zeolites with nSi /nAl > 1 requires a further improvement of the NMR spectroscopic techniques since the presence of Si−O−Al and Si−O−Si bridges leads to a superposition of at least eight signals, which are, in addition, broadened by a distribution of chemical shifts. At present, the most widely applied technique for solid-state NMR investigations of base sites in solid catalysts is the application of probe molecules, which are summarized in Table 5. Bosch et al. [109] studied the strength of base sites in zeolites Li-Y, Na-Y, K-Y, Rb-Y and Cs,Na-Y by adsorption of trichloro- and trifluoromethane

1175

Probe molecules applied for the characterization of base sites. 1 H and 13 C NMR shifts are referenced to tetramethylsilane (δ1 H = 0 ppm, δ13 C = 0 ppm)

Tab. 5

Probe molecule Trichloromethane

Trifluoromethane

Pyrrole

Chloroform

13 C-Chloroform

13 C-Methyl

iodide

13 C-Nitromethane

Resonance/effect

Refs.

hydrogen-bonded trichloromethane at δ1 H = 7.55 (Li-Y) to 8.23 ppm (Cs;Na-Y–90) 1 H: hydrogen-bonded trifluoromethane at δ1 H = 6.62 (Li-Y) to 7.6 ppm (Cs,Na-Y–90) 1 H: hydrogen-bonded pyrrole at δ1 H = 8.4 (Li-Y) to 11.5 ppm (K-X) 1 H: hydrogen-bonded chloroform at δ1 H = 7.45 (H-Y) to 8.70 ppm (Na,Ge-X) 13 C: hydrogen-bonded 13 C-chloroform at δ13 C = 77.9 (H-Y) to 81.7 ppm (Na,Ge-Y) 13 C: methoxy groups occurring at δ13 C = 58.5 (Na-ZSM-5) to 54.0 ppm (Cs,Na-X) 13 C: δ13 = 102 to 112 ppm C for nitromethane at mixed magnesium– aluminum oxides

[109]

1 H:

[109]

[110]

[111]

[111]

[112–116]

[117]

as probe molecules. Depending on the type of the univalent cation in zeolite Y, they observed 1 H MAS NMR signals at resonance positions of 7.55 (Li-Y) to 8.23 ppm (Cs,Na-Y) for adsorption of trichloromethane and of 6.62 (Li-Y) to 7.6 ppm (Cs,Na-Y) for adsorption of trifluoromethane. S´anchez-S´anchez et al. [110, 111] utilized pyrrole and chloroform as NMR probes for basic zeolites. Figure 11 shows 1 H MAS NMR spectra of pyrrole adsorbed on various alkali metal-exchanged zeolites X and Y [110]. The hydrogen atoms at the rings of the pyrrole molecules are not influenced by the different zeolites and cause the two signals at 6–7 ppm. The 1 H NMR shift of the hydrogen atoms at the ring nitrogens, however, covers a range between 8.4 (Li-Y) and 11.5 ppm (K-X) and indicates the different base strengths of zeolite oxygen atoms contributing to the hydrogen bondings with the pyrrole molecules. In these experiments, a stronger low-field shift corresponds to a higher base strength. An important advantage of pyrrole as probe molecule for the strength of base sites is the remarkable sensitivity (large shift range) and the good resolution of the 1 H MAS NMR spectra. References see page 1176

1176

3.2 Chemical Properties

Increasing base strength

11.5

60

H-ZSM-5

K-X

d13C / ppm

59

Na-X

Cs-Y

Na-ZSM-5 Na-Mor

58

Mg,Na-Y Ca,Na-Y H-Y

57 Na-Y 56 Na-X

55

Li-X

Cs-X

54

K-Y 2.2

2.4

2.6

2.8

3.0

3.2

Sm Na-Y

Fig. 12 Dependence of the 13 C MAS NMR shift of methoxy groups bound on framework oxygens in bridging positions on the mean Sanderson electronegativity Sm of the zeolite framework [116].

8.4 Li-Y 14

12

10

8

6

4

2

0

d1H / ppm 1H

MAS NMR spectra of pyrrole adsorbed on various alkali metal-exchanged zeolites X and Y [110].

Fig. 11

Independently of the spectroscopic method applied, important disadvantages of the above-mentioned probe molecules are that most of them are not totally unreactive in the presence of strong base sites and that they form different adsorption structures, complicating the evaluation of the spectra. An interesting approach, which is free from this problem, is the use of methoxy groups directly formed at the basic framework oxygen atoms by a conversion of methyl iodide or methanol as NMR spectroscopic probes. Applying 13 C MAS NMR spectroscopy, Bosacek et al. [113, 116] found a correlation between the isotropic chemical shift of surface methoxy groups bound to zeolite oxygens in bridging positions and the mean electronegativity S m of the zeolite framework (Fig. 12). According to this correlation, a low 13 C NMR shift of methoxy groups corresponds to a high base strength of the framework oxygen atoms. Methoxy groups bound at framework oxygens of alkali metal-exchanged zeolites Y and X cover a range of 13 C NMR shifts from 54.0 to 56.5 ppm. For zeolites X and Y impregnated with basic guest compounds, such as alkali metal hydroxides, additional 13 C MAS NMR signals of methoxy groups occur at 50 ppm. A stepwise increase in the methyl iodide adsorption on zeolites X and Y impregnated with alkali metal hydroxides leads to the preferential formation of methoxy groups at

50 ppm. Therefore, this signal indicates the presence of strongly basic compounds. For zeolites X and Y impregnated with alkali metal hydroxides, two high-field signals of methoxy groups bound at strongly basic guest compounds were found at ca. 50 and 52 ppm [115]. According to Krawietz et al. [118], guest compounds formed by impregnating a support with cesium hydroxide or acetate are a mixture of cesium oxide (Cs2 O), peroxide (Cs2 O2 ) and superoxide (CsO2 ). Therefore, the 13 C NMR signals observed at ca. 50 and 52 ppm could be an indication of the presence of different basic guest compounds on impregnated zeolites X and Y [115]. References 1. (a) H. Pfeifer, D. Freude, M. Hunger, Zeolites 1985, 5, 274; (b) H. Pfeifer, Colloids Surf. 1989, 36, 169; (c) H. Pfeifer, in Handbook of Heterogeneous Catalysis, 1st Ed., G. Ertl, H. Knoezinger, J. Weitkamp (Eds.), Vol. 2, Wiley-VCH, Weinheim, 1997, p. 732. 2. G. Engelhardt, D. Michel, High-Resolution Solid-State NMR of Silicates and Zeolites, Wiley, New York, 1987, 485 pp. 3. J. F. Haw, J. B. Nicholas, T. Xu, L. W. Beck, D. B. Ferguson, Acc. Chem. Res. 1996, 29, 259. 4. M. Hunger, Solid State Nucl. Magn. Reson. 1996, 6, 1. 5. M. Hunger, Catal. Rev. Sci. Eng. 1997, 39, 345. 6. M. Hunger, J. Weitkamp, in In-Situ Spectroscopy of Catalysts, B. M. Weckhuysen (Ed.), American Scientific Publishers, Stevenson Ranch, 2004, p. 177. 7. M. Hunger, M. W. Anderson, A. Ojo, H. Pfeifer, Microporous Mater. 1993, 1, 17. 8. M. Kalwei, H. Koller, Solid State Nucl. Magn. Reson. 2002, 21, 145. 9. S. Ganapathy, R. Kumar, L. Delevoye, J.-P. Amoureux, Chem. Commun. 2003, 2076.

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3.2.5

Active Phase–Support Interactions Agnieszka M. Ruppert and Bert M. Weckhuysen∗ 3.2.5.1

Metal–Support Interactions

Supported metal catalysts are comprised of small metal nanoparticles dispersed on the surface of a support material. Previously, the role of the support oxide was commonly thought to be limited to provide large metal surface areas. However, it is now generally accepted that the activity and selectivity of these catalysts are dependent on the type and composition of the support used [1–3] or even on thermal pretreatment conditions of catalysts consisting of the same metal and support [4]. In other words, the role of the support goes well beyond strictly physical phenomena such as increasing the surface area, and the specific interaction of the support with the active phase often controls the course of a catalytic reaction. Understanding the effect of the support oxide on the properties of metal nanoparticles is a challenging subject because it opens a way towards modeling and tuning of the catalytic properties by a deliberate choice of the support oxide [5–8]. As a consequence, gaining an insight into the interfacial structure between metal nanoparticles and their support, particularly under relevant reaction conditions, has been an important topic in the field of heterogeneous catalysis for decades [9, 10]. Therefore, ∗

Corresponding author.

3.2.5 Active Phase–Support Interactions

one would like to gather information on the nature of the chemical bonding at the surface, the direction of the ‘‘electron flow’’ at the metal–support interface, how the support influences the sorption and catalytic properties of the metal nanoparticles and how the support affects the bonding within adsorbed molecules. Due to the progress in experimental techniques, in combination with surface science studies, gas-phase cluster studies and theoretical calculations, detailed insight into these fundamental questions has been obtained and we refer the interested reader to several excellent review papers on this topic [11–18]. However, a unifying picture is still lacking and one of the major objectives for future research clearly is to move from qualitative ideas on the effect of the support oxide on adsorption and catalysis for supported metal nanoparticles to a more quantitative understanding of these metal–support effects. Only then we will be able to design supported metal nanoparticles towards a required conversion and selectivity for a specific catalytic reaction. The first step in the design of such ‘‘fine-tuned’’ supported metal catalysts requires a fundamental understanding of the chemistry of catalyst preparation. Indeed, the preparation method will ultimately dictate the surface composition, size and morphology of the supported metal nanoparticles and as a consequence determine the metal–support interface. Catalyst preparation procedures aim to make and stabilize small and uniform metal nanoparticles at the surface of a support material. The obvious reason is that one likes to have a catalytically active phase exposing as large fraction of active sites as possible to the reactant molecules present in the gas or liquid phase surrounding the catalyst material. This implies that decreasing the metal nanoparticle size results in an increased active phase surface area, leading further to an increased number of active sites. Interestingly, when the particle size is smaller than a few nanometers, quantum size effects will occur and the electronic structure changes from one comprised of valence and conduction bands to one made up of individual molecular states, thus leading to a metal–insulator transition. In a typical preparation procedure, a catalytically active precursor complex is first dissolved in a solution and then contacted with the catalyst support [19]. After impregnation, the catalyst is dried and subsequently activated, which is usually a reduction with either H2 or CO. In the case of metal precursor solutions with a suspended oxide support, the oxide is typically covered with terminal hydroxyl groups of varying basicity and acidity. The choice of the catalyst precursor complex relates to whether basic hydroxyl groups or acidic groups cover the surface of the support oxide, i.e. basic hydroxyl groups dominate alumina supports, whereas silica supports contain only weak acid hydroxyl groups

1179

(see Chapter 2.4.2). For example, to prepare a highly dispersed Pt on alumina catalyst, a negatively charged Pt complex is typically used, such as PtCl2− 4 . This complex can be ion exchanged onto the basic hydroxyl groups of the alumina support. In contrast, in preparing a Pt/SiO2 catalyst, Pt(NH3 )2+ 4 is typically used to carry out the ionexchange reaction with protons of the hydroxyl groups. In subsequent catalyst drying and activation steps, a complex set of surface reactions will take place, which depend on the type of support and also the metal loading. For easily reducible metals and a fairly strong metal precursor–support interaction, small metal nanoparticles are formed at the surface of the support (e.g. Pt/Al2 O3 ). In contrast, on silica often large and weakly interacting metal nanoparticles are formed. Furthermore, when non-easily reducible metals, e.g. Co2+ , are used they may react with silica or alumina and make a Co aluminate or Co silicate layer. In the case of titania supports, it is even possible that TiOx particles partially cover the reduced metal particle and as a consequence influence the catalytic performance, and also block active sites. This led Bond to propose a classification scheme in which the metal–support interactions (MSI) are divided into strong, medium and weak [20]. Strong metal–support interactions (SMSI) have been ascribed to metals supported on reducible oxides, such as TiO2 , whereas weak metal–support interactions (WMSI) are associated with noble metals supported on non-reducible oxides, such as SiO2 and Al2 O3 . The example for which Co aluminate or Co silicate are formed can be considered as an example of medium metal–support interaction (MMSI). Other examples of MMSI are those in which mostly zeolites are involved as supports. MSI effects are also a function of the type of metal, in the case of supported Ag catalysts the effect decreases in the order SiO2 > Al2 O3 , whereas for supported Pt catalysts the sequence is the reverse. The behavior of high-temperature-reduced noble metal catalysts supported on reducible oxides has been examined in a broad range of carbonyl group hydrogenations, including reductions of CO [21, 22], acetone [23], acetophenone [24], crotonaldehyde [25, 26], cinnamaldehyde [27], benzaldehyde [28], phenylacetaldehyde [29], citral [30] and acetic acid [31]. In all cases the turnover frequency (TOF) of the C=O bond hydrogenation was much higher for the high-temperature-reduced Pt/TiO2 catalysts (being in the SMSI state) than for the same catalysts reduced at low temperature or catalysts supported on non-reducible oxides, such as SiO2 and Al2 O3 . The above examples indicate that the support oxide can actively ‘‘shape’’ the catalytic properties of the active metal phase. However, in some cases the support oxide may even show some catalytic reactivity itself. Whereas in the case of References see page 1187

1180

3.2 Chemical Properties

an alumina support the basic hydroxyl groups dominate its surface, the few highly acidic hydroxyl groups also present may dramatically affect a catalytic reaction. An example is the selective oxidation of ethene catalyzed by silver supported on alumina. Ethene oxide, which is produced by the catalytic reaction of oxygen and ethene over Ag, can be isomerized to acetaldehyde via the acidic protons of the alumina. The acetaldehyde can then be rapidly oxidized over Ag to CO2 and H2 O. This combustion reaction is an example of bifunctional catalysis and is of course unwanted. Another interesting phenomenon to be discussed is the different behavior of Pt nanoparticles supported on zeolite Na-Y and Cs-Y, as shown in Fig. 1. The zeolite Y pore structure can be seen as a physical boundary, which limits the growth of the metal nanoparticles to its supercage size (1.3 nm). As a consequence, it was found with H2 chemisorption, HRTEM and EXAFS that welldefined Pt nanoparticles of about 1 nm were formed in zeolite Y when a mild calcination treatment was used in which the temperature was increased with a heating rate as low as 0.2 ◦ C min−1 , prior to reduction in pure H2 [32]. This is illustrated for Pt/Na-Y in Fig. 1a and no

large clusters with a diameter higher than 1 nm could be observed in the related HRTEM picture. However, in the case of the Pt/Cs-Y sample, a significant portion of the Pt nanoparticles are in the size range 1–5 nm. This is evidenced by the larger Pt−Pt coordination number from EXAFS experiments, and also from the HRTEM and 3D-TEM pictures, shown in Fig. 1b and c. Figure 1c shows that larger Pt nanoparticles are not on top of the zeolite Y crystals, but most probably grown in the Pt/Cs-Y sample exceeding the supercage size. This led to the creation of mesopores. Evidence for similar effects has been reported for CdS and ZnS clusters in zeolite Na-X, Ir clusters in zeolite Na-X and Pt clusters in ZSM-5 zeolites [33, 34]. The origin of these effects remains unclear, however. The role of the support oxide, in this particular case a zeolite support, on these supported Pt nanoparticles is therefore not limited to providing large Pt surface areas. The two different Pt-Y samples possess not only different catalytic conversion in e.g. CO oxidation because they have different Pt dispersions, but also because the support oxide influences the surface chemistry of the Pt nanoparticles; and as a consequence the catalyst performances, as will be discussed below.

20 nm

20 nm

(a)

(b)

(c)

Fig. 1

(a) High-resolution TEM image of Pt/Na-Y; (b) high-resolution TEM image of Pt/Cs-Y; (c) three-dimensional TEM image of Pt/Cs-Y.

3.2.5 Active Phase–Support Interactions

To the best of our knowledge, the first mention of the enhancement of the catalytic activity of supported metal nanoparticles goes back to the 1960s in a study on the hydrogenolysis of neopentane in which Pt nanoparticles were supported on zeolite Y [35]. Since then, hydrogenolysis has been considered as a sensitive probe reaction to reveal the effect of the support on the catalytic properties of supported metal nanoparticles. More specifically, hydrogenolysis appears to be an ‘‘electronic-sensitive’’ reaction and the TOF may vary by several orders of magnitude when going from one support oxide to another. For example, Homeyer et al. [36] showed that the TOF of Pd/H-Y in neopentane hydrogenolysis is about two orders of magnitude higher than that of Pd/SiO2 and much higher than that of non-acidic Pd/Na-Y zeolites. Moreover, in addition to the enhanced hydrogenolysis activity, a lower activation energy for neopentane hydrogenolysis was found for Rh/H-Y zeolites containing ‘‘electrondeficient’’ Rh particles, unlike the behavior of ‘‘neutral’’ Rh particles in Rh/SiO2 catalysis [37]. Careful analysis of the experimental data made it possible to exclude other possible explanations for these enhancements of reaction rate. Furthermore, it is now generally accepted that similar metal–support effects can be observed for many other catalytic reactions and related phenomena, such as hydrogenation, oxidation, aromatization and sulfur and coke resistance, and for a large variety of metals, including Pt, Pd, Rh, Ru, Au, Ir, Ni, Cu and Co. Table 1 discriminates between several reactions, which lead to an enhancement or a suppression of the reaction rate when using either ‘‘electron-rich’’ or ‘‘electron-poor’’ supported metal nanoparticles. Another interesting example includes the change in selectivity in the catalytic hydrogenation of 1,3-butadiene towards the 1,4-addition reaction over ‘‘electron-deficient’’ Pt nanoparticles [38, 39]. As a result of this observation, the formation of 2-butene has even been considered as a sensitive probe for monitoring the presence of Ptδ+ species. The changing of the electronic properties of the active metal phase by a support oxide was also described by Touroude and co-workers, who examined Pt/ZnO catalysts with the support partially reduced to metallic Zn, Tab. 1

1181

for the selective hydrogenation of crotonaldehyde to crotyl alcohol [40, 41]. Thanks to electronic interactions with Zn, the C=C adsorption on Pt was disfavored because of the loss of 5d electrons, while the C=O bond could be selectively adsorbed on Ptδ− −Znδ+ entities, leading to a higher selectivity towards unsaturated alcohol [42]. Additionally, the presence of Cl− (from the Pt precursor) further polarized those sites, leading to increased Lewis acidity and even more selective adsorption of the carbonyl oxygen. It is beyond the scope of this chapter to be exhaustive on all catalytic reactions influenced by the support material and we refer the reader to many interesting books describing the different effects a support may induce on the catalytic performances of supported metal nanoparticles [2, 6, 7]. Metal–support effects can be revealed by (1) changes due to metal nanoparticle charging (generally referred to as electronic effects), (2) effects related to variations in metal nanoparticle shape and crystallographic structure (generally referred to as geometric effects) and (3) the appearance of specific active sites at the metal nanoparticle–support boundary. Electronic effects are best defined as phenomena arising from electronic perturbations at the interface between metal nanoparticles and the support and include an ‘‘electron charge transfer’’ between metal and support and a changing of electric field at the metal–support interface. In contrast, geometric effects imply that the exposed surface of metal nanoparticles and the related active species are a function of the support type and reaction environment. More specifically, changes in the shape or size of the metal nanoparticle will control the relative ratios of edge, corner and terrace sites and therefore have an influence on both the catalytic activity and selectivity. Further, geometric effects are related to the decoration or encapsulation of the active metal phase and the sintering of the metal particles on the support. In addition, the support may induce strain on the metal–metal bonding at the surface, giving rise to a different type of supported metal particle. However, geometric and electronic effects mostly cannot be treated as completely independent phenomena. For instance, increasing the size of metal nanoparticles References see page 1187

Experimental results relating to metal–support effects and their relation to the presence of electron-rich and electron-poor metal

clusters Reactions or related phenomena

‘‘Electron-rich’’ metal nanoparticle

‘‘Electron-poor’’ metal nanoparticle

Alkane hydrogenolysis Aromatics hydrogenation Sulfur resistance Coke resistance Hydrogenation of α,β-unsaturated aldehydes

Decrease in activity Decrease in activity Decrease Increase Increase in C=O hydrogenation selectivity

Increase in activity Increase in activity Increase Decrease Decrease in C=O hydrogenation selectivity

1182

3.2 Chemical Properties

results in an increase in the electron bandwidth, but the nature of the exposed planes and the topology of the surface sites also change. Hence the ambiguity in discriminating electronic and geometric factors is well known to the catalysis community. Furthermore, it is far from trivial to provide unambiguous experimental evidence for the presence of ‘‘interfacial’’ active sites at the metal–support boundary, although an increase in the CO hydrogenation reaction rates has been associated with the formation of interfacial active sites when making reduced metal nanoparticles on a support oxide. The aforementioned hydrogenation of α, β-unsaturated aldehydes is an example of a catalytic reaction where all those effects meet. The now well-recognized SMSI effect, responsible for increased unsaturated alcohol selectivity in noble metal catalysts supported on reducible oxides, has been attributed to both the formation of metal–support interfacial active sites and the change of metal electronic properties by cations of the partially reduced support [4, 43]. Englisch et al. [44] describe two complementary effects – of type (2) and (3) in the above classification, respectively – contributing to the favoring of C=O hydrogenation. Over Pt/SiO2 (non-reducible support), the allyl alcohol selectivity increased with Pt particle size. For large metal particles, the high fraction of Pt(111) surfaces was concluded to favor the adsorption of crotonaldehyde via the carbonyl bond, which has also been theoretically calculated by Delbecq and Sautet [45]. Over Pt/TiO2 (reducible support), the selectivity depended also on the degree of TiOx decoration of the metal surface, controlled by the catalyst pretreatment temperature. Larger Pt particles also accounted for increased selectivity. However, smaller particles facilitated TiOx migration, leading to the formation of interfacial active sites. Bowker et al. described a mathematical model of a reaction in which the rate is proportional to the length of the metal–support phase boundary, which contains the active sites [46]. According to this model, the rate increases with the metal loading up to some maximum value, after which it begins to drop (when metal particles start to overlap). At high loadings, the rate is reduced to zero due to complete coverage of the support with the active phase. This model worked very well in practice with the photocatalytic degradation of methanol over a Pd/TiO2 catalyst. The same group suggested in another study that the SMSI effect in a Pd/TiO2 catalyst can involve the formation of PdTi and PdTi2 alloys with distinctive atomic structures [47]. Table 2 summarizes some spectroscopic and microscopic techniques that have been shown to provide detailed information on geometric, electronic and/or interfacial active site properties of supported metal nanoparticles [48]. The potential and limitations of these techniques are also included. In the following, we will

A selection of spectroscopic and microscopic techniques for studying supported metal nanoparticles, including their potential and limitations

Tab. 2

Technique

EXAFS/XANES IR EPR XPS Electron microscopy

Information obtained Particle size

Particle shape

Electronic properties

In situ capabilities

+ − + ± +

± ± − ± +

+ + + + −

+++ +++ ++ + +

discuss their application in the field of supported metal catalysis in more detail. Infrared (IR) spectroscopy in combination with, e.g., carbon monoxide as probe molecule remains one of the most informative techniques for characterizing supported metal catalysts [49]. The spectra of adsorbed CO are known to be the result of the interplay of the interaction between metal d-orbitals and σ -bonding and π-antibonding orbitals of adsorbed CO. This may result in different types of bonding to the metal surface. In a detailed study, Visser and co-workers [50, 51] analyzed well-defined Pt nanoparticles with an average diameter of 1 nm supported on a series of zeolite Y samples containing different monovalent (H+ , Na+ , K+ , Rb+ and Cs+ ) and divalent (Mg2+ , Ca2+ , Sr2+ and Ba2+ ) cations. These different cations introduced into the zeolite structure resulted in an altered ionicity of the catalyst material due to a change in the negative charge density of the zeolite oxygen atoms. Time-resolved IR spectroscopy measurements allowed the study of the temperatureprogrammed desorption of C≡O from supported Pt nanoparticles to monitor these changes by probing the local environment of the adsorbed C≡O. An example of such set of data is shown in Fig. 2 for zeolite Pt/H-Y. It was found that the red shift from the linear Pt-coordinated C≡O vibration compared with that of gas-phase C≡O increases with increasing cation radius-to-charge ratio. In addition, a systematic shift from linear- to bridgebonded C≡O was observed for decreasing Lewis acidity of the introduced cation as expressed by the Kamlet–Taft parameter α [52], as illustrated in Fig. 3. A decreasing α value results in an increasing electron richness on the zeolite framework oxygen atoms. This observation was confirmed with X-ray absorption spectroscopy, and the intensity of the experimental Pt atomic XAFS intensity correlates with the Lewis acidity of the cation introduced in the zeolite matrix. Furthermore, it was found on the basis of a semi-quantification of the obtained IR spectra

3.2.5 Active Phase–Support Interactions

1183

Absorbance/ a.u.

0.14

0.07

160 120 in m e/ m i -t

80 40

0.00 2300

2100

1900

D

TP

1700

Wavenumber/cm−1 1.0 nm

Time-resolved IR spectra obtained during temperature-programmed CO desorption from a Pt/H-Y zeolite sample, including a schematic representation of a Pt cluster encaged in the supercage of zeolite Y.

Fig. 2

IR L:B intensity ratio

8

6

4

2 0

1

2

3

4

5

Lewis acidity of cation (a) Relationship between IR linear-to-bridged CO species intensity ratios and Lewis acid properties of the cations introduced in zeolite Y as expressed by the Kamlet–Taft parameter α.

Fig. 3

that the C≡O coverage increases with increasing electron density on the Pt nanoparticles. Similarly, Ji et al. found that hydrogen chemisorption on zeolite Y-supported Pt nanoparticles was strongly influenced by the support ionicity and decreased with decreasing electron richness of the support framework oxygen atoms [53]. These results also indicate that one should be cautious in interpreting H2 chemisorption results in terms of only Pt dispersion, since it is applicable only when comparing Pt particles dispersed on the same support oxide. Indeed, hydrogen chemisorption experiments revealed that the determined H/Pt values have no relationship with the Pt dispersion calculated from independent high-resolution transmission

electron microscopy (HRTEM) results and the hydrogen chemisorption capacity of Pt is directly influenced by the ionicity of the support oxygen. Therefore, serious deviations in the real Pt dispersions occur if Pt is supported on oxide materials with different ionicities. In a more systematic study, van der Eerden et al. [54] studied 14 supported 1 wt.% Pt catalysts with different chemical compositions and support pore structures (microporous, mesoporous and macroporous supports) and compared the linear-to-bridged (L : B) Pt-coordinated C≡O as obtained with IR spectroscopy with the atomic XAFS intensities measured on the same set of samples after adsorption of C≡O at room temperature in the same spectroscopic in situ cell. The results are shown in Fig. 4. The IR L : B ratios span a range of more than 12 units, whereas the atomic XAFS intensities change between 0.8 and 2.6 × 10−2 A˚ −2 . It is clear that the atomic XAFS peak decreases with decreasing IR L : B ratio and both the atomic XAFS peak intensity and the IR L : B ratio can be considered as complementary measures for the electron charge of the Pt nanoparticles. In other words, the atomic XAFS peak intensity and the IR L : B ratio decrease with increasing ionization potential of the Pt atoms in the supported nanoclusters. The advantage of the atomic XAFS method, however, is that no probe molecule is needed and could be used in the future to probe the electronic properties of supported metal nanoparticles under reaction conditions in real time, delivering mechanistic insight into the working catalyst. References see page 1187

Atomic XAFS intensity (× 10−3)

1184

3.2 Chemical Properties

20

10

0 2

4

6

8

10

12

14

IR L:B intensity ratio Atomic XAFS peak intensity of supported Pt nanoparticles as a function of the corresponding IR linear-to-bridged CO species intensity ratio.

Fig. 4

One of the simplifying assumptions in the description of catalytic reactions has often been to consider the catalyst surface as static, thus providing a fixed ensemble of active sites exposed at a surface of a catalytic solid. However, there is no guarantee that the catalyst nanostructures observed remain stable during catalytic action. Indeed, in situ characterization studies have shown that supported metal nanoparticles may undergo structural transformations in response to changes in the reaction conditions and that the transformations can have a significant impact on the catalytic performances [55, 56]. A powerful technique, which can provide such detailed information, is in situ HRTEM [57]. In a seminal paper, Hansen et al. [58] have shown that in situ HRTEM may provide detailed insight into the dynamics of supported metal nanoparticles when exposed to various gas environments. The system under investigation was a Cu/ZnO-based catalyst used for the industrial production of methanol. It was found that when the catalyst was exposed to more oxidizing conditions by adding water vapor to hydrogen gas, the Cu nanoparticles obtained a more spherical shape, i.e. the Cu nanoparticles are terminated by a higher fraction of (110) and (100) facets, relative to the more close-packed (111) facets, than in pure hydrogen. Hence the more open (110) and (100) facets turned to stabilize relative to the (111) facets. At the interface, however, the contact area between Cu and ZnO does not change significantly, pointing to the fact that water adsorption on the different exposed Cu facets is the main driving force for the gas-induced surface reconstructions and the resulting reshaping of the supported Cu nanoparticles. On the other hand, addition of a more reducing gas, such as carbon monoxide, to the hydrogen gas resulted in more marked changes. It was found that the Cu nanoparticles transform into disc-like structures caused by an increased wetting of the ZnO support. On average, the interfacial area increased by

about 50%, indicating a large decrease in the interfacial energy. Previously, it was suggested that Cu mainly affects the ZnO surface energy, whereas the Cu surface energies are not perturbed [59–61]. This is probably a good approximation because the coverage of hydrogen and carbon monoxide on the Cu surfaces is fairly low under the conditions used in the study. It was proposed that a change in the oxidation potential of the gas phase may change the oxygen content in the ZnO surface and therefore the interface energy. The intimate structure of the metal–support interface determines many of the qualitative phenomena discussed so far. However, the complexity of the metal–support interface makes it difficult to obtain structural and electronic information and as a consequence experimental data on the metal–support interface are almost lacking. Important information, has however been obtained with EXAFS, as illustrated in Fig. 5. It was found that for supported Pt catalysts the metal–support interface is characterized by two metal–support oxygen contributions, one at 2.18 A˚ and another in the range 2.7–2.9 A˚ [62–66]. Desorption of hydrogen at elevated temperature results in a shortening of the Pt−O distance to 2.05 A˚ and some sintering of the metal particles occurred. Based on EXAFS, H2 TPD and FTIR data, it was proposed that the long Pt−O distance originates from neutral hydrogen atoms located at the interface between the metal nanoparticles and the support oxide. Thus, after reduction three-dimensional Pt particles are formed, which are separated from the support surface by a layer of hydrogen atoms. Treatment at higher temperature leads to desorption of hydrogen, bringing the metal atoms in direct contact with the oxygen atoms of the support. Removal of hydrogen increases the strength of the metal–support interaction and the shape of the Pt nanoparticles transforms from hemispherical into ‘‘raftlike’’ (Fig. 5). Moreover, better uniformity of the Pt–O distances after high-temperature reduction suggested epitaxial growth between platinum particles and the support oxide. Similarly, EXAFS allowed the interaction of a metal cluster with support lattice cations to be elucidated, an example being Ir clusters supported on MgO [67]. It was found that Ir exists as a mixture of three-dimensional and raft-like Ir4 and Ir6 clusters. The Ir−MgO interface is characterized by a single Ir−Mg bond with a length ˚ It was of 1.6 A˚ and four-fold Ir−O coordination at 2.6 A. suggested that the Ir atom is located on top of the Mg in the MgO (100) crystal face. The long Ir−O distance was attributed to the presence of hydrogen at the interface. Another interesting spectroscopic technique, which can provide insight into the size, structure, charge and local environment of small metal nanoparticles

3.2.5 Active Phase–Support Interactions

∼2.7 Å

∼2.2 Å

Pt atom (a) Fig. 5

1185

H atom

O atom

Me atom, (Me-metal of the support oxide) (b)

Schematic representation of the metal–support interfacial structure after reduction with hydrogen at (a) 300 and (b) 450 ◦ C.

is electron paramagnetic resonance (EPR). EPR is a non-invasive technique that yields information on supported paramagnetic metal clusters. Roduner’s group has studied with continuous-wave and pulsed EPR methods platinum clusters immobilized on L- and faujasite-type zeolites [68, 69]. It was found that an icosahedral magic cluster of 12 equivalent platinum nuclei and a 13th atom in the center that is invisible in EPR could be immobilized in the cages/channels of the zeolites. Following reduction with hydrogen, each of the 12 surface platinum atoms is bound to one hydrogen atom, giving rise to a cluster of the form Pt13 H12 , probably with a small positive charge. Small clusters are normally assumed to occur in various structures and with a broad distribution regarding the number of atoms. Surprisingly, these EPR studies indicate that only a single well-defined EPR-active cluster was observed apart from a minor unidentified second species. These two species account for ca. 10% of the total Pt content of the sample and therefore no information could be obtained for the remaining 90% of the sample, which may be diamagnetic or paramagnetic but EPR silent. There are strong indications that a reservoir of EPR-silent but structurally similar clusters exists which can be partially converted to EPR-visible species by H−D exchange or by gas adsorption. Interestingly, clusters with 13 ± 5 Pt atoms were also found in Pt/Na-Y zeolites, as established with EXAFS [70]. In the case of the combustion of toluene over supported Pd catalysts, it has been shown that the reactivity of Pd is related to the acid–base properties of the support, which influence the surface oxidation state of the supported Pd nanoparticles [71, 72]. It was found that Pd loaded on MgO and WO3 , which have strong basic and acidic character, respectively, were relatively inactive to the reaction, whereas metal oxides with weak acid-based properties, such as Al2 O3 , SiO2 and Nb2 O5 , were much more active. In order to elucidate the effect of the support oxide on the surface oxidation state of Pd, XPS can be used.

It was found that the Pd 3d5/2 XPS peak position shifted towards higher binding energy with increasing acidity of the support. These facts mean that the surface of Pd is easily oxidized when Pd is supported on acidic supports, whereas Pd supported on a basic support became more difficult to be oxidized. The observation can be explained through the electronic character of the support, which is related to the acid–base properties. Acidic supports with electrophilic character result in the electron-deficient state of Pd, hence the Pd surface is easily oxidized to generate the surface PdO. On the other hand, a basic support makes the Pd particle more electron rich, hence the Pd surface becomes more difficult to oxidize. However, XPS may not always provide clear evidence for a change in oxidation state of the surface metals, an illustrative example being the support effects reported for the hydrogenation of cinnamaldehyde over carbon nanofibersupported Pt catalysts. Toebes and co-workers [73–75] studied the influence of the support surface composition on the performance of the catalyst by changing the number of oxygen-containing surface groups on the carbon nanofiber support. In this manner the overall catalytic activity could be increased by a factor of 25. The differences in intrinsic activities are even much larger since in the catalyst treated at higher temperatures internal diffusion limitation was encountered. The enhanced activity is mainly caused by a strong increase in the hydrogenation rate of the C=C bond, whereas only a slight increase in the C=O bond hydrogenation is observed. XPS and H2 chemisorption experiments provided no clear evidence for a change in the electronic structure of the Pt nanoparticles induced by the oxygen-containing group present on/in the surface of the carbon nanofibers, although a linear decrease in the hydrogenation activity with an increase in the number of acidic groups on the carbon nanofiber support was found. In other words, a modification of the electronic properties of the metal is not the prime cause of the large References see page 1187

1186

3.2 Chemical Properties

HOOC O

COOH O

O

N2, ∆T O

O

PtCNF Polar CNF surface low activity, weak adsorption of reactant on support

O

PtCNF973 Non-polar CNF surface high activity, optimal support-assisted adsorption of reactant

Schematic representation of cinnamaldehyde adsorption explaining the enhanced activity for the carbon nanofiber-supported platinum catalyst after the removal of the majority of the oxygen-containing groups, by support-assisted catalysis.

Fig. 6

changes in the catalytic behavior of the catalysts. Alternatively, it was proposed that hydrogenation is assisted by adsorption of cinnamaldehyde on the carbon support after removal of the oxygen-containing surface groups. The suggested mechanism is illustrated in Fig. 6. It is assumed that on a CNF- or carbon nanofiber-supported Pt catalyst with a large amount of oxygen-containing surface groups adsorption and dissociation of hydrogen and adsorption of the organic reactants all have to take place on the Pt particles. In other words, adsorption of cinnamaldehyde on the polar surface is weak. On the Pt particles hydrogenation sites are present and apparently the hydrogenation of cinnamaldehyde on Pt is rather difficult. On the other hand, when hardly any oxygen groups are present, reactant molecules adsorb with the benzene ring on the non-polar surface of the carbon nanofibers. It was speculated that on the carbon surface other sites are present that can very efficiently hydrogenate C=C bonds. Finally, we should also discuss the formation of specific active sites at the metal–support interface since there is a growing body of experimental data available which provide evidence for the existence of such active sites at the metal nanoparticle periphery composed of the metal site and the site on the support surface. Figure 7 illustrates the adsorption of a molecule containing a polar group on such an ‘‘interfacial active site’’. Several groups have attributed the enhanced activity of supported metal catalysts, consisting of a noble metal phase supported on a reducible oxide, such as TiO2 and Nb2 O5 , in CO hydrogenation reactions, to the presence of those active sites. For example, Somorjai and co-workers [76–78] proposed that the effect of a transition metal oxide species, such as TiOx , deposited on Rh surface in hydrogenation of the C=O bond is associated with carbonyl bond activation through simultaneous adsorption of the carbon end of the C=O bond to the metal site and the oxygen end of the C=O bond to the Lewis acid site on the support oxide. This interesting concept provides an adequate explanation for the support effect of reducible metal oxides on the activity and selectivity of the metal phase in the hydrogenation of CO and hydroformylation of ethene.

C O

Schematic representation of the adsorption of a reactant molecule containing a polar group on an ‘‘interfacial active site’’ present in supported metal catalysts.

Fig. 7

Similarly, Vannice and co-workers [79–81] proposed a model to explain the extremely high activity of a Pt/TiO2 catalyst in the hydrogenation of acetophenone and the enhanced selectivity towards crotyl alcohol in crotonaldehyde. The model also implies the creation of special active sites at the metal–support interface, which can coordinate the oxygen end of the C=O bond and thereby specifically activate the bond. However, excessive decoration of Pt surface with TiOx particles may lead to the loss of catalytic activity. A last example to illustrate this interesting phenomenon can be found in the enhanced selectivity of Ru/ZrO2 towards cinnamyl alcohol in cinnamaldehyde hydrogenation, which was ascribed to the formation of Ru−Zrn+ sites at the periphery of the particles [82]. The presence of mixed Ru−Zrn+ sites appears to decrease the strength of the C=O bond, thus facilitating the hydrogenation reaction. This concept of ‘‘interfacial’’ active sites coupling a metal center and a Lewis acid–base function of adjacent centers stabilized by the support makes it possible to predict that these sites may also be active in the conversion of other molecules with polar functional groups, such as CN, CS and NH.

References

Acknowledgment

B.M.W. acknowledges financial support from NWO-CW VICI for carrying out part of the research work described in this chapter. The authors thank Dr. U. Ziese (UU) for measuring the 3D-TEM picture shown in Figure 1c. References 1. A. T. Bell, Science 2003, 299, 1688. 2. R. A. van Santen, M. Neurock, Molecular Heterogeneous Catalysis, a Conceptual and Computational Approach, Wiley-VCH, Weinheim, 2006, pp. 47, 57, 242, 248–249. 3. J. Hagen, Industrial Catalysis, a Practical Approach, WileyVCH, Weinheim, 1999. 4. S. J. Tauster, S. C. Fung, R. L. Garten, J. Am. Chem. Soc. 1978, 100, 170. 5. G. C. Bond, Metal-Catalyzed Reactions of Hydrocarbons, Springer Science and Business Media, New York, 2005. 6. J. A. Anderson, M. F. Garcia (Eds.), Supported Metals in Catalysis, Imperial College Press, London, 2005, pp. 60–73, 134, 145, 251–253, 334–337. 7. J. M. Thomas, W. J. Thomas, Principles and Practice of Heterogeneous Catalysis, VCH, Weinheim, 1997, p. 519. 8. H. F. Rase, Handbook of Commercial Catalysts, CRC Press, Boca Raton, FL, 2000. 9. S. J. Tauster, S. C. Fung, J. Catal. 1978, 55, 29. 10. G. L. Haller, Adv. Catal. 1989, 36, 173. 11. M. Che, C. O. Bennett, Adv. Catal. 1989, 36, 55. 12. D. R. Rainer, D. W. Goodman, J. Mol. Catal. A 1998, 131, 259. 13. Z. Xu, F. S. Xiao, S. K. Purnell, O. Alexeev, S. Kawi, S. E. Deutsch, B. C. Gates, Nature 1994, 372, 346. 14. D. W. Goodman, Chem. Rev. 1995, 95, 523. 15. W. M. H. Sachtler, Acc. Chem. Res. 1993, 26, 383. 16. W. M. H. Sachtler, Z. C. Zhang, Adv. Catal. 1993, 39, 129. 17. B. C. Gates, Chem. Rev. 1995, 95, 511. 18. A. Y. Stakheev, L. M. Kustov, Appl. Catal. A 1999, 188, 3. 19. G. Ertl, H. Kn¨ozinger, J. Weitkamp (Eds.), Preparation of Solid Catalysts, Wiley-VCH, Weinheim, 1999. 20. G. C. Bond, Metal–Support and Metal–Additive Effects in Catalysis, Elsevier, Amsterdam, 1982, p. 1. 21. M. A. Vannice, R. L. Garten, J. Catal. 1979, 56, 236. 22. M. A. Vannice, C. C. Twu, J. Catal. 1983, 82, 213. 23. B. Sen, M. A. Vannice, J. Catal. 1988, 113, 52. 24. S. D. Lin, D. K. Sanders, M. A. Vannice, Appl. Catal. A 1994, 113, 59. 25. C. G. Raab, J. A. Lercher, Catal. Lett. 1993, 18, 99. 26. P. Claus, S. Schimpf, R. Sch¨odel, P. Kraak, W. M¨orke, D. H¨onicke, Appl. Catal. A 1997, 165, 429. 27. A. B. de Silva, E. Jordano, M. J. Mendes, P. Fouilloux, Appl. Catal. A 1997, 148, 253. 28. M. A. Vannice, D. J. Poondi, J. Catal. 1997, 169, 166. 29. D. J. Poondi, M. A. Vannice, J. Catal. 1997, 124, 79. 30. U. K. Singh, M. A. Vannice, J. Mol. Catal. A 2000, 163, 233. 31. W. Rachmady, M. A. Vannice, J. Catal. 2000, 192, 322. 32. B. M. Weckhuysen, unpublished results. 33. G. Schulz-Ekloff, Stud. Surf. Sci. Catal. 1991, 69, 65. 34. M. Wark, J. N. I. Schul-Ekloff, A. Zukal, Stud. Surf. Sci. Catal. 1991, 69, 189. 35. R. A. Dalla Betta, M. Boudart, in Proceedings of the 5th International Congress on Catalysis, J. W. Hightower (Ed.), North-Holland, Amsterdam, 1973, p. 1329.

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3.2.5.2

Oxide–Support Interactions

Jaap A. Bergwerff and Bert M. Weckhuysen∗

Solid catalysts, in which a metal oxide phase acts as the active component and is supported on a support oxide, are applied in a broad range of industrially relevant reactions. Some important examples are listed in Table 1. The active phase is most often a transition metal oxide, such as Fe, V, Cr, Mo or W, because of their ability to adapt different oxidation states and coordination environments depending on the reaction conditions applied. Porous oxides, such as Al2 O3 , SiO2 , SiO2 /Al2 O3 and TiO2 , and different types of zeolites, such as ZSM-5 and mordenite, are commonly used as support oxides. In a large number of reactions in which metal oxide catalysts are used, the nature of the support that is applied has a pronounced effect on the observed activity and selectivity. Whereas the active sites of a metal catalyst are located on the surface of a metal nanoparticle, in a metal oxide catalyst they consist of metal ion complexes in close interaction with the support. Due to the intimate contact between support and active phase, a strong influence of the support on the nature of the active sites can be expected. However, as was already discussed for metal catalysts in Chapter 3.2.5.1, it is often difficult to establish the exact nature of this support effect. ∗

Corresponding author.

Characterization studies by Berlier et al. elegantly show how the support material influences the coordination and reactivity of supported metal ions to a great extent and how spectroscopy can be used to evaluate this effect [11]. In this study, the coordination of Fe2+ ions in different siliceous matrices was monitored using the detection of adsorbed NO by IR spectroscopy. Absorption by excitation of different vibrational modes of NO is observed at different frequencies when the number of adsorbed NO molecules per Fe2+ center changes. Hence the number of coordinatively unsaturated sites around the Fe2+ ions in the different samples can be determined. This can yield important information on the active sites in catalytically interesting materials such as Fe-ZSM-5, a promising material for the selective catalytic reduction of NOx [12] and selective oxidation reactions [13, 14]. Samples that were studied included FeO supported on amorphous silica and Fe-containing silicalite and ZSM-5. Surface Fe2+ centers of small FeO clusters deposited on amorphous silica can accommodate only one NO molecule, as they are hardly exposed. As a result, only the characteristic NO stretch vibration of [Fe2+ (NO)] complexes is observed in IR spectra recorded on this material after NO dosing. After heat treatment of the Fe/SiO2 sample at 1073 K, dispersion of Fe2+ ions takes place and isolated Fe2+ centers are generated that can form [Fe2+ (NO)2 ] complexes after exposure to NO. In the same way, the coordination of Fe in the crystalline structures of Fe-silicalite and Fe-ZSM-5 was examined. The Fe atoms that are present inside the zeolite framework in Fe-silicalite samples, treated at 773 K, are capable of accommodating up to three NO molecules. Heat treatment at 1073 K was found to lead to the migration of Fe species from framework positions and the formation of extra-framework FeO clusters. This is expressed in the IR spectra by an increase in the intensity of the nitrosyl band. In addition, the bands become much broader since the formation of a number of different [Fe2+ (NO)x ] complexes becomes possible in this amorphous material. Such an agglomeration is not observed after heat treatment of Fe-ZSM-5 at 1073 K. The presence of Al3+ in the framework of the Fe-ZSM-5 sample prevents the aggregation of extra-framework Fe2+ since individual ions, expelled from the framework by the heat treatment, are stabilized by the negative charge of the Al[OSi]− 4 entities [11]. The example described above clearly illustrates the flexibility of oxidic support materials to accommodate transition metal ions in different configurations. However, the difficulties that are connected with the description of supported metal oxide catalysts become obvious when a comparison is made with metal complexes applied in homogeneous catalysis. An example of such a catalytic

3.2.5 Active Phase–Support Interactions Examples of industrial applications of supported metal oxide catalysts

Tab. 1

Application

System

Reduction of NOx by NH3 Synthesis of phthalic anhydride Polymerization of alkenes Dehydrogenation of light alkanes Metathesis of alkenes Precursor for hydrotreating catalysts

Refs.

(W)V/TiO2 [1–3] [4–6] V/TiO2 [7] Cr/SiO2 Cr/Al2 O3 [8] [9] Re/Al2 O3 Mo/Al2 O3 , W/Al2 O3 [10]

R N R

O N

N Cr Cl

Cr

O

R SIO2

Cl Cl

(a)

(b)

(a) Molecular structure of a Cr–triazacyclohexane complex applied as a homogeneous ethene trimerization catalyst. (b) Schematic representation of the structure of the active site in a heterogeneous Phillips Cr/SiO2 ethene polymerization catalyst.

Fig. 1

system is presented in Fig. 1 a, in which a triazacyclohexane complex of chromium is shown, which is used for the trimerization of α-alkenes, such as ethene [15]. In these homogeneous complexes, the coordination environment of the metal ion is determined by the ligand that is used for complexation. The nature and number of the coordinating atoms can be controlled, in addition to the electronic properties and the steric environment of the metal ion. Characterization of these complexes is relatively straightforward as solutions can be obtained that solely contain the complex of interest. Furthermore, X-ray crystallographic data often provide a detailed picture of the active sites in these catalysts, although this information does not necessarily reflect the state of the catalyst material under reaction conditions. In supported metal oxide catalysts, the oxidic support can be regarded as the ligand, controlling the coordination and electronic properties of the metal ions and hence their reactivity. The exact nature of the supported metal complexes depends on the structure and surface composition of the support. To evaluate the exact nature of the metal oxide–support interaction, one needs to obtain an insight into the structure of the active species on a molecular scale. Hence, whereas the characterization of metal catalysts is often limited to textural analysis and the determination of the dispersion of the active phase, molecular spectroscopy is an essential supplementary tool for the characterization of metal oxide catalysts. However, characterization is more difficult,

1189

due to the heterogeneous nature of the support surface, resulting in the presence of a variety of metal sites on the support. A well-known example is the Phillips Cr/SiO2 catalyst used for ethene polymerization [7, 16], where only a small percentage of the Cr centers are actually active in polymerization. The lack of knowledge and the resulting challenges in the characterization of supported metal oxide catalysts are clearly illustrated when the schematic representation of the active site in a Phillips polymerization catalyst, as depicted in Fig. 1b, is compared with the well-defined structure of the homogenous Cr catalyst in Fig. 1a. Furthermore, due to the high mobility and reactivity of the surface metal complexes, their exact nature is very much a function of reaction temperature, pressure and atmosphere. It is therefore not surprising that the need to characterize catalysts under working conditions is most urgently felt in this field of catalysis and most in situ spectroscopic studies are carried out on supported metal oxide catalysts [17–21]. The information on metal oxide–support interactions that can be derived from such studies will therefore be the main topic of this chapter. However, to gain an insight into the molecular structure of the as-synthesized material, the preparation of supported metal oxide catalysts and the application of spectroscopy at this stage of a catalyst’s life will first be addressed. For the preparation of supported metal oxide catalysts, impregnation of the oxidic support material is often carried out with aqueous solutions of Cr6+ , Mo6+ , V5+ and W6+ salts. Transition metal oxo-anions are formed in these impregnation solutions, where monomeric 2− 3− 2− complexes (CrO2− 4 , MoO4 , VO4 , WO4 ) are present at high pH and low concentrations, whereas larger entities, (6−x)− such as Cr2 O2− and H2 W12 O6− 40 are 7 , Hx Mo7 O24 formed at low pH and high concentrations. The speciation of these complexes in solution can be calculated from their formation constants [22] and monitored using Raman and UV/visible spectroscopy [23–25]. Interaction with the support surface takes place via Coulomb interaction or a condensation reaction with surface hydroxyl groups in which adsorbed metal oxide groups are formed. Protonation of the support oxide surface will lead to a Coulomb interaction with the negatively charged transition metal oxo-anions and a higher affinity towards the formation of surface adsorbed species. Hence the use of acidic solutions generally leads to a stronger interaction between the metal precursor in solution and the support oxide and is therefore often preferred [26]. During calcination, the metal complexes are converted into a metal oxide phase covalently connected to the support oxide. As water is removed from the system the hydrated metal oxide complexes present in the References see page 1196

1190 O

3.2 Chemical Properties

O

O

M

M

O O

O M

M

MxOy

Support

Metal-oxide loading

Schematic representation of the metal oxide phase, present on the surface of a support metal oxide catalyst under dehydrated conditions support as a function of metal oxide loading.

Fig. 2

dried material are forced to react with surface hydroxyl groups. This can be observed using IR spectroscopy, in which the intensity of the absorption bands resulting from the stretch vibrations of support OH groups is found to decrease with increasing metal oxide loading [27]. The different MOx phases present on the surface of dehydrated supported metal oxide catalysts are represented schematically in Fig. 2. At low metal oxide loadings, dehydration will result in the formation of isolated metal oxide entities. When the metal oxide loading increases, these groups will link up and a twodimensional polymer is formed on the support oxide. Only when all OH groups have reacted and the surface is fully covered will three-dimensional bulk oxide clusters start to form, which can be observed with XRD or more sensitively with Raman spectroscopy. For SiO2 , the OH group density is typically low and the formation of bulk oxide is already observed at low loadings for SiO2 -supported systems. For SiO2 -supported catalysts, the preparation method is found to be the most critical factor. Therefore, a direct titration of surface hydroxyl groups by reactive transition metal complexes, such as metal (oxo)chlorides or alkoxides, is found to give the best dispersion for these catalyst systems [28]. When the supported metal oxide catalysts are exposed to the ambient atmosphere or even traces of water after calcination, rehydration will take place instantaneously and support–metal oxide bonds are broken again. A system will be formed with transition metal oxo complexes dissolved in a thin film of water on the support oxide surface. The nature of these metal complexes is similar to that of the anions found in solution, which allows for a reasonably straightforward identification, as ample reference spectra are available. Monomeric complexes are stable on supports with a high zero point charge and at low metal oxide loadings, whereas isopolyanions are found for high metal oxide loadings and on more acidic supports, i.e. those with a low zero point of charge [29]. Although there is general consensus on the surface processes sketched above, the exact molecular structure of the metal oxide complexes, present on the support surface after calcination, is still a matter of debate. Possible

molecular structures for metal oxide groups vary in the number of metal–oxygen–support bonds and the number of terminal M=O bonds per metal atom. Numerous workers have tried to answer these questions by applying a whole variety of spectroscopic techniques to study these systems under dehydrated conditions. In an extensive study, Raman spectroscopic measurements were carried out on a series of ZrO2 -supported metal oxide catalysts at different metal oxide loadings, before and after exchange of surface oxygen by 18 O [30]. The most intense band in the Raman spectra recorded on supported metal oxide catalysts under dehydrated conditions is ascribed to the stretch vibrations of terminal M=O bonds, as they can be expected to be accompanied by the largest change in polarizability. A second band in the Raman spectra of metal oxide catalysts is often observed at lower frequency and is reported to be the result of M−O−M or O−M−O bending vibrations in polymeric MOx species. 18 O isotope exchange leads to a ca. 40 cm−1 red shift of both Raman bands. The observation of only one M=18 O band shows that only mono-oxo metal complexes are present, even at higher metal oxide loadings, when polymeric metal oxide domains exist on the support oxide surface. A number of possible molecular structures for supported metal oxide groups on a ZrO2 surface are presented in Fig. 3, along with the position of the vibration bands observed by Raman spectroscopy. However, this study did not provide any insight into the exact way in which these metal oxide species are connected to the ZrO2 surface. To illustrate the complexity of the matter, the debate on the molecular structure of vanadium oxides in lowloaded vanadium oxide catalysts is summarized in this paragraph. In these systems, a band at ∼1030 cm−1 is observed in Raman spectra for different supports. It is traditionally ascribed to the V=O stretch vibration of an isolated VO4 unit symmetrically attached to the surface through three V−O−support bonds in a pyramid-like geometry [30–32]. The molecular structure of this configuration is depicted in Fig. 4a. This assignment was supported by the coordination of the V atom derived from EXAFS analysis, which showed three oxygen atoms at a distance of 1.82 A˚ and one at a distance of 1.62 A˚ [33]. The high natural abundance and large magnetic moment of the 51 V isotope allows one to derive the coordination of V atoms in these systems from solid-state NMR studies [34–37]. Two different types of V centers were observed in Al2 O3 - and TiO2 -supported V catalysts at different vanadium oxide loadings [36]. A comparison with crystalline reference compounds shows that one type, which is predominant at low vanadium oxide loadings, indeed represents V atoms in a four-coordinated V−O environment. Another contribution stems from octahedrally coordinated V sites and is present only at higher vanadium oxide loadings [38]. In the case of V catalysts,

3.2.5 Active Phase–Support Interactions

O

O

O

O

850 cm−1

O

O

O

O O

930 cm−1

O

O

O

O

Nb

V O

O

O

1005

O W

O

O

O

O

790 cm−1

O

O

O V

O

O

O

W O

920 cm−1 1030 cm−1

O

O

Mo O

O

O

O

1005cm−1

O

Mo

Cr

O

O O

996 cm−1

O

O

O Cr

Cr O

880 cm−1

O

1010 cm−1

1030 cm−1

cm−1

O

890 cm−1

1191

980 cm−1

Nb O

O

O

O

O Re

Re O

O O O

O O

O O

Schematic drawings of the molecular structures of MoO3 , Nb2 O5 , WoO3 , V2 O5 and Re2 O7 present on ZrO2 surfaces under dehydrated conditions, including the observed vibrational frequencies belonging to the molecular bonds.

Fig. 3

the position of the M−O−M band is at around 920 cm−1 . The observation of this band at extremely low V loadings makes this assignment at least questionable and alternative explanations have been suggested. For instance, in the case of V/Al2 O3 catalysts, it was assigned to a V−O−Al vibration [6, 39]. A recent investigation has shown that both Raman bands can just as well be explained by assuming a VO4 group to be present connected to the support by only one oxygen bridge, as depicted in Fig. 4b [40–42]. In this model, one terminal V−O bond is much shorter that the other two, which are of similar length to the V−O bond going to the support, in accord with the observations from EXAFS and NMR studies. Furthermore, the single V=O terminal bond in this ‘‘umbrella’’ model is in agreement with conclusions from the isotope exchange experiments mentioned above [30]. This example of discussion highlights that it is very difficult to link spectroscopic observations unambiguously with molecular structures under dehydrated conditions and often no firm conclusions can be drawn. One of the problems stems from the fact that, in contrast to hydrated complexes, no suitable reference compounds are available for these complexes. Theoretical calculations on relevant cluster compounds could help overcome this problem. However, for the construction of these models, insight

into the structure of the support surface is essential. Despite the huge amount of research dedicated to the characterization of especially SiO2 and γ -Al2 O3 surfaces using predominantly IR spectroscopy [43] and theoretical methods [44, 45], the description of these surfaces is still rather rudimentary. In general, it can be stated that a careful characterization of the catalyst material under study is essential for the interpretation of any catalytic test data. However, in the case of transition metal oxide catalysts, the exposure to reaction conditions can lead to severe structural rearrangements. More specifically to the supported vanadium oxide catalysts case study, small traces of water present in the reaction mixture may result in the formation of other molecular structure configurations as illustrated in Fig. 4c and d. In this case, water may result in a partial detachment of the supported vanadium oxides, resulting in the formation of one or two V−OH groups. Regardless of this important observation, the thermodynamically most stable state of a supported metal oxide catalyst is determined by the equilibrium of surface free energies of the support oxide material and the active metal oxide phase [46]. For a supported References see page 1196

1192

3.2 Chemical Properties

988

(g) O2 /He, 1 h ~969

(f) O2 /He, 0.5 h (a)

(e) MeOH /O2 /He 5 h

(c)

(d) MeOH /O2 /He 3 h

(c) MeOH /O2 /He 1 h

814

X 0.25

(b) MeOH /O2 /He 0.3 h (a) O2 /He 1100

1000

Raman (b)

900

800

shift/cm−1

(d)

In situ Raman spectra of a 4% MoO3 /TiO2 physical mixture in different environments at 230 ◦ C (a) flowing O2 –He; (b) methanol oxidation 20 min; (c) methanol oxidation 1 h; (d) methanol oxidation 3 h; (e) methanol oxidation 5 h; (f) flowing O2 –He, 0.5 h; (g) flowing O2 –He, 1 h (Copyright 1999 American Chemical Society).

Fig. 5

Silicon

Oxygen

Vanadium

Hydrogen

Molecular structures of the cluster models for an isolated VO4 group attached to an SiO2 surface. Representations of the pyramid model (a), umbrella model with peroxo group (b), umbrella model with one V−OH group (c) and umbrella model with two V−OH groups (d), as proposed in literature.

Fig. 4

metal oxide catalyst, the thermodynamic tendency will generally be towards complete coverage of the support oxide surface by the active phase and the formation of so-called monolayer catalysts. At temperatures above the Tammann temperature, surface migration of metal oxide complexes becomes possible and spreading of one oxide phase over the other can proceed. It has been shown, for instance, that treatment at elevated temperatures of physical mixtures of active oxides such as MoO3 , WO3 and V2 O5 and supports such as Al2 O3 and TiO2 can lead to the spreading of the MOx phase over the support oxide [47]. Transport of the metal oxide phase can also take place via the formation of volatile metal complexes, such as metal–alkoxy complexes that are formed during the oxidation of different alcohols over transition metal oxides [48]. To illustrate this reaction-induced spreading of transition metal oxides, a physical mixture of 4% MoO3 and TiO2 was exposed to a mixture of methanol and O2 at 230 ◦ C for several hours. Raman spectra were recorded on the sample during reaction and are presented in Fig. 5. Upon exposure to the reactant stream, the intensity of the

bands due to crystalline MoO3 (characterized by vibrations at 814 and 988 cm−1 ) that were present in the spectrum of the physical mixture, was found to decrease with time. At the same time, bands due to the presence of surface molybdate methoxy groups started to appear (969 cm−1 ). With increasing time on-stream, the activity of the sample in methanol oxidation steadily increased. In fact, after approximately 4 h, the activity and selectivity towards formaldehyde of the sample that was introduced into the reactor as a physical mixture were similar to those of an MoO3 /TiO2 catalyst prepared by impregnation. Reoxidation of the sample in O2 leads to the reappearance of MoO3 bands albeit of decreased intensity, as compared with the starting material. Apparently, a fraction of the MoO3 crystals was (partially) reduced during reaction and therefore not observed in Raman spectra, but reoxidized after exposure to O2 . Note that the reaction temperature at which this metal oxide spreading was observed was below the Tammann temperature (270 ◦ C) of MoO3 . Volatile metal–methoxy complexes provide an additional pathway for the spreading of the metal oxide phase at low temperatures. Similar experiments showed the reactioninduced spreading of V2 O5 , CrO3 , Cr2 O3 and Re2 O7 over TiO2 and SnO2 supports. However, when the same

3.2.5 Active Phase–Support Interactions

experiment was repeated for 4% MoO3 /SiO2 and 4% V2 O5 /SiO2 physical mixtures, no increase in catalytic activity was observed with time on-stream, although a decrease in intensity of the MoO3 and V2 O5 Raman bands was observed. Volatile metal–methoxy complexes were again formed, but spreading did not take place. Instead, deposition of metal oxide was observed in cooler parts of the reactor system [48]. The lack of interaction between SiO2 and catalytically active metal oxides has the effect that SiO2 -supported samples often show divergent behavior in catalytic studies. The drastic reorganization of the active phase under a reactive atmosphere once again illustrates the need to study catalysts under reaction conditions. In homogeneous catalysis, the influence of a certain ligand on the catalytic properties of a transition metal ion can be investigated in a relatively straightforward manner. Stoichiometric complexation of the transition metal ions can be brought about in solution and turnover numbers for metal complexes with different ligands can be derived with relative ease. To determine the effect of the support ligand on the catalytic properties of surface metal oxide groups is a much more intricate task. Ideally, one would like to obtain turnover numbers of similar metal oxide surface groups deposited on different support materials. One way to achieve this is to prepare catalysts with low transition metal oxide loadings. Hence it can be assumed that isolated metal oxide groups are created. However, in this case, a large part of the support oxide surface is still exposed and the catalytic activity of the support oxide material may hinder interpretation of the catalytic data. Furthermore, the catalytic activity of these low-loaded systems will be low and characterization of the active phase is difficult because of the low density of active sites. Another option is to prepare systems in which the support oxide surface is fully covered with a transition metal oxide phase. Since different support oxide materials show different specific surface areas, the metal oxide loading needs to be adjusted in each case. This approach was adopted to investigate systematically the influence of the support cation on the activity of supported vanadium oxide catalysts in a number of different reactions (methanol oxidation [49], butane oxidation [50] and the selective catalytic reduction of NO with NH3 [51]) for a number of oxidic support materials (Al2 O3 , TiO2 , ZrO2 , CeO2 , SiO2 , Nb2 O5 ). In determining the turnover frequency (TOF) per V atom, it was assumed in these studies that all V atoms are active in the reaction. Raman spectra recorded on these materials under dehydrated conditions indicate that the vanadium oxide phase consists of isolated and polymeric vanadium oxide surface groups and confirm the absence of any crystalline V2 O5 . The position of the main V−O vibration band at 1030 cm−1 , generally attributed to the V=O stretch

1193

vibration of isolated VO4 groups, was found to be the same for all materials. Furthermore, during reaction, the position of this band was found to remain at the same position, indicating that the molecular structure of the active site is nearly identical for all samples. It was concluded that reduction of the metal oxide phase did not take place to any great extent, as the intensity of the bands in the Raman spectra remained more or less constant. Despite the observed similarity in molecular structure of the active phase, turnover numbers for these systems in the reactions mentioned above were found to vary considerably from support to support. For instance, in methanol oxidation, a 104 -fold increase in the activity of the VOx phase going from CeO2 through ZrO2 , TiO2 , Nb2 O5 to Al2 O3 was established, as is illustrated by the turnover numbers for the various support materials presented in Table 2 [49]. This trend was explained by taking into account the Sanderson electronegativity of the support cation, which follows the trend Ce > Zr > Ti > Nb > Al. It can be imagined that a lower electronegativity of the support cation results in a higher electron density of the support−O−V bond. The higher basicity of the bridging oxygens renders them more reactive towards the adsorption of methanol and the rate-determining formation of a methoxy-intermediate on the catalyst surface. It is generally accepted that during oxidation reactions, the supported transition metal oxide phase undergoes a redox cycle in a Mars–van Krevelen mechanism. From this point of view, a number of workers have studied the influence of the degree of reduction of the MOx phase on the activity in different oxidation reactions [52–55]. Reduced transition metal oxide centers are generally not observed in Raman spectroscopy (the most powerful technique for the characterization of fully oxidized metal oxide catalysts) because of their low Raman cross-section. Although, the degree of reduction could in principle be derived from the decrease in the signal of the M−O bands in Raman spectra, coloring of the sample often makes it difficult to determine quantitatively the Turnover frequencies for methanol oxidation over supported vanadium oxide catalysts at monolayer coverage of surface vanadium oxide species [49]

Tab. 2

Catalyst 25% V2 O5 /Al2 O3 7% V2 O5 /Nb2 O5 6% V2 O5 /TiO2 4% V2 O5 /ZrO2 3% V2 O5 /CeO2 References see page 1196

TOF/s−1 0.068 0.4 1.1 1.7 10

1194

3.2 Chemical Properties

6

(a) 4

F(R∞)

(b) 2

0 −2

(c)

35000

25000

15000

Wavenumber / cm−1 UV/visible/NIR diffuse reflectance spectra of 1% V2 O5 /SiO2 (a) in the oxidized state and (b) after H2 reduction at 550 ◦ C for 1 h. (c) Difference spectrum (b–a) (Copyright 1999 Academic Press). Fig. 6

extent of reduction using this spectroscopic technique. UV/visible/NIR spectroscopy applied in the diffuse reflectance mode can be very informative in these cases. Fully oxidized metal oxide catalysts only exhibit intense ligand-to-metal charge-transfer (LMCT) bands. Reduction leads to a decrease in the intensity of these LMCT bands and the appearance of far less intense and often broad d–d transitions in the visible and NIR region of the spectrum. These changes are illustrated in Fig. 6, in which the spectra of a 1% V/SiO2 catalyst before and after reduction in H2 are presented [55]. Determination of the degree of reduction can now be accomplished by quantification of the decrease in intensity of the LMCT bands (>30 000 cm−1 ) or the increase in the intensity of the d–d transitions (10 000–30 000 cm−1 ). Both approaches have their advantages and disadvantages and in all cases an extensive calibration procedure is required to obtain reliable information. In an explorative study to evaluate the possibilities of UV/visible diffuse reflectance spectroscopy for obtaining quantitative structure–activity relationships, a series of supported Cr catalysts were studied in the dehydrogenation of butane [56]. Their activity was evaluated as a function of (a) Cr loading, (b) temperature, (c) time onstream, (d) zero point of charge (ZPC) of the support material and (e) isobutane partial pressure. A statistical model was constructed based on the use of ‘‘design of experiments’’ (DOE), in which the catalytic activity was expressed as a function of the experimental factors a–e. Activity was found to increase with increasing chromium oxide loading, reaction temperature, ZPC and time onstream and to decrease at higher butane partial pressures. Furthermore, during reaction, UV/visible spectra were

recorded on the catalyst bed. Reduction of Cr6+ was indeed observed when the catalyst was exposed to the isobutane feed at elevated temperature. All UV/visible spectra were found to consist of two pure component spectra, one containing contributions of the support and Cr6+ LMCT bands and the other the bands of the d–d transitions of Cr3+ and Cr2+ reduced centers. From analysis of the UV/visible spectra, the activity was found to increase linearly with the extent of reduction at a given time for certain reaction conditions. The highest amount of reduced surface sites and consequently the highest activity were found for a 7.5 wt.% Cr/Al2 O3 sample. In an in situ UV/visible study on supported vanadium oxide catalysts, it was found that the reducibility of the vanadium oxide phase is a function of the specific oxide used as the support material [55]. At similar vanadium oxide loadings, the degree of reduction after exposure of the catalysts to a reductive C2 H6 –He stream at 450 ◦ C follows the trend ZrO2 > Al2 O3 > SiO2 . This trend again parallels that in Sanderson electronegativity, mentioned above. It can easily be envisaged that an increased electron density of the V−O−support bonds makes for an easier reduction of the V cation. When the V loading on a specific support material is increased, the ease of reduction also increases. Apparently larger V2 O5 domains are more easily reduced. It was further observed that under ethane oxidation conditions (450 ◦ C, 2.4 : 8 : 81.6 C2 H6 : O2 : He), only a small amount (104

Possible pathways for formation of soot-like matter (from Refs. [24, 28]).

Fig. 1

Growth of precursor molecules

References see page 1214

Formation of primary particles

Nucleation, chemical reactions

0.5 nm

1 – 3 nm

Growth, coagulation, coalescense, aggregation

5 nm

Agglomeration

Fuel

Oxidation etc.

Ring formation, growth

10 – 50 nm

Reaction time

Fig. 3

104

2.0

0

CX

106

Atomic mass/amu

(CX HY AZ)N

Soot aggregates 108

Approximate time/ms

Dehydrogenation and loss of heteroelement A

Concurrent polymerization and dehydrogenation

1010

10

Dehydrogenation and loss of heteroelement A

Polymerization

CX HY AZ

and their resistance to regeneration. The formation of such cokes may include educt molecules, intermediates, impurities and product molecules. Different mechanisms

Particle diameter/nm

feedstock molecules to carbonaceous entities consisting of more than 104 carbon atoms. Figures 2 and 3 give an impression of the time-scale of the formation and growth of such grades of soot-like carbons. In addition to the formation and deposition of sootlike carbon species in thermally driven processes, the interactions of carbons from catalytically driven coking and their catalytic conversion are at the focus of catalyst research, since the chemical nature of the various grades of carbonaceous entities and the strength of interaction with particular surface sites of catalysts are crucial with respect to a detrimental impact on activity

1199

Growth of particles from fuel to soot-like entities over time (adapted from Ref. [27]).

>100 nm

>1µm

1200

3.2 Chemical Properties

Type A: Clean Pd surface

Type C: Partial carbon overlayer and bulk carbide

Type B: Complete carbon overlayer on Pd particle

Type D: Partial carbon overlayer, no bulk carbide

Type E: Clean Pd surface with interstitial carbon

Turbostratic carbon support Fig. 4

The varieties of Pd or Pd/C particles that may form on high surface area carbon supports (from Ref. [33]).

and reaction sequences have to be considered in the individual process, as has been highlighted [1–11, 27–32]. Examples of more strongly adherent carbons which interact more strongly with the catalyst are depicted in Fig. 4 [33]. Numerous investigations of model catalysts, and also industrial catalysts, have shown that distinct transformations of carbonaceous deposits may occur during catalyst operation. A prominent example is the Fischer–Tropsch synthesis. Studies on unsupported Fe and Fe/K catalysts led to a deactivation model: Simultaneous conversion of atomic to polymeric to graphitic carbon and of active carbon-rich carbides to inactive carbon-poor carbides was reported [34] in the context of a competition model [35] with several reaction pathways of carburization, synthesis and deactivation competing for the carbon which is deposited on the surface by the dissociation of CO. The rates of carbide formation, the carburization behavior, coking and the decomposition of Ni3 C as a function of temperature in Ni/SiO2 and 4Fe : Ni/SiO2 catalysts have been revealed [36]. Also for nickel catalysts applied in steam reforming in the production of synthesis gas and hydrogen, detailed information on the dynamic changes at the catalyst surfaces are available. During investigations of methane steam reforming on Ni/αAl2 O3 catalysts, it was observed by XPS depth profiling that carbon diffuses and dissolves in the bulk and is partly carbidic. The slow diffusion of carbon back to the surface causes a slow combustion of carbon in the catalyst [37]. The interaction of hydrogen with carbidic carbon deposited on Ni(100) in isothermal hydrogenation experiments by decomposing CH4 and

by H2 -TPD (temperature-programmed desorption) at carbon coverages of 0–0.34 monolayers was studied. In the hydrogenation curves carbon states with different reactivities were observed. The effective hydrogen binding energy depends on both the grade of carbon and the hydrogen coverage [38]. In addition to the formation of carbides, the catalytic transformation of amorphous carbon into filamentous carbon and related sp2 -type entities such as graphite platelets is an important phenomenon. Transmission electron microscopy (TEM) examination of the formation of carbon whiskers from acetylene on nickel showed that the growth rate of the whiskers is related to the rate of diffusion of carbon in Ni. The growth rates of whiskers of different diameters (15–36 nm) were determined as 52–90 nm s−1 . The influence of temperature and of hydrogen exposure prior to the exposure to acetylene was studied together with the nucleation of amorphous carbon around the nickel particles and the formation of graphite platelets at the edges of larger Ni particles [39]. In situ TEM of Ni/MgAl2 O4 catalysts which were heated at 720 ◦ C in a gas mixture of nH2 /nCH4 = 1 : 1 at 5 mbar (1 mbar = 100 Pa) in the TEM have shown that most of the Ni particles were observed to be encapsulated in layers of graphitic carbon but some of the metal particles were active in the growth of whiskers. Lower growth rates of

C7 H16 > cyclohexane > trimethylbutane > C4 H10 was reported [89]. In catalytic cracking, the basicity of the hydrocarbon has a great influence on coke formation (see Fig. 8 in Ref. [14]). Hydrocarbons with structures like cyclopentadiene, phenanthrene and others were recognized as potent coke precursors [16]. In the reforming of naphtha, the absence of alkali metals and of silica was advantageous in increasing of catalyst 3.2.6.3

activity and lifetime [89]. Conversely, effects of potassium and other electropositive elements on the microstructure and the three-dimensional secondary structure of carbons are systematically used in modifying and adjusting the structure of carbons on a technical scale [90]. Alkali and alkaline earth elements can also contribute to the specific morphology of carbonaceous deposits and, therefore, the adhesion of the carbons concealing a catalyst surface. CO2 reactivity of carbons is correlated with the alkali metal content. The reactivity may increase with increasing content of, particularly, calcium, potassium or sodium (see, e.g., Ref. [91]). The influence of coke graphitization can be compensated for by alkali enrichment. Coking may be reduced by alkali metals due to reduced acidity of the support [92]. In graphitic matter, alkali metal entities can be intercalated between the sp2 layers. Intercalation was observed to enhance the dihydrogen uptake of carbons 10-fold due to an enlargement of the interplanar spacing of the graphite sheets facilitating the access of H2 molecules [93]. Coke Distribution, Porosity and Pore Blocking An inhomogeneous distribution of carbonaceous deposits in porous catalysts can be of major concern in catalyst deactivation and regeneration. The local maximum temperature rise during regeneration of a catalyst in which coke is distributed non-uniformly may greatly exceed that obtained for a uniformly coked pellet [94]. In the temperature rise in pellets during gas–solid diffusion-controlled reactions, effects of heat- and masstransfer resistances have to be considered. Sintering of catalyst pellets in which reactions such as combustion of coke occur can be induced [95]. Pore blocking, coke and gas-phase concentration profiles inside a catalyst particle due to tortuosity and diffusion phenomena have been discussed [96]. Low coking rates could be achieved by proper feedstock blending to prevent pore mouth blockage [16]. In studying fractured or cleaved catalyst pellets, optical microscopy, scanning electron microscopy (SEM) together with energy-dispersive X-ray microanalysis (EDX), electron probe microanalysis (EPMA) and ion microprobe analysis with secondary ion mass spectrometry (SIMS) are used for detecting the carbonaceous deposits and recording carbon distribution profiles across the catalyst particles by electron beam or ion beam line scans [97]. The irregularities of the carbon concentration throughout the pellets of reforming catalysts were noted. Rapid deactivation of H-mordenite catalysts is generally attributed to its one-dimensional structure. Even small amounts of carbonaceous deposits in the pore mouth or low molecular weight coke in a pore channel may block 3.2.6.4

3.2.6 Carbonaceous Deposits

most of the active sites in the channel. In n-heptane cracking, even 1% of coke could block the access of n-heptane to a pore volume 10 times greater than the volume really occupied by the coke. Models of the location of carbon at low and high coke content were given as derived from adsorption data on the accessibility of molecules of different size [98]. For ZSM-5, offretite and mordenite zeolites, the formation of ‘‘external coke’’ deposited under hydrogen-deficient conditions and of ‘‘intracrystalline coke’’ deposited in the presence of hydrogen released during aromatization reactions were highlighted together with an impact of shape-selective effects in the removal of coke from the different kinds of zeolite channel networks [99]. Nuclear magnetic resonance (NMR) techniques are frequently employed to reveal the influence of coke on the deactivation of zeolites [99, 100]. Characterization of carbonaceous deposits in zeolites is carried out by 13 C magic angle spinning (MAS) NMR and cross-polarization (CP) techniques [101]. For example, the mechanisms in alkylation of isobutene with 1-butene on LaY zeolite were revealed. At typical alkylation temperatures (80 ◦ C), only paraffinic carbon was observed, whereas at higher temperature aromatics and olefins were detected by means of 13 C CP/MAS NMR [102]. The technique gives detailed information on the chemical nature of carbon atoms but cannot discriminate between carbon inside the pores and on the outer surface of zeolites. 29 Si and 27 Al NMR are used to monitor changes of the zeolite framework during coke deposition, such as a structural distortion by the growth of carbonaceous deposits in the confined geometry of the zeolite channels and the influence of extra-framework Al species. 1 H MAS NMR is able to focus on Brønsted acidic sites as a function of the degree of coking and on the hydrogen content of the carbons. 129 Xe NMR provides information on the location of the carbon by adsorbing Xe in porous materials (Ref. [100] and literature reviewed therein). Complementary to NMR techniques for in situ studies of coke deposition, e.g., on H-ZSM-5 catalysts, in situ Fourier transform infrared (FTIR) spectroscopy is used to characterize the chemical composition of carbonaceous deposits. The properties of coke formed during the conversion of ethylbenzene were studied in the C−H deformation vibration range, including the ‘‘coke’’ band at 1610 cm−1 [103], the position of which shifts to lower wavenumbers with increasing sp2 character of the carbonaceous deposits with increasing time on-stream. The band is assigned to the C=C stretch in aromatic and highly conjugated structures. A band at about 1630 cm−1 is assigned to stretching vibrations of double bonds in polyolefinic structures rather than C=C vibrations in aromatic rings [104]. Vibrational bands at 1610 and 1490–1510 cm−1 were used to evaluate the amount of

1207

coke deposited with time of catalyst operation [103]. In the transformation of m-xylene on USHY, zeolite coke formed at 723 K was more polyaromatic than coke formed at 523 K [105]. Time-resolved in situ FTIR permits the study of the transformation of olefinic molecules into aromatic/graphitic structures with increasing reaction time or temperature [106] and the evolution of these structures of varying nH /nC ratio. Also, the influence of temperature and hydrogen pressure on the regeneration of zeolite catalysts coked in, e.g., isobutene/butane alkylation has been studied to identify suitable conditions for restoring the activity [107]. FTIR together with capillary gas chromatography was used to evaluate the extent of deactivation by propylene oligomerization and coke deposition in the acid-catalyzed alkylation of benzenes with propylene to cumene on Beta-, NU-87, EU-1 and MCM-22 zeolites with regard to propylene turnover and coke formation [108]. Another in situ technique is the measurement of mass changes during coking in correlation with deactivation by means of a tapered element oscillating microbalance (TEOM) in combination with gas chromatography of product mixtures. This was demonstrated in the acidcatalyzed conversion of 2-propanol to propene and diisopropyl ether on H-Y, H-ZSM-5, H-ZSM-22 and H-mordenite zeolites, which have different pore systems and, therefore, different sensitivities towards carbon deposition [109]. Combining surface analytical methods such as XPS and SIMS with cycles of surface erosion by ion sputtering of a catalytic material complements electron microscopy in order to localize and to study the in-depth distribution of carbonaceous deposits with enhanced depth resolution and to discriminate between pore blocking on the one hand and physical shielding or poisoning of active surface sites on the other [110]. Deactivation of a weakly acidic borosilicate of the MFI type (B-MFI) during use in the vapor-phase Beckmann rearrangement of cyclohexanone oxime to ε-caprolactam could not be attributed to an enhanced deposition of carbons such as coke or degradation products at the catalyst surface. Only about 16% of the surface area of the most strongly deactivated catalyst was covered with carbon. An enrichment of nitrogen-containing linear polymeric species formed by ring-opening reactions of unsaturated nitriles preferentially formed inside of the zeolite was responsible for the deactivation by a pore blocking mechanism. A pronounced and irreversible deboration effect in the surface regions of the zeolite as a reason for catalyst deactivation could be ruled out. SiO+ /SiOH+ SIMS fragment ion profiles indicated varying relative amounts of silanol groups and changes in silanol group References see page 1214

1208

3.2 Chemical Properties

distribution gradients in the near-surface regions of the pore system of the zeolite at different grades of deactivation and carbon penetration into the porous solid. Graphiticity The bulk structural properties of carbons of varying aromatic/graphitic contributions are usually evaluated by means of X-ray diffraction (XRD). Graphitic ordering, paracrystallinity or amorphous entities can be detected and information on the crystallite dimension (Lc) and turbostratic or lubricostratic mismatch in the stacking of graphite sheets (see Ref. [24], p. 91) is accessible. The sharper the 002 Bragg peak, the higher is the degree of ordering of the carbon structure. The use of the Scherrer equation allows the calculation of the size of the crystalline regions [98]. It may be preferable to evaluate the graphiticity of the carbonaceous matrix without interference from background contributions by adsorbed aliphatics, green oils and other amorphous matter. Therefore, carbonaceous deposits may be studied before and after appropriate extraction pretreatments similar to the case of elemental hydrogen analysis (see Section 3.2.6.6). In many cases, additional information on the fine structure of the carbons is required and the complementary use of transmission electron microscopy techniques (bright field, dark field, phase contrast and high-resolution imaging [111]) is highly important in order to obtain a coherent picture of the paracrystallinity or of the degree of aromaticity or graphiticity of the carbonaceous matter. The combination of data from various methods provides a representative impression on the sp2 character of the carbon atoms. A typical analytical procedure for characterizing the structure of carbonaceous deposits on Pt/alumina reforming catalysts has been reported [45], combining the information from XRD with results from TEM and other techniques. For obtaining an overview of the morphological peculiarities of low-volatility carbonaceous deposits on catalysts at a mesoscopic scale, scanning electron microscopy is an additional valuable tool [44]. In order to evaluate the electronic structure of, e.g., gas oil-derived coke deposits in zeolites, the carbon K-absorption edge has been measured in a special version of X-ray absorption spectroscopy (XAS): fluorescence yield near-edge spectroscopy (FYNES). In combination with 13 C MAS NMR, XPS and elemental analysis, location and characterization of the coke is possible. Energies and relative intensities of the C1s → π ∗ and σ ∗ resonances can give information on the degree of polyaromatic annellation and graphiticity in the cokes in comparison with reference data on well-defined carbons [112]. 3.2.6.5

Similar information is accessible with much higher spatial resolution by means of electron energy loss spectroscopy (EELS) in the transmission electron microscope [113, 114]. Combination with elemental analysis, surface-specific techniques (XPS, SIMS etc.) and bulk techniques, such as NMR, provides the maximum information on carbons at the surface and inside a porous catalytic material. Complementary to the use of XAS or EELS, the XPS signal of carbonaceous materials can be evaluated. The C1s XPS signal of purely aliphatic carbon exhibits a fairly symmetrical shape. The signals of graphitic carbons are narrow but asymmetrically broadened towards higher binding energy, because of interactions between the emitted photoelectron and exciton states near the Fermi energy in a final state effect [115]. The transition from a symmetrical to an asymmetric C1s line envelope has been studied on polynuclear aromatic compounds such as coronene, ovalene and pentacene. It occurs when the aromatic structures exceed dimensions in the range 1–2 nm [116, 117]. The asymmetry of the signal represents the mean area over which the π-electrons of the aromatic system are delocalized [116]. This structure sensitivity of the C1s signal shape has been used to evaluate the degree of order and disorder in graphitic/aromatic structures and cokes [118]. Not only the shape but also the position of the C1s peak of sp2 and sp3 carbon may show variations [119, 120]. The specific properties of a precious metal at the catalyst surface seem, along with variations in the feed-stream nH /nC ratio, to be highly relevant with respect to the nH /nC ratio of deposited carbon. SIMS measurements revealed pronounced differences in the hydrogen-related properties of carbonaceous deposits produced from ethylene on the surfaces of polycrystalline Pt, Rh, Ir and on PtRh and PtIr alloys [121–124]. Examination of various carbonaceous materials shows that the SIMS fragmentation patterns, and also the depth dependences of the secondary ion yields, depend on the crystallinity and the hydrogen content of the material [44]. The CH− fragmentation channel is indicative of finely crystalline, microcrystalline or amorphous material, whereas carbon of higher structural order is represented by enhanced − relative intensities of the C− 2 and C2 H fragment ions and, especially, the acetylene-like fragment ion C2 H− 2 under the conditions of quasistatic/dynamic SIMS. The matrix dependence of the secondary ion yields of carbons can be exploited to obtain qualitative analytical fingerprints associated with the sp2 character of coke and, hence, its micromorphology. A suitable method to complement NMR and FTIR for the evaluation of the internal sp2 /sp3 structure of cokes from the hydrogen point of view at a bulk level is IINS [125]. Each IINS spectrum in Fig. 10 gives an

3.2.6 Carbonaceous Deposits

S (Q,w)

0.12 0.08 (a) 0.04

(b)

0.00 0

500

1000

1500

2000

0.5

S (Q,w)

0.4 0.3 0.2

(c)

0.1

(d)

0.0 0

200

400

600

800

1000

Wavenumber/cm−1 Fig. 10 IINS spectra of cokes from large-scale chemical reactors as collected from the catalysts surfaces. (a) CVD-grade, low-temperature coke deposited at the surface of a Pd/SiO2 catalyst (hydrogen content 44.100 ppm H). (b) Graphite-type, high-temperature coke deposited at the surface of a Pt/Al2 O3 catalyst (residual hydrogen content: 43 ppm H). (c) Low-temperature coke from the surface of a Pd/SiO2 catalyst. (d) Includes results of a computer simulation of the spectrum of [Fe(H2 O)Cl5 ]2− (dashed line).

impression of the bulk carbon and the hydrogen-related properties (proton dynamics) of up to 20 g per sample [77, 79]. Completely different mechanisms of carbon deposition were involved. Figure 10a shows the IINS spectrum of carbonaceous deposits collected from an industrial catalyst operating at low temperature in the HCl recycle gas stream of the vinyl chloride process (overall hydrogen content: 44 100 ppm). The presence of residual water and its potential contribution to the analytical data can be excluded by IINS with high sensitivity. Principally, the spectrum of polymer-like carbon originating from copolymerization of acetylene and ethylene was expected. However, Fig. 10a perfectly matches the spectrum of a well-defined CVD carbon [126]. X-ray fluorescence (XRF) analysis indicated that more than 50 wt.% of this material consists of inorganic contaminants including Fe and Cl, which are closely linked to the finely divided, highly absorbing carbon component of the deposits. XPS also confirmed the presence of reduced Fe. In the IINS spectrum, the inorganic component of the coke is largely suppressed (by virtue of the small scattering cross-sections of the inorganic materials) and the hydrogen-containing part of the coke is selectively revealed as a pure CVD material.

1209

The low-energy graphite band at 112 cm−1 indicates the presence of sp2 entities originating from catalytically driven coke transformation of amorphous carbon into aromatic/graphitic entities of enhanced size under the influence of iron contaminants. The band at 270 cm−1 is a measure of the degree of cross-linking between sp2 -type clusters and sp3 -type carbon chains containing CH2 (bands at 1300–1500 cm−1 [125, 126]). It contains information on the interactions between different carbonaceous components of the bulk of the carbonaceous network. The Robertson model for a-C : H (amorphous hydrogenated carbon) is applicable [127, 128]. Related information on the internal structure of carbonaceous deposits has been obtained by studying coke formation in hydrotreating catalysts. The coke precursors were identified by HPLC and gas chromatography coupled with mass spectroscopy (GC/MS). It was concluded that the initial adsorption of the molecules in the polyaromatic fraction takes place via heteroaromatic nitrogen or the π-electrons in the aromatic compounds. Aromatic and aliphatic fractions were evaluated by 13 C MAS NMR. It was also shown that the coke consisted of condensed aromatic clusters with short terminal or bridging aliphatic species. The coke did not act as a poison but caused deactivation by physical blocking of the catalyst surface [129]. In the case of highly graphitic carbonaceous deposits, the IINS spectrum of which is depicted in Fig. 10b (overall hydrogen content 43 ppm [70]), the use of NMR or FTIR was prevented by the electrical conductivity and strong absorption of electromagnetic radiation by the carbon atoms. The residual traces of hydrogen terminate the edges of the graphite sheets. With respect to the operating temperature of the catalyst, spectra similar to the soot-like material were expected for this kind of methane-derived coke. However, the spectra strongly resemble the IINS features of pure polycrystalline graphite and of graphitized carbon black as formed at about 3000 ◦ C [125, 126, 130]. The high purity and sp2 character of the material are indicated by the sharp graphite modes at 112 cm−1 and the fact that the other vibrational features are confined to the region below 1700 cm−1 (graphite has no fundamental modes above 1700 cm−1 ). Evidence for sp3 carbon is largely missing. This indicates an accidentally very low partial pressure of hydrogen during the deposition and growth of this clean species of coke. The presence of orientation effects which are known for the anisotropic structure of well-defined graphites are also indicated by the double structure of the graphite band inside this coke. In contrast to Fig. 10a, the spectrum of a carbon from the same process is shown in Fig. 10c; this strongly resembles the signals of the well-defined species References see page 1214

1210

3.2 Chemical Properties

[Fe(H2 O)Cl5 ]2− [77, 79], the simulated spectrum of which is also included in Fig. 10d. Alternative structures would show different vibrational spectra. The strongest band at 386 cm−1 is assigned to the Fe−OH2 torsional mode, indicating another main reason for catalyst deactivation present. These examples illustrate that under certain conditions, extensive growth of pure CVD-like carbon occurred, which, according to XRF, XPS and REM/EDX measurements, contained significant amounts of welldispersed iron, iron chlorides and other contaminants. On the other hand, the presence of traces of moisture can lead to some dew point corrosion in the HCl recycle gas stream and to the formation of a distinct [Fe(H2 O)Cl5 ]2− species dominating the whole INS spectrum of this coke. In spite of the industrial origin of the cokes, Fig. 10a–d reveal the presence of well-defined components dominating the IINS spectra of these apparently ill-defined samples: pure oriented graphite, CVD carbon and one iron complex of defined H2 O content. Carbonaceous deposits of porous shape were formed by clustering of polymer droplets under the influence of acetylene during steam cracking of ethane (nC2 H6 : nH2 O = 3 : 1, T = 850 ◦ C). Electron microscopy revealed an influence of the presence of sulfur on the raggedness of the surface of these materials [131]. Also 15–50-nm size metal particles were involved. It was concluded that a certain fraction of the deposits at the outlet of the steam cracker were formed due to the presence of acetylene. Hydrogen Content and Hydrogen Analysis The specific properties of different types of carbonaceous deposits and cokes can have an impact on chemical reactions. For example, coke deposited on a reforming catalyst affects the hydrogenolysis more than the isomerization of n-pentane [132]. According to observations reviewed in Sections 3.2.6.2 and 3.2.6.5, parameters that govern the sp2 /sp3 character and, hence, structure, morphology and adhesion of carbonaceous deposits to the catalyst surface or their impact in deactivation include the presence of additional components catalytically relevant in coke-forming and -transformation processes, impurities, debris, catalyst poisons, etc. However, the key analytical information is the hydrogen content of and its chemical bonding to the carbonaceous entities. This is extremely valuable information for a fast assessment of its sp2 /sp3 character. It allows a rough, but nevertheless safe, estimate of the chemical structure of the deposits. Limiting cases in the nH /nC ratio are (i) aliphatic deposits (sometimes referred to as ‘‘white’’ or ‘‘soft’’ or ‘‘smooth’’ coke, predominantly low-temperature carbons) with a structure of, e.g., a polyolefin running through the catalyst pore system and an nH /nC ratio close to 2 and (ii) condensed polyaromatic/graphitic deposits (referred 3.2.6.6

to as ‘‘black’’ or ‘‘hard’’ or ‘‘high-temperature’’ coke) which is strongly deficient in hydrogen (nH /nC around 0.5 or even much lower). Type (ii) carbonaceous deposits mainly from high-temperature processes may resemble the structure and properties of purely graphitic materials with hydrogen contents down to V 0 ). The quantity  (c − c0 ) dV (8) na = na (app) + c0 V s = Vg

w

p

Schematic arrangement of a simultaneous volumetric and gravimetric adsorption experiment.

Fig. 1

V /v g is the amount of gas which would be contained in the volume V if the gas concentration were uniform throughout the volume. That the amount actually present is n0 means that local variations in gas concentration must occur: the gas concentration within the bulk of the solid is zero, but is greater than c0 in the interfacial layer. The difference between n0 and V /v g may be called the apparent adsorption na (app) = n0 −

V vg

(5)

and is a directly observable quantity. The precision with which it can be measured is controlled only by the experimental precision in T, p and V, and by the reliability with which v g (and c0 ) can be calculated from the equation of state of the gas. The apparent adsorption may, alternatively, be defined by measuring the amount of gas which has to be added to the system at constant T, p to increase the volume V back to V 0 . The apparent adsorption is then equal to the extra amount of gas which can be accommodated in a volume V 0 at a given T, p when the solid is introduced. It can, therefore, be expressed in terms of the local deviations of the concentration, c, of adsorptive molecules, from the bulk concentration c0 :  a (c − c0 ) dV (6) n (app) = V0

is thus equal to the Gibbs adsorption (surface excess amount of adsorbed substance, cf. [4]) when the surface of the solid is taken as the Gibbs dividing surface: it is the difference between the amount of substance actually present in the interfacial layer and that which would be present at the same equilibrium gas pressure in a reference system in which the gas-phase composition is constant up to the Gibbs surface, and in which no adsorptive penetrates into the surface layer or the bulk of the solid. The operational definition of na is thus  na = (c − c0 ) dV = n − c0 V g (9) Vg

where n is the total amount of gas admitted. The precision with which na can be determined depends, not only on the precision of T , p, V and v g but also on the precision with which V s (and hence V g ) is known. The volume of the solid (i.e. the volume enclosed by the Gibbs surface) is often defined experimentally as that volume which is not accessible to a nonsorbable gas (e.g. helium–leading to the helium dead space). In making this identification it is assumed that the volume available to He atoms is the same as that for molecules of the gas under investigation (which is not true, for example, if the solid acts as a molecular sieve, or if the molecules of the gas are significantly larger than the He atom), and that the solid does not swell under the influence of the adsorbate. Helium adsorption may occur if the solid contains very fine pores (or pore entrances) and the only proof that the adsorption is zero is that the apparent value of V s is independent of temperature. It is usual to take the high temperature limit of V s as being the correct value, but if V s is determined at a temperature widely different from

3.3.1 Reporting Physisorption Data for Gas/Solid Systems

that used in an adsorption experiment, a correction for the thermal expansion of the solid may be required. Alternatively, V s may be estimated from the known density of the bulk solid with the implied assumption that this is the same as that of the material of the adsorbent. The above discussion in terms of the volumetric technique when applied to gravimetric measurements gives for the apparent change in weight
 Vs w = w − w0 = na − g M (10) ν where M is the molar mass of the adsorptive. Thus na =

Vs w + g M ν

(11)

The second term on the right-hand-side is the buoyancy correction which has the same origin as the deadspace correction in a volumetric determination. An alternative but less useful definition of adsorption is  c dV (12) ns = Va

where Va = τ As is the volume of the interfacial layer and c is the local concentration. Va has to be defined on the basis of some appropriate model of gas adsorption which gives a value of τ the layer thickness. Provided that the equilibrium pressure is sufficiently low and the adsorption not too weak, then ns = na

(13)

the surface excess amount (na ) and total amount (ns ) of substance in the adsorbed layer become indistinguishable and the general term amount adsorbed is applicable to both quantities. 3.3.1.4

Experimental Procedures

Outgassing the Adsorbent Prior to the determination of an adsorption isotherm all of the physisorbed species should be removed from the surface of the adsorbent. This may be achieved by outgassing, i.e. exposure of the surface to a high vacuum–usually at elevated temperature. To obtain reproducible isotherms, it is necessary to control the outgassing conditions (temperature program, change in pressure over the adsorbent and the residual pressure) to within limits which depend on the nature of the adsorption system. Instead of exposing the adsorbent to a high vacuum, it is sometimes expedient to achieve adequate cleanliness of the surface by flushing the adsorbent with an inert gas (which may be the adsorptive) at elevated temperature. With certain microporous solids reproducible isotherms are only obtained after one or

3.3.1.4.1

1221

more adsorption–desorption cycles. This problem can be overcome by flushing with the adsorptive and subsequent heating in vacuum. Where physisorption measurements are to be employed for the determination of surface area and/or porosity, the rigorous surface cleanliness required in chemisorption studies is unnecessary and outgassing to a residual pressure of ca. 10 mPa is usually considered satisfactory. Such conditions are readily achieved with the aid of conventional vacuum equipment–usually a combination of a rotary and diffusion pump in conjunction with a liquid nitrogen trap. The rate of desorption is strongly temperature dependent and to minimize the outgassing time, the temperature should be the maximum consistent with the avoidance of changes in the nature of the adsorbent and with the achievement of reproducible isotherms. Outgassing at too high a temperature or under ultra-high vacuum conditions (residual pressure 200, say). It should be appreciated that the BET analysis does not take into account the possibility of micropore filling or penetration into cavities of molecular size. These effects can thus falsify the BET surface areas and in case of doubt their absence should be checked by means of an empirical method of isotherm analysis or by using surface area reference samples (see Section 3.3.1.6.2). For the determination of small specific surface areas ( CH4 > N2 , and they are also reasonably close to the permselectivity reported by Kusakabe et al. [97] for the binary mixture CO2 /N2 (≈5–7). It should be noted that permeability ratios measured on a given membrane practically coincide with the corresponding permeance ratios; that is, the effect of the membrane thickness is largely factored out in taking the ratio. Although the predicted permeability ratios for the three pure gases through silicalite-1 under low occupancy conditions are in reasonable agreement with experimental data, the absolute values of predicted permeability are, by a factor of 120 to 340, higher than the macroscopically measured experimental values. These discrepancies in absolute values are most probably attributable to the fact that silicalite-1 has been modeled as a perfect, fully siliceous crystal. On the other hand, in macroscopic measurements, defects and aluminum (acid) sites are inevitably present, leading to a decrease in the mobility of sorbates. Intercrystalline boundaries, gel remaining References see page 1675

1674

5.5 Computer Simulations

Tab. 2

Predicted and experimental permeability (Pe ) ratios in MFI membranes

Sorbate CH4 /N2 CO2 /CH4 CO2 /N2 CH4 /N2 CO2 /CH4 CO2 /N2

Pe -ratio

Technique

Sorbent

T/K

Reference

1.2 2.3 2.8 1.6 2.2 3.4

Membrane permeation Membrane permeation Membrane permeation MD MD MD

ZSM-5 ZSM-5 ZSM-5 Silicalite-1 Silicalite-1 Silicalite-1

303 303 303 300/298 303/300 303/298

[96] [96] [96] [40] [40] [40]

7.5 × 10−3

Ds / (m2 s−1)

6.0 × 10−3 4.5 × 10−3 3.0 × 10−3 1.5 × 10−3 0.0 10−3 10−2 10−1 100 101 102 103 104 105 l/µm Fig. 12 Transition of self-diffusivity from the molecular to the Knudsen diffusion regime as a function of mean free path λ for ethane in the reconstructed bed of NaX crystals of Fig. 2 at 298 K. (Reproduced from Ref. [7].)

where p is the pressure and σ is the critical diameter of ethane. Upon collision with a crystal surface, a new direction is selected according to the cosine law [99]. These results confirmed the evidence for different apparent tortuosity factors in the Knudsen and bulk regimes, which was first observed via PFG NMR measurements in the same system [33, 34]. Further analysis of the kinetic Monte Carlo simulation results revealed that the latter difference may be attributed to the dependence of the molecular paths on the geometry of the particles and the porosity of the zeolitic bed [7, 98]. Intracrystalline transport barriers in MFI-type crystals due to intersections and intergrowths between the elementary structural crystal units play a significant role. Indeed, such a role has been studied experimentally by K¨arger and coworkers [100, 101], using PFG NMR, interference microscopy, and Monte Carlo simulation for methane, n-butane, and i-butane diffusion. Conclusions As a result of the continuing rapid growth of the modern computer, computer-aided modeling of transport phenomena under confinement has revealed fundamentally interesting and practically relevant aspects of the physics of sorbate molecules in micropores. The driving force behind this type of study in zeolites has been to gain insight into the molecular-level processes that shape sorption equilibria and transport rates governing the performance of these materials as catalysts and sorbents. Simulation at the molecular level has proven to be a prerequisite for the modeling of such processes, as conventional macroscopic techniques cannot address the physics of nanoscopically confined guest molecules in a predictive manner. Important sampling by Monte Carlo methods, mainly in the grand canonical ensemble, have been used to capture the sorption equilibria of a variety of fluids in zeolites. Configurational bias techniques have significantly enhanced the efficiency of insertions of long flexible (e.g., alkanes) and bulky (e.g., aromatic hydrocarbons) sorbates into the confining zeolite framework. The competition between adsorption enthalpy and entropic effects as a 5.5.2.5

between the crystals, or non-idealities in the crystals are additional sources of resistance to mass transport [97]. The intercrystalline diffusivity of ethane in the reconstructed bed of NaX (see Section 5.5.2.2.2B and Ref. [98]), as predicted over a wide range of mean free paths, from the Knudsen to the molecular regime via kinetic Monte Carlo simulation [7, 98], is illustrated graphically in Fig. 12. Random trajectories are generated in the void space of the medium in such a way so that in the bulk gas the lengths l between successive collisions follow an exponential distribution expected from the Poisson process character of intermolecular collisions:
 −l (9) l f (l) = exp l

where f (l) is the conditional probability of having a collision-free trajectory length between l and l + dl, with a mean value of l being denoted by l . In the bulk gas phase, √ l is the molecular mean free path λ = kB T /pσ 2 π 2,

References

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5.5.3

Computer Simulations of Shape Selectivity Effects Berend Smit∗ and Theo M. Maesen

Introduction Today, almost every gasoline molecule has seen the interior of a zeolite and experienced the effect of shape selectivity. Given the economical importance [1–5] implied by this statement, one would expect that we have a very good understanding of the mechanisms underlying shape selectivity. Clearly, we do have a very good understanding of the chemical reactions that take place inside the pores of a zeolite. In addition, many mechanisms have been proposed to explain the various product distributions that have been observed experimentally. However, we are still very far away that by looking at a zeolite structure we can provide a prediction of the products. Let us consider as a simple example the conversion of n-decane in an acid zeolite. The chemistry tells us that isomerization reactions take place giving mono-, di-, and tribranched isomers. These isomerization reactions compete with cracking, and this results in a large number 5.5.3.1

∗

Corresponding author.

5.5.3 Computer Simulations of Shape Selectivity Effects

of different types of shorter alkanes. Whereas the reactant is a simple linear n-alkane, the product is a complex distribution of many different components. So, the principle aim is to use shape selectivity in such a way that the optimal product distribution for a given application can be obtained. What should the appearance of the zeolite be in order to obtain as much dibranched products as possible? Can one envision a structure that selectively removes C5 from the product distribution? Answering these questions requires a molecular understanding of shape selectivity, the various forms of which are reviewed in this chapter. Weisz and Frilette coined the term ‘‘shape selective catalysis’’ during the 1950s, when they discovered that only molecules permeable into an LTA-type zeolite were catalytically converted, to the exclusion of others [6–9]. In a 1971 review Venuto pointed out that, in addition to shape selectivity, there are ‘‘. . . reactions on external surfaces and special effects’’ [10]. Whereas progress has been made in establishing and understanding external surface reactions [11–14] (variously called ‘‘nest effects’’ [15] or ‘‘pore mouth’’ [10, 11, 16–18] and ‘‘key-lock’’ [19–22] catalysis), the (other) ‘‘special effects’’ have mushroomed into a myriad of phenomena each with their individual name. An early example of these special effects is the ‘‘window’’ or ‘‘cage effect’’, coined by Chen et al. [10, 23]. However, this was later merged into the classical shape selectivity model – initially as a form of mass transport shape selectivity [23–25] and more recently as a form of sorption shape selectivity [26]. Santilli and Zones coined the term ‘‘secondary shape selectivity’’ to describe the selective hydroconversion of n-C6 instead of n-C16 by AFX-type zeolites [27]. Derouane postulated that ‘‘molecular traffic control’’ might occur when small n-CN

n-C=N

Me-CN−1

diMe-CN−2

Me-C=N−1

molecules can diffuse through small and large molecules through large channels of one and the same molecular sieve [28, 29]. Van Nostrand and coworkers coined the term ‘‘inverse shape selectivity’’ to denote the accelerated formation rate of reaction intermediates that have a shape more commensurate with the framework topology than others [30–32]. At this point, is it important to note that most of these mechanisms have been proposed with very little knowledge on the thermodynamics and diffusion properties of the adsorbed molecules. Molecular simulations have provided us with such thermodynamic and transport data on well-defined model systems. Here, we demonstrate that the availability of these thermodynamic data has given us new insights into the mechanism of shape selectivity. Conventional Hydroconversion Mechanisms Before addressing the effect of shape selectivity on these reactions, it is worthwhile first to discuss what occurs in the absence of shape selectivity. 5.5.3.2

5.5.3.2.1 Basic Mechanism In alkane hydroconversion, a metal site dehydrogenates alkanes into alkenes, an acid site converts the alkenes into isomers or cracking products, whereupon the metal site hydrogenates the converted alkenes back into alkanes [33–35]. When starting with an n-alkane, the hydroconversion can be described as a series of consecutive hydroisomerization steps, with each step increasing the degree of branching [35–37]. If this process is simplified by only considering methyl group branches, the hydroisomerization of an n-alkane of N carbon atoms can be described as in Fig. 1. In References see page 1690

triMe-CN−3

diMe-C=N−2

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tetraMe-CN−4

triMe-C=N−3

tetraMe-C=N−4

......

Me-CN−1−M+n-CM Me-CN−1−M+Me-CM−1 etc.

Schematic reaction mechanism of an n-alkane hydroconversion reaction on, for example, an amorphous aluminosilicate (i.e., in the absence of shape selectivity.) n-Alkane feed and hydroisomerization products (top) dehydrogenate into alkene intermediates (vertical ← →, e.g., Pt-catalyzed). Alkenes hydroisomerize in a chain of acid-catalyzed hydroisomerization reactions (horizontal ← → ). With increasing degree of branching it is increasingly more likely that isomers crack (vertical →, acid-catalyzed) and hydrogenate into smaller alkanes (vertical ← → , e.g., Pt-catalyzed).

Fig. 1

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5.5 Computer Simulations

addition to the hydroisomerization reactions that change the degree of branching, there are also those that change the distribution of branching towards thermodynamic equilibrium [21, 38–40]. None of the hydroisomerization reactions equilibrates completely because they compete with consecutive hydrocracking reactions that decompose the isomers [21, 36, 38–42]. The probability of a molecule undergoing a hydrocracking reaction increases with increasing degree of branching, because more extensively branched isomers afford the formation of more stable carbocationic hydrocracking transition states [21, 37–40]. For as long as the molecules are adsorbed, we refer to them as reaction intermediates, and they can either desorb to become a product or continue to react. Those intermediates that have two or more methyl groups sufficiently close to each other such that they can hydrocrack relatively fast, are termed cracking precursors. Figure 1 also illustrates that, even in an ideal hydroconversion experiment in which a single component is utilized as feed, the product is a complex mixture of products that originate from the hydroisomerizing and hydrocracking of reaction intermediates. Shape selectivity is used to optimize the product distribution for a given application. To illustrate the type of product distribution to be expected in the absence of shape selectivity, let us consider the hydroconversion reaction of an n-alkane in more detail. Figure 1 suggests that the hydroisomerization reactions allow the formation of any branched isomers. At high and intermediate alkene coverage of the acid

2,2diMe-C8

n-C6 + i-C4

3/

10

sites α, α, γ -trimethylalkene hydrocracking dominates the hydrocracking product slate [26]. When this is the case, the product slate consists of a histogram with a single maximum indicative of preferential hydrocracking at the center of the chain irrespective of the n-alkane feed length [26, 34–36, 43, 44]. This is because the probability of formation of α, α, γ -trimethylalkene hydrocracking precursors is dependent on the proximity of the methyl groups to the center of the chain [26, 34, 36, 37, 43]; for reasons of symmetry there are fewer permutations of their precursor transition state closer to the center [26]. For the system shown in Fig. 1 we can therefore expect a product distribution dominated by products originating from 3,3,5-trimethylheptane, and from 3,5- and 3,3- and 4,4-dimethyloctane (see Fig. 2). In addition to the traditional kinetic n-alkane hydroconversion network discussed here, an alternative network was proposed [32]. Both networks agree with the older literature [45, 46] in that the cationic dialkylcyclopropyl transition states play a key role in hydroisomerization, but they disagree as to the role of these transition states in hydrocracking. The traditional network postulates that the dialkylcyclopropyl transition states only play a role in hydroisomerization, and that they have to open and form a fully fledged branched alkene before hydrocracking occurs [36, 37]. The alternative network postulates that the dialkylcyclopropyl transition states do not have to open before hydrocracking sets in, and that they themselves can initiate molecular scission [32]. Both kinetic networks are equally suitable for explaining

2,4,4triMe-C7

2,4diMe-C8 C3 + i-C7

2/

10

4,4diMe-C8 2

n-C5 + i-C5

3

2/

3

i-C5 + i-C5

1/

3

2,2,4triMe-C7

3,3,5triMe-C7

n-C4 + i-C6

i-C6 + i-C4

/10

3,5diMe-C8

3,3diMe-C8

/10

The hydrocracking precursors and products of C10 . The boxes contain the probabilities for forming hydrocracking products, assuming that all α, γ - and α, α-dimethyloctanes are available in equal amounts and that there is no preference for hydrocracking. The same was done for α, α, γ -trimethylheptane hydrocracking. Only hydrocracking routes involving at least one tertiary carbocation transition state are included because only these routes are fast enough to make an impact. When there are only secondary carbocation transition states involved (as in monomethylalkane hydrocracking), hydrocracking occurs at a significantly lower rate.

Fig. 2

5.5.3 Computer Simulations of Shape Selectivity Effects

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the hydrocracking product distributions when α, α, γ trimethylalkanes or α, γ - and α, α-dimethylalkanes dominate and the hydrocracking product slates resemble Gaussian or flat histograms [32, 37]. Here, we use the traditional network as the starting point. 5.5.3.2.2 Shape Selectivity: General Principles The original notion of shape selectivity [26] is simply the observation that, if a molecule cannot permeate through the pores of a zeolite, it will not adsorb as a reactant or desorb as a product [6]. If the adsorption of a molecule is inhibited, it will show up intact in the product slate. If the desorption of a molecule is inhibited, it could still form as a reaction intermediate in the adsorbed phase, but only molecules that originate from this reaction intermediate through consecutive reactions will show up in the product distribution. For example, the FAU-type topology exhibits large cavities in which di- and tribranched hydrocarbons form easily, whereas the TON-type topology exhibits much smaller pores in which only the monobranched isomers form easily. As a consequence, if we use the reaction scheme in Fig. 1, we deduce that the product distribution obtained from TON-type zeolites comprises many monobranched (and some dibranched) isomers and their hydrocracking products, whereas the distribution obtained from FAU-type zeolites comprises many dibranched (and some tribranched) isomers and their hydrocracking products [16]. Clearly, differently sized and shaped zeolite pores will interact differently with differently sized reactants, reaction intermediates and products. As a result, the zeolite topology can leave its ‘‘signature’’ on a particular product distribution. Although the original concept of shape selectivity is appealingly simple, it can only explain relatively few product distributions. Therefore, forms of shape selectivity other than the originally proposed reactant and product shape selectivity have been proposed. Examples include transition state, reaction intermediate, and exterior surface shape selectivity. In the following we will provide a short discussion on these five forms of shape selectivity. In particular, the reaction scheme shown in Fig. 1 will be used to illustrate how these various forms of shape selectivity influence product distribution.

A Transition-State Shape Selectivity Transition-state shape selectivity occurs when a zeolite topology influences the reaction rates of the adsorbed molecules by modifying the relative Gibbs free energies of formation of the corresponding transition states [47, 48] (see Fig. 3). It is the only form of shape selectivity that occurs irrespective of the extent of mass transfer limitations between gas and adsorbed phase [14]. In the reaction scheme of Fig. 1, the transition state for hydroisomerization is a

Transition-state (top) and reaction-intermediate shape selectivity (bottom). In transition-state selectivity the zeolite modifies the ease of formation of a reaction intermediate by modifying the ease of generation of the transition state required for its formation; in reaction-intermediate shape selectivity the zeolite modifies the ease of formation of the reaction intermediate directly. When a zeolite impedes formation of a reaction intermediate, the Brønsted–Evans–Polanyi (BEP) principle is likely to apply, so that reaction-intermediate and transition-state shape selectivity occur simultaneously [52]. When a zeolite facilitates formation of a reaction intermediate the BEP principle is likely not to apply [52], so that there can be reaction-intermediate shape selectivity without transition-state shape selectivity.

Fig. 3

dialkylcyclopropyl cation. An example of transition-state selectivity is the inability of TON-type zeolites to form (and hydrocrack) α, α-dimethylalkanes [49–51]. Since the transition states in alkane hydroisomerization occur late in the reaction path, the transition state is sterically similar to the α, α-dimethylalkane products [49–51], and similar van der Waals forces will similarly increase the Gibbs free energy of formation of both the transition states and the products [52]. This linear relationship between the Gibbs free energy of formation of transition states and that of products is an example of the semi-empirical Brønsted–Evans–Polanyi (BEP) relationship [52–54]. According to this relationship, the prohibitively high Gibbs free energy of formation (and adsorption) of α, α-dimethylalkane products inside TON-type zeolites [49–51, 55] is an indication for a similarly high Gibbs free energy of formation of the transition state for α, α-dimethylalkane formation, so that transition-state shape selectivity will inhibit α, α-dimethylalkane product formation inside TON-type zeolite channels, irrespective of the absence or presence of mass transfer limitations [56]. B Reactant Shape Selectivity If any reactant in a feed is too large to permeate a zeolite this reactant may References see page 1690

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5.5 Computer Simulations

Reactant shape selectivity. The zeolite leaves the dibranched alkane intact because this isomer is too large to fully permeate into the zeolite pores. Instead, the zeolite selectively transforms the monobranched isomers that can fully adsorb.

Fig. 4

reach the product slate virtually intact. This would be an extreme form of reactant shape selectivity (see Fig. 4) [57], in which only those reactants undergo catalytic reactions that can fully adsorb so as to form reaction intermediates, convert, and desorb as products. In general, reactant shape selectivity occurs when reactant conversion is inversely proportional to the extent they exhibit mass transfer limitations. Accordingly, the reactants that adsorb with the smallest mass transfer limitation will be converted most and will be the least prevalent in the product slate. An example of reactant shape selectivity is the selective combustion of exclusively the linear isomers from a mixture of branched and linear butanes and butenes on Pt-loaded, Ca,Na-exchanged LTA-type zeolite [57]. Whereas, these early experiments involved reactants that are categorically excluded from the LTA-type zeolite pores [57], later experiments included reactants that are excluded by different degrees and, therefore, focused at differences in mass transfer rates as a dominant cause for reactant shape selectivity [58]. C Product Shape Selectivity If reaction intermediates are too large to desorb intact from a zeolite, only their consecutive reaction products can end up in the product slate. This would be an extreme form of product shape selectivity (Fig. 5). In general, product shape selectivity occurs when some reaction intermediates exhibit higher mass transfer limitations than others, so that they remain in the adsorbed phase and continue to react for a longer period of time than other, less mass transfer-limited reaction intermediates. Accordingly, the products that desorb with the smallest mass transfer limitation will be the most prevalent in the product slate. A typical example of this form of shape selectivity is the cracking of hexane isomers in Ca-exchanged LTAtype zeolites [6]. This process yields only linear and no branched cracking products, because only the former can desorb from the zeolite pores. D Exterior Surface Shape Selectivity In some instances the exterior surface of zeolites process reactants that are

Product shape selectivity. The zeolite does not contribute any geminal dibranched isomers to the product slate because these isomers react much more rapidly than they desorb. Therefore, geminal dibranched isomers cannot leave the zeolite intact. Instead, the zeolite selectively yields products that can desorb rapidly.

Fig. 5

either too large to adsorb completely [26] or diffuse too slowly [11–13] to fully permeate the adsorbate. Whether the exterior zeolite surface has a sufficiently regular structure to yield product distributions different from amorphous aluminosilicates remains a subject of debate. In reflecting the lack of agreement on the relevance of the exterior surface to shape selective catalysis, the process has been given a plethora of names, including pore mouth catalysis [11], key-lock mechanism [19, 20], nesteffect [5, 14], and exterior surface shape selectivity [14]. E Other Forms of Shape Selectivity In the previous sections we have listed some of the published mechanisms for shape selectivity. Within the literature, many other types of shape selectivity can be identified, including the ‘‘concentration’’ or ‘‘solvent effect’’ [59–63], the ‘‘confinement’’ or ‘‘solvent effect’’ [64–67], molecular traffic control [28, 29], secondary shape selectivity [27], inverse shape selectivity [32], and the ‘‘cage’’ or ‘‘window effect’’ [23, 24]. Molecular Simulations At this point it is important to recall that the mechanisms described in the previous section are based on the product distribution observed in the desorbed (often gas) phase. Comparatively little is known about the thermodynamics and diffusion properties of the molecules in the adsorbed phase, even though these molecules determine the product distribution observed in the desorbed phase. Molecular simulation provides an attractive alternative to compute these missing properties. A short review on the main simulation techniques is provided here, but for a more extensive description the reader is referred elsewhere [68]. 5.5.3.3

5.5.3 Computer Simulations of Shape Selectivity Effects

5.5.3.3.1 Simulation Techniques The idea of a molecular simulation is simple – that is, to provide a model in the form of an intermolecular potential that describes the interaction between the molecules adsorbed in the zeolite. This model provides the input for a molecular dynamics simulation or a Monte Carlo simulation from which the corresponding thermodynamic and/or transport properties are obtained. For the successful application of these techniques it is important that the intermolecular potentials give a sufficiently realistic description of the experimental systems, and that the simulation techniques are sufficiently powerful such that, for the molecules of interest, accurate properties can be computed within a reasonable amount of CPU time. As will be demonstrated below, both issues are non-trivial.

A Molecular Dynamics of Adsorbed Molecules Suppose we use a molecular dynamics simulation for a system of N particles for which we solve Newton’s equations of motion [69]. If there are no external forces working on the system, the total energy is conserved in a molecular dynamics simulation. We therefore perform a simulation in which the energy, E, number of particles, N , and the volume, V , are imposed. A molecular dynamics simulation therefore samples the microcanonical ensemble [68]. Ideally, the real system should be mimicked as well as possible. If successful, the experimental data would also be reproduced, including the diffusion coefficients. Hence it would also be observed that, for the long-chain hydrocarbons that are of interest to hydrocracking, the diffusion coefficient is so small that the corresponding simulations to obtain reliable statistics become prohibitively long. As a consequence, standard molecular dynamics can only be used for those systems that diffuse sufficiently fast that accurate thermodynamic or transport properties can be obtained. B Monte Carlo Simulation of Adsorbed Molecules Let us consider the experimental set-up required to measure adsorption isotherms. The aim is to measure the number of adsorbed molecules as a function of the pressure of the gas or liquid that is in contact with the zeolite. Experimentally, the most common system is a zeolite in a container that contains a gas or liquid which is maintained at a constant temperature and pressure (or partial pressure in the case of a mixture). At equilibrium, the adsorbed gas molecules have the same temperature and chemical potential as the molecules in the container. The container can be seen as a reservoir that fixes the temperature and chemical potentials of the adsorbed components. In principle, the experimental set-up could be mimicked by simulating a gas or liquid in contact with a reservoir.

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However, such a system is not very convenient from a simulation point of view. This experimental set-up closely resembles the grandcanonical ensemble [68] wherein the temperature, volume, and chemical potentials are imposed. An important advantage of the Monte Carlo technique is that a simulation in the grand-canonical ensemble can be performed. In this situation, the reservoir and zeolite are not in direct physical contact, but the Monte Carlo procedure guarantees that the adsorbed molecules have an equal temperature and chemical potential in both the reservoir and in the zeolite. Hence, periodic boundary conditions can be used for the entire zeolite, and the presence of an interface is avoided. The input of the simulation is temperature and chemical potential of the molecules in the reservoir, and the average number of adsorbed molecules is a result of the simulation. The grand canonical Monte Carlo technique functions best if the acceptance of trial moves by which particles are added or removed is not too low. For atomic fluids, this condition effectively limits the maximum loading in a zeolite at which the method can be used. Although the grand canonical Monte Carlo technique can be applied to simple models of non-spherical molecules, special techniques are required as the method converges very poorly for all but the smallest polyatomic molecules. Although both molecular dynamics and Monte Carlo are very efficient for atoms or small molecules, for large molecules both methods require significant amounts of CPU time. For example, June et al. [70] studied the relaxation of n-butane and n-hexane in MFI using molecular dynamics, and concluded that the zeolite slowed down the relaxation of these molecules by several orders of magnitude – the longer the chains, the slower the relaxation. Hence, the CPU requirements increase significantly for MD simulations of these long-chain alkanes. The diffusion coefficients of linear alkanes in MFI are sufficiently high that these can be simulated using MD [71], but for the mono-branched alkanes MD can only be used at very high temperatures [72, 73]. Branched alkanes in MFI preferentially adsorb in the intersections between the zig-zag and straight channels [49], and the diffusion is therefore an activated process in which the molecule jumps from one intersection to another [74]. This very slow diffusion path could be avoided via a Monte Carlo simulation in which a new configuration is generated at a random position in the zeolite. The probability that such a move will be accepted depends on the energy difference between the new and the old configurations. Clearly, if a new position is generated on top of a zeolite atom, the attempt will be rejected. For a chain molecule this implies that none of the atoms References see page 1690

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5.5 Computer Simulations

should overlap with the zeolite atoms. If for methane this probability is 1 out of 1000 attempts, then for ethane this will be of the order 1 out of 106 , and for n-octane 1 out of 1024 . The conventional Monte Carlo method is therefore very efficient for noble gases or small molecules, but for long-chain alkanes it is equally inefficient as molecular dynamics. C Configurational-Bias Monte Carlo Technique The configurational-bias Monte Carlo (CBMC) technique has been developed to make possible the insertion of longchain molecules in moderately dense liquids. The original configurational-bias Monte Carlo technique has been developed for lattice models [75, 76] and has been extended to continuous models [77]. Here, we show how this method can be used to simulate the adsorption of long-chain hydrocarbons in zeolites. In a CBMC simulation the molecules are not inserted at random but rather are grown atom by atom. This growing process introduces a bias which can be removed exactly by adjusting the acceptance rule [68]. In a CBMC simulation a molecule is grown atom by atom using a method based on an algorithm developed by Rosenbluth and Rosenbluth [78]. In this growing scheme, overlap is avoided with the zeolite atoms, and the corresponding bias is removed exactly by adjusting the acceptance rules [68]. The basic CBMC scheme has been extended to branched molecules [49, 79, 80], cyclic molecules [81, 82], and all-atoms models which explicitly include the hydrogen atom [83, 84]. Several ‘‘tricks’’ have been devised to increase the efficiency of a CBMC simulation [85]. Compared to ordinary Monte Carlo simulations, CBMC can be up to 10 orders of magnitude more efficient, which has made these types of simulation for long-chain hydrocarbons possible. In addition, as the growing step in a CBMC scheme provides information on the free energy of adding a molecule to the system, an important application of CBMC is to compute the free energy of a molecule inside the pores of a zeolite. D Rare Event Simulations A very small diffusion coefficient often is the result of molecules which are trapped in low (free) energy sites and once in a while ‘‘hop’’ from one to another adsorption site. In order to compute a diffusion coefficient reliably, a sufficient number of hops must be observed. Most of the CPU time is, however, spent on molecules that ‘‘wait’’ close at an adsorption site until a fluctuation gives them sufficient kinetic energy to take the barrier between adsorption sites. The higher the barrier, the longer the molecules remain trapped and – on the time scale of a molecular dynamics simulation – such hopping becomes a very rare event.

Special techniques have been developed to simulate such rare events [68]. The basic idea is to compute the hopping rate in two steps [86, 87]. First, the probability that a molecule can be found on top of the barrier is computed. This calculation is followed by a separate simulation in which the probability is computed that a molecule that starts on top of the barrier ends up in the next adsorption site and does not recross the barrier. The probability of finding a molecule on top of the barrier can be computed directly from the free energy profile, which is the free energy as a function of the position of the molecule in the zeolite. The second step involves the average time that it takes a molecule to cross the barrier. The simplest approach is to assume that transition state theory (TST) holds. A molecule arriving at the top of the barrier is assumed to be in equilibrium with its surroundings, and as a consequence the velocity distribution is given by the Maxwell distribution corresponding to the temperature of the system. TST assumes that half of the molecules that reach the barrier also cross the barrier. TST theory ignores the possibility that such a particle recrosses the barrier and returns into the cage from which it has originated; this may, for example, be due to collisions with the zeolite atoms. This recrossing can be computed directly from a molecular dynamics simulation in which the molecules start on top of the barrier, and the recrossing probability is computed directly. As this involves a simulation that begins on top of the barrier, it is much faster than simulating the time it takes a molecule to climb the free energy barrier. These rare event methods have been applied to zeolites at low [26, 74, 88–90] and even at high loadings [91]. Intermolecular Potentials Most simulation studies follow the assumptions pioneered by Kiselev and coworkers [92] for the adsorption of non-polar molecules. The zeolite is assumed to be rigid and purely siliceous. The adsorbate–zeolite interactions are dominated by the dispersive interactions with the oxygen atoms of the zeolite. The smaller silicon atoms contribute little to the dispersive interaction, and are taken into account implicitly via the oxygen atoms. Further refinements involve the use of a flexible lattice, or the effect of charges for systems that involve polar or Coulombic interactions [93].

5.5.3.3.2

A Zeolite–Zeolite Interactions The starting point here is the crystal structure of the zeolite. (Details of most of these structures can be found on the Internet [38].) In case of a rigid zeolite structure, the crystal structure can be used to generate a simulation box containing the desired number of unit cells of the zeolite crystal. Computer

5.5.3 Computer Simulations of Shape Selectivity Effects

packages are available to perform this procedure both quickly and conveniently [94]. If a rigid lattice is assumed, there is no need for a model of the zeolite–zeolite interactions. In case a flexible zeolite is essential, a variety of models that describe the zeolite–zeolite interactions have been published, the accuracy of which can partly be assessed via a comparison of the calculated vibrational infrared (IR) spectra with the experimental spectra. These models have been discussed in detail by Demontis and Suffritti [95], and the interested reader is referred to this review for details on these models and further references. In order to limit the CPU requirements of a fully flexible zeolite, methods have been developed in which the normal vibrational modes and harmonic crystal approximation are used [96]. B Adsorbate–Adsorbate Interactions For the simulation of hydrocarbons, a variety of models has been proposed, the most realistic being all-atom models in which both the carbon and hydrogen atoms are considered explicitly. In united-atom models, the CH3 , CH2 , or CH groups are considered as a single atom. From a computational point of view, the united-atom model is more efficient and has fewer parameters to be determined; consequently, most studies use a united-atom model. A comparison of the results of such simulations with experimental data for the adsorption and diffusion shows a satisfactory description of the experimental data. In addition, the scatter in these currently available experimental data makes it very difficult to prove that an all-atom model is essential. For adsorbate–adsorbate interactions it is convenient to distinguish between the intramolecular and intermolecular interactions. The former group are very important to arrive at a realistic representation of the conformation of the adsorbate molecules. Fortunately, these potentials can be based on quantum chemical or spectroscopic data, and therefore for most molecules these models provide a sufficiently accurate description of the intramolecular interactions. In addition, comparisons of various models of, for example, the torsion or bond-bending, show little influence on the thermodynamic properties, such as the vapor-liquid curve [97]. For the hydrocarbon special force fields, both united-atom and all-atom models have been developed that provide an accurate description of the entire vapor–liquid coexistence curve [80, 81, 84, 97–103]. C Zeolite–Adsorbate Interactions Often, zeolite– adsorbate interaction parameters are obtained from fitting to experimental data, and therefore these parameters depend on which accuracy and type of data are used in the fitting procedure. For example, the parameters obtained from fitting to diffusion coefficients can be different from

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those fitted to, for example, the heats of adsorption or Henry coefficients. Beerdsen et al. have shown that a very accurate fitting of the parameters of a united-atom model can be obtained through fitting on experimental isotherms with inflection points [91]. This procedure uniquely determined the adsorbent–adsorbate interaction parameters, and is very sensitive to the size parameter as the inflection points in the isotherms are often related to a subtle interplay between different adsorption sites. Several alternative force fields have been developed that also provide an accurate description of the experimental adsorption isotherms [104, 105]. Free Energy Model In this section, it will be shown that molecular simulations afford computation of the thermodynamic and diffusion properties of the reaction intermediates in the adsorbed phase. Knowledge of these properties of the relevant reaction intermediates affords a reinvestigation of the basic assumptions for kinetic models that had been in vogue for a long time. 5.5.3.4

5.5.3.4.1 Introduction In the most general case, a feed mixture of many reactants must be considered. The first step is that these molecules must adsorb; indeed, for a given partial pressure of the various components, the concentration of the molecules depends on the free energies of adsorption. In the absence of diffusion limitations, a relatively low free energy of adsorption of a particular feed component implies a high adsorbed phase concentration, which often indicates a high reactivity of this reactant. Once a reactant has been adsorbed and has – by definition – formed a reaction intermediate, it can undergo a sequence of reactions forming various other reaction intermediates. The relative importance of these reaction intermediates depends on their free energy of formation. Depending on their relative free energy of formation and adsorption, and depending also on the presence of kinetic barriers to desorption, reaction intermediates will either desorb as a product or will continue to transform into other reaction intermediates. In the free energy model the free energies of adsorption and free energy of formation play an essential role. The free energies of formation of most molecules are tabulated for the gas phase. If molecules are formed in a zeolite, a contribution of the zeolite must be added to this free energy of formation, and this correction corresponds to the free energy needed to bring a molecule from the gas phase into the zeolite. Both the free energy of formation in References see page 1690

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the adsorbed phase and the free energy of adsorption must be known at reaction conditions. This implies that reaction intermediates do not approach gas-phase thermodynamic equilibrium, but rather adsorbed-phase thermodynamic equilibrium. As was observed using in-situ NMR [106], free energy contributions by the zeolite can cause major shifts to the gas-phase equilibrium distribution. At this point it should be emphasized that, in practice, thermodynamic equilibrium among the reaction intermediates often will not be reached, and other (kinetic) factors contribute to the product distribution. Therefore, this free energy model will provide an idealized reference, which can be a useful starting point to further investigate the details of a product distribution. 5.5.3.4.2 Conventional Shape Selectivity The traditional definition of shape selectivity is related to the observation that reactants which are too large to fit inside the zeolite pores do not go on to form products, that transition states which are too large to fit inside the pores of zeolites do not form, and that reaction intermediates that are too large to desorb intact continue to undergo consecutive reactions [14]. In the free energy model, a high and positive contribution of a zeolite to the free energy of formation is indicative of a bad fit of a transition state or reaction intermediate. Transition states that fit badly exhibit a high free energy of formation, and therefore will contribute little to the product slate. Similarly, reaction intermediates that fit badly exhibit

a high free energy of formation and will therefore maintain a low concentration inside the pores. If a minimal concentration of a reaction intermediate entails a minimal contribution to the desorbed phase depends on the reactivity of the thermodynamically impeded species. Thus, contributions from the highly reactive trimethylalkanes can dominate the hydrocracking product slate [44], even though they have a relatively high free energy of formation [26]. A comparison of n-decane hydroconversion on FAUand TON-type zeolites provides a good illustration of the relevance to conventional shape selectivity of the contributions of the zeolite to free energy of formation of reaction intermediates in the adsorbed phase. Figure 6 shows these contributions for FAU- and TON-type zeolites to selected reaction intermediates in the hydroconversion reaction of n-C10 , as discussed in Fig. 1. In FAU, the contribution of the zeolite is relatively small, as the cages of FAU are sufficiently large to accommodate all reaction isomers. For this system gas-phase thermodynamic data are therefore a good approximation. The free energies of formation of alkanes with the same degree of branching are similar, and the resulting ideal Gaussian product distribution is simply determined by the number of pathways that generate a particular reaction intermediate; the product slates obtained on amorphous aluminosilicates [34, 35] and on crystalline aluminosilicates with a FAU-type topology [36, 37] are therefore virtually identical. The TON-type pores are

40 FAU

30

20

10

0

−10

∆Gi - ∆Gn-decane/[KJ mol−1]

∆Gi - ∆Gn-decane/[KJ mol−1]

40

30

123.6

95.8

TON

20

10

0

−10

3,3,5-trimethylheptane 4,4-dimethyloctane 2,4-dimethyloctane 5-methylnonane 2-methylnonane

Contributions of the zeolite to the free energy of formation of typical reaction intermediates in the hydroconversion of n-decane. The figure illustrates the free energy difference between the reaction isomer and n-decane at T = 415 K in the zeolite FAU (left) and TON (right).

Fig. 6

5.5.3 Computer Simulations of Shape Selectivity Effects

much smaller, and the free energy calculations clearly show that this topology inhibits the formation of isomers with proximate branches. 5.5.3.4.3 Transition-State and Reaction-Intermediate Shape Selectivity Molecular simulations clearly show that the free energy of formation of alkane isomers with proximate branches are prohibitively high inside TON-type zeolites (Fig. 6), and that – therefore – these isomers will not form in the adsorbed state. Traditionally, the inhibition of formation of these isomers has been attributed to a prohibitively high free energy of formation of the transition state preceding formation of these adsorbed isomers; that is, it has been attributed to kinetic instead of thermodynamic factors [49]. In acid zeolite chemistry this transition state is usually associated with some type of carbocation. Transition-state selectivity can be defined as a change of free energy of this carbocationic transition state relative to a reaction intermediate as induced by the zeolite framework. Van der Waals interactions between a transition state and the zeolite framework can decelerate the reaction by increasing the free energy of formation of transition states that are incommensurate with the particular zeolite topology [47, 48, 58, 107, 108]. Alternatively, ionic interactions between a transition state and the zeolite framework can accelerate the reaction by decreasing the free energy of formation of transition states that are commensurate with the particular zeolite topology [52]. In order to quantify these effects it is necessary to perform a detailed quantum chemical calculation in the pores of the zeolite, and also to determine the transition state in the pores. These are often very time-consuming calculations and are usually limited to a small region of the zeolite; hence, only a few studies have been published that take the full zeolite structure into account [52]. However, methods are being developed to integrate these quantum chemistry calculations using embedded methods with force field-based methods. At present, it is not yet feasible to perform a full quantum-chemical calculation to determine unambiguously the free energy of the transition state [109]. However, for many components the BEP relationship holds, which states that the activation energy and reaction energy are related linearly [52–54, 56, 110–112]. Hence, if the zeolite increases the relative free energy of one of the reaction intermediates, the corresponding transition state increases similarly. The consequence of this relationship is that differences in the free energies of the transition states of competing reactions can be estimated from the differences in free energy of formation of the corresponding reaction intermediates (see Fig. 7). If it is assumed that the semi-empirical BEP relationship holds for the

1685

Gas-phase

Inside a zeolite The Brønsted–Evans–Polanyi (BEP) relationship. In the gas phase the free energies of formation of two products are virtually the same, and so are the free energies of formation of the transition state. In the gas phase, the zeolite increases the free energy of formation of one product relative to that of the other product. According to the semi-empirical BEP relationship, this increases concomitantly the free energy of formation of the transition state for this product – the alkane isomer with geminal methyl groups in this example.

Fig. 7

entire reaction scheme, it can be deduced which transition states have an increased free energy of formation in the adsorbed state from the computed free energies of the reaction intermediates. A recent suggestion is that the BEP relationship might not hold for instances where adsorbent–adsorbate van der Waals interactions decrease the reaction energy in the adsorbed phase [52]. This would suggest that evaluating the acceleration of reactions due to the zeolite topology-induced facilitated formation of a transition state requires a full-fledged quantum chemical evaluation. Even if the BEP relationship were always to hold, contributions by the zeolite to either the free energy of formation of the transition state or to the free energy of formation of adsorbed reaction intermediates will result in two distinct forms of shape selectivity [113]. Transition-state shape selectivity will occur irrespective of the presence of diffusion limitations [58], whereas reaction-intermediate shape selectivity will only occur when a reaction is diffusion-limited [114]. Alkane hydroconversion in MFI-type pores provides an illustration of reaction-intermediate shape selectivity. Molecular simulations have shown (see Fig. 8) a large and positive contribution of these zeolites to the free energy of formation of α, α, γ -trimethylalkanes [56]. By comparison, these materials impede the adsorption and formation of α, γ -dimethylalkanes to only a small extent [56]. They do not impede the formation of α, αdimethylalkanes, monomethylalkanes, and n-alkanes, as the shape of these isomers is commensurate with that of the MFI-type intersections, so that all have a similar Gibbs free energy of formation in the adsorbed References see page 1690

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40

30

MFI

20

10

0

−10

∆Gi-∆Gn-decane/[kJ mol−1]

∆Gi-∆Gn-decane/[kJ mol−1]

40

30

20

MEL 3,3,5-trimethylheptane 4,4-dimethyloctane 2,4-dimethyloctane 5-methylnonane 2-methylnonane

10

0 −10

Comparison of MFI and MEL types of pore. Top: artist’s impression of the structures. Bottom: free energies of formation of selected reaction intermediates relative to that of adsorbed n-decane.

Fig. 8

phase [56]. Accordingly, these free energies indicate that the consecutive hydroisomerization of n-alkanes into monomethylalkanes and α, α-dimethylalkanes dominates the shape selectivity in MFI-type pores. Since the favored α, α-dimethylalkanes are hydrocracking precursors which have a low diffusion coefficient, the net result would be that the shape selectivity inside MFI-type pores enhances the hydrocracking rate at the cost of the hydroisomerization rate. Now, we can compare the product distributions obtained from n-decane conversion in MFI- and MELtype pores. Figure 8 shows that the structures of these two zeolites are very similar; the main difference is that MFI has both sinusoidal and straight channels, whereas MEL has only straight channels. Despite these similarities, the n-decane hydrocracking product distribution is very different; the iso-butane yield of MEL is twice that of MFI-type zeolites [56]. Figure 8 shows also that the computed free energies of formation for α, α, γ -trimethylalkanes are highly repulsive, to the extent that their Gibbs free energy of formation in the adsorbed phase effectively prohibits the formation of these hydrocracking precursors in both zeolites. In the absence of α, α, γ -trimethylheptanes, α, α- and α, γ -dimethyloctanes are the most likely hydrocracking precursors [56]. Interestingly, the free energy calculation indicates that α, α-dimethyloctanes have a relatively low free energy at MFI-type intersections, suggesting that these intermediates are commensurate with the shape of the MFI-type intersections, whereas α, γ -dimethyloctanes

are commensurate with the shape of one of the MEL-type intersections. Thus, 4,4-dimethyloctane fits snugly when it has its octane backbone in the straight MFI-type channel and the two methyl-groups in the zigzag channel, while 2,4-dimethyloctane has a perfect fit in the large MEL-type intersection because the distance between the two branches matches the distance between the two intersecting channels. The commensurate isomers have the lowest free energy of formation [56]. Due to the relatively large zeolite crystals and high acid site densities used [29, 115], alkanes with the lowest free energy of formation are preferentially formed but cannot diffuse out of the zeolite without being hydrocracked [56]. Since the hydrocracking of 2,4-dimethyloctane yields isobutane, while hydrocracking of 4,4-dimethyloctane yields n-butane, the preferential formation of commensurate isomers suggests an explanation for the twice as high iso-butane yield of MEL- as compared to MFI-type zeolites [56]. This example nicely illustrates the concept of reactionintermediate shape selectivity; the zeolites preferentially form reaction intermediates that have a low free energy of formation. In the case of MEL- and MFI-type zeolites the important reaction intermediates are those that are commensurate with the zeolite structure and therefore have an unusually low free energy. They can only form at the intersections, for they have a very large free energy at any other position. As a consequence, these molecules cannot diffuse in the pores of the zeolite and are a typical example of ‘‘ship-in-the-bottle’’ molecules. So far, these

5.5.3 Computer Simulations of Shape Selectivity Effects

reactions have only been studied for severely diffusionlimited MFI- and MEL-type zeolite crystals. As an example of reaction-intermediate shape selectivity induced by adsorbate–adsorbate intermolecular interactions, we can consider the effect of pore size on the hydroisomerization selectivity of the C6 hydrocracking products formed during n-hexadecane hydroconversion. Figure 9 shows that the ratio between the 2,3-dimethylbutane and n-hexane yields as a function of pore size forms a bell-shaped curve [26, 30–32, 116]. These results indicate that there exists an optimal pore diameter for the formation of branched alkanes; this phenomenon is often referred to as inverse shape selectivity [32], whereby the zeolite is favoring – rather than inhibiting – the formation of the bulkiest isomers, the dibranched alkanes. This phenomenon was explained in terms of an optimal fit of the branched alkene reaction intermediate with the zeolite pores. However, at low pressures, recent simulations do not reproduce the optimum fit, as was originally reported [114, 116]. At intermediate pressures [117, 118], recent experiments have not reproduced the height of the optimum. Rather, only at high pressures, where adsorbate–adsorbate interactions are important, could such an optimum be reproduced quantitatively [116].

1.2

Normalized yield ratio

1.0 0.8 0.6 0.4 0.2

FAU

LTL

DON

MAZ

AFI

MOR

SSZ-31

BEA

MTW

0.0

The effect of pore size on the yield ratio of 2,3-dimethylbutane and n-hexane during n-C16 hydrocracking (increasing pore size from MTW to FAU). So as to facilitate a comparison between simulated (left bar), experimental adsorption ratios (middle bar) and hydroconversion yield ratios (right bar), all ratios were divided by the values obtained for AFI. When the pores are small (as with MTW-type zeolites), repulsive adsorbent–adsorbate van der Waals interactions impede dimethylbutane (DMB) formation; when the pores increase in size these impeding interactions disappear and inter-adsorbent interactions favor formation of the better-packing DMB; when the pore size increases above the 0.74 nm, differences in packing efficiencies disappear because the adsorbents no longer have to line up head-to-tail but can pack in an increasingly more random, liquid-like fashion.

Fig. 9

1687

If the pores are too narrow for the bulky dibranched alkane to fit, this will be reflected in a high and positive contribution of the zeolite to the free energy of formation of 2,3-dimethylbutane relative to n-hexane. When the pore size increases towards an optimum size, more 2,3-dimethylbutane compared to n-hexane can fit in the tubular channels as the dibranched molecule is more compact. Because of these differences in effective size at sufficiently high pressure, the more compact molecule has the lower free energy [26, 116, 119]. This size entropy effect is also responsible for differences in adsorption behavior of these isomers. Figure 9 shows that the tubular MAZ- and AFI-type pores share this optimum size for adsorbing and forming 2,3-dimethylbutane instead of n-hexane [116]. When the pores are still larger, the molecules no longer stack linearly so that the adsorbed phase approaches a liquid phase and the entropic size effect vanishes. This difference in packing efficiency leaves its mark on the product slate, because slowly diffusing n-C16 locks up the initial C6 hydrocracking products sufficiently long to have them approach adsorbed phase chemical equilibrium. Once desorbed, C6 is unlikely to compete with C16 for adsorption in the zeolite, so that no chemical equilibration towards the gas phase will occur. Thus, severe mass transfer limitations between the gas and adsorbed phases are a prerequisite for this type of reaction-intermediate shape selectivity to occur. In the absence of mass transfer limitations, FAU-, MAZ-, and MOR-type zeolites yield a virtually identical C6 isomer product slate [120]. Whether the optimum branched isomer yield reported for MAZ- and AFI-type zeolites is indeed only a result of reaction-intermediate shape selectivity remains the subject of debate. Recently, it was suggested that transitionstate shape selectivity might also contribute [118]. If there is indeed a contribution, it must be small as an optimized MAZ-type zeolite yields the same yield ratio of 2,3-dimethylbutane and n-hexane in the hydrocracking of n-hexadecane as do FAU- or MOR-type zeolites [120]. In addition, it was found that 2,3-dimethylbutane diffuses faster than other hexane isomers in MOR- as compared to FAU-type zeolites [121]. If this faster diffusion at high loading for the more efficiently packing molecule can be extrapolated to all topologies mentioned in Fig. 9, then product shape selectivity might contribute to the observed shape selectivity. 5.5.3.4.4 Reactant Shape Selectivity Zeolites can shapeselectively process more of one reactant than of another reactant because: (i) the former has a higher diffusion rate; or (ii) the former has a lower free energy of adsorption. References see page 1690

5.5 Computer Simulations

In the former instance of reactant shape selectivity, the reaction needs to be adsorption rate-limited before reactant shape selectivity occurs [58], whereas in the latter instance it is not a prerequisite for shape selectivity. When differences in free energy of adsorption determine the shape selectivity, they do so by affecting the relative concentration of different molecules inside a zeolite. These can be remarkably different from that outside the zeolite. A lower free energy of adsorption is a measure of a higher concentration inside the zeolite, and a higher reactivity. Experimentally, insights into reactant shape selectivity due to differences in free energy of adsorption have been obtained in studies on the chain lengthdependence of the reactivity of n-alkanes [26, 122, 123], so that the discrimination of zeolites between n-alkanes of various lengths depends on the pore topology. To the extent that the reactivity of n-alkanes as a function of chain length varies with zeolite topology, it is – by definition [14] – an example of (reactant) shape selectivity. At low pressure and loading the Henry coefficient is directly proportional to the free energy of adsorption [123, 124]. The Henry coefficient is a measure of the pressure required to adsorb a given amount of molecules in the pores of the zeolite. Most zeolites (such as, FAU- [19, 42], OFF- [125] and MFI- [126] types [122, 127]) are similar to amorphous aluminosilicates, in that the Henry coefficient increases monotonically with the chain length [123]. Hence, a monotonic increase in Henry coefficient with chain length implies a monotonic increase in reactant concentration with chain length. As a result, in the adsorbed phase, the reactivity of n-alkanes in hydroconversion increases monotonically with chain length [19]. Accordingly, the products originating from the longer n-alkanes dominate the product slate. This selectivity for processing longer alkanes is an example of shape selectivity only in as much as it depends unambiguously on the zeolite pore topology. Figure 10 shows that, for some zeolites, the Henry coefficient decreases as a function of chain length. It is interesting to see how such chain length dependence relates to the chain length-dependence of the n-alkane hydroconversion rate of ERI. The ERI-type pore topology exhibits small (diameter ≈0.4 nm) openings (or ‘‘windows’’) providing access to somewhat larger cages [128]. Only from n-C4 to n-C6 ERI-type zeolites exhibit the usual increase in reactivity due to a combination of a lower free energy of adsorption and a higher intrinsic reactivity of the n-alkane with chain length [19, 129]. For n-alkanes longer than n-C6 , hydrocracking diffusion limitations set in [25] and the reactivity becomes increasingly abnormal. Figure 11 shows the chain length dependence of the diffusion coefficient of the n-alkanes in ERI. It is apparent that

Henry Coefficient KH / (mmol/g/Pa)

1688

10−2 10−3 10−4 10−5 10−6 10−7 10−8 10−9 10−10 10−11 10−12 10−13 10−14

ERI CHA LTA OFF

2

4

6

8 10 12 14 16 18 20 22 24 26 28

Carbon number of n-alkane Henry coefficient as a function of n-alkane chain length in various zeolites. When the window size is above a certain diameter (as in OFF-type zeolites), attractive van der Waals interactions between the pore wall and the adsorbed alkane decrease the adsorption enthalpy more than they increase the adsorption entropy; hence, the net effect is a linear decrease in free energy of adsorption (and an exponential increase in Henry coefficient) with increasing n-alkane chain length. When the window size is below a certain diameter (as in ERI-type zeolites), repulsive van der Waals interactions between the window and the adsorbed alkane increase the adsorption entropy more than they decrease the adsorption enthalpy, so that the net effect is a linear increase in free energy of adsorption (and an exponential decrease in the Henry coefficient) with increasing n-alkane chain length.

Fig. 10

diffusion limitations are severe, for they reduce the reactivity from n-C6 to n-C8 [25], and thereby offset both an increase in intrinsic reactivity [19, 26] and a decrease in free energy of adsorption [26]. A similar reduction in reactivity has been reported only once for FER-type zeolites [130]. Surprisingly, the monotonic decrease in reactivity with n-alkane chain length in ERI-type zeolites is interrupted at n-C10 , for n-C10 is more reactive than n-C8 [25]. Figure 11 shows that this increase correlates well with an increase in diffusion rate [26], suggesting that the strong diffusion limitations remain the dominant cause for changes in reactivity from n-C6 to n-C10 . Interestingly, molecular simulations fully support the traditional model, which postulates that diffusion rates are the cause for both the decrease and the increase in reactivity. For n-alkanes longer than n-C10 , the reactivity as a function of n-alkane chain length changes direction again and now decreases monotonically with the n-alkane chain length. This cannot be related exclusively to changes in the diffusion rates, for both experimental and simulated diffusion data indicate that the diffusion rates are still increasing with increasing n-alkane chain length and do not peak before n-C12 –n-C13 [26]. Only when considering both the simulated non-monotonic variation in Henry

5.5.3 Computer Simulations of Shape Selectivity Effects

Exp Sim.

D / (nm 2/s)

109

106

103

100

2

4

6

8

10

12

14

16

Carbon number of n -alkanes

Diffusion coefficients (experiments: gray = experimental results [24]; black = simulations) in ERI as a function of chain length at 600 K. Due to the high density of windows with a repulsive van der Waals energy between the n-alkane and the window in ERI-type zeolites, the diffusion rates in the ERI-type zeolites are several orders of magnitude smaller than those in MFI-type.

Fig. 11

coefficient and the simulated non-monotonic variation in diffusion coefficient with n-alkane chain length can the experimentally reported monotonic decrease in reactivity with n-alkane chain length from n-C10 to n-C16 be reproduced [123]. A fully quantitative comparison is difficult due to the likely onset of catalysis at the exterior surface of the ERI-type crystals [123]. 5.5.3.4.5 Product Shape Selectivity Zeolites can, in a shape-selective manner, yield more of one product than of another product because: (i) the former has a higher diffusion rate; or (ii) the former has a higher free energy of adsorption. In (i), the reaction needs to be desorption ratelimited before product shape selectivity occurs [58], while in (ii) desorption rate limitation is not a prerequisite for shape selectivity. In both cases the zeolite yields more of a particular product because such a product desorbs faster than other products and therefore escapes the reaction cycle at a relatively early stage. An example of the preferential formation of products that combine the highest free energy of adsorption with the lowest free energy barrier to diffusion (product shape selectivity) is the product isomer distribution observed in n-alkane hydroconversion on TON-type zeolites. Recent experimental [42] and simulated [49–52, 73, 131, 132] adsorption data are in agreement that both linear and monobranched alkanes can fully adsorb into TONtype zeolite pores. Experimental [55] and simulation [49] methods also agree that dibranched alkanes with geminal dimethyl groups cannot adsorb in TON-type zeolite pores. By extension, dimethylalkanes should also adsorb into TON-type pores, provided that the methyl groups are far enough apart. However, closer scrutiny reveals that

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dimethylalkanes can be divided into two groups: those that are commensurate with the periodicity of the TONtype zeolite wall; and those that are incommensurate. The commensurate dimethylalkanes combine a low free energy of adsorption with a high free energy barrier to diffusion, whereas the incommensurate molecules combine a high free energy of adsorption with a low free energy barrier to diffusion [56]. Therefore, only the latter type of dimethylalkanes are found in the hydrocracking product slate [56, 133]. However, this nice example of the importance of the Frenkel–Kontorowa effect to catalysis remains the subject of debate [22, 56]. Concluding Remarks The results of these studies have demonstrated the importance of a proper knowledge of the thermodynamic and transport properties of molecules adsorbed in zeolites at reaction conditions. As experiments at these conditions are very difficult to perform, molecular simulations offer an attractive alternative. In particular, as recent developments in novel simulation techniques and force field have allowed for computing these properties to a sufficient degree of accuracy, they may be considered as a good alternative for real experimental data. It is proposed that shape selectivity be analyzed using a simple concept based on the contribution of the zeolite on the free energies of formation – that is, the free energy of transferring a molecule from the gas phase into the zeolite. The lower this contribution to the free energy, the greater the probability that these molecules are formed as reaction intermediates in the pores of the zeolite. Depending on the diffusion coefficient, these reaction intermediates may leave the zeolite to become products, or they may continue to react. In a similar context, understanding the contribution of the zeolite to the free energies of adsorption of the reactants, provides us with direct information on the contribution of the zeolites to the reaction rates. We show, by re-examining some well-known examples of various forms of shape selectivity that, once the contributions of the zeolites to the relative free energies of the reaction intermediates are known, a very different explanation of shape selectivity becomes apparent compared to that observed experimentally. 5.5.3.5

Acknowledgments

As this chapter summarizes the work of many, the authors are especially indebted to Edith Beerdsen, David Dubbeldam, R. Krishna, Bei Liu, Sofia Calero, Martijn Schenk, and Thijs Vlugt. These studies were funded in References see page 1690

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References 61. J. A. Rabo, in Proceedings from the Sixth International Zeolite Conference, D. Olson, A. Bisio (Eds.), Butterworth, Guildford, 1984, pp. 41. 62. J. A. Rabo, in New directions in molecular sieve science and technology, J. Basset (Ed.), NATO ASI Series, Series C: Mathematical and Physical Sciences, Kluwer Academic Publishers, Dordrecht, 1988, Vol. 231, pp. 245. 63. J. A. Rabo, R. D. Bezman, M. L. Poutsma, Acta Phys. Chem. 1978, 24, 39. 64. E. G. Derouane, J. Mol. Catal. A: Chemical 1998, 134, 29. 65. E. G. Derouane, NATO ASI Series, Series B: Physics 1990, 221, 225. 66. E. G. Derouane, J. B. Nagy, Appl. Catal. 1989, 52, 169. 67. E. G. Derouane, J. B. Nagy, C. Fernandez, Z. Gabelica, E. Laurent, P. Maljean, Appl. Catal. 1988, 40, L1. 68. D. Frenkel, B. Smit, Understanding Molecular Simulations: from Algorithms to Applications, 2nd edn. Academic Press, San Diego, 2002, 69. M. P. Allen, D. J. Tildesley, Computer Simulation of Liquids, Clarendon Press, Oxford, 1987. 70. R. L. June, A. T. Bell, D. N. Theodorou, J. Phys. Chem. 1992, 96, 1051. 71. R. C. Runnebaum, E. J. Maginn, J. Phys. Chem. B 1997, 101, 6394. 72. D. Schuring, A. P. J. Jansen, R. A. van Santen, J. Phys. Chem. B 2000, 104, 941. 73. E. B. Webb, G. S. Grest, M. Mondello, J. Phys. Chem. B 1999, 103, 4949. 74. T. R. Forester, W. Smith, J. Chem. Soc., Faraday Trans. 1997, 93, 3249. 75. J. Harris, S. A. Rice, J. Chem. Phys. 1988, 88, 1298. 76. J. I. Siepmann, D. Frenkel, Mol. Phys. 1992, 75, 59. 77. D. Frenkel, G. C. A. M. Mooij, B. Smit, J. Phys.: Condens. Matter 1992, 4, 3053. 78. M. N. Rosenbluth, A. W. Rosenbluth, J. Chem. Phys. 1955, 23, 356. 79. M. Dijkstra, J. Chem. Phys. 1997, 107, 3277. 80. M. G. Martin, J. I. Siepmann, J. Phys. Chem. B 1999, 103, 4508. 81. C. D. Wick, M. G. Martin, J. I. Siepmann, J. Phys. Chem. B 2000, 104, 8008. 82. Z. Chen, F. A. Escobedo, J. Chem. Phys. 2000, 113, 11382. 83. M. D. Macedonia, E. J. Maginn, Mol. Phys. 1999, 96, 1375. 84. B. Chen, J. I. Siepmann, J. Phys. Chem. B 1999, 103, 5370. 85. T. J. H. Vlugt, M. G. Martin, B. Smit, J. I. Siepmann, R. Krishna, Mol. Phys. 1998, 94, 727. 86. C. H. Bennett, in Algorithms for chemical computations, R. E. Christoffersen (Ed.), ACS Symposium Series, American Chemical Society, Washington, DC, 1977, pp. 63. 87. D. Chandler, J. Chem. Phys. 1978, 68, 2959. 88. D. Dubbeldam, E. Beerdsen, T. J. H. Vlugt, B. Smit, J. Chem. Phys. 2005, 122. 89. F. Jousse, S. M. Auerbach, J. Chem. Phys. 1997, 107, 9629. 90. T. Mosell, G. Schrimpf, J. Brickmann, J. Phys. Chem. B 1997, 101, 9476. 91. E. Beerdsen, B. Smit, D. Dubbeldam, Phys. Rev. Lett. 2004, 93, art. no 248301. 92. A. G. Bezus, A. V. Kiselev, A. A. Lopatkin, P. Q. Du, J. Chem. Soc., Faraday Trans. II 1978, 74, 367. 93. S. Calero, D. Dubbeldam, R. Krishna, B. Smit, T. J. H. Vlugt, J. F. Denayer, J. A. Martens, T. L. M. Maesen, J. Am. Chem. Soc. 2004, 126, 11377. 94. Cerius2, 2000. 95. P. Demontis, G. B. Suffritti, Chem. Rev. 1997, 97, 2845.

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6

Macrokinetics and Transport Processes 6.1

Rate Procurement and Kinetic Modeling1 Freek Kapteijn, Rob J. Berger, and Jacob A. Moulijn∗

6.1.1

Introduction

Knowledge of the rate and selectivity of a catalytic reaction is essential for its application in practice. More precisely, the kinetics of the reaction are required in the form of a rate expression, i.e. the function that tells how the reaction rate varies with temperature, pressure and composition of the reacting system. The kinetics determine the size of the catalytic reactor for a given overall production rate and, without it, design would be highly speculative. Kinetic studies focus on the selection of an adequate rate expression and determination of the unknown rate parameters that it contains, as expressed in the equation r = f (pi , . . . , T , NT , ki , . . . , Ki , . . . , Keq )

(1)

measurements can be disguised by slow mass or heat transport phenomena inside and outside the catalyst particle or by the reactor configuration, so that not intrinsic rates are determined, thus rendering the data and efforts worthless. A simple test to exclude external mass-transport limitations consists in simultaneously varying the flow rate and the amount of catalyst, while keeping their ratio constant. If no external limitations exist, the resulting conversions should be the same. Internal mass-transport limitations can be revealed by variation of the catalyst particle size. For a reaction to be chemically controlled, the conversion should be independent on the particle size. Catalyst testing and the influence of transport phenomena on the determination of reaction kinetics are dealt with in more detail in Chapter 9.1. In this chapter, the emphasis will be on the methods to obtain data relevant for kinetic modeling and comparison of catalyst activities for the ultimate purpose of engineering applications. The procedure for kinetic modeling is outlined together with the related parameter estimation, the determination of the rate constants in the rate expression. It should serve as a comprehensive reference to this type of activity in catalysis. Most of it comes from authoritative reviews, the reader is referred to for detailed information [2–15].

Generally, the rate is not measured directly but is derived from a measured quantity, conversion or concentration, under given operating conditions such as catalyst amount and feed rate. Apart from kinetic studies to determine the rate equation, other purposes of measuring rates are (i) comparison of various catalyst formulations in screening of new catalysts, (ii) the timedependent behavior of the catalyst activity to predict its long term performance and (iii) to characterize catalysts in e.g. temperature-programmed reduction (TPR) or sulfiding (TPS) studies. In this sense, rate procurement is one of the core businesses of catalysis. The importance of reaction kinetics was confirmed in an industrial survey in 1996 [1, 2]. Because of its importance, rate procurement has to be carried out in the correct way. Rate

Rates of catalytic reactions are obtained by measurement of the conversion of a key component, often the ratelimiting reactant, in laboratory reactors and relating this to the amount of catalyst used and the amount or flow rate of reactants used, to obtain an intrinsic quantity, mol s−1 (amount)−1 . For practical applications, the mass or volume of a catalyst is most relevant as amount, but

1 A list of symbols used in the text is provided at the end of the chapter. ∗ Corresponding author.

References see page 1712

Handbook of Heterogeneous Catalysis, 2nd Ed. .. .. Edited by G. Ertl, H. Knozinger, F. Schuth, and J. Weitkamp Copyright  2008 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim ISBN: 978-3-527-31241-2

6.1.2

Rate Procurement – Types of Laboratory Reactors

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6.1 Rate Procurement and Kinetic Modeling

for comparison studies the amount of active phase on a supported catalyst, its specific surface area or the number of active sites may be preferred. In the latter case, this yields the turnover frequency (TOF) [4], which is relevant for fundamental studies. The number of active sites is, however, often difficult to determine and the mass of the catalyst, W , will be used throughout this chapter, resulting in a rate dimension of mol s−1 kg−1 . Other quantities are easily derived from this. Three ideal reactor types are relevant from reactor theory [16], the two continuous flow types, the plug-flow reactor (PFR) and the continuous-flow stirred-tank reactor (CSTR) and the well-stirred batch reactor. The first two are generally operated under steady-state conditions, yielding time-independent relations for the conversion, whereas the batch reactor operates under transient conditions and the conversion level changes as a function of the elapsed reaction time. The relations that hold for component A in a reaction such as − νP P + νQ Q + . . . νA A + νB B + . . . ← −−− −− → in these reactors are given by PFR: d

dx
A  = −νA × r W

CSTR:


(2)

FA0 xA

 = −νA × r

(3)

dxA = −νA × r dt

(4)

W FA0

Batch: NA

All expressions contain the reaction rate r, but only in Eq. (3) can the rate be calculated directly from the observed conversion. The differential equation of Eq. (2) describes the conversion as a function of the space-time (a kind of residence time), W/FA0 , which can be considered as the reactor coordinate for a given feed rate. Only the exit conversion is measured and not a reaction rate, which varies with the conversion over the bed length. Only for low conversion levels, where the rate can be assumed constant over the bed length, does Eq. (2) reduce to Eq. (3). Therefore, the CSTR and the PFR at low conversion are referred to as differential reactors, in contrast to the integral PFR. Often people calculate rates for the latter on the basis of Eq. (3), but this will be an average value depending on the conversion level. In Eq. (4), the rate is the time derivative of the conversion curve, which can be constructed from the observed conversion–time behavior by mathematical treatment such as differentiation equations or polynomial

or spline interpolation, provided that the product analysis is fast enough to follow the reaction. The same approach can be followed in principle for the plug-flow reactor if data are collected at various space-time values. In kinetic studies, it is not necessary to use rates for determination of the rate parameters. In Eqs. (2)–(4), r represents the rate expression for the reaction under consideration, whatever its mathematical form. This can be inserted into Eq. (2) and subsequently integrated after separation of variables, leading to Eq. (5). The result may be an implicit expression, containing the rate parameters, describing the relation between space-time and conversion. It will be shown later that this relation can also be used for parameter determination.  xA dx W (5) = − 0 νA r FA 0 The choice of the reactor type also influences the type of kinetic information that can be obtained, particularly concerning the effects of the reaction products or intermediates on the reaction rate. A PFR operated at low conversion will provide little information on these effects unless intermediates or products are added to the feed. More information can be obtained in a PFR operated at higher conversion. The CSTR configuration, however, is the most suitable for this purpose since in the whole reactor volume the reactions take place at the same (high) conversion at which the concentration of reaction products is highest. More information about reactor types and the suitability for obtaining accurate and reliable experimental data can be found in Chapter 9. Obviously, a stable catalyst activity is required to measure the reaction kinetics in the normal way. However, catalyst deactivation may be unavoidable in some studies and may even be the major goal of the study. The extent and the type of catalyst deactivation taking place are also important for the choice of the reactor type. A PFR operated at low conversion will be very sensitive to a poison in the feed whereas a CSTR operated at high conversion will be more strongly affected by a poison formed from the reaction products. The batch reactor is a rather unsuitable reactor type in the case of catalyst deactivation, because in a batch reactor catalyst deactivation cannot be observed unless the experiment can be repeated with the same catalyst sample. For rapidly deactivating catalysts a special configuration exists, in which the catalyst is fed to the reactor together with the feed (see Chapter 9.1). In addition to the reactor types described in Chapter 9.1, several dedicated reactor types have been developed for a more in-depth investigation of the reaction mechanism and reaction kinetics. A large group is formed by plug-flow reactors used under transient conditions, in

6.1.3 Kinetic Modeling

which the kinetic data are obtained by analyzing the combined reactor and catalyst response upon a stimulus. Mostly used are a small reactant pulse (e.g. in temporal analysis of products (TAP) [17–20] and positron emission profiling (PEP) [21, 22]) or a concentration step change (in step-response experiments (SRE) [23]). Isotopically labeled compounds are used, which allow operation under overall steady-state conditions, but under transient conditions with respect to the labeled compound (SSITKA) [22, 24–28]. In this type of experiment, both time- and position-dependent concentration profiles will develop, which are described by sets of coupled partial differential equations (PDEs). These include the concentrations of proposed intermediates at the catalyst. The mathematical treatment is more complex and more parameters are to be estimated; for more details see, e.g., Ref. [21]. Another type of reactor that was particularly developed for obtaining much kinetic information in a short experimentation time is the temperature-scanning reactor (TSR), developed by Wojciechowski and Rice [29, 30]. This system allows measurement of the reaction kinetics including temperature dependences by temperature scans at various feed flow rates, resulting in a very dense data collection grid. This reactor does not require isothermicity along the reactor axis during the scans. Nevertheless, there are some requirements for allowing adequate use of this technique: (i) a constant catalyst activity along the bed, (ii) plug-flow behavior, (iii) isothermal conditions at the start of the runs, (iv) high accuracy of the effluent analysis, (v) rapid stabilization of the catalyst performance on variation of the temperature and (vi) absence of (irreversible) slow catalyst deactivation. By performing simulations of the heat transport phenomena in the temperature-scanning plug-flow reactor, Kolkowski et al. [31] found that one must take special care to prevent the occurrence of significant radial temperature gradients and significant temperature differences between the bulk gas phase and catalyst surface, which is often difficult to avoid. There are numerous other ways to reduce the experimental effort, besides reducing the time and effort needed per experimental data point by applying, e.g., automation, fast analysis methods and parallelization of reactors [2, 32–35]. A further reduction of the experimental effort can be obtained by sequential experimental design, i.e. carefully planning the new experimental conditions [10, 11], as will be discussed in Section 6.1.4.5. Basically, kinetic studies consist of (i) the acquisition of rate versus space-time data or of conversion versus space-time data sets for various conditions of pressure, temperature and composition, (ii) derivation and selection of adequate rate expressions and (iii) determination of the

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unknown parameters in the rate expression by nonlinear least-squares methods for which the basic Eqs. (2)–(5) are applied. This will form the main body of what follows. 6.1.3

Kinetic Modeling

There exist many different types of rate expressions [36]. Well known are the power rate law expressions, which are in principle purely empirical and valid only in the range of conditions for which they were established. A better choice for rate expressions is the Langmuir–Hinshelwood–Hougen–Watson (LHHW) model [10, 16], which allows more extrapolation freedom. It is based on a sequence of elementary steps, the (micro)kinetic model, that constitute the overall catalytic process. The breakdown to elementary processes should preferably be based on known information on the reaction mechanism. From the kinetic model, a rate expression can be derived where several assumptions are introduced. We focus here on the kinetics first developed by Hinshelwood, based on Langmuir adsorption, and later by Hougen and Watson [37] to become a systematized tool for process applications. The rate expressions have become known as LHHW expressions. Much has been discussed about the fundamental validity of these kinetic models [5, 14, 15], in view of the underlying assumptions, but their usefulness in engineering applications is undoubted. In the next subsection the procedure is presented for the derivation of the rate expressions. It is emphasized here that in the area of biocatalysis (e.g. enzymatic and microbial conversions), Michaelis and Menten took a similar approach, nowadays known as Michaelis–Menten kinetics. The reproduction of whole cells in microbial conversions yields an extra dimension (autocatalysis) in the kinetic description (Monod kinetics). In contrast to heterogeneous and homogeneous catalysis, a restricted number (∼200) of kinetic models are identified for enzyme catalysis, which makes rapid screening of the resulting different rate expressions (∼100) feasible [38]. Straathof [39] developed the computer code Encora to find the most suitable enzyme kinetic rate expression by progress curve analysis and to fit the kinetic parameters. This parallel development lasted for more than half a century before scientists acknowledged each others’ activities [40]. A comparison of the nomenclature from the fields of heterogeneous catalysis and biocatalysis is given in Table 1. References see page 1712

1696

6.1 Rate Procurement and Kinetic Modeling

Nomenclature used in heterogeneous catalysis kinetics and biocatalysis kinetics

Tab. 1

Heterogeneous catalysis Kinetics

0=

Biocatalysis

Linearization Catalytic center Turnover number

Langmuir–Hinshelwood Michaelis–Menten kE0 CA kNT KA pA v= r= 1 + KA pA kM + CA KA = adsorption constant kM = Michaelis constant Hougen–Watson Lineweaver–Burke ‘‘Active site’’ Enzyme k/s−1

Turnover frequency

r /s−1 NT

Number of turnovers

No. of molecules converted/No. of active sites

Rate expression

6.1.3.1

the steady-state assumption, viz., the assumption that the concentrations of these species do not vary in time: dθ∗ = k−1 θN2 O∗ + 2k3 sθO2 ∗ dt − k1 pN2 O θ∗ − 2k−3 spO2 θ∗2 dθO∗ = k2 θN2 O∗ + 2k−3 spO2 θ∗2 − 2k3 sθO2 ∗ dt dθN2 O∗ 0= = k1 pN2 O θ∗ − k−1 θN2 O∗ − k2 θN2 O∗ dt 0=

(10)

Equations (10) are equivalent to those of Eq. (9), in which only two are independent. Using this steady-state assumption, generally one or more of these unknowns can be eliminated, but one remains. The second assumption is the site balance. The total concentration of active sites is constant and equal to NT : NT = [N2 O∗ ] + [O∗ ] + [∗ ]

Rate Expression

or Derivation of the Rate Expression The welldocumented decomposition of nitrous oxide into nitrogen and oxygen [41] is used as an example: 6.1.3.1.1

2N2 O −−−→ 2N2 + O2

(6)

This irreversible reaction is proposed to proceed according to Eq. (7) over many transition metal oxides, whereby steps 1 and 3 are reversible and step 2 is irreversible. 1. N2 O+∗ ← |2| → N2 O∗ ∗ (7) 2. N2 O → N2 + O∗ |2| ← 3. 2 O∗ → O2 + 2∗ |1| Since these steps are all elementary processes, the rate expression can be based directly on the rate equation, resulting in Eq. (8). In these expressions s represents the nearest neighbors of an active site.

This generally yields a too complex expression for the rate [5, 42] and it is therefore assumed that some of the steps are in quasi-equilibrium; the forward and backward rates are much larger than their difference, e.g. ri  r+i , r−i . The other step(s) that remain(s) are called the rate-determining step(s). In the example (see Figure 1), the adsorption of N2 O and desorption of O2 , steps 1 and 3, are considered in quasi-equilibrium and step 2 rate determining. For the former two it holds that: K1 =

k1 θN2 O∗ = ===⇒ θN2 O∗ = K1 pN2 O θ∗ k−1 pN2 O θ∗

K3 =

 k3 pO2 θ∗2 = ===⇒ θO∗ = pO2 /K3 θ∗ 2 k−3 θO∗

r = r2

r2 = r+2 = k2 NT θN2 O

(12)

(13)

(8) r+1 r−1

Since steps 1 and 2 must proceed twice per overall reaction and step 3 only once, the stoichiometric numbers of the steps equal two, two and one, respectively, and net rate of each step equals the overall net rate r of N2 O conversion as in r = r1 = r2 = 2r3

(11)

and the overall rate is now given by

r1 = r+1 − r−1 = k1 NT pN2 O θ∗ − k−1 NT θN2 O r3 = r+3 − r−3 = k3 NT sθO2 ∗ − k−3 NT spO2 θ∗2

1 = θN2 O∗ + θO∗ + θ∗

r+2

r+3 r−3

(9) r

The three unknown fractional occupancies of empty sites ∗ , oxidized sites O∗ and sites occupied by N2 O must be eliminated from Eq. (8). This is achieved by

Rate determining

Quasi equilibrium

Visualization of the rate-determining step and quasi-equilibrium steps.

Fig. 1

6.1.3 Kinetic Modeling

It is implicitly assumed that all adsorption sites are energetically uniform and can accommodate only one species, as was first suggested by Langmuir. In most cases, it is not necessary to consider more complex adsorption isotherms, such as Freundlich (exhibiting a linear decrease in the adsorption energy with increasing coverage) and Temkin isotherms (exhibiting a logarithmic decrease in the adsorption energy with increasing coverage), since Corma et al. [43] demonstrated that kinetic modeling is often unable to distinguish between the basic Langmuir adsorption isotherm and the alternative adsorption isotherms. The use of the Langmuir adsorption isotherm simplifies the treatment. Inserting Eqs. (11) and (12) in Eq. (13) yields the relation between the rate and the partial pressures of N2 O and O2 : r=

k2 NT K1 pN2 O  pO2 /K3

1 + K1 pN2 O +

(14)

The denominator of Eq. (14) represents the distribution of the sites that are empty, oxidized and occupied by N2 O. The dissociative adsorption of O2 (backward step 3) gives rise to a square root dependence on the O2 partial pressure. Further simplification is obtained by the initial reaction rate consideration, i.e. the product concentrations are that small that they can be neglected in the rate expression. In our example this eliminates the term with pO2 and since typically K1 pN2 O  1, this yields a first order reaction in N2 O: r = k2 NT K1 pN2 O

where the numerator is a product of an apparent rate constant factor, containing the rate constant of the rate determining step and the active site concentration and the driving force for the reaction, a measure how far the overall reaction is from thermodynamic equilibrium. The overall equilibrium constant Keq can be calculated from thermodynamics. The denominator is a kind of inhibition term, always lowering the rate. It expresses the distribution of the active site over the different surface species, including components j that simply adsorb on the sites but which do not take part in the reaction, called inhibitors. The power n of the denominator is generally 0, 1 or 2 and indicates how many surface species are involved in the rate-determining step. Pure fitting values up to 8 have been reported for n. Froment and Bischoff [16] provide a useful step-by-step scheme to build LHHW rate expressions. Considering some limiting cases of the example rate expression Eq. (14), which are shown in Fig. 2, one can recognize three different power rate dependences: (i) Weak adsorption of N2 O and O2 (or low partial pressures of both): first order in N2 O and zero in O2 : r = k2 NT K1 pN2 O

r=

k2 NT K1 pN2 O  pO2 /K3

r = k2 NT

Characteristics of the LHHW Rate Expression The LHHW approach demonstrated above has been followed successfully for many reactions. Extensive examples are described in, e.g., Refs. [10, 15, 16]. One of the most elegant illustrations is the cracking of nalkanes in ZSM-5 [44], where the sites can be considered to approach much better uniform energetics than in other amorphous porous catalysts. A general rate expression of the LHHW type for a reaction A + B ↔ C is of the form

=

rate factor × driving force inhibition term

(18)

(iii) Strong adsorption N2 O or high partial pressures, weak adsorption of oxygen: order in N2 O and O2 both zero, the observed rate is constant:

6.1.3.1.2

krds NT Ki . . . {pA pB − pC /Keq } r= n  1 + KA pA + KB pB + KC pC + j Kj pj

(17)

(ii) Weak adsorption of N2 O (or low partial pressures) and strong dissociative adsorption of O2 : first order in N2 O and −0.5 order in O2 :

(15)

Various rate expressions may be derived from a kinetic model, depending on the assumptions of quasiequilibrium and rate-determining step. Experimental data should provide information on which expression describes the rate dependence best.

1697

(19)

Weak adsorption

Strong adsorption N2O

Strong adsorption O2

Visualization of some limiting cases in the N2 O decomposition. Light spheres, N2 O; dark spheres, O; N2 is not represented.

Fig. 2

(16)

References see page 1712

1698

6.1 Rate Procurement and Kinetic Modeling

Power rate expressions may therefore be considered as special cases of the LHHW models, but valid over a limited range of conditions. In fact, it can be derived from Eq. (14) that the reaction orders depend on the surface occupancies. For the given example: nN2 O = 1 − θN2 O 1 nO2 = − θO2 2

(20)

Some remarks can also be made about the temperature dependence of the reaction rate, as expressed in an app apparent activation energy, Ea . It is assumed that the rate constant has an Arrhenius-type dependence and that the equilibrium constants follow van’t Hoff and the active site concentration is constant:  Eai ki = k0i exp − RT  Hi (21) Ki = K0i exp − RT For the three cases above, the following can be distinguished: (i) The apparent activation energy is lower than that of the rate-determining step since an adsorption enthalpy, H1 , has a negative value thermodynamically: app

Ea

= Ea,r2 + Hads,N2 O

(22)

(ii) The apparent activation energy is given by app

Ea

1 = Ea,r2 + Hads,N2 O + Hdes,O2 2

(23)

and its magnitude will be larger than in case (i) because the desorption enthalpy is a positive quantity. Oxygen has to desorb first before N2 O can adsorb and react. (iii) In this case the real activation energy of the rate determining process is observed: app

Ea

= Ea,r2

(24)

A salient example of case (i) is the cracking of n-alkanes over ZSM-5. The apparent activation energy of the initial reaction rate decreases as a function of the chain length and for C16 and higher the apparent activation energy becomes even negative, due to the larger negative value of the adsorption enthalpy [44]. As a result of the temperature dependences of the parameters, the rate expression may change depending on the operating temperature. Strongly adsorbing components occupy fewer sites with increasing temperature and the apparent rate expression may change from Eq. (18) to Eq. (17) or from Eq. (19) to Eq. (17).

Adaptation of LHHW Models and Other Models The example above showed the usefulness of applying fundamental knowledge in the construction of the kinetic model. As mentioned, kinetic models based on such a (micro)kinetic analysis are often valid in a broader window of conditions than simple power-law equations. A potential disadvantage is that they typically contain more parameters, as a result of the large number of elementary reactions and the adsorption constants of all adsorbing compounds. Particularly with feeds containing many different compounds and with complex reaction networks, this easily becomes unsolvable without adapted approaches. In addition to the often applied lumping of groups of compounds with similar properties, there are more sophisticated approaches to avoid this problem. A successful example is the single-events approach in which the reaction scheme is simplified in terms of elementary steps of the functional groups only, reducing the number of parameters [45, 46]. Of course, one should also considerer techniques allowing the estimation of unknown parameters independently (e.g. adsorption isotherms and thermodynamic properties). Also, moredetailed treatments of the microkinetics of catalytic reactions, as described in Chapter 5.2.1 and in Refs. [3, 10, 16], may be helpful. In all approaches a certain number of unknown parameters remain to be estimated. Alternatively, models created using algorithms based on neural networks are nowadays sometimes applied in kinetic modeling, particularly in industrial research. Such models are called ‘‘black models’’, in contrast with the models based on first principles, which are called ‘‘white models’’. These black models have the advantage of easy use and not requiring any microkinetics modeling. However, these have the disadvantage of requiring more experimental input data and a very poor reliability outside the window of experimental conditions covered. The advantages of both types of models can be combined to form hybrid or gray models [47]. This type of model can be used if only part of the process or reaction kinetics is known [48]. The gray model separates the model into systematic variation due to known sources (white model), systematic variation due to unknown sources (black model) and residual variation. 6.1.3.1.3

Deactivation Kinetics Deactivation of a catalyst negatively affects the concentration of active sites on the catalyst, NT . This should not be confused with the reversible inhibition of the active sites by competitive adsorption, which is treated above. Deactivation may be caused by various phenomena, such as sintering, irreversible adsorption and fouling (for example, coking or metal depositions in petrochemical conversions, see Chapter 7.1). 6.1.3.2

6.1.3 Kinetic Modeling

All these deactivation mechanisms may have different effects on NT as a function of time, position in the bed and even the position inside the catalyst pellet. This complicates the ways in which deactivation can be accounted for, but also allows characterization of the deactivation mechanism to a large extent. The kinetics of deactivation of porous catalysts can be divided into four categories according to their relation to the main catalytic reaction, generalized by A → B [49, 50], where (see the equations below)  is a deactivation function (i.e. the catalyst activity at time t relative to fresh catalyst), kd is the deactivation rate constant, p is the concentration of reactant A, product B or poison P, m is the order of the deactivation and n is the order of its concentration dependence. 1. Parallel deactivation, which depends on the concentrations of reactants, e.g. poisoning or fouling due to the deposition of side-products from the main reaction: d = −kd pAn m dt

(25)

2. Consecutive deactivation, which depends on the concentrations of reaction products, e.g. poisoning or fouling due to the consecutive decomposition of products of the main reaction: d = −kd pBn m dt

(26)

d = −kd pPn m dt

(27)

4. Independent deactivation, which is not affected by fluid composition; e.g. some structural transformations or sintering processes that occur at reaction temperature: d = −kd m dt

can give much information about the deactivation mechanism. Segmental analysis of sectioned pellets can provide evidence of the deactivation profile throughout the pellet. Janssens et al. [51] demonstrated both for the hydrodemetallization of vanadyltetraphenylporphyrin over an Mo/SiO2 catalyst. Differential reactors, preferably CSTRs, are required for more detailed studies of most deactivation kinetics. This reactor is also suitable for testing whole pellets under realistic fluid velocity and isothermal conditions. An important cause of deactivation in industry is coking, which often arises from a side path of the main catalytic reaction or from a precursor that adsorbs strongly on the active sites, but which cannot be related to a measurable gas-phase concentration. For example, for the reaction A → B, the site balance also contains the concentration of blocked sites C∗ . The deactivation function C will be a function of the quantity of coke deposited on the catalyst (CC ): C ≡

NT − [C∗ ] = f (CC ) NT

(28)

The first three deactivation types depend on the concentrations of reactants, products or impurities and the local impact is influenced by concentration gradients within pellets and throughout the catalyst bed. Depending on the rates of the main reaction and that of the deactivation reaction, concentration profiles may develop in the catalyst pellets as a result of pore diffusion limitations and these can result in various deactivation profiles such as pore mouth, throughout or core deactivation. Post mortem characterization of segmentally discharged samples

(29)

Various correlations have been reported for this function [16]: f (CC ) = exp(−αCC )

(30)

f (CC ) = 1 − αCC

(31) 2

3. Side-by-side deactivation, which depends on the concentrations of other components, not involved in the main reaction, e.g. poisoning or fouling by feedstock impurities:

1699

f (CC ) = (1 − αCC ) f (CC ) =

1 1 + αCC

(32) (33)

If the deactivation follows from a consecutive reaction of product B into C∗ according to B → C∗ , an LHHW approach could lead to the following rate of formation of C∗ : rCC = C

kC0 NT KB pB 1 + KA pA + KB pB

(34)

The parallel reaction of A → C∗ can be treated similarly. Note that the rate expression for the reaction A → B together with Eq. (31) indicates that the deactivation depends on the composition of the reaction mixture and may vary along the length of the reactor. Under integral reactor conditions this will therefore lead to a non-uniform catalyst deactivation in the packed bed. By using the linear dependence on the amount of coke formed, i.e. Eq. (29), it follows from Eq. (34) that rCC ≡

kC0 NT KB pB dCC 1 dC (35) =− = C dt α dt 1 + KA pA + KB pB

References see page 1712

1700

6.1 Rate Procurement and Kinetic Modeling

By integration, the following equation results for the deactivation function as a function of time-on-stream:

  t αkC0 NT KB pB dt (36) C (t) = exp − 0 1 + KA pA + KB pB Since Eq. (36) is difficult to handle without doing a thorough simulation, the deactivation function for coking is sometimes not expressed as an empirical function of the amount of coke formed, but directly as a function of time, thus independent of the gas composition. This would allow a direct prediction of the catalyst’s lifetime. Various empirical activity functions have been reported [16]: C = 1 − α t

(for 0 ≤ t ≤ α −1 )

C = exp(−α t) C =

1 1+αt

(37) (38) (39)

Despite their convenience, the use of such timedependent empirical expressions must be discouraged since the fitting ‘‘constant’’ α in Eqs. (37)–(39) will often be a function of the operating conditions determining the coke deposition. By ignoring this, these time-dependent correlations in most cases will only be valid in a narrow window of operation conditions. Therefore, empirical correlations as a function of coke deposited are strongly preferred over empirical correlations as a function of time [16]. In some cases, the catalyst does not deactivate completely after a long time-on-stream but only to a certain residual activity level. A typical example is sintering of the metal crystallites on a catalyst support material. A simple way to accommodate residual activity is to include a term for the residual activity SS in the empirical power law decay expressions. Fuentes [52] reported the generalized power-law equations of Eqs. (40) and (41) to describe such behavior: d = −kd pAn ( − SS )m dt d = −kd pAn (m − m SS ) dt

(40)

6.1.4

Parameter Estimation – Model Discrimination

Rate expressions contain a number of unknown parameters with a physical meaning. Their values are estimated by using on the one hand the experimental data and on the other the calculated values predicted from the rate expressions in the reactor and optimizing a certain objective function. This is called data regression [9]. Several techniques exist to achieve this goal. For a given rate expression, this yields the optimum parameter values, but it has to be decided whether the rate expression is the most adequate one. This selection can be based on statistical analysis and on physical significance of the parameter values (e.g. they should be positive in many cases). Further model discrimination can be achieved by carefully planning additional experiments, i.e. experimental design methods. For simplicity, the treatment in the following subsections will be restricted to single-response models, i.e. models considering only one dependent variable. Section 6.1.4.6 explains how multi-response models can be treated. 6.1.4.1

Data Regression

6.1.4.1.1 Maximum Likelihood and Variance Models In the parameter estimation theory, it is generally assumed that the experimental errors are normally distributed with zero mean and a constant variance σ 2 . The parameter values can then be estimated by maximizing the likelihood function of the parameters. If the errors are also statistically independent, this likelihood function reads  n 1 1 (42) L(β|y) = √ exp − 2 (yi − ηi )2 2σ ( 2π)n i=1

Under these conditions, this is identical with minimizing the sum of squares of residuals Resi , the SSR defined in Eq. (43). This is called the objective function for the parameter estimation. SSR = S(β) =

n

Resi 2

i=1

(41)

It is noted that this approach with deactivation until a residual activity can also be applied for the initial fast decay of a catalyst down to a more stable state, often referred to as ‘‘lining out’’. Of course, the deactivation may also be a function of other compounds than reactant A. The rate of sintering, for example, often depends on the concentration of specific species, such as steam, hydrogen or oxygen.

=

n

β1 , β2 , . . . , β p

(yi − ηi )2 −−−−−−−−→ Min

(43)

i=1

where βi represents the true, but unknown, parameter value and ηi the true, but unknown, value of the response variable, the measured variable. The latter is also called the dependent variable, representing conversion or product concentrations, in contrast to the independent variables xi , which represent the experimental settings such as temperature, pressure and concentrations. Since the true

6.1.4 Parameter Estimation – Model Discrimination

values of βi and ηi are unknown, the estimated or calculated values will be denoted by bi and yˆi . Although in most cases the assumption of a constant variance is appropriate, one might consider the variance to be related with the absolute value of the response. One option is to assume a constant relative variance, in which it is assumed that the variance is proportional with the value of the response. Also, heteroscedastic variance models, a combination of both, can be applied: σ 2 = ω2 (y 2 )γ

(44)

where γ represents the heteroscedasticity of the variance (0 < γ < 1); γ = 0 represents a constant variance model and γ = 1 represents a constant relative variance model. In order to avoid numerical problems at response values close to zero, y 2 should be increased with a very small number. These alternative variance models are useful in cases in which the variance behaves heteroscedastically, the response varies over more than one order of magnitude and where it is important to fit accurately the small response values. Linear Regression If the model for the response variable is linear in the parameters for each observation at chosen settings of the independent variables xi , then for the set of n observations and p parameters one can write [9] 6.1.4.1.2

y1 = β1 x11 + β2 x12 + . . . . + βp x1p = η1 + Res1

and the estimate of the error variance σ 2 is given by the mean SSR, the minimum sum of squares of residuals divided by its corresponding number of degrees of freedom (d.f .): n

ResT Res σ 2 ≈ s2 = = n−p

. . (45)

In matrix notation, this reads (46)

and the least-squares criterion as β

ResT Res −−−→ Min

(47)

The solution of this linear least-squares problem is exact and the estimate vector b for the parameter values βi follows, in matrix notation [9]: b = (XT X)−1 · XT y

(48)

Under the conditions mentioned above, b is an unbiased estimate of the real parameter values and its variance–covariance matrix V(b) is given by V(b) = (XT X)−1 σ 2

i=1

n−p

(50)

When the errors are normally distributed with zero mean but the error variance is not constant and the errors are interdependent, represented by a variance–covariance matrix V, i.e. Res ∼ N (0, Vσ 2 )

(51)

then a weighted least-squares minimization has to be applied, with the following results: SSR = S(b) = ResT V−1 Res −−−→ Min

(52)

with vector of residuals Res having elements of Resi = yi − yˆi

(53)

The maximum likelihood estimates are then given by b = (XT V−1 X)−1 · XT V−1 y

(54)

and the variance–covariance matrix by V(b) = (XT V−1 X)−1 σ 2

s2 =

y = Xβ + Res

(yi − yˆi )2

(55)

An estimate s 2 of the unknown factor σ 2 is now obtained from

y2 = β1 x21 + β2 x22 + . . . . + βp x2p = η2 + Res2

yn = β1 xn1 + β2 xn2 + . . . . + βp xnp = ηn + Resn

1701

(49)

=

S(b) ResT V−1 Res = n−p n−p (y − Xb)T V−1 (y − Xb) n−p

(56)

Based on the (symmetric) variance–covariance matrix of the parameter estimates V(b), one can determine confidence limits of the parameter estimates. The diagonal elements vii contain the parameter estimate variances and the off-diagonal elements the covariances between the parameter estimates. The confidence range of the parameters (i.e. the interval of parameter values that are statistically not significantly different from the estimated value bi ), at a selected probability level (1 – α), is defined by √ √ bi − tn−p,1−α/2 vii < βi < bi + tn−p,1−α/2 vii (57) where tn−p,1−α/2 represents Student’s t-value at n − p degrees of freedom and a confidence level of 100(1 – α)%. References see page 1712

1702

6.1 Rate Procurement and Kinetic Modeling

Usual values of α are 0.05 and 0.01. This relation is valid if all other parameters are kept at their optimum estimate bj . Since the parameters are in fact correlated, a joint confidence region of the estimates b can be defined, which accounts for the simultaneous variation of all the parameters:     p S(β) ≤ S (b) 1 + F (p, n − p ; 1 − α) (n − p) (58) where F (p, n − p ; 1 − α) is the α-percentage point of Fischer’s F -distribution at p and (n − p) degrees of freedom. The boundary of the joint confidence region is defined by all values of β which satisfy the hyperellipsoid in the p-dimensional parameter space around b: (b − β)T XT X (b − β) = ps 2 F (p, n − p ; 1 − α)

(59)

All parameter combinations enclosed by the ellipsoidal surface do not deviate significantly from the maximum likelihood estimates b at the probability level (1 − α). Examples for a two-dimensional case are given in Figure 3, to be discussed later. Nonlinear Regression Most kinetic expressions, however, are not linear in the parameters and two approaches can be followed. The first is to rewrite the expression in a linear form and apply the linear least squares minimization to obtain parameter values. Equation (14) can be reformulated as  1 1 1 1 1 + = + √ r k2 NT k2 NT K1 pN2 O k2 NT K1 K3 √ pO 2 × (60) = b0 + b1 x1 + b2 x2 pN2 O 6.1.4.1.3

Z2

Exact

Z1

b2

+

b1 (a)

+

b1 (b)

Confidence contours for two parameters, b1 and b2 . The optimum value indicated by the + sign. (a) The linearized contour, Eq. (59), and the banana-shaped exact contour, Eq. (58); (b) examples of parameter improvement designs.

Fig. 3

bk = ak XkT Resk = ak XkT (y − yˆ k ) xij =

∂ yˆi ∂bj

for i = 1, . . . , n and

j = 1, . . . , p

(61) The methods differ in the determination of the steplength factor ak at the kth iteration, since the direction of the steepest descent is, due to nonlinearities, not necessarily the optimum one, but only for quadratic dependences. Some methods therefore use the secondderivative matrix of the objective function with respect to the parameters, the ‘‘Hessian’’ matrix, to determine the parameter improvement step-length and its direction: H = ∇ 2 SSR where hij =

∂ 2 SSR ∂bi ∂bj

i, j = 1, . . . , p

−1 T T ˆk ) bk = H−1 k · Xk Resk = Hk · Xk (y − y

Approximate

b2

where three parameters are present and two ‘‘independent variables’’. This approach was used by Hougen and Watson in heterogeneous catalysis and Lineweaver and Burke in biocatalysis. Similarly, exponential relations are ‘‘linearized’’ by taking the logarithm. It is noted that the error distributions change due to these types of transformations and the error limits provided by the equations given above are incorrect [11]. Therefore, it is better to use the nonlinear model directly in a nonlinear regression of the observed variable, the nonlinear least-squares method. Because of the nonlinearity, minimization is an iterative process. Various more or less efficient optimization strategies have been developed [53–55] and can be classified as direct search methods and gradient methods. The direct search methods, such as those of Powell [56], Rosenbrock and Storey [57] and Nelder and Mead [58] (‘‘simplex’’) start from initial guesses and vary the parameter values individually or combined thereby searching for the direction to the minimum SSR. The gradient methods, such as the Newton, Gauss– Newton, Fletcher and Levenberg–Marquardt methods, use the derivative vector of the SSR with respect to the parameter directions to determine the direction where this gradient changes most, the steepest descent direction:

(62)

In least-squares minimization, where it is assumed that the residuals are small, this Hessian can be approximated by the earlier encountered first derivative matrix multiplication:  ∂ 2 SSR ∂ T ∂ Res = hij = 2 Res ∂ bi ∂ bj ∂ bi ∂ bj =2

∂ 2 Res ∂ ResT ∂ Res + 2 ResT ∂ bi ∂ bj ∂ bi ∂bj

or, in matrix notation, H ≈ XT X.

(63)

6.1.4 Parameter Estimation – Model Discrimination

bk = (XkT Xk + λk I)−1 · XkT Resk

(64)

For large values, this takes the direction of the steepest descent, whereas for λ → 0 it has the Gauss–Newton direction. During the parameter estimation procedure, the value of λ is adapted (generally decreased finally to zero), so this Levenberg–Marquardt method is an optimum compromise between the steepest descent, efficient far from the minimum, and the Gauss–Newton method, efficient when close to the minimum [54]. Generally, the models are so complex that analytical derivatives cannot be provided, so they are calculated numerically by a small variation of the parameter value. It is noted that the gradient techniques directly provide the parameter variance–covariance matrix, Eq. (49) or (55), since this is already used in the calculation of the parameter step changes at each iteration. Also, the relations for the estimates of the parameter error limits, Eq. (57), and the joint confidence region, Eqs. (58) and (59), can be used. These are only approximations due to the nonlinearity of the model. Hence Eq. (59) gives an approximate joint confidence region, a hyperellipsoid, whereas Eq. (58) defines the contour of the real confidence region at an approximate confidence level of 100(1 – α)%. This often gives a banana-shaped contour, as shown in Fig. 3. Sometimes a very strong correlation exists between the parameters, resulting in a confidence contour having a strongly elongated shape [11]. This makes convergence during minimization difficult. A well-known example is reaction rate constants and their temperature dependence, expressed by the Arrhenius relation such as in Eq. (21). It is strongly recommended to reduce this correlation by applying a reparametrization as shown in Eq. (65). This yields a much more spherical shape of the confidence contour and facilitates convergence [11].    Ea 1 1 k = kref exp − − R T Tref  Ea where kref = k0 exp − (65) RTref and Tref represents a reference temperature, which should be chosen within the experimentally covered temperature window, preferably the average temperature in this window. A similar reparametrization is recommended for equilibrium constants.

Also, an analysis of the source of information concerning the parameters to be estimated can help to avoid difficult convergence and strong correlations between parameters. This can be demonstrated using the following simple sequential reaction: r1

r2

A −−−→ B −−−→ C the kinetics of which are described using the following LHHW equations: r1 =

k1 KA pA 1 + KA pA + KB pB + KC pC

r2 =

k2 KB pB 1 + KA pA + KB pB + KC pC

(66)

The concentration of the desired product, B, as a function of time is shown in Fig. 4. On closer inspection, it appears that the crucial information needed to estimate the five parameters involved originates from different parts of the curve, as illustrated in the figure. Without any experimental data in the region of the first ellipse on the left-hand side, it will be difficult to estimate k1 accurately and the parameter estimation will end up with a strong correlation between k1 and KA . The shape of the curvature in the area of the second ellipse gives much information about KA and KB . If KA pA 1 and KA pA KB pB + KC pC , the curve would be a straight line up to full conversion of A. If KA pA  1, there will be a curve like the one in the figure. Similar explanations can be given for the information concerning the other parameters. Since the information for estimating KA and KB , and also for KB and KC , originates from the same part of the curve, it will be difficult to estimate both accurately without a strong correlation. To overcome this problem,

1.0

CB / mol L–1

However, the Hessian or its approximation has to be inverted to determine the parameter step-change for the next iteration and, especially far from the real minimum of SSR, the matrix is not positive definite, a requirement for inversion. Levenberg [59] and Marquardt [60] therefore added a diagonal matrix to it and allowed this contribution to vary according to a parameter λ, the Marquardt parameter. For the kth iteration, this yields

1703

K2

0.8

KA and KB

0.6 0.4

KB and KC K1

0.2 0.0 0

100

200

300

400

500

Time /min

Illustration of the locations of the source of information concerning the parameters in the LHHW equations (open triangles, experimental data; curve, calculated data using the optimum parameter estimates).

Fig. 4

References see page 1712

1704

6.1 Rate Procurement and Kinetic Modeling

one should carry out some additional experiments with co-feeding of component B. In addition to strong correlation with other parameters, another problem that often occurs with the adsorption constants in LHHW expressions is that these are too small or too large to be estimated accurately. For allowing an accurate estimation of, e.g., KA , the value of KA pA should vary around unity whereas KB pB + KC pC may not be much larger than unity. Otherwise, the parameter estimator can only indicate that the value must be very high or very low, but without a reliable confidence range. In that case, it is recommended to rearrange the model after elimination of this parameter. In any parameter estimation, one should strive to limit the number of parameters to be estimated and to start from good initial guesses for the parameters, not only to avoid difficult convergence due to strong correlations between parameters, but also to avoid the risk of ending up in a local minimum that is far away from the global minimum. The best way to achieve this is to estimate required parameters from other types of experiments such as adsorption measurements or using microkinetics approaches, as discussed in Section 6.1.3. The linearization technique is also useful to provide good initial guesses for the parameters. In nonlinear regression the parameter estimation iteration continues until a certain criterion is satisfied or the maximum number of function evaluations is exceeded. These criteria may be that the relative change in the SSR value or in the parameter values is below a preset value or the norm of the gradient is less than a certain value (in the minimum this gradient norm vanishes), etc. Kinetic Data Handling The observations made in kinetic studies are usually concentration or conversion measurements. CSTRs and differential PFRs yield rate data directly, so the experimentally observed rate can be directly fitted to the rate expression and its parameters estimated. Integral PFR data cannot be used directly in this way, since one has a differential equation that describes the conversion profile along the catalyst bed [Eq. (2)]. For simple cases this can be integrated analytically, yielding an implicit expression in the observed variable [Eq. (5)]. Sometimes the independent variable W/Fi0 is now used as observed variable and its SSR minimized [10], but this interchange of dependent and independent variables destroys the error properties and the parameter error limits are not correct. Often the parameter estimates correspond well [10] and can be used as starting guesses for more robust minimization to determine the real error bounds. 6.1.4.2

When the integration of Eq. (2) cannot be performed analytically, the integration should be carried out numerically, e.g. using robust routines such as the Runge–Kutta, Adams–Moulton predictor–corrector or Bulirsch–Stoer methods with step size and error control [61–63]. A similar approach holds for the interpretation of batch reactor data. Here, the spatial integration has to be replaced by a time integration of Eq. (4). A nice example of such an application is the catalyzed hydrodemetallization of Ni porphyrins [64], a reaction of the type −−  −−  A −− − −B −− − − C −−−→ D which is a multi-response system (see Section 6.1.4.6). In combining the numerical integration with nonlinear parameter estimation, one should be aware of the nesting of errors in the procedures. One should not demand a high accuracy in the SSR determination if the numerical integration does not have sufficient accuracy. The innermost level, the integration, requires the highest numerical accuracy, the next level, the numerical derivatives calculation, the next highest and the stopping criteria the lowest, otherwise the procedure will fail to converge. It is clear that one has to play around with the numerical accuracies to develop a feeling for the proper settings of these values. Most challenging are kinetic studies by transient techniques. Here, one has a set of coupled partial differential equations (PDEs) to be solved that describe the concentration profiles along the catalyst bed as a function of time, embedded in the parameter estimation routine. Good results have been obtained by the relatively easy implementable numerical method of lines (NUMOL) [65]. In this method, the spatial derivatives are replaced by an algebraic approximation in ng grid points along the catalyst bed length. For each PDE this results in a set of ng coupled ordinary differential equations (ODEs). These initial value problems can be solved by the robust explicit integrators mentioned above. The use of implicit integrators, such as LSODE [66], can have certain advantages. These use the Jacobian of the dependent variables, i.e. the derivative with respect to all grid point positions. This Jacobian is rather sparse and a reduced calculation effort can be obtained by application of programs that take this sparsity into account, such as LSODES [66]. After initialization and set-up of the Jacobian, the implicit methods are fairly fast in their time integration. However, when the parameter values change substantially during the parameter estimation, a new Jacobian has to be calculated after each parameter iteration, thus losing this advantage. Therefore, in parameter estimation explicit routines have some preference and have much lower memory

6.1.4 Parameter Estimation – Model Discrimination

requirements. A more detailed discussion and examples can be found in Ref. [21]. Model Testing A kinetic expression that is selected should satisfy certain statistical tests and physicochemical constraints before it can be indicated as an adequate model. Often these tests cannot be applied all together due to lack of information, but should be considered as far as possible in the evaluation and selection of the best rate expression. When still a set of competitive expressions is left, some model discrimination techniques may be applied or additional experiments should be conducted, based on careful, efficient planning. 6.1.4.3

6.1.4.3.1 Statistical Testing – Model Adequacy The residual sum of squares, SSR, contains contributions of a pure error, PE, due to pure experimental errors, and a lack-of-fit, LF, due to the inadequacy of the model. The pure error sum of squares PE can be obtained by, e.g., ne replicated experiments at a number of, at least one, experimental settings. The relations that hold are as follows: ne

(yr − y) ¯ 2

degrees freedom ne − 1

Physicochemical Testing – Kinetic Parameters The kinetic parameters in the rate models all have a positive value and negative values resulting from the parameter estimation have no physical meaning. 6.1.4.3.4

r=1

se2 =

PE ne − 1

LF = SSR − PE slf2 =

6.1.4.3.3 Statistical Testing – Residual Distribution The error model used in the minimization is based on the hypothesis that the residuals have zero mean and are normally distributed. The first is easily checked, but the latter is only possible when sufficient data points are available and a distribution histogram can be constructed. An adequate model also follows the experimental data well, so if the residuals are plotted as a function of the dependent or independent variable(s) a random distribution around zero should be observed. Nonrandom trends in the residuals mean that systematic deviations exist and indicate that the model is not completely able to follow the course of the experimental data, as a good model should do. This residual trending can also be evaluated numerically by correlation calculations, but visual inspection is much more powerful. An example is given in Fig. 5 for the initial rate data of the metathesis of propene into ethene and 2-butene [67]. One expression was based on a dual-site Langmuir–Hinshelwood model, while the other was based on the carbene mechanism, which is nowadays generally accepted as the reaction mechanism. From the residual plots the latter model is clearly the best.

degrees freedom n − p − ne + 1

LF n − p − ne + 1

(67) Residual rate

PE =

These variances of PE and LF, se2 and slf2 , can be used in an F -test for lack-of-fit. If Fc =

slf2 se2

1705

> F (n − p − ne + 1, ne − 1; 1 − α)

(68)

p propene

0

p propene

(a)

Residual rate

then there is a chance (1 – α) that the model is inadequate. The model is retained if the calculated value does not exceed the F -value. This test is only rarely applied [67], since often there are hardly replicated data available. Statistical Testing – Parameter Significance Certain parameters have badly determined values and it is questionable whether they should be retained in the rate expression. This is simply achieved by the ttest, implicitly present in Eq. (57). If the interval of the estimated parameter contains zero then it does not deviate statistically from zero, at the applied confidence level (1 – α).

0

6.1.4.3.2

(b)

Residual rate distribution for the metathesis of propene. (a) Langmuir–Hinshelwood model; (b) model based on the carbene mechanism [59].

Fig. 5

References see page 1712

1706

6.1 Rate Procurement and Kinetic Modeling

The same holds for the temperature dependence. Activation energies should be positive, unless a combination of parameters is fitted and apparent values are obtained. Due to their composition, even negative values may be obtained [see Eqs. (22)–(24)]. Moreover, from thermodynamics it follows that adsorption enthalpies are negative and desorption enthalpies positive. The same holds for the associated entropy changes. A molecule loses degrees of freedom upon adsorption, at least one translational degree when mobile adsorption occurs. These entropy losses can be estimated fairly well [3, 68, 69]. In an analysis of kinetic modeling parameters, Boudart et al. [70, 71] found a correlation between the enthalpy and entropy of adsorption that may serve as a guideline, rather than a constraint. Table 2 gives a summary of the physicochemical constraints. It should be emphasized that the LHHW kinetic rate expressions are derived with assumptions of energetic uniformity and if this is violated then these constraints should be used with caution. In a transient kinetics study, Dekker et al. [25] nicely showed that an occupancydependent CO adsorption enthalpy on Pt results in very low values of the reaction activation energy and might even become negative. Much more can be said about the magnitude of preexponential factors and activation energies of elementary processes based on statistical thermodynamics applied to collision and reaction-rate theory [3, 68], but in view of the remark above one should be cautious in their application and limit it to well-defined model reactions and catalyst surfaces. Physicochemical Testing – Site Density Steadystate kinetic modeling does not yield separately the value of the site density NT , but always in combination with rate parameters [see Eqs. (14)–(19)]. Only selective poisoning experiments and transient kinetic techniques can yield this information [72, 73]. Maatman [74] analyzed many catalytic systems and gave a (broad) range of acceptable 6.1.4.3.5

Physicochemical [69–71, 74]

Tab. 2

Feature Adsorption Unimolecular surface reaction Rule I Rule II Rule III Guideline I Guideline II Site density

constraints

of

kinetic

values for site densities L (sites cm−2 ), included in Table 2. Discrimination Between Rival Models After fitting data to the various rate expressions that have been derived, it will appear that often several models do not differ much in their mean SSR values. Statistical tests, such as an F -test, cannot be performed directly on these mean SSR values since they are based on the same data set and are therefore statistically not independent. Therefore, other approaches have been devised to allow further selection between the possible candidates, some of which are considered below. 6.1.4.4

Model Discrimination – Initial Rate Expressions Hougen and Watson [37] suggested analyzing the rate dependence on the partial reactant pressure or the total pressure at low conversion levels where the product concentrations can be neglected and so-called initial rates are measured. Depending on the assumed ratedetermining step in the kinetic model, a different pressure dependence is predicted, as exemplified in Fig. 6. This allows a direct discrimination between possible rate expressions of different models. 6.1.4.4.1

Model Discrimination – Diagnostic Parameters Suppose a selection has to be made between two rival models, both of which have been fitted already to 6.1.4.4.2

T3 r0

T3

r0

T2 T2 T1

T1

parameters

Constraint

p0

p0

adsorption

surface

T3

r0

◦ Kads = exp(−Sads /R) exp(−H◦ads /RT) k = k0 exp(−Ea /RT)

T2 T1

◦ ◦ < Sgas 0 < −Sads ◦ −Hads > 0 Ea > 0 ◦ >∼ 40 (J mol−1 K−1 ) −Sads ◦ < 51.2 − 0.0014 H◦ads (J mol−1 K−1 ) −Sads −9 10 < L < 10 (sites nm−2 )

p0 desorption Examples of initial rate dependences on the partial reactant pressure for different rate determining steps.

Fig. 6

6.1.4 Parameter Estimation – Model Discrimination

the experimental data and their parameters have been estimated. Then a new dependent variable is defined [75] containing a non-intrinsic diagnostic parameter λ and the calculated predictions of the two models: yˆ =

1 (yˆ1 + yˆ2 ) + λ (yˆ1 − yˆ2 ) 2

(69)

Now a linear least-squares minimization is used to determine λ. If λ takes the value 0.5 then model 1 is preferred, whereas for −0.5 model 2 is the best. The confidence limits should be evaluated to be taken into account since it should not include the value of the other model. For m models, Wilks [76] proposed the following new dependent variable: yˆ = λ1 yˆ1 + λ2 yˆ2 + . . . . . . + λm yˆm with

0 ≤ λi ≤ 1 and

m

λi = 1

(70)

i=1

The diagnostic parameters λi are scaled and their estimation is also a (constrained) linear least-squares problem. These parameter values represent the fraction of y that is accounted for by model i. 6.1.4.4.3 Model Discrimination – Bartlett’s Chi-Squared Test The m rival models with their mean SSR or error variances, si 2 , are pooled to a set of variances that are tested upon homogeneity. The following random variable is therefore defined [77]:

ln s 2 χc 2 =

m

(d.f.)i −

m

(d.f.)i ln si 2

i=1 i=1

m 

1 1 1 1+ − m 3(m − 1) (d.f.)i i=1 (d.f.)i i=1

m

with s 2 =

(d.f.)i si 2

i=1 m

1707

Sequential Experimental Design Optimum planning of kinetic experiments is crucial for an efficient experimental effort in this time-consuming activity. Initially, a series of experiments will be conducted to determine the effect of the various independent variables (experimental conditions), which can be carried out according to a classical factorial design (Fig. 7). If linear relations are to be tested, then only two conditions for each independent variable should be tested: at the lowest and at the highest possible setting and this for each independent variable. Also, cross-correlations are taken into account in this way. For q independent variables this leads to 2q experimental settings. If nonlinearity is assumed to be involved, at least three conditions should be taken per independent variable, leading to a 3q -factorial design. A picture of these designs is given in Fig. 7. Intermediate designs can be devised to reduce the rapidly increasing number of experiments, for example by the introduction of a center point, through which nonlinearities can be determined with less effort than a complete 3q -factorial design. After this initial stage, the selection of the best model and accurate parameter values are the two major goals in kinetics. Subsequent experiments are carried out under the optimum conditions for further model discrimination or parameter improvement. In both goals one should distinguish the criterion to achieve the goal, i.e. the discrimination criterion or parameter accuracy criterion, and the design criterion that determines the experimental settings. The experiments are limited by the experimental range of conditions that can be achieved in the equipment. Usually the experimental conditions are represented by a grid of combinations. In each grid point the design criterion is evaluated to determine the optimum experimental setting. This grid search is easier and faster than application of optimization routines described above. After each new experiment, the parameter estimation procedure is repeated, the new information is extracted and a new design is made until a satisfactory goal is 6.1.4.5

(71) (d.f.)i

i=1

This variable is distributed like χ 2 with m − 1 degrees of freedom (d.f.). If its value exceeds χ 2 (m − 1, 1 – α) the m models do not form a homogeneous set at a probability of (1 – α). Now the model with the highest si 2 value is rejected and the test is repeated with the remaining m − 1 models until the χ 2 is no longer exceeded. The remaining set of si 2 is homogeneous and the models are competitive. As these models are not yet guaranteed to be adequate, this should be based on the statistical tests given above.

(a)

22

(b)

23(+1)

(c) 23 + 3 × 2 + 1

Factorial designs for initial sets of experiments. (a), (b) 2q designs for linear models of two and three independent variables, respectively; (b), (c) reduced designs for nonlinear models of three independent variables. Fig. 7

References see page 1712

1708

6.1 Rate Procurement and Kinetic Modeling

reached. With all the power present nowadays in PCs it will be clear that the most time-consuming step is the conduction of the experiment, which motivates this approach.

No overlap confidence limits

Largest divergence in value

Model 1

Discrimination Between Rival Models – Design Criterion The various approaches to discriminate between models and to test models’ adequacies have been treated above and focus is given to the design criterion. A simple approach can be followed to determine the maximum divergence between model predictions. Based on statistical considerations for two models, a simple expression was derived to measure the divergence D(xk ) at the experimental settings xk of the experimental grid: 6.1.4.5.1

Rate

Model 2

Experimental range

p0

D(xk ) = {yˆ (1) (xk ) − yˆ (2) (xk )}2

(72) Illustration of discrimination between two rival models. Models (solid lines) are plotted as a function of pressure, together with their confidence areas and the experimental data (stars). At the highest pressure the largest difference in rate values exists, but confidence regions overlap. At the lowest pressure these regions do not overlap, an optimum location for the next experiment.

Fig. 8

The set of conditions that yields the largest value of the divergence is chosen at which to conduct the next experiment. The criterion is easily extended to m models, as follows: D(xk ) =

m−1

m

{yˆ (i) (xk ) − yˆ (j ) (xk )}2

(73)

i=1 j =i+1

The double summation is used to take each of the models as a reference in order to avoid mislocations of optimum discrimination conditions. Since the model adequacy criteria and the design criterion are independent of one another, any type of design criterion can be used. An alternative one is given by [10] D(xk ) =

m−1

m

|yˆ (xk ) − yˆ (i)

(j )

(xk )|

(74)

For a nonlinear model, the matrix X represents the matrix of partial derivatives of the dependent variable with respect to the parameters in all data points, Eq. (61), and vk the vector of the partial derivatives taken at settings xk of the independent variables. Box and Hill [78] derived the following expression for the divergence between m rival models: D(xk ) =

m−1

m

 πi,n πj,n

i=1 j =i+1



−{yˆ (xk )− yˆ (xk )} (i)

(i)

i=1 j =i+1

2

(σki 2 − σkj 2 )2 (σ 2 + σki 2 )(σ 2 + σkj 2 )

1 1 + σ 2 +σki 2 σ 2 +σkj 2



(76) The divergence is determined not only by the value of the response variable y, ˆ however, but also by their statistical uncertainties. Large differences between the model predictions are not necessarily important if their associated uncertainty is large. This is illustrated in Fig. 8, which shows the predictions and their confidence limits. An experiment that is conducted under conditions where the confidence limits do not overlap has more power than one where the difference between the models is the largest, but where the confidence regions overlap. The variance of a model yˆ at settings xk is given by [9]: for linear models: V [y(x ˆ k )] = σk 2 = xTk (XT X)−1 xk σ 2

(75a)

for nonlinear models: V [y(x ˆ k )] = σk 2 = vkT (XT X)−1 vk σ 2

(75b)

After n experiments the data are analyzed and the next (n + 1)th experiment has to be conducted under those conditions xk which maximize D(x); πi,n is the probability or degree of adequacy reached after n experiments. Initially all models start with equal probabilities. After the (n + 1)th experiment this probability is updated by the Bayesian posterior probabilities of the models: πi,n+1 =

πi,n pi m

πj,n pj

i = 1, 2, . . . , m

j =1

1 in which pi = pi (yn+1 ) =  2π(σ 2 + σn+1,i 2 )  (i) [yn+1 − yˆn+1 ]2 (77) × exp − 2(σ 2 + σn+1,i 2 )

6.1.4 Parameter Estimation – Model Discrimination

is the probability density of the (n + 1)th observation yn+1 under model i. Discrimination is achieved as soon as the posterior probability of a model approaches the value of 1. Inspection of Eq. (76) shows that:

A-optimality

1709

Z1

E1-optimality (Shape design)

• Experimental conditions that result in large divergence in model predictions contribute strongly to D due to the squared term. • Predictions which are highly unreliable, with large variances contribute less due to the inverse proportionality. • Greater weight is given to the models which approximate the results best.

E2-optimality

b2

Z2 D-optimality (Minimum volume) b1

Several application examples were given by Froment and Hosten [10]. Recently, Chen and Asprey [79] presented methods for the optimum design of dynamic experiments for model discrimination in multi-response experiments. 6.1.4.5.2 Parameter Improvement Once a model has been selected, it is often important to improve the precision of the estimated parameters. The cornerstone of the theory is the covariance matrix of the parameter estimates. The covariance matrix defines a hyperellipsoid around the optimum parameter combination; the joint confidence region, Eq. (59), can be written as

(b − β)T XT X (b − β) = δ

(78)

which can be transformed to 

z1 √ δ/λ1

2
2  2 zp z2 + √ + ··· +  = 1 (79) δ/λ2 δ/λp

where λi represent the eigenvalues of the matrix XT X and zi the coordinates of the eigenvectors, i.e. the axes of the hyperellipsoid. This is indicated in Fig. 3. Various criteria can be used to improve the parameter estimates [10], which are illustrated in Fig. 9. 6.1.4.5.3 Parameter Improvement – D Optimality (Minimum Volume Design) The volume of the hyperellipsoid represents the total parameter uncertainty, so its minimization is a valid goal. The volume can be determined by

Vol. = constant ×

p 

 δ/λk

Geometric representation of the divergence criteria for parameter improvement. The shaded area represents the approximate joint confidence region between two parameters, b1 and b2 .

Fig. 9

set of existing data settings and the determinant of the matrix XT X is evaluated. This is done for each grid point to determine the highest value. 6.1.4.5.4 Parameter Improvement – E Optimality (Shape An alternative objective is to render the Design) confidence region as spherical as possible; this means that the longest axis of the hyperellipsoid has to be reduced (E1 optimality). Mathematically, the length of the axes is inversely proportional to the eigenvalues, so the smallest eigenvalue of XT X has to be maximized. Alternatively, the ratio of the lowest to the largest eigenvalue can be maximized (E2 optimality). Two interesting case studies in which the D optimality in addition to the E optimality is used for parameter improvement can be found in Ref. [10].

Improved Parameter Estimation – A Optimality The variance of the parameter estimates can be minimized, which means the minimization of the trace of (XT X)−1 : 6.1.4.5.5

p

xk

(XT X)−1 ii −−−→ Min

(81)

i=1

Geometrically, this means that the dimensions of the enclosing box around the joint confidence region is minimized, as visualized in Fig. 9.

k=1

= constant × 

δ p/2 1 √ λ1 λ2 . . . λp det XT X

(80)

The experimental conditions are now chosen by a grid search such that the det XT X is maximized. For this purpose each new condition is individually added to the

Improved Parameter Estimation – Other Criteria Alternatively, one might try to improve the least accurate parameter estimate or minimize the correlation between the parameters by minimizing the sum of squares of the 6.1.4.5.6

References see page 1712

1710

6.1 Rate Procurement and Kinetic Modeling

correlation coefficients between the parameters which are also calculated from the elements of (XT X)−1 : p p

xk

ρij 2 −−−→ Min

i=1 j =1

(XT X)ij −1 {ρij } =  T −1 (XT X)−1 ii (X X)jj

(82)

Flowsheet for Sequential Experimental Design A typical flow diagram of the application of sequential experimental design is given in Fig. 10. It should be noted that in the case of both model discrimination and parameter improvement, the design criteria often select experimental settings at the borders of the realizable conditions. In a way this is not surprising since at these locations the models often diverge most from each other or from the data, so they have the highest new information content. For one independent variable this location is easy to recognize, but with several independent variables it will be more easily achieved by the techniques described above. 6.1.4.5.7

Multi-response Models All treatments above have been given for single response data, but in most applications several dependent variables have to be considered simultaneously [10]. Multi-response data are familiar outputs of experiments and processes involving multicomponent mixtures, multiple streams or multiple methods of observation. The errors in these 6.1.4.6

Define experimental region (grid) Initial experiments: Factorial design

Parameter estimation of the rival models that are left

Discrimination criterion: One model left?

Y

Parameters of sufficient accuracy?

N

Y

STOP

N

Design criterion: Maximize divergence Search for grid point

Design criterion: Apply divergence criterion for parameter improvement Search for grid point

Perform experiment Add result to data set

Fig. 10 Flowsheet of sequential experimental design in the process of model selection, followed by parameter improvement.

outputs will almost always contain interdependencies, which should be in some way taken into account in the objective function for minimization for parameter estimation and experimental design. Most ways of incorporating the interdependencies are based on the following assumptions, applicable to the case that in each of the n experiments v dependent variables are determined having the associated error εij (i = 1, . . . , n, j = 1, . . . , v): • All errors are normally distributed with zero mean. • The n errors corresponding to the n observations of the j th response are statistically independent and have a constant variance σjj . • The v errors within the ith experiment are statistically interdependent with a covariance matrix . If the errors are independent, this matrix reduces to a diagonal matrix. • The covariance matrix  is the same for all n experiments. Depending on the application and the knowledge concerning the interdependencies, there are various ways to incorporate the interdependency of the errors. If the covariance matrix  is known, the following objective function can be used: v v

k=1 l=1

σ kl

n

(yik − ηik )(yil − ηil )

(83)

i=1

where σ kl are elements of the inverted matrix  −1 . Unfortunately, in most cases little knowledge exists concerning the interdependencies of the errors. Therefore, one often ignored these interdependencies. However, since this will often result in inferior parameter estimates, more sophisticated algorithms have been developed that extract the interdependencies from the experimental data in combination with the model to be fitted [80–82]. These algorithms are based on a Bayesian estimation [83, 84]. These algorithms also contain features to deal with missing data, e.g. as a result of analyses that were only partly successful. 6.1.5

Computer Codes and Software Packages

For more than a decade, minimization techniques have been well implemented in academic computer codes [61, 85]. Although many companies and universities have developed their own in-house codes for these purposes, there are nowadays also many commercially available software packages containing capabilities for parameter estimation and a few of these also contain features for sequential experimental design. Most of

List of Symbols

the packages are supplied with user-friendly graphical user interfaces, which facilitate use by non-experienced researchers and produce much statistical information concerning the estimates obtained. They can be found in several process modeling packages and in advanced programming languages/platforms [86]. Examples are Aspen Custom Modeler [87], Athena Visual Studio [88], gPROMS [89], Matlab [90], Mathematica [91], Parametra [92] and Presto-Kinetics [93].

b CC d.f . D(xk )

Eai F (p, n − p, 1 − α)

6.1.6

Conclusion

In this chapter, focus has been given to the derivation and application of the LHHW models to describe catalyzed reactions. In spite of objections that can be made against their underlying assumptions, these expressions have been proven to be fairly successful. An important reason is probably the fact that it contains intrinsically a capacity, a limited number of catalytically active centers where the reaction takes place and which are distributed among the different adsorbing components. A completely similar approach applies to biocatalytic reactions, indicated here by Michaelis–Menten kinetics. Detailed knowledge of the reaction kinetics is important, particularly if high conversions and/or high selectivities are to be achieved in industrial reactors. Recycles are often present in processes and some components will accumulate. The effect of such components on activity and selectivity should be carefully taken into account, otherwise reactor sizes or deactivation rates may be tremendously underestimated. Kinetic studies are laborious, the most time-consuming part nowadays being the experimental work. Due to the increasing availability and user-friendliness of software packages for doing the parameter estimation, the required modeling effort is reduced. A significant reduction in the costly experimental activities may be achieved by applying optimum design of the experiments. Although the mathematical expressions are relatively simple and can be easily incorporated in parameter estimation programs, their application is still not frequently encountered in the literature. However, this may change with the increasing availability of software capable of doing sequential experimental design and the continuous effort to increase efficiency. List of Symbols

A ak

bi

component or species step length improvement factor at kth parameter iteration (steepest descent) estimated value parameter i

– –

–

FA0 hij H Hi I ki k0i ki0 kref Ki Keq L L(β|y) m NA NT n p pA pi r ri R Resi

Res s2 SSR S(b) Si

a

vector of estimated parameters coke content of catalyst number of degrees of freedom divergence between rival models at experimental settings vector xk activation energy of step i F -value of Fischer’s distribution at p and n – p degrees of freedom and confidence level (1 − α) molar flow of A element of Hessian matrix Hessian matrix enthalpy change for step or process i unity matrix rate constant of step or process i pre-exponential factor of ki rate constant value without deactivation rate constant defined through Eq. (65) equilibrium constant of step or process i overall equilibrium constant of a reaction site density (Table 2) likelihood function of parameter vector β in model y number of rival models amount of A site concentration number of experiments number of parameters partial pressure of A probability density under model i, Eq. (77) reaction rate reaction rate of step i universal gas constant residual, observed minus calculated value of dependent variable at ith experiment vector of residuals estimated variance sum of squares of residuals SSR function as a function of parameter vector b entropy change for step or process i

References see page 1712

1711

mol kg−1 – – a J mol−1 –

mol s−1 – a J mol−1 – – – –

a a a

–

a

–

a

–

a

nm−2 – – mol mol kg−1 – – Pa – mol s−1 kg−1 mol s−1 kg−1 J mol−1 K−1 – a

– –

a a

J mol−1 K−1

1712

6.1 Rate Procurement and Kinetic Modeling

t time tn−p,1−α/2 t-value of Student’s distribution at n – p degrees of freedom and (1 – α) probability level T temperature reference temperature within Tref the experimental range W amount of catalyst element i, j of variance matrix vij V(b) variance–covariance matrix of parameter estimates b V [y(x ˆ k )] variance of predicted value of y at experimental settings vector xk of grid point k xA conversion of A xi setting of independent variable i xk vector of independent variables at the kth grid point X matrix of independent variable settings of experiments value of dependent (observed) yi variable of ith experiment y vector of dependent variable values eigenvector coordinate zp α deactivation constant or α-percentage point in F - or t-distribution real (unknown) value of βi parameter i β vector of real (unknown) values of parameters δ constant εij error of the j th dependent (observed) variable at the ith response (experiment) γ heteroscedasticity of the variance ηi real (unknown) value of dependent variable yi πi,n probability of model i after n experiments θA surface occupancy of species A  hyperellipsoid function (confidence contour) Wilks’ diagnostic parameter in λi model discrimination λk Levenberg–Marquardt parameter after the kth iteration the pth eigenvalue λp  catalyst deactivation function (0 ≤  ≤ 1)

s

C – SS

K K kg –

ρij σ2  a

νA –

a

χC2 ∂

– –

∧ a

–

–

a

s−1 – –

a

– –

a

– – – – –

– –

– – – –

a

– – – –

a The

dimension of this variable depends on the systems considered.

References

–

– –

catalyst deactivation function due to C catalyst deactivation function at steady state correlation coefficient between parameters i and j real variance of a model matrix of error variances in multi-response experiments stoichiometric number of A in reaction function defined for model discrimination, Eq. (71) operator for partial derivatives referring to calculated value

a

1. A. N. R. Bos, L. Lefferts, G. B. Marin, M. H. G. M. Steijns, Appl. Catal. A 1997, 160, 185. 2. R. J. Berger, E. H. Stitt, G. B. Marin, F. Kapteijn, J. A. Moulijn, CatTech 2001, 5, 30. 3. M. Boudart, G. Dj´ega-Mariadassou, Kinetics of Heterogeneous Catalytic Reactions, Princeton University Press, Princeton, NJ, 1984, 222 pp. 4. M. Boudart, in Handbook of Heterogeneous Catalysis, Vol. 3, G. Ertl, H. Kn¨ozinger, J. Weitkamp (Eds.), VCH, Weinheim, 1997, p. 958. 5. M. Boudart, Kinetics of Chemical Processes, ButterworthHeinemann, Boston, 1991, 242 pp. 6. E. G. Christoffel, Laboratory Studies of Heterogeneous Catalytic Processes, Elsevier, Amsterdam, 1989, 263 pp. 7. E. Alpay, L. S. Kershenbaum, N. F. Kirkby, Chem. Eng. Sci. 1995, 50, 1063. 8. L. K. Doraiswamy, D. G. Tajbl, Catal. Rev. Sci. Eng. 1974, 10, 177. 9. N. Draper, H. Smith, Applied Regression Analysis, Wiley, New York, 1981, pp. 709. 10. G. F. Froment, L. Hosten, in Catalysis; Science and Technology, J. R. Anderson, M. Boudart (Eds.), Springer-Verlag, Berlin, 1981, p. 97. 11. J. R. Kittrell, in Advances in Chemical Engineering, T. B. Drew, G. R. Cokelet, J. W. Hoopes, T. Vermeulen (Eds.), Vol. 8, Academic Press, New York, 1970, p. 97. 12. K. C. Pratt, in Catalysis; Science and Technology, J. R. Anderson, M. Boudart (Eds.), Springer, Berlin, 1980, p. 174. 13. J. R. Anderson, K. C. Pratt, Introduction to Characterization and Testing of Catalysts, Academic Press, Sydney, 1985, 457 pp. 14. S. W. Weller, Catal. Rev. Sci. Eng. 1992, 34, 227. 15. M. I. Temkin, Adv. Catal. 1979, 28, 173. 16. G. F. Froment, K. B. Bischoff, Chemical Reactor Analysis and Design, Wiley, New York, 1979, 765 pp. 17. J. T. Gleaves, J. R. Ebner, T. C. Kuechler, Catal. Rev. Sci. Eng 1988, 30, 49. 18. J. T. Gleaves, G. S. Yablonsky, P. Phanawadee, Appl. Catal. A 1997, 160, 55. 19. S. O. Shekhtman, G. S. Yablonsky, S. Chen, J. T. Gleaves, Chem. Eng. Sci. 1999, 54, 4371.

References 20. M. Soick, D. Wolf, M. Baerns, Chem. Eng. Sci. 2000, 55, 2875. 21. S. C. van der Linde, T. A. Nijhuis, F. H. M. Dekker, F. Kapteijn, J. A. Moulijn, Appl. Catal. A 1997, 151, 27. 22. K. A. Vonkeman, Exhaust Catalysis Studies Using In-Situ Positron Emission, Ph.D Thesis, Eindhoven University of Technology, 1990. 23. F. H. M. Dekker, Transient Kinetics in Heterogeneous Catalysis, PhD Thesis, University of Amsterdam, 1995. 24. J. Happel, Isotopic Assessment of Heterogeneous Catalysis, Academic Press, Orlando, FL, 1986, 196 pp. 25. F. H. M. Dekker, J. G. Nazloomian, A. Bliek, F. Kapteijn, J. A. Moulijn, P. L. Mills, J. J. Lerou, Appl. Catal. A 1997, 151, 247. 26. C. Mirodatos, Catal. Today 1991, 9, 83. 27. R. H. Nibbelke, J. Scheerova, M. H. J. M. de Croon, G. B. Marin, J. Catal. 1995, 156, 106. 28. H. A. J. van Dijk, J. H. B. J. Hoebink, J. C. Schouten, Top. Catal. 2003, 26, 111. 29. B. W. Wojciechowski, N. M. Rice, Chem. Eng. Sci. 1993, 48, 2881. 30. B. W. Wojciechowski, Catal. Today 1997, 36, 167. 31. M. Kolkowski, J. Malachowski, F. J. Keil, C. Liebner, D. Wolf, M. Baerns, Chem. Eng. Sci. 2003, 58, 4903. 32. J. A. Moulijn, J. P´erez-Ram´ırez, R. J. Berger, G. M. Hamminga, G. Mul, F. Kapteijn, Catal. Today 2003, 81, 457. 33. R. Hoogenboom, M. A. R. Meier, U. S. Schubert, Macromol. Rapid Commun. 2003, 24, 16. 34. U. S. Schubert, E. J. Amis, Macromol. Rapid Commun. 2004, 25, 19. 35. M. A. R. Meier, R. Hoogenboom, U. S. Schubert, Macromol. Rapid Commun. 2004, 25, 21. 36. R. Mezaki, H. Inoue, Rate Equations of Solid-Catalysed Reactions, University of Tokyo Press, Tokyo, 1991, 399 pp. 37. O. A. Hougen, K. M. Watson, Chemical Process Principles, Vol. III, Wiley, New York, 1947, p. 1. 38. A. J. J. Straathof, J. Mol. Catal. B 2001, 11, 991. 39. A. J. J. Straathof, Software Code Encora 1.2, 2005; freely available from http://www.bt.tudelft.nl/live/pagina.jsp?id= ca553f24-64c7-4d5e-a344-ce30830894ab&lang=en. 40. F. Kapteijn, J. A. Moulijn, R. A. van Santen, R. Wever, in Catalysis: an Integrated Approach R. A. van Santen, P. W. N. M. van Leeuwen, J. A. Moulijn, B. A. Averill (Eds.), Studies in Surface Science and Catalysis, Vol. 123, Elsevier, Amsterdam, 1999, p. 81. 41. F. Kapteijn, J. J. Rodriguez-Mirasol, J. A. Moulijn, Appl. Catal. B 1996, 9, 25. 42. F. Kapteijn, J. A. Moulijn, R. A. van Santen, in Catalysis: an Integrated Approach to Homogeneous, Heterogeneous and Industrial Catalysis, J. A. Moulijn, P. W. N. M. van Leeuwen, R. A. van Santen (Eds.), Elsevier, Amsterdam, 1993, p. 69. 43. A. Corma, F. Llopis, J. B. Montonand, S. W. Weller, Chem. Eng. Sci. 1988, 43, 785. 44. J. Wei, Ind. Eng. Chem. Res. 1994, 33, 2467. 45. G. G. Martens, J. W. Thybaut, G. B. Marin, Ind. Eng. Chem. Res. 2001, 40, 1832. 46. G. Froment, Catal. Rev. Sci. Eng. 2005, 47, 83. 47. S. P. Gurden, J. A. Westerhuis, S. Bijlsma, A. K. Smilde, J. Chemom. 2001, 15, 101. 48. H. J. Ramaker, E. N. M. van Sprang, S. P. Gurden, J. A. Westerhuis, A. K. Smilde, J. Proc. Control 2002, 12, 569. 49. O. Levenspiel, Chemical Reaction Engineering, 2nd Ed., Wiley, New York, 1972, p. 537. 50. O. Levenspiel, J. Catal. 1972, 24, 265. 51. J. P. Janssens, R. M. de Deugd, A. D. van Langeveld, S. T. Sie, J. A. Moulijn, in Catalyst Deactivation 1997, C. H.

52. 53. 54. 55. 56. 57. 58. 59. 60. 61.

62. 63. 64.

65. 66.

67. 68. 69. 70. 71. 72. 73. 74. 75. 76. 77. 78. 79. 80. 81. 82. 83.

84. 85. 86.

87. 88.

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Bartholomew, G. A. Fuentes (Eds.), Studies in Surface Science and Catalysis, Vol. 111, Elsevier, Amsterdam, 1997, p. 283. G. A. Fuentes, Appl. Catal. 1985, 15, 33. Y. Bard, Nonlinear Parameter Estimation, Academic Press, New York, 1974, 341 pp. D. M. Himmelblau, Applied Nonlinear Programming, McGrawHill, New York, 1972, 498 pp. T. F. Edgar, D. M. Himmelblau, Optimization of Chemical Processes, McGraw-Hill, New York, 1988, 652 pp. M. J. D. Powell, Comput. J. 1964, 7, 155. H. H. Rosenbrock, C. Storey, Computational Techniques for Chemical Engineers, Pergamon Press, New York, 1966, 328 pp. J. A. Nelder, R. Mead, Comput. J. 1965, 7, 308. K. Levenberg, Q. Appl. Math. 1944, 2, 164. D. Marquardt, J. Soc. Ind. Appl. Math. 1963, 11, 431. W. H. Press, S. A. Teukolsky, W. T. Vetterling, B. P. Flannery, Numerical Recipes; the Art of Scientific Computing, Cambridge University Press, Cambridge, 1992, 994 pp. R. Bulirsch, J. Stoer, Num. Math. 1966, 8, 1. C. W. Gear, L. R. Petzold, SIAM J. Numer. Anal. 1984, 21, 367. R. L. C. Bonn´e, Hydrodemetallization of Ni-TPP and VO-TPP over Sulfided Molybdenum and Vanadium Catalysts, PhD Thesis, University of Amsterdam, 1990. W. E. Schiesser, The Numerical Method of Lines, Academic Press, San Diego, 1991, 326 pp. A. C. Hindmarsh, in Scientific Computing, 10th IMACS World Congress on Systems Simulation and Scientific Computation, Montreal, Canada, 8–13 August 1982, R. S. Stepleman, M. Carver (Eds.), NHPC, Amsterdam, 1983, p. 55. F. Kapteijn, L. H. G. Bredt, E. Homburg, J. C. Mol, Ind. Eng. Chem. Prod. Res. Dev. 1981, 20, 457. R. A. van Santen, J. W. Niemantsverdriet, Chemical Kinetics and Catalysis, Plenum Press, New York, 1995, 280 pp. M. A. Vannice, S. H. Hyum, B. Kalpacki, W. C. Liauh, J. Catal. 1979, 56, 358. M. Boudart, D. E. Mears, M. A. Vannice, Ind. Chim. Belge 1967, 32, 281. M. Boudart, AIChE J. 1972, 18, 465. F. Kapteijn, R. Meijer, J. A. Moulijn, Energy Fuels 1992, 6, 494. P. Biloen, J. Mol. Catal. 1983, 21, 17. R. W. Maatman, Adv. Catal. 1980, 29, 97. R. Mezaki and J. R. Kittrell, Can. J. Chem. Eng. 1966, 44, 285. S. S. Wilks, Ann. Math. Stat. 1946, 17, 257. M. S. Bartlett, Proc. R. Soc. London, Ser. A 1937, 160, 268. G. E. P. Box, W. J. Hill, Technometrics 1967, 9, 57. B. H. Chen, S. P. Asprey, Ind. Eng. Chem. Res. 2003, 42, 1379. W. E. Stewart, M. Caracotsios, J. P. Sorensen, AIChE J. 1992, 38, 641. W. E. Stewart, M. Caracotsios, J. P. Sorensen, AIChE J. 1992, 38, 1302. W. E. Stewart, Y. Shon, G. E. P. Box, AIChE J. 1998, 44, 1404. G. E. P. Box, G. C. Tiao, Bayesian Inference in Statistical Analysis, Addison-Wesley, Reading, MA, 1973; reprinted Wiley, New York, 1992, 588 pp. G. E. P. Box, N. R. Draper, Biometrika 1965, 52, 355. Netlib, http://www.netlib.org, 2005. R. J. Berger, J. Hoorn, J. Verstraete, J. W. Verwijs, in Reaction Kinetics and the Development of Catalytic Processes, G. F. Froment, K. C. Waugh (Eds.), Studies in Surface Science and Catalysis, Vol. 133, Elsevier, Amsterdam, 2001, p. 631. Aspen Technology, Aspen Custom Modeler, http://www.aspen tech.com/product.cfm?ProductID=54; accessed January 2006. Athena Visual Software, Athena Visual Studio, http:// athenavisual.com; accessed January 2006.

1714

6.2 Determination of Diffusion Coefficients in Porous Media

89. Process Systems Enterprise, gPROMS, http://www. psenterprise.com/gproms/index.html; accessed January 2006. 90. The Mathworks, Matlab, http://www.mathworks.com; accessed January 2006. 91. Wolfram Research, Mathematica, http://wolfram.com; accessed January 2006. 92. Hydromantis, Parametra, http://www.hydromantis.com/ software05.html; accessed January 2006. 93. Computing in Technology, Presto-Kinetics, http://www.citwulkow.de/tbbpres.htm; accessed January 2006.

6.2

Determination of Diffusion Coefficients in Porous Media .. .. Jorg Karger∗

6.2.1

Definitions

Diffusion is the process of molecular transport associated with the stochastic movement of the individual diffusants [1–4]. Diffusion coefficients in porous media may be defined on the basis of the generalized Fick’s first law: jz,i = −

j

Dij

∂cj ∂z

(1)

where jz,i and ci denote, respectively, the flux density in the z direction and the concentration of the ith component. The parameters Dij denote the generalized diffusion coefficients (or diffusivities). They depend on the given porous medium and the temperature, as well as on the diffusants involved and their concentrations. In using Eq. (1), the porous medium is considered as quasi-homogeneous. Concentrations and fluxes must be defined, therefore, with respect to sufficiently large volume elements and areas, so that the averages over all volume elements and areas are identical [5]. In the case of intracrystalline diffusion in porous crystallites, this is ensured by considering volume elements substantially exceeding the unit cell dimensions. In amorphous materials and beds of porous crystallites, much larger volume elements must be considered. The structure of such volume elements can clearly only be compared in a statistical sense. Diffusivities defined on the basis of Eq. (1) for inhomogeneous systems are sometimes called ‘‘apparent’’ diffusivities [6], since with this definition the system under study appears to be homogeneous with respect to its transport properties, even though it is not really homogeneous. ∗

Corresponding author.

Equation (1) says nothing about the underlying mechanisms, nor does it imply anything about the dependence of the diffusivities on the concentration of the relevant components. For diffusivities, which vary considerably with concentration, the concentration range covered within a sample during diffusivity measurement should be kept as small as possible. Otherwise only ‘‘integral’’ diffusivities may be determined, from which it is difficult to extract the ‘‘differential’’ diffusivities (see Ref. [7], pp. 279–282). In a complex system such as a catalyst particle, a variety of processes may contribute to the diffusion fluxes [5–8] (see Chapter 6.3). In general, the diffusants within the pore system of a catalyst particle co-exist in two phases, a gas phase and an adsorbed phase. In the vicinity of the external surface, the gas-phase concentration in the catalyst particle is determined by that of the surrounding atmosphere. If there are no additional resistances (e.g. film resistances), both concentrations are identical. Since the external gas-phase concentration is easily accessible experimentally, molecular fluxes through the catalyst particle are often related to the gas-phase concentration, even though this is at variance with the understanding of Fick’s first law as formulated by Eq. (1). The ‘‘effective’’ diffusivities [6] defined in this way are clearly different from those introduced by Eq. (1) [5]. The definition of effective diffusivities from Fick’s laws evokes divergences in the procedure to determine tortuosities, including the risk of misunderstandings. If, following Fick’s laws, diffusivities are referred to the concentration of the diffusants in the system under consideration, the tortuosity simply results as the ratio between the coefficients of self-diffusion in the unrestricted (free) fluid, i.e. in the reference system, Dref , and of the fluid in the system under consideration, D. This is the procedure how, by e.g. PFG NMR [9–13], the tortuosities of porous catalysts may be determined in a most efficient way. If, however, the catalysts under study are described by an effective diffusivity, i.e. by referring their intrinsic transport properties to external concentrations, rather than by using their genuine, intrinsic diffusivities, the ratio of the diffusivities must be additionally multiplied by the porosity (implying equal concentration of the fluid inside and outside the pores) [5]. Depending on the conditions of the diffusion experiment, the diffusivities defined by Eq. (1) are given particular names. For one-component diffusion under the influence of a concentration gradient, Eq. (1) reduces to jz = −D

dc dz

(2)

where D denotes the transport diffusivity. Single-component diffusion under equilibrium conditions can be monitored either by labeling some of the

6.2.2 Measurement of Transport Diffusion

molecules or by following their trajectories. Considering the diffusion flux of the labeled molecules, again a proportionality relation of the type of Eq. (2) may be established. The factor of proportionality is called the coefficient of selfdiffusion (or tracer diffusion). In a completely equivalent way (see Ref. [7], pp. 23–30), the self-diffusion coefficient may be determined on the basis of Einstein’s relation z2 (t) = 2Dt

(3)

as the factor of proportionality between the mean-square displacement of the diffusants and the observation time, provided that, for sufficiently large observation times, this proportionality is established (see Chapter 5.5.2). Self-diffusion under equilibrium conditions may also be monitored in multicomponent systems. Again, with Eqs. (2) and (3) a self-diffusion coefficient (of a particular component) may be defined. This coefficient depends on the nature and the concentration of all molecular species involved as well as on the nature of the catalyst particle. Mass transfer in catalysis proceeds under nonequilibrium conditions with at least two molecular species (the reactant and product molecules) involved [14–16]. Under steady-state conditions, the flux of the product molecules out of the catalyst particle is stoichiometrically equivalent (but in the opposite direction) to the flux of the entering reactant species. The process of diffusion of two different molecular species with concentration gradients opposed to each other is called counter-diffusion, and if the stoichiometry is 1 : 1 we have equimolar counterdiffusion. The situation is then similar to that considered in the case of self or tracer diffusion, the only difference being that now two different molecular species are involved. Tracer diffusion may be considered, therefore, as equimolar counter-diffusion of two identical species. The process of diffusion of (generally two) components with concentration gradients pointing in the same direction is called codiffusion. According to Eq. (1), the flux density of each individual component may depend on the concentration gradients of both species. The factors of proportionality are called coefficients of codiffusion. Obviously, any component in a two-component system is described by two coefficients of codiffusion, relating the flux of this component to the concentration gradient of the same component and of the other component, respectively. The diffusivities depend, in general, on the concentrations of both components. 6.2.2

Measurement of Transport Diffusion Steady-State Measurements Steady-state experiments have their origin in Darcy’s classical permeation measurements [6, 7]. In a typical 6.2.2.1

1715

experimental set-up a pressure drop is imposed across a porous plug (often a single cylindrical pellet) of catalyst by maintaining a constant (known) pressure on one side while the pressure on the other side is maintained at a lower constant value or allowed to increase with passage of the diffusant through the plug. In the latter case, monitoring the downstream pressure provides a simple and convenient way of measuring the flux. More recently, this technique has even been successfully applied to transport diffusion studies in zeolite crystallites by incorporating them into impermeable membranes [17–21]. As a consequence of the cost of preparing such membranes, the most recent activities focus on the fabrication of membranes composed of randomly intergrown zeolite crystals which may be prepared by direct in situ zeolite synthesis on appropriately chosen supports [20, 22–24]. In general, however, the method suffers from the disadvantage that the permeability rather than the diffusivity is measured. The two quantities are directly correlated only in the regimes of intracrystalline and Knudsen diffusion [6, 8, 25, 26]. This limitation is overcome in the Wicke–Kallenbach method [27]. This method differs from the original permeability experiment in that, by applying a carrier gas, the flux is measured under constant pressure conditions with a known concentration difference maintained across the membrane or pellet. In this way, any non-diffusive contribution (Poiseuille flow) to the observed fluxes may be excluded. Remarkably, in the procedure of ‘‘effectiveness factor analysis’’, the phenomenon of catalysis itself may be exploited to gain information about intrinsic diffusivities under steady-state conditions [28–31]. This method relies on intentional modification of the Thiele modulus √ φ = R k/D (with R, k and D denoting, respectively, the radius of the catalyst particle, the intrinsic rate constant and the diffusivity to be determined) and monitoring the effect on the measured rate of reaction. In the intermediate range between the limits of sufficiently fast diffusion (φ  1, no influence of diffusion) and slow diffusion (φ 1, measured rate constant increases proportional to the square root of the diffusivity), the observed dependence may directly yield the respective diffusivities [28–31] (see Section 6.2.4). Time Lag Measurements Information about transport diffusion in catalyst particles can also be deduced during the initial, unsteady-state period of a permeation experiment. In this stage, the number of molecules passing the plug of catalyst per unit time will increase from zero until the rate of permeation 6.2.2.2

References see page 1723

1716

6.2 Determination of Diffusion Coefficients in Porous Media

characterizing the steady state behavior is attained. In the limit t → ∞, the total amount of molecules which have permeated in the time interval 0 . . . t is given by the relation [6, 7, 32] Q(t −−−→ ∞) =

  L2 cD × t− L 6D

(4)

with c, D, and L denoting, respectively, the concentration and diffusivity of the diffusants within the porous system under study and its extension in the direction of the concentration gradient. The time intercept of a plot of Q(t → ∞) versus time gives a time lag τ=

L2 6D

(5)

which may be used for the determination of the transport diffusivity. Procedures to remove the restrictions of the permeation technique, also inherent in the time lag method, have been described by Grachev et al. [33] and Gibilaro et al. [34]. As with the Wicke–Kallenbach method, they are based on the application of a carrier gas. Details of these methods may be found in Ref. [6]. Sorption Rate Measurements The most widely used unsteady-state method for determining diffusivities in porous solids involves measuring the rate of adsorption or desorption when the sample is subjected to a well-defined change in the concentration or pressure of sorbate. The experimental methods differ mainly in the choice of the initial and boundary conditions and the means by which progress towards the new position of equilibrium is followed. The diffusivities are found by matching the experimental transient sorption curve to the solution of Fick’s second law. Detailed presentations of the relevant equations may be found in the literature [6, 7, 12, 35–38]. For spherical particles of radius R, for example, the fractional uptake after a pressure step obeys the relation 6.2.2.3

γ (t) =

m(t) − m(0) m(∞) − m(0)

 ∞ 6 1 2 2 Dt × exp −n π 2 =1− 2 π n2 R

(6)

n=1

where the concentration in the surrounding atmosphere after the pressure step is assumed to be constant. This is a valid approximation if the system volume is large and the step is small enough so that the diffusivity within the covered range of concentrations remains constant. The series converges rapidly for large values of time, and for

γ > 0.7 it is sufficient to retain the first term only:  Dt 6 γ (t) ≈ 1 − 2 × exp −π 2 2 π R

(7)

Alternatively, in the short-time range √ [γ (t) < 0.3], the uptake may be approximated by the t law  Dt γ (t) ≈ 6 (8) πR 2 For sufficiently short times, Eq. (8) holds for any particle size. In this case, R is understood as an equivalent radius being equal to three times the ratio between the particle volume and the external surface. If the uptake measurements are carried out in a small volume, the concentration in the surrounding atmosphere will decrase after the pressure step. Equation (6) is then replaced by γ (t) = 1 −

∞

6λ(λ + 1)

9 + 9λ + qn 2 k2  Dt × exp qn 2 2 R n−1

(9)

with λ = p(∞)/[p(0) − p(∞)] and qn denoting the nonzero roots of tan qn =

3qn 3 + λqn 2

For sufficiently short times, the applied, with an additional factor

(10) √

t law may again be

p(0) − p(−0) p(∞) − p(−0) on the right-hand side of Eq. (8). The quantities p(−0), p(0) and p(∞) denote, respectively, the sorbate pressure (and hence the concentration) in the surrounding atmosphere before the uptake experiment, at the beginning of the uptake experiment, and after attaining the new equilibrium value. Exactly the same formalism may be used for considering diffusion in biporous catalyst particles, if sufficiently fast molecular exchange between the various ranges of microporosity may be implied [5]. In this case, Eqs. (6)–(9) may be preserved with the understanding that D has to be replaced by the expression Dinter εp /[εp + (1 − εp )K], with εp and Dinter denoting, respectively, the macroporosity and the diffusivity in the macropores and with the equilibrium constant K between the gaseous and adsorbate phases. Conventionally, molecular uptake is recorded gravimetrically [39–41]. A particularly charming version of gravimetric measurement is brought about by the ‘‘tapered element oscillating microbalance’’ (TEOM) [42–44],

6.2.2 Measurement of Transport Diffusion

where mass changes are recorded by following the corresponding changes in the frequency of the mechanical oscillation of the sample. Alternatively, for a limited supply of adsorbate, molecular uptake may also be calculated from a knowledge of the time dependence of the pressure (piezometric method [45, 46]) or composition of the gas phase. Changing the sorbate pressure by a step change of the gas volume has proved to be a very efficient method for following fast sorption processes (single-step method [47, 48]). The sorption uptake may also be measured volumetrically by means of a gas burette arrangement [49]. In principle, the determination of molecular uptake may be based on any experimentally accessible quantity which is a function of the amount adsorbed. Being directly sensitive to a certain molecular species, in this respect the application of spectroscopic methods is particularly suitable. IR spectroscopy has been successfully applied to studying molecular uptake by beds of zeolite catalysts [50–52] as well as–in combination with IR microscopy [53–57]–on individual crystallites. Similarly, NMR spectroscopy has also been used to monitor the time dependence of the sorbate concentration within porous media [58]. Moreover, recent progress in NMR imaging allows the observation of concentration profiles within porous media with spatial resolution below the millimeter region [59–69]. Similarly, positron emission profiling (PET) [70] may also be applied to in situ observation. In this technique, information about concentration profiles is gained by injecting small amounts of radiochemically labeled molecules marked with positron-emitting isotopes. Unambiguous information about molecular transport may clearly only be deduced from sorption experiments if the applied model correctly reflects the multitude of processes accompanying molecular uptake. In this way, also by exploiting the option of ‘‘temporal analysis of products (TAP)’’ reactor systems [42, 71–73], estimates of the different rates of diffusion in complex systems have been given. The validity of such model assumptions [74, 75] should be checked by varying a characteristic parameter of the system, for example the crystallite size for the study of intracrystalline diffusion in beds of crystallites [76, 77]. As a necessary condition for the validity of the model, the measured dependency must coincide with the theoretical prediction. Among the external processes possibly influencing the rate of overall adsorption, the access of the diffusants into the sorption vessel and the dissipation of the adsorption heat deserve special attention. A substantial number of models have been developed to quantify these influences, which are generally referred to as the valve effect [78–80] and the heat effect [7, 81–83]. In turn, in Ref. [84] a novel method for uptake measurements has

1717

been based on the heat effect during molecular sorption. By IR monitoring of the surface temperature it has become possible to acquire a second, independent source of information about the internal processes within the sample, yielding useful information in particular for fast processes. (see Section 6.2.2.4). A decisive breakthrough in unveiling the controlling steps during molecular uptake and release by nanoporous catalysts has recently been achieved by developing an option for monitoring transient intracrystalline concentration profiles by means of interference microscopy, as explained in more detail in Section 6.2.2.6. Frequency Response Measurements As an alternative approach to conventional uptake measurements, in the frequency response technique [85–90] one follows the response of the sample to a regular periodic perturbation, e.g. a sinusoidal variation of the system volume. Using complex notation, one may write for the time dependence of the system volume 6.2.2.4

V (t) = V0 (1 − v exp{iωt})

(11)

For sufficiently small perturbations, both the induced pressure variation P (t) and the amount adsorbed A(t) are also sinusoidal functions of time: P (t) = P0 {1 + p exp[i(ωt + ϕ)]}

(12)

A(t) = A0 {1 + a exp[i(ωt + φ)]}

(13)

The pressure variation, P (t) = P0 p exp(iωt) × exp(iϕ), and the variation of the amount adsorbed, A(t) = A0 a exp(iωt) × exp(iφ), must be correlated, therefore, by a proportionality relation of the type A(t) = (αc − iαs )

V0 P (t) RT

(14)

with αc and αs representing the real and imaginary parts of the complex factor of proportionality. They may be interpreted, therefore, as the in-phase and out-of-phase components of the adsorbed species with respect to the pressure variation and are commonly called the inphase and out-of-phase characteristic functions. By using the equation of material balance, d/dt[P (t)V (t)/RT + A(t)] = 0, the in-phase and out-of-phase characteristic functions may be easily shown to be determined through the experimentally accessible quantities p and ϕ by the relationships v cos ϕ − 1 p v αs = sin ϕ p

αc =

References see page 1723

(15) (16)

1718

6.2 Determination of Diffusion Coefficients in Porous Media

Theoretical expressions for the frequency dependence of the characteristic functions may be calculated by solving the relevant diffusion equation under oscillating boundary conditions [85–89]. The diffusivities are determined by matching the experimental curves to the theoretical expressions for the given model. As a rule of thumb, the out-of-phase characteristic function may be expected to pass through a maximum at ω ≈ D/ l 2 (where l is the characteristic diffusion length for the particles under study), approaching zero for frequencies both much larger and much smaller than this ‘‘resonance’’ frequency. A fundamental advantage of the frequency response method is its ability to yield information concerning the distribution of molecular mobilities. For example, a bimodal distribution of diffusivites, which is difficult to detect by conventional sorption measurements, leads to two different resonances [91]. Moreover, from an analysis of the frequency response spectrum it is even possible to monitor molecular diffusion in combination with chemical reactions [86]. Thus, in addition to a large number of investigations with monodisperse microporous materials [88, 89, 92–95], detailed studies concerning more complex situations such as bidisperse porous media [96], acid site distributions [97], isomerization reactions [98] and carbon nanotubes [99] have also been performed. As in conventional sorption experiments, however, the intrusion of heat effects limits the information provided by this technique for fast adsorption–desorption processes [100], leading to a bimodal form of the frequency response characteristic curves [90, 101]. This limitation may be circumvented in the thermal frequency response method [102–105]. By applying the IR temperature monitoring technique [84], in this method one is able to record the temperature response of the sample as a consequence of the adsorption heat released or consumed during the adsorption–desorption process. Thus, in addition to the pressure response a second, independent quantity becomes experimentally accessible. Since the sample temperature is influenced in a completely different way by the rates of molecular uptake and heat release, any confusion between these two influences is excluded. Moreover, by referring the temperature response to the variation of the pressure rather than to the variation of the volume, the influence of disturbing effects like adsorption by the vessel walls or temperature changes accompanying the changes of the volume may be excluded. Chromatographic and Flow Methods The influence of heat effects and of external mass transfer resistances (film resistance) may be significantly reduced by the application of flow methods. In the usual chromatographic experiment, a steady flow of 6.2.2.5

an inert (non-adsorbing) carrier is passed through a small column packed with the catalyst under study. A small pulse of sorbate is injected at the column inlet, and the effluent concentration is monitored continuously. While the retention volume is a measure of the adsorption equilibrium, the dispersion of the response peak is determined by mass transfer resistance and axial mixing within the column. A particularly convenient way of analyzing the experimental data may be based on the moments of the response curves of the chromatographic column. The first and second moments, µ and σ 2 , are defined by the relationships  ∞ ct dt µ ≡ t¯ = 0 ∞ and c dt 0

 2

σ =

∞ 0



c(t − µ)2 (17)

∞

c dt 0

for the pulse response, and by  ∞ c 1− dt and µ ≡ t¯ = c0 0  ∞ c 2 σ = 1− t dt − µ2 c0 0

(18)

for the step response, where c denotes the concentration of the adsorbate in the effluent. For the step response, c approaches the equilibrium value c0 . For a biporous adsorbent model [6–8], when the equilibrium constant K is large, the second moment is given by the relation
 Rp Rp 2 DL Rc 2 εv σ2 = + + + 2µ2 vL L(1 − ε) 3 kf 15εp Dp 15KDc (19) with the following notations: Dc and Dp are the diffusivities in the micro- and macroparticles with the respective radii Rc and Rp ; ε is the external void fraction of the adsorbent bed; εp the porosity of the adsorbent particle; DL the axial dispersion coefficient; L the column length; v the interstitial fluid velocity; kf the external mass transfer coefficient of the macroparticle; K the dimensionless Henry’s law constant. Equation (18) includes the important result that the second moment is simply the sum of the moments of the different transport resistances (see Chapter 6.1). The method of moments has also been successfully applied to the analysis of sorption rate [32, 106] and steady-state [107] measurements. It follows from Eq. (19) that by varying

6.2.2 Measurement of Transport Diffusion

characteristic parameters of the experiment such as the flow rate v or the size Rp of the macroparticles, the contributions of the different mechanisms may, in principle, be discriminated. In reality, however, various deviations from the ideal behavior which was assumed in the derivation of Eq. (19) may become relevant. Such deviations include the influence of a pressure drop over the column [108, 109], nonlinearities in the adsorption isotherm, heat transfer resistances [110] and the influence of the finite pulse width [111]. More detailed accounts of the application of chromatography to the study of diffusion in porous media may be found in several reviews [6, 7, 112, 113]. While in chromatographic methods external heat and mass transfer resistances can be eliminated more easily than in a static system, the response may be additionally affected by axial dispersion, as represented by the first term on the right-hand side of Eq. (19). In the zero length column (ZLC) method the basic advantages of the chromatographic method are retained while eliminating the limitations imposed by axial dispersion [114–120]. In this technique, which has proved to be particularly useful for the observation of fast sorption phenomena, a small sample of adsorbent (in general only consisting of microparticles of radius R) is equilibrated at a uniform sorbate concentration and then desorbed by purging with an inert gas at a flow rate high enough to maintain essentially zero sorbate concentration at the external surface of the particles. By following the composition of the effluent, it is possible to measure the rate of desorption which may be attributed to the transport diffusion within the particles. Analytical solutions for various models may be found in the literature [115, 120–123], including the application to liquid-phase counter-diffusion measurements [124]. For sufficiently large purge flow rates and small sorbate concentrations (Henry’s law range), the adsorbate concentration c(t) in the outgoing gas flow obeys the following relation in the long-time limit  −π 2 Dt c(t) ∝ exp (20) R2 This expression may be easily deduced from the corresponding relation for the desorption experiment (Eq. (7)) by differentiation with respect to time. Microscopic Measurement Since all techniques of diffusion measurement presented so far are based on the measurement of either mean concentrations within the catalyst particles or of the vapor pressure in the surrounding atmosphere, they are generally referred to as macroscopic techniques of measurement (or mesoscopic ones, if only one crystal is concerned) [125–129]. Most recently, however, two novel 6.2.2.6

1719

techniques have been introduced into heterogeneous catalysis which are likewise able to record phenomena of transport diffusion over regions which might be much smaller than the extensions of the individual zeolite crystallites and which, hence, are referred to as microscopic techniques. By choosing probe molecules containing nuclei with a large coherent scattering cross-section for neutrons (e.g. deuterium), in quasi-elastic neutron scattering (QENS) (see also Section 6.2.3.2) the fluctuation of molecular concentrations and hence – following Onsager’s regression theorem [130–132] – the transport diffusivities as the factors of proportionality between (microscopic) gradients and fluxes may be directly recorded [119, 133–139]. In fact, QENS represents the only technique by which both transport diffusion (viz. the rate of molecular fluxes by analyzing coherent scattering) and self-diffusion (the rate of molecular redistribution by analyzing incoherent scattering) may be determined on microscopic scales for one and the same sample [136, 139]. Alternatively, interference microscopy has been introduced to monitor microscopically the time dependence of intracrystalline concentration profiles during molecular uptake and release [140–143]. In this technique, molecular concentrations may be monitored owing to their influence on the optical density and hence on the interference pattern of the host material (e.g. a zeolite crystallite) under study. First systematic investigations of transient uptake and release with a number of zeolite framework types such as MFI [143–145], AFI [142, 146–148] and FER [149] revealed substantial deviations of their real structure from the ideal textbook pattern. In most cases, the permeation through transport resistances on the outer surface (‘‘surface barriers’’ [149]), through internal barriers [144, 145] and/or cracks [149] rather than intracrystalline diffusion turned out to represent the rate-determining mechanism of molecular exchange between the intracrystalline space and the surrounding [150]. Since, so far, most of these measurements have been performed with specially synthesized large zeolite crystallites, it is not yet possible to assess the value of these findings for genuine zeolite catalysts, the crystallite diameters of which are typically below the range of micrometers and thus below the space scale accessible by interference microscopy. With respect to spatial resolution, interference microscopy notably exceeds IR microscopy [53–57]. In special applications in heterogeneous catalysis, however, in particular during a simultaneous observation of diffusion and conversion, this advantage of interference microscopy might be compensated by the potentials of References see page 1723

1720

6.2 Determination of Diffusion Coefficients in Porous Media

IR spectroscopy for selective measurement. First applications of IR microscopy to studying intracrystalline concentration profiles [146, 149, 151] and recent technical progress in IR micro-imaging [152, 153], however, open new perspectives also in this field. 6.2.3

Measurement of Self-Diffusion Elementary Steps of Diffusion The process of molecular diffusion may be viewed conceptionally as a sequence of jumps with statistically varying jump lengths and residence  times. Information about the mean jump length l 2  and the mean residence time τ , which might be of particular interest for a deeper understanding of the elementary steps of catalysis, may be provided by spectroscopic methods, in particular by quasi-elastic neutron scattering (see Section 6.2.3.2) and nuclear magnetic resonance (NMR). NMR spectroscopy has been applied to the determination of the mean residence time of adsorbed molecules in essentially two different ways, namely by proton magnetic relaxation studies and deuteron magnetic resonace line shape analysis. In the case of adsorbed molecules, proton magnetic relaxation is mainly determined by magnetic dipole interactions with protons of the same molecule (intramolecular proton–proton coupling), with protons of other molecules (intermolecular proton–proton coupling), and with paramagnetic iron. Separating the different contributions by a relaxation analysis [154, 155] it is, in principle, possible to determine the correlation time for any of these interactions, and from this, under appropriate conditions, the mean residence time between successive jumps. The deuteron magnetic line shape is determined by the gradient of the electric field brought about by the surroundings. Molecular reorientation with rate constants of the order of or larger than the linewidth of the rigid molecule leads to an averaging of the effective electric field gradient which is reflected in characteristic changes of the measured NMR spectrum. Attributing the observed spectra to the results of model calculations allows an estimate of the rates of reorientation and with them of the mean residence time [156, 157]. By assuming a molecular jump length comparable to the separation between adjacent large cavities as a reasonable estimate of the molecular mean jump length, deuteron NMR has been used in Refs. [158–164] to predict zeolitic diffusivities on the basis of the relation 6.2.3.1

D=

l 2  6τ

(21)

Eq. (21) is the microscopic equivalent of the Einstein equation (3) and implies that succeeding jumps are

uncorrelated. In general, however, backward jumps will take place with a higher probability (the correlation effect [7]), so that Eq. (21) provides an upper limit of the diffusivity. Alternatively, with known values for D and τ , Eq. (21) allows an estimate of the lower limit of the mean-square jump length [156, 157]. With the advent of two-dimensional NMR spectroscopy [165], a novel access to monitoring molecular interchange has become possible. In two-dimensional exchange NMR [166, 167], it is possible to monitor changes in the NMR frequencies of the nuclei considered between two subsequent time intervals, which may be separated by milliseconds up to seconds (this period being limited by the longitudinal nuclear magnetic relaxation time T1 ). Since in many cases these changes in the resonance frequencies may be assigned to molecular interchanges between well-defined positions, from the knowledge of their separation and the rate of interchange as provided by two-dimensional exchange NMR, diffusivities may be estimated [164, 168–170]. The displacements covered in these experiments may be of the order of interatomic distances. With Eq. (21), the lower limit of accessible diffusivities may thus be estimated to be on the order of 10−18 m2 s−1 , thus allowing the determination of diffusivities much smaller than accessible by, e.g., QENS or PFG NMR, as discussed in the subsequent sections. Quasi-elastic Neutron Scattering Diffusion measurement by thermal neutron scattering is based on an analysis of the (quasi-elastic) broadening in the energy distribution of the outgoing neutron beam (for the application of neutron scattering for the investigation of structural properties of catalysts, see Chapter 3.1.3.9). This broadening may be conceptionally understood as a Doppler broadening, caused by energy transfer between the incident wave and the scattering centers (in general the protons of the molecules under study). Since the energy transfer corresponding to a diffusive motion is very small, the term quasi-elastic neutron scattering (QENS) is used for such studies. For a diffusive motion with infinitely small diffusion steps, the half width at half maximum of the energy distribution of the outgoing beam (quasi-elastic and incoherent part [171, 172]), is determined by 6.2.3.2

E ≡ ¯hω(κ) = ¯hκ 2 D

(22)

with κ and ω denoting, respectively, the differences in the wavevectors and the frequencies between the incident and scattered neutron beams. On the basis of Eq. (22), the self-diffusivity may be easily determined from a plot of ω vs. κ 2 . For diffusion steps of finite length, one obtains [173] E ≡ ¯hω(κ) = ¯h

1 − exp(−κ 2 l 2 /6) τ

(23)

6.2.3 Measurement of Self-Diffusion

which allows the determination of both the mean residence time (for large values of κ) and the selfdiffusivity (for small values of κ, where Eqs. (22) and (23) coincide). Moreover, using Eq. (21) the mean jump length may also be estimated from these two quantities. QENS has been applied to several zeolitic adsorbate-adsorbent systems [135, 136, 157, 160, 171, 174–178]. Instrumental and methodological progress in QENS, in particular the advent of neutron spin-echo instruments [119, 138, 179, 180], have notably enhanced the range of measurement [136]. Today, the observation of displacements up to 10 nm and of diffusivities down to 5 × 10−14 m2 s−1 are possible. Moreover, first extensive measurements of transport diffusivities in zeolite catalysts with molecular species containing nuclei with sufficiently large coherent scattering cross-sections have been performed (see section 6.2.2.6). Pulsed Field Gradient NMR In the pulsed field gradient (PFG) NMR method [often also referred to as the PGSE (pulsed gradient spin echo) technique], molecular transport is studied by making use of the spatial dependence of the nuclear magnetic resonance frequency in an inhomogeneous magnetic field [59, 60, 155, 157]. Superimposing the constant magnetic field over two short time intervals of duration δ and separation t by a well defined inhomogeneous magnetic field B = gz (the field gradient pulses), the NMR signal (spin echo) generated by an appropriate sequence of rf pulses is attenuated by a factor  (t, γ δg) = P (z, t) cos(γ δgz) dz (24) 6.2.3.3

Here, P (z, t) (the mean propagator [59, 181]) denotes the probability (density) that during the time interval t a molecule within the sample is shifted over a distance z in the direction of the applied field gradient. The quantity γ denotes the magnetogyric ratio, a characteristic quantity of the given nucleus. In a homogeneous system of diffusivity D, the propagator is given by the relation  2 −z (25) P (z, t) = (4πDt)−1/2 exp 4Dt In this case Eq. (24) becomes (t, γ δg) = exp(−γ 2 δ 2 g 2 Dt)

(26)

and D results straightforwardly form a plot of ln  versus (γ δg)2 . Molecular displacements covered in PFG NMR measurements are typically of the order of a micrometer. Owing to the large magnetogyric ratio, protons offer the best conditions for PFG NMR diffusion studies with respect to both signal intensity and sensitivity

1721

towards molecular displacements. However, PFG NMR studies in porous media have also been carried out using 13 C [182], 15 N [183], 19 F [184] and 129 Xe NMR [182]. The lower limit of diffusivities accessible by 1 H PFG NMR is of the order of 10−14 m2 s−1 . However, such low diffusivities may only be measured under suitable conditions, in particular for large nuclear magnetic relaxation times T1 and T2 [59, 185, 186]. Owing to the quite general correlation between the primary experimental data and molecular propagation as provided by Eq. (24), PFG NMR is particularly suitable for diffusion studies in heterogeneous systems. It allows the determination of a variety of parameters characterizing molecular transport in porous media. In the investigation of beds or granules of zeolite crystallites it is possible to measure directly the rates of intra-crystalline diffusion and of long-range diffusion, i.e. of the rate of molecular propagation through the bed or the granule [155, 157]. So far, it has been possible to observe five different patterns of concentration dependence for intracrystalline diffusion [187]. Further transport-related phenomena observable in such studies include diffusion anisotropy [188], the formation of transport resistances (surface barriers) by coke deposits [157] or deterioration of the crystal structure [189], deviations from ordinary diffusion following Einstein’s relation (e.g. single-file diffusion [190–193] or diffusion in fractals [194]) and the influence of the pressure of compaction and of inert gases [195]. Being able to distinguish between intracrystalline and external coke deposits [196], PFG NMR may be very helpful in tracing the microdynamics of catalyst deactivation and regeneration (see Chapter 7.1). As a non-invasive technique, sensitive to a particular nucleus or even (by applying Fourier transform NMR [197]) to a particular molecular species, PFG NMR may be applied to study the diffusivity of any individual component in a multicomponent system [198, 199]. Moreover, by this method the in situ observation of the diffusion of the reactant and product molecules during catalytic reactions has become possible [200]. PFG NMR diffusion measurements with varying observation time have been performed for both an elimination of the disturbing influence of the finite size of the zeolite crystallites under study [201, 202] following the conception of time-dependent diffusivities by Mitra and co-workers [203, 204] and the discrimination of internal transport barriers [205–207]. Recent methodological progress includes the combination of pulsed field gradients with magic angle spinning (MAS) NMR, which may contribute considerably to an enhancement of the spectral resolution in multicomponent diffusion measurements [208, 209] and also to an enhancement of the observation time. References see page 1723

1722

6.2 Determination of Diffusion Coefficients in Porous Media

Tracer Techniques The methods described so far for studying self-diffusion are essentially based on an observation of the diffusion paths, i.e. on the application of Einstein’s relation (Eq. (3)). Alternatively, molecular self-diffusion may also be studied on the basis of Fick’s laws by using isotopically labeled molecules. As in the case of transporter diffusion, the diffusivities are determined by comparing the measured curves of tracer exchange between the porous medium and the surroundings with the corresponding theoretical expressions. As a basic assumption of the isotopic tracer technique for studying self-diffusion, the isotopic forms are expected to have identical properties. For small percentage differences in the atomic masses, this is a reasonable assumption. Differentiation between the isotopes may be based on the differences in their mass (gravimetry [210] and mass spectroscopy [211]), in their nuclear magnetization (NMR spectroscopy [155]), or in their IR frequencies [50]. If one of the isotopes is radioactive, the activity provides a particularly simple and sensitive measure of the labeled species [212]. In principle, by applying labeled components, the various concepts of measuring transport diffusion may be correspondingly applied to the measurement of selfdiffusion, provided that the exchange process between the labeled and unlabeled molecules is determined by diffusion rather than being masked by other phenomena. For the investigation of fast processes, a tracer variant of the ZLC method [116, 120, 213–217] has proved to be of particular relevance. 6.2.3.4

6.2.4

Diffusion in Multicomponent Systems

In principle, any of the techniques described for studying self-diffusion may be applied to both single- and multicomponent systems. So far, however, most self-diffusion measurements of multicomponent systems have been carried out by PFG NMR. Such measurements are possible by applying a set of samples with deuterated compounds with only one species (namely that to be measured) in the protonated form [218], by simultaneously considering different nuclei (e.g. 1 H and 19 F [184, 219], 1 H and 129 Xe [220] and even 1 H, 13 C and 15 N [221–223] in three-component diffusion studies) and by Fourier transform PFG NMR [198, 199, 208]. Diffusion measurements under non-equilibrium conditions are more complicated due to the difficulties in ensuring well defined initial and boundary conditions. IR spectroscopy has proved to be a rather sensitive tool for studying simultaneously the intracrystalline concentration of different diffusants, including the occupation density of catalytic sites [52]. By choosing appropriate

initial conditions, in this way both co- and counterdiffusion phenomena may be followed [55]. Information about molecular transport diffusion under the conditions of multicomponent adsorption may also be deduced from flow measurements [224]. As in the case of singlecomponent adsorption, the diffusivities are determined by matching the experimental data (i.e. the time dependence of the concentration of the effluent or the adsorbent) to the corresponding theoretical expressions. A rather special possibility to attain information about molecular diffusion is provided by catalytic reactions if they proceed in the range of medium Thiele moduli (i.e. in the transition range between intrinsically and transport controlled reactions) [225]. By analyzing the dependence of the effective reaction rate on the catalyst particle size [28, 30] and/or the intrinsic rate constant [226] it is possible to determine the rate of counter-diffusion of the reactant and product molecules. Similarly, as with the flow techniques (see Chapter 6.3 of this Handbook), also in this case the diffusivities are determined under steady-state conditions (see Section 6.2.2.1). 6.2.5

Correlation Between the Different Diffusivities

Diffusivities are often measured under conditions which are far from those of catalytic reactions. Moreover, corresponding to their different nature, the various measuring techniques are limited to special ranges of application. The possibility of a mutual transformation of the various diffusivities would therefore be of substantial practical relevance. Since each of the coefficients of selfdiffusion and transport diffusion in single-component and multicomponent systems refers to a particular physical situation, one cannot expect that the multitude of information contained in this set of parameters can in general be adequately reflected by a smaller set of parameters. Any correlation which might be used in order to reduce the number of free parameters must be based on certain model assumptions. To a first-order approximation, the coefficients of transport diffusion (D) and self-diffusion (D ∗ ) under the conditions of single-component adsorption are correlated by the equation D = D∗

∂ ln p(c) ∂ ln c

(27)

where p(c) denotes the external gas-phase pressure of the adsorbate that is necessary to maintain the adsorbate concentration c. Eq. (27) may be shown by both thermodynamic [227, 228] and statistical [229] arguments to be valid only under the assumption that the correlation of the motion of different molecules may

References

be assumed to be negligibly small. For a number of systems this assumption seems to be justified, since there are various comparative studies of transport and self diffusion [230–232] as well as MD simulations [229, 233] which are in reasonable agreement with Eq. (27). The generalization of Eq. (27) to multicomponent adsorption yields for the transport diffusivities (see Eq. (1)) Dij =

Di∗ ci ∂ ln pi × cj ∂ ln cj

(i, j = 1 . . . n)

(28)

with pi denoting the set of partial pressures necessary to maintain the set of sorbate concentrations cj . Hence, if one knows the corresponding multicomponent adsorption isotherms, the elements of the n × n matrix of transport diffusivities can be determined from the multicomponent self-diffusivities by using Eq. (28). In general and following, in particular, the evidence provided by QENS experiments simultaneously recording transport diffusion and self-diffusion [136, 139, 234], one has to be aware of deviations from Eq. (27) and, hence, from Eq. (28). In this case, the formalism of the generalized Maxwell–Stefan equations [235–242] may serve as a useful means of correlating different diffusivities. It is based on the conception that the driving force evoked by the gradient of the chemical potential of a molecular species i is balanced by the momentum exchange between this species and all other species, including the host framework [129, 235, 243]. MD simulations with binary mixtures adsorbed in zeolites of type MFI and FAU confirm the expediency of such approaches to correlate the binary transport diffusivities with the self-diffusivities [31, 234, 239, 242, 243]. Under industrial conditions, zeolite catalysts are generally run at high temperatures and low concentrations. Cross-effects between the various molecular species are therefore negligibly small in general and, again, Eqs. (27) and (28) result as useful approaches for correlating selfand transport diffusivities. 6.2.6

A Look into the Future

Diffusion research with nanoporous materials has continued to provide us with surprising new results. These have resulted from the advent of new measuring techniques and the expanding range of material classes and originate from the complexity of the systems, the real structure of which is repeatedly found to deviate notably from the ideal case. The persistently stormy evolution in the field suggests a continuation of this tendency. With the most recent, powerful developments in quasielastic neutron scattering today molecular propagation may be covered over distances that are long enough

1723

to ensure the measurement of genuine translational diffusion, but which – on the other hand – are still so small that the influence of deviations from the ideal pore structure may be neglected. Quasi-elastic neutron scattering has, therefore, become the method of choice to confirm the results of molecular modeling of diffusion in a given structure type by experimental evidence [137, 244]. Simultaneously, the elucidation of the steps really limiting molecular transport during catalytic processes has turned out to be a much more complicated problem. Being able to record the rate of molecular propagation over displacements from about 100 nm up to hundreds of micrometers, pulsed-field gradient NMR is ideally suited to trace the hierarchy of transport resistances that possibly occur during heterogeneous catalysis [5]. In addition to the drag exerted by the genuine pore structure, transport resistances may also result from intracrystalline stacking faults, structure collapse or other structural peculiarities, in particular close to the crystal surface and various types of deposits, with coke as its most important representative. The novel options of interference microscopy, allowing the direct observation of transient concentration profiles, make this technique particularly powerful and informative for exploring the different intracrystalline transport phenomena, including the quantification of surface resistances [245]. The identification of these resistances and their quantitation as a function of the reactant and product molecules under study, and also of their evolution with temperature and time, are among the great challenges of future experimental work. The resulting data sets represent an inevitable input for any strategy of performance enhancement striving to exploit fully the potential of transport optimization. References 1. P. Heitjans, J. K¨arger, Diffusion in Condensed Matter: Methods, Materials, Models, Springer-Verlag, Berlin, 2005, 965 pp. 2. J. K¨arger, F. Grinberg, P. Heitjans (Eds.), Diffusion Fundamentals, Leipziger Universit¨atsverlag, Leipzig, 2005, 615 pp. 3. M.E. Glicksman, Diffusion in Solids, Wiley, New York, 1999, 498 pp. 4. R.H. Doremus, Diffusion of Reactive Molecules in Solids and Melts, Wiley, New York, 2001, 312 pp. 5. J. K¨arger, S. Vasenkov, Microporous Mesoporous Mater. 2005, 85, 195–206. 6. H.W. Haynes, Catal. Rev. Sci. Eng. 1988, 30, 563–627. 7. J. K¨arger, D.M. Ruthven, Diffusion in Zeolites and Other Microporous Solids, Wiley, New York, 1992, 605 pp. 8. C.S. Satterfield, Mass Transfer in Heterogeneous Catalysis, MIT Press, Cambridge, MA, 1970, 267 pp. 9. S. Olayinka, M.A. Ioannidis, Transp. Porous Media 2004, 54, 273–295. 10. P.N. Sen, Concepts Magn. Reson., Part A 2004, 23, 1–21.

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6.3

Simultaneous Heat and Mass Transfer and Chemical Reaction Roland Dittmeyer∗ and Gerhard Emig

6.3.1

Introduction

Heterogeneous catalytic reactions, by their nature, involve a separate phase of catalyst, embedded in a phase of reacting species. Therefore, the chemical transformation relies References see page 1781 ∗ Corresponding author.

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6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

on a number of physical transport processes which may have a strong influence on the rate of the overall process and which may introduce an additional dependence on the operating conditions. In the industrially important situation that the catalyst is a porous solid and the reactants form either a gaseous or a liquid phase, the following seven steps can be observed (Fig. 1): 1. diffusion of the reactants through a boundary layer or film adjacent to the external surface of the catalyst (film diffusion or interphase diffusion) 2. diffusion of the reactants through the porous interior of the catalyst to the point at which the chemical transformation occurs (pore diffusion or intraparticle diffusion) 3. adsorption of the reactants on the inner surface 4. surface reaction at specific active sites 5. desorption of the products from the inner surface 6. diffusion of the products through the porous structure to the external surface (intraparticle diffusion) 7. diffusion of the products through the external boundary layer into the bulk fluid phase (interphase diffusion). Adsorption, surface reaction and desorption are sequential steps; these constitute the chemical transformation. Interphase and intraparticle diffusion also occur in a serial manner. However, intraparticle diffusion and chemical transformation (adsorption, surface reaction and desorption) take place simultaneously. Most chemical reactions are accompanied by heat effects. Thus, in addition to film and pore diffusion, interphase and intraparticle heat transfer also occurs. The net kinetics of the overall reaction (steps 1–7) are normally called effective kinetics or

macrokinetics, in contrast to the kinetics of the chemical transformation (steps 3–5), which are termed intrinsic kinetics or microkinetics. Heat and mass transfer processes always proceed with finite rates. Hence, even when operating under steady state conditions, more or less pronounced concentration and temperature profiles may exist across the phase boundary and within the porous catalyst pellet as well (Fig. 2). As a consequence, the observable reaction rate may differ substantially from the intrinsic rate of the chemical transformation under bulk fluid phase conditions. Moreover, the transport of heat or mass inside the porous catalyst pellet and across the external boundary layer is governed by mechanisms other than the chemical reaction, a fact that suggests a change in the dependence of the effective rate on the operating conditions (i.e. concentration and temperature). Normally, the intrinsic chemical rate is an exponential function of temperature, according to the Arrhenius law, whereas the mass transfer rate is less strongly influenced by a temperature change. The intraparticle effective diffusivity De is proportional to T 3/2 when molecular diffusion dominates and proportional to T 1/2 for the case of governing Knudsen diffusion. The interphase mass transfer coefficient kf exhibits basically the same temperature dependence as the molecular diffusivity (kf ∼ T 3/2 ). boundary catalyst layer particle

Bulk phase A1

A2

7

Boundary layer

1

A2

A1

1

A1 2

2

A1

4

A2

3 5 Active site at catalyst surface

A2 4

ci

(a)

(b)

T

Porediffusion

ci

Adsorption/ 5 desorption

(c)

(d)

Chemical reaction

Individual steps of a simple, heterogeneous catalytic fluid–solid reaction A1 → A2 carried out on a porous catalyst.

Fig. 1

ci

ci

6

A1

T

T

Porous catalyst 6 A 2

3

Filmdiffusion

7

T

Stationary concentration (reactant) and temperature profiles inside and around a porous catalyst pellet during an exothermic, heterogeneous catalytic fluid–solid reaction: (a) without transport influence; (b) limited only by intraparticle diffusion; (c) limited by interphase and intraparticle diffusion; (d) limited only by interphase diffusion (dense pellet).

Fig. 2

6.3.1 Introduction

When plotting the natural logarithm of the effective rate constant ln ke against the inverse of temperature (Arrhenius diagram), the general curve depicted in Fig. 3 is obtained. At a lower temperature the chemical reaction is slow and thereby rate controlling. Concentration and temperature remain constant over the entire cross-section of the catalyst pellet (Fig. 2a). In this region, the slope of the curve in Fig. 3 is proportional to the intrinsic activation energy EA . By increasing the temperature, the intrinsic chemical rate is accelerated more strongly than the rate of intraparticle diffusion. When both of these simultaneous steps occur with a rate of about the same order of magnitude, a transition region is passed where the slope changes with temperature. As the intraparticle diffusion finally becomes rate determining, the concentration inside the porous catalyst pellet drops markedly (Fig. 2b). The effective activation energy in this regime is roughly half of the true value, derived in detail in Section 6.3.4.1. Upon further increasing the temperature, a second transition regime is passed. Finally, interphase mass transfer becomes the rate-controlling step. As a consequence, the reactant concentration already falls off steeply across the external boundary layer (Fig. 2c). Here, the slope of the curve in Fig. 3 tends towards zero, corresponding to an effective activation energy in the range of less than 5–10 kJ mol−1 . In general, not only the effective activation energy but also the effective order of reaction is changed during the transition from kinetic to diffusion control. According to Fick’s first law, the rate of diffusion (interphase and intraparticle) is proportional to the concentration gradient, i.e. it is first order. The effective reaction order observed under severe intraparticle diffusion control approaches a value of (n + 1)/2, where n is the intrinsic order of

Film diffusion Transition Pore diffusion Transition region II region I control control

Kinetic control

In ke

Slope ~ 0

Slope =

− EA 2R

Slope =

−E A R

1/T

Transition from the kinetic regime to the diffusioncontrolled regime of a heterogeneous catalytic fluid–solid reaction carried out on a porous catalyst.

Fig. 3

1729

reaction (cf. Section 6.3.4.1). In the case that interphase mass transfer controls the overall rate, the effective concentration dependence of the reaction rate is close to first order (cf. Section 6.3.4.2). Changes in the effective reaction rate, the apparent activation energy and the apparent order of reaction during the transition from the kinetic regime to the diffusion-controlled regime are of great importance for the technology of heterogeneously catalyzed reactions. These effects must be considered in practical situations, as otherwise wrong predictions with respect to selectivity and yield of the catalyst may result. This chapter is concerned with the mathematical modeling of coupled chemical reaction and heat and mass transfer processes occurring in porous catalysts. It focuses primarily on steady-state catalyst operation, which is beyond any doubt the preferred industrial practice. Nonstationary operation may be important for the startup and shutdown of an industrial reactor or with respect to dynamic process control. However, these effects are not discussed here in great detail because of the limited space available. Instead, the interested reader is referred to the various related monographs and articles available in the literature [1–6]. The engineering implications of the interaction between diffusional transport and chemical reactions were first identified by Damk¨ohler [7], Zeldovich [8] and Thiele [9] in the late 1930s and then by Wheeler [10] and Weisz and Prater [11] in the early 1950s. Since then, the authors of numerous chemical engineering textbooks have addressed this problem within one or two separate chapters [12–16]. The most comprehensive treatment of the subject has been given by Aris [1, 17], although this is a more theoretical treatment. Practical aspects have been ably treated by, for example, Satterfield [18, 19] and Carberry [20, 21]. More recent textbooks with a good discussion of the topics of this chapter have been authored by, for example, Baerns et al. [22], Davis and Davis [23] and L¨owe [24]. Published in 2005, the books by Fogler [25], Schmidt [26] and Emig and Klemm [27] may also be recommended. Finally, Klemm et al. have treated the same subject in Chapter 5.2 of the Handbook of Porous Solids [28]. Additional information with respect to certain aspects of this chapter may be also found in Chapters 6.2, 5.5 and particularly 9.1. Note that Chapter 6.2 is concerned in detail with the modeling of diffusion and with the determination of diffusivities. In this special field, a good textbook was recently published by Keil [29]. The contents of the present contribution can be outlined as follows. Section 6.3.2 introduces the basic principles of coupled heat and mass transfer and chemical reaction. Section 6.3.3 covers the classical mathematical treatment References see page 1781

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6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

of the problem by example of simple reactions and some of the analytical solutions which can be derived for different experimental situations. Section 6.3.4 is devoted to the point that heat and mass transfer may alter the characteristic dependence of the overall reaction rate on the operating conditions. Section 6.3.5 contains a collection of useful diagnostic criteria available to estimate the influence of transport effects on the apparent kinetics of single reactions. Section 6.3.6 deals with the effects of heat and mass transfer on the selectivity of basic types of multiple reactions. Section 6.3.7 gives an overview of available classical macroscopic models to describe the diffusional transport in porous catalysts. Section 6.3.8 focuses on a practical example, namely the utilization of shape selectivity effects in the disproportionation of ethylbenzene to benzene and diethylbenzenes on an H-ZSM-5 zeolite catalyst. Finally, Section 6.3.9 provides some advice on an appropriate strategy to apply the methods described to reaction engineering problems. The authors’ objective is to make the reader familiar with the basic concepts for a quick and effective study of the phenomena of simultaneous heat and mass transfer and chemical reaction. The reader should then be in a position to decide whether or not these phenomena play a decisive role in an actual catalytic system at hand. 6.3.2

Mathematical Description

The mathematical description of simultaneous heat and mass transfer and chemical reaction is based on the general conservation laws valid for the mass of each species involved in the reacting system and the enthalpy effects related to the chemical transformation. The basic equations may be derived by balancing the amount of mass or heat transported per unit of time into and out of a given differential volume element (the control volume) together with the generation or consumption of the respective quantity within the control volume over the same period. The sum of these terms is equivalent to the rate of accumulation within the control volume: accumulation within control volume/time = in/time – out/time + generation or consumption/time The region over which this balance is invoked is the solid porous catalyst pellet, which, for the sake of simplicity, is described as a pseudo-homogeneous substitute system with regular pore structure. This virtual replacement of the solid catalyst pellet by a fictitious continuous phase allows a convenient representation of the mass and enthalpy conservation laws in the form of differential equations. Moreover, the three-dimensional shape of the

catalyst pellet is replaced by assuming a one-dimensional model geometry (sphere, infinite cylinder, infinite flat plate) in order to simplify the mathematical treatment. For the case of a sphere, the control volume is given by a thin spherical shell of thickness dr and radius r. If we assume that the complex diffusion process inside the porous structure can be represented by Fick’s first law and if we additionally suppose that the volume change due to reaction is negligible (i.e. the total number of moles is constant), we arrive at the following form of the mass conservation law for the reacting species i:  ∂ci  2 ∂ci 2 4πr δr = 4πr Di,e ∂t ∂r r  ∂ci  − 4π(r + δr)2 Di,e + 4πr 2 δrRi (1) ∂r r+δr where Ri is the pseudo-homogeneous rate of production or disappearance of species i per unit volume, which in the general case of a network of multiple reactions may be the result of M different chemical transformations. Therefore, Ri =

M

νi,j rj

(2)

j =1

In Eq. (1) Di,e is the effective diffusivity of species i in the reaction mixture, which can be determined on the basis of various models of the diffusion process in porous solids. This aspect is discussed more fully in Section 6.3.7 and in Chapter 6.2. Di,e is affected by the temperature and the pore structure of the catalyst, but it may also depend on the concentration of the reacting species (Maxwell–Stefan diffusion [30], cf. Section 6.3.7.2). As Di,e is normally introduced on the basis of more or less empirical models, it may not be considered as a physical property, but rather as a model-dependent parameter. The concentration gradient dci /dr at the position r + δr in Eq. (1) may be expanded as a Taylor series:    ∂ci  ∂ 2 ci  ∂ci  = + δr (3) ∂r r+δr ∂r r ∂r 2 r The resulting approximation is substituted in Eq. (1). The terms which then include the expression δr 2 in the numerator can be neglected (δr is already small). After rearranging, the following mass balance is finally obtained:  2 2 ∂ci ∂ci ∂ ci + (4) = Di,e + Ri ∂t ∂r 2 r ∂r Equation (4) is equivalent to Fick’s second law (nonstationary diffusion), expanded by an additional source term which accounts for the production or consumption

6.3.2 Mathematical Description

of species i caused by chemical transformations. Similarly to this mass balance, an enthalpy balance may be also derived: ρcp

∂T = λe ∂t



2 ∂T ∂ 2T + 2 ∂r r ∂r

+ HR

(5)

Equation (5) is an expansion of Fourier’s second law (non-stationary heat conduction), where the additional term HR , in analogy with Ri in Eq. (4), specifies the total heat of reaction produced or consumed per unit volume and time. This gives HR =

M

rj (−HR,j )

(6)

j =1

Although it is clear from Eq. (6) that HR can be negative (endothermic) as well as positive (exothermic), in the following it is always considered as the heat release term, for notational simplicity. In the general case, Eqs. (4) and (5) constitute a system of nonlinear coupled second-order partial differential equations. To specify the boundary conditions for this problem, it is necessary to include the external (interphase) heat and mass transfer, as both the concentration and the temperature at the external surface of the catalyst pellet may differ from the corresponding values in the bulk of the surrounding fluid phase. To avoid unnecessary complications, the view is restricted here to the symmetric case. This means that the catalyst pellet is assumed to be located in a homogeneous concentration and temperature field, i.e. the surrounding fluid phase is perfectly mixed. A consequence of this assumption is that uniform reaction conditions prevail at the pellet surface. Concentration or temperature gradients, which might occur on the surface in extreme situations [2], are neglected here. At the phase boundary no discontinuity of the mass and heat fluxes can occur. Therefore, the mass flux density of species i from the bulk fluid phase to the external catalyst surface, determined by the interphase mass transfer, must equal the mass flux density of this species attributed to intraparticle effective diffusion. The external mass flux density is given by the product of the mass transfer coefficient kf and the concentration difference between the bulk fluid phase and the external surface (ci,b − ci,s ), whereas the mass flux density towards the center of the porous pellet is obtained by multiplying the effective diffusivity Di,e by the concentration gradient at the external pellet surface (∂ci /∂r)r=R . Equating both expressions (pellet surface: r = R) yields the first boundary condition [Eq. (7)]. An analogous equation [Eq. (8)] can be derived for the

1731

heat transfer:

 ∂ci  kf (ci,b − ci,s ) = Di,e ∂r r=R  ∂T  hf (Tb − Ts ) = λe ∂r r=R

(7) (8)

The second boundary condition is dictated by the symmetry of the (idealized) pellet geometry, i.e. the concentration and temperature gradients must disappear at the pellet center [r = 0; Eqs. (9) and (10)]:  ∂ci  =0 (9) ∂r r=0  ∂T  =0 (10) ∂r r=0 If the external heat and mass transfer resistances are negligible, Eqs. (7) and (8) can be simplified by replacing the unknown surface values ci,s and Ts with the known conditions in the bulk fluid phase and then transform to the following simple boundary conditions at the external pellet surface: ci,s = ci,b

(11)

Ts = Tb

(12)

The condition of symmetry at the center of the pellet remains uninfluenced by a change of the boundary condition at the external pellet surface. The solution of the above system of partial differential equations [Eqs. (4)–(12)] yields the concentration and temperature profiles inside the catalyst pellet and if necessary across the external boundary layer, as a function of time. However, there are only few cases of practical importance where this complete solution is required, such as startup and shutdown periods, dynamic process control options such as the so-called ‘‘Matros concept’’ with flow reversals (for redox processes) or situations where the catalyst is rapidly deactivated. In most day-to-day-operations, only the stationary solution will be of interest. Hence, all partial derivatives with respect to time in Eqs. (4)–(12) can be set to zero, which leads to the following equations for the stationary case: Mass balance:

2 dci −Ri d2 ci + = 2 dr r dr Di,e

Enthalpy balance:

2 dT d2 T −HR + = 2 dr r dr λe

(13) (14)

This is a system of ordinary differential equations. Its solution is subject to the same boundary conditions References see page 1781

1732

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

[Eqs. (7)–(12)] as the solution of Eqs. (4) and (5). However, the variables involved are no longer time dependent, which means that the partial derivatives ∂ci /∂r and ∂T /∂r can be replaced by the ordinary derivatives dci /dr and dT /dr. For simultaneous treatment of interphase and intraparticle transport resistances, the boundary conditions at the external pellet surface are now given by Eqs. (15) and (16), whereas Eqs. (11) and (12) still hold for the case of negligible interphase transport resistance. The boundary conditions at the pellet center, for both cases, are defined by Eqs. (17) and (18):  dci  kf (ci,b − ci,s ) = Di,e (15) dr r=R  dT  (16) hf (Tb − Ts ) = λe dr r=R  dci  =0 (17) dr r=0  dT  =0 (18) dr r=0 However, for a basic design or a detailed simulation of chemical reactors, profiles of the concentrations and the temperature inside the catalyst pellet are not of primary interest, but rather the effective rates of production or disappearance of the reacting species and the effective heat release or consumption as well. Both are defined according to Eqs. (19) and (20) as averaged values, related to the pellet volume:  1 Ri,e = Ri dV (19) Vp Vp  1 HR,e = HR dV (20) Vp Vp To calculate Ri,e and HR,e , normally not their definition as given in Eqs. (19) and (20) will be used. Instead, it is more convenient to utilize the condition that, during stationary operation, the amount of heat and mass converted inside the catalyst pellet per unit time must meet the flux of the respective quantity across the external pellet surface. For the mass of the reacting species i, this gives the following relationship:  4 dci  3 2 (21) πR (−Ri,e ) = 4πR Di,e 3 dr r=R When both sides of Eq. (21) are divided by the pellet volume, we arrive at the final equation for determining Ri,e :  3 dci (22) Ri,e = − Di,e R dr r=R

In the same way, for HR,e :  3 dT HR,e = − λe R dr r=R

(23)

According to Eqs. (22) and (23), the concentration and temperature gradients at the external pellet surface must be known to evaluate Ri,e and HR,e . As soon as the concentration and temperature profiles over the pellet radius are available, the gradients at the pellet surface can be determined by differentiation. However, full analytical solution of the coupled mass and enthalpy balances is achieved only when the system can be described properly by just a single key reaction (the conversion problem), which, in addition, obeys simple power law or Langmuir–Hinshelwood-type kinetics. In contrast, when several key reactions must be considered simultaneously (the selectivity problem) or when it is necessary to use a complex kinetic expression, then the solution in general requires the use of numerical methods. For didactic reasons, a detailed discussion of the first case, the analytically solved (simple) conversion problem, is given initially. The main purpose of this is to demonstrate how the reaction rate under certain conditions is fundamentally influenced by heat and mass transfer. Attention is then given to the more complicated case in which several key reactions have to be considered and selectivity in addition to conversion becomes an issue. 6.3.3

Single Reactions (Conversion Problem)

Whenever the kinetics of a chemical transformation can be represented by a single reaction, it is sufficient to consider the conversion of just a single reactant. The concentration change of the remaining reactants and products is then related to the conversion of the selected key species by stoichiometry and the rates of production or consumption of the various species differ only by their stoichiometric coefficients. In this special case, the combined influence of heat and mass transfer on the effective reaction rate can be reduced to a single number, termed the catalyst efficiency or effectiveness factor η. From the pioneering work of Thiele [9] on this subject, the expressions ‘‘pore-efficiency concept’’ and ‘‘Thiele concept’’ have been coined. As a typical example of this type of reaction, here the transformation A1 → products may be considered, where the kinetics are described by a simple power rate law of the order n. Since this reaction is completely characterized by specifying the conversion of reactant A1 , the above system of differential equations [Eqs. (11)–(18)] may be readily expressed in a convenient, non-dimensional form. For this purpose, the reactant concentration and the temperature

6.3.3 Single Reactions (Conversion Problem)

are related to their corresponding values in the bulk fluid phase [Eqs. (24) and (25)], and the radius coordinate r is divided by the pellet radius R to introduce a dimensionless coordinate [Eq. (26)]: f =

c cb

(24)

T Tb r x= R θ=

(25) (26)

With these definitions, Eqs. (11)–(18) yield a structure which suggests a grouping of the different variables to several non-dimensional numbers: 1. The well known Thiele modulus φ of the reaction. This is defined as the ratio of the intrinsic chemical rate, calculated at bulk fluid phase conditions, to the maximum rate of effective diffusion at the external pellet surface. For spherical catalyst pellets, the Thiele modulus is given by  k(Tb )cbn−1 (27) φ=R De 2. The Biot number Bim for mass transport. This can be interpreted as the ratio of internal to external transport resistance (intraparticle diffusion versus interphase diffusion): Bim =

kf R De

(28)

3. The Biot number Bih for heat transport. Analogously to Bim , this is defined as the ratio of the internal to external heat transfer resistance (intraparticle heat conduction versus interphase heat transfer): Bih =

hf R λe

(29)

4. The Arrhenius number, which is a dimensionless representation of the intrinsic activation energy, related to the bulk temperature: γ =

EA RTb

(30)

5. The Prater number, defined as the maximum observable temperature difference between the pellet surface and the pellet interior, related to the bulk temperature: β=

cb De (−HR ) λe Tb

(31)

1733

Substituting these non-dimensional numbers into Eqs. (11)–(18) and after some rearrangement, the general dimensionless representation of the problem is obtained, as depicted in Table 1. These equations are valid not only for spherical pellet geometry, but also for an infinite cylinder and an infinite flat plate (slab). The dimensionless numbers x, φ, Bim and Bih must then be calculated on the basis of the respective characteristic length, i.e. the cylinder radius or the plate thickness. Moreover, the parameter b in Eqs. (35) and (36) is a factor depending on the pellet geometry: it is 2 for the sphere, 1 for the cylinder and 0 for the flat plate. An overall catalyst effectiveness factor of the reaction may now be defined as the ratio of the observed effective rate, averaged over the pellet volume, divided by the intrinsic chemical rate which would be expected in absence of concentration and temperature gradients in the system (i.e. under bulk fluid phase conditions): η=

re rb

(32)

When Eqs. (22), (24) and (26) are substituted into this relationship, the following expression for the effectiveness factor of a spherical catalyst pellet is obtained:  df 3De dx x=1 (33) η= 2 R kcbn−1 Using the definition of the Thiele modulus according to Eq. (27) finally yields  3 df (34) η= 2 φ dx x=1 This equation is valid for the general case, where the effective rate is influenced by both external and internal heat and mass transfer. In real situations, however, one or more of the transport steps involved frequently proceed at a rate substantially above that of the chemical reaction. If this happens, these steps may be neglected without significantly affecting the observable reaction rate. Hence the system may be simplified. The extent of simplification depends on which of the transport steps can be neglected. In the following, the different cases that occur in practical situations are discussed separately. Pore Diffusion in an Isothermal Pellet The simplest case is given when the effective reaction rate is influenced by pore diffusion only, whereas the interphase heat and mass transfer resistances as well 6.3.3.1

References see page 1781

1734

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction Dimensionless representation of the stationary mass and enthalpy balance equations for combined interphase and intraparticle transport and reaction (single, nth-order irreversible reactions)

Tab. 1

Parameter

Mass balance Enthalpy balance Boundary conditions Pellet center (x = 0)

Relationship γ d2 f b df + − φ2 f n e dx 2 x dx



Equation 1 1− θ

γ d2 θ b dθ + − βφ 2 f n e 2 dx x dx





1 1− θ

=0  =0

 df  =0 dx x=0  dθ  =0 dx 

(35) (36)

(37) (38)

x=0

Pellet surface (x = 1) with interphase gradients

   df  = Bim 1 − fs dx x=1  dθ  = Bih (1 − θs ) dx 

(39) (40)

x=1

without interphase gradients

fs = 1

(41)

θs = 1

(42)

as the effective heat conduction resistance inside the porous pellet may be neglected. Despite its simplicity, this case is of great practical importance, as it is often found when the reaction is associated only with a small heat release (or consumption), and/or when the catalyst exhibits good thermal conductivity (e.g. metals). The pellet temperature is assumed to be constant and equal to the temperature of the surrounding fluid phase. Hence the solution of the enthalpy balance is trivial. Moreover, the concentration at the external pellet surface is identical with the concentration in the bulk fluid phase. Therefore, the mass balance [Eq. (35)] and the corresponding boundary conditions [Eqs. (37) and (41)] are governed by a single parameter, namely the Thiele modulus φ. For single, irreversible reactions obeying simple, integer-order power rate laws, this problem can generally be solved analytically. In the case of a first-order reaction in a spherical pellet, the following mass balance is found: d2 f 2 df = φ2f + 2 dx x dx

(43)

Introducing a new variable y = f x allows this differential equation to be rewritten in a very simple form: d2 y = φ2y dx 2

(44)

A general solution of Eq. (44) is y = f x = C1 eφx + C2 e−φx

(45)

From the condition of symmetry at the pellet center [Eq. (37)], it can be easily deduced that the absolute values of the constants C1 and C2 must be equal: C1 = −C2

(46)

Therefore, f =

2C1 sinh(φx) x

(47)

The constant C1 is obtained from the boundary condition at the external pellet surface [Eq. (41)]: C1 =

1 2 sinh(φ)

(48)

Combining Eqs. (47) and (48) leads to the final relationship for the change of the dimensionless concentration f across the pellet radius: f (x) =

sinh(φx) x sinh(φ)

(49)

To illustrate the decrease of the reactant concentration towards the pellet center, Fig. 4 shows a plot of the concentration profiles for few selected values of the Thiele modulus.

6.3.3 Single Reactions (Conversion Problem)

1

1

1

f=1

Sphere

f=2 h

f

0.1 Cylinder

0.4 f=5 0.2

0

2 3

n=1 T = const.

0.8

0.6

1735

Flat plate f = 10 0.01 0.01

0

0.2

0.4

0.6

0.8

0.1

1

x

Normalized concentration profiles of reactant A1 versus the dimensionless pellet radius, calculated from Eq. (49) for different values of the Thiele modulus φ (isothermal, first-order irreversible reaction in a sphere).

Fig. 4

By differentiation of Eq. (49), the concentration gradient at the external pellet surface is obtained as  φ df  = −1 (50)  dx x=1 tanh(φ) Substituting Eq. (50) into Eq. (34) yields the final expression [Eq. (51)], which gives the effectiveness factor η of a first-order irreversible reaction in a spherical pellet as a unique function of the Thiele modulus:   3 1 1 η= − (51) φ tanh(φ) φ Equation (51) has been derived for a spherical catalyst pellet. The corresponding solution for the cylinder is η=

2 I1 (φ) φ I0 (φ)

(52)

and that for the flat plate is η=

tanh(φ) φ

(53)

where I1 (φ) and I0 (φ) denote the modified Bessel functions of first and zero order, respectively [31]. Figure 5 shows the dependence of the effectiveness factor on the Thiele modulus for the different pellet shapes. At small values of φ, the effectiveness factor approaches unity in all cases. Here, the chemical reaction constitutes the rate determining step – the corresponding concentration profiles over the pellet cross-section are flat (Fig. 4). This situation may occur at low catalyst activity (k is small), large pore size and high porosity (De is large)

1 f

10

100

Effectiveness factor η as a function of the Thiele modulus φ for different pellet shapes. Influence of intraparticle diffusion on the effective reaction rate (isothermal, first-order irreversible reaction).

Fig. 5

and/or small catalyst pellets (R is small, i.e. in fluidized bed reactors R is typically around 50 µm). At large values of φ, the function η(φ) for the sphere approaches the asymptotic solution η = 3/φ which, on the double logarithmic scale of Fig. 5, is a straight line with a slope of −1. The asymptotes for the cylinder and the flat plate obey the same slope (logarithmic scale), but do not cross the line η = 1 at φ = 3, as for the spherical case, but at values of φ = 2 and φ = 1, respectively [Eqs. (51)–(53)]. In this range of the Thiele modulus, pore diffusion controls the overall reaction rate. As a consequence, the concentration of reactant A1 tends to approach zero towards the center of the pellet (Fig. 4). This situation is typical for very active catalysts or catalysts with small pore size, low porosity and/or large pellet diameter. The observation that the slope of the asymptotic solution for η(φ) in Fig. 5 seems to be unaffected by the catalyst geometry suggests that the dependence of the effectiveness factor on the Thiele modulus might be described by a single asymptotic relationship, valid for arbitrary pellet shapes. In fact, Aris [32] showed that the curves of η(φ) for a sphere, infinite cylinder and infinite flat plate fall together almost perfectly when a generalized Thiele modulus φp is introduced, which is related to the ratio of the pellet volume to the external pellet surface (corresponding to the Sauter diameter) as a characteristic length scale of diffusion:  Vp kcbn−1 φp = (54) Sp De References see page 1781

1736

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

1

1 Flat plate

n=1 T = const.

Flat plate T = const.

8

Cylinder

n=2

6

Sphere

n=0

n=1

h

h

4

0.1

2

0.01 0.01

0.1

1

10

100

fp

Effectiveness factor η as a function of the generalized Thiele modulus φp for different pellet geometries. Influence of intraparticle diffusion on the effective reaction rate (isothermal, first-order irreversible reaction). Fig. 6

Figure 6 shows the effectiveness factor for any of the three different pellet shapes as a function of the generalized Thiele modulus φp . It is obvious that for larger Thiele moduli (i.e. φp > 3) all curves can be described with acceptable accuracy by a common asymptote η = 1/φp . The largest deviation between the solutions for the individual shapes occurs around η ≈ 1/φp . However, even for the extremely different geometries of the flat plate and the sphere, the deviation of the effectiveness factor is still below 10%. From a physicochemical point of view, this can be explained by the fact that in the range of strong intraparticle diffusion limitation only a small part of the outer pellet shell participates in the chemical reaction. In this situation, the local surface curvature is relatively unimportant [2]. The above considerations can also be extended in a simple way to different reaction orders, if the modulus φp is further modified. In this sense, Petersen [33] defined a generalized Thiele modulus φpn which takes into account the effects of the pellet shape as well as the influence of the reaction order:  n+1  n−1  Vp  2 kcb n+1 = φp φpn = (55) Sp De 2 In Fig. 7, the effectiveness factor is shown as a function of the generalized Thiele modulus φpn for different reaction orders (flat plate). It is obvious that, except for the case of a zero-order reaction, the curves agree fairly well over the entire range of interest. The asymptotic solution η = 1/φpn is valid for any reaction order and for values of the modulus φpn > 3.

0.1 0.1

2

4

6

8

1 fpn

2

4

6

8

10

Effectiveness factor η as a function of the generalized Thiele modulus φpn for different reaction orders. Influence of intraparticle diffusion on the effective reaction rate (isothermal, first-order irreversible reaction in a flat plate).

Fig. 7

However, in many practical situations the problem exists that effective rate constants and activation energies have been derived on the basis of laboratory experiments. The question then arises as to whether or not these parameters are influenced by transport effects. With the relations given so far, this question cannot be answered yet, since according to its definition by Eq. (27) the Thiele modulus is based on the intrinsic rate constant k. This problem can be solved by introducing a new modulus, which in contrast to φ contains only observable (effective) quantities and therefore can be determined without any knowledge of the intrinsic kinetics. This modulus is known as the Weisz modulus ψ. It is defined as the ratio of the effective pseudo-homogeneous reaction rate versus the maximum effective rate of diffusion at the external pellet surface [Eq. (56)]. Physically, the Weisz modulus can be looked at as the ratio of the time constants of diffusion τD = R 2 /De and reaction τR = cb /re . ψ=

R 2 re D e cb

(56)

Equation (56) can be used only for spherical catalyst pellets and first-order irreversible reactions. However, for convenience and in analogy with the Thiele modulus, a generalized modulus ψpn can also be defined which applies to arbitrary pellet shape and arbitrary reaction order. This is defined as  2 Vp n + 1 re ψpn = (57) Sp 2 D e cb Comparing Eqs. (56) and (57) and recalling the definition of the effectiveness factor according to Eq. (32)

6.3.3 Single Reactions (Conversion Problem)

yield the following simple relationship between the Thiele modulus and the Weisz modulus: ψ = ηφ 2 ;

2 ψpn = ηφpn

(58)

A plot of the effectiveness factor from Eq. (53) against the Weisz modulus ψpn from Eq. (58) gives the curve depicted in Fig. 8 for a first-order reaction. On the basis of this diagram, the effectiveness factor can be determined easily once the effective reaction rate and the effective diffusivity are known. Film and Pore Diffusion in an Isothermal Pellet The situation becomes more complicated when the reaction rate is already high enough to cause the reactant concentration to drop significantly across the external boundary layer. In this case, the combined effects of intraparticle and interphase mass transfer must be considered. If we again restrict our considerations to reactions that are accompanied only by minor heat effects, we may still assume a constant temperature inside the catalyst pellet and within the boundary layer. Hence, no solution of the enthalpy balance [Eq. (36)] is required. Moreover, the mass balance [Eq. (35)] takes exactly the same form as in the previous case of negligible interphase concentration gradients, thus giving rise to an identical general solution (see Section 6.3.3.1). Again, for the simplest case of a first-order irreversible reaction, the solution is given by Eq. (47). However, the boundary condition at the external pellet surface is now defined by Eq. (39) instead of Eq. (41). As a consequence, a different expression for the integration constant C1 results, which is not only a function of 6.3.3.2

1 8

Flat plate T = const.

6

n=0 n=1

n=2

h

4

2

0.1 0.1

2

4

6

8

1 ypn

2

4

6

8

10

Effectiveness factor η as a function of the generalized Weisz modulus ψpn for various reaction orders. Influence of intraparticle diffusion on the effective reaction rate (isothermal, first-order irreversible reaction in a flat plate). Fig. 8

1737

the Thiele modulus φ, but also depends on the Biot number for mass transport Bim . Hence a complete characterization of this problem already requires two parameters. After determining C1 from Eq. (39) and rearranging, the following solution for the normalized concentration inside the pellet is obtained:   sinh(φx) Bim tanh(φ) f = (59) x sinh(φ) φ + (Bim − 1) tanh(φ) Differentiation of this expression with respect to x and substitution of the derivative at the point x = 1 into Eq. (34) yields the final relationship for determining the catalyst efficiency:    1 3 1 Bim tanh(φ) − η= (60) φ tanh(φ) φ φ + (Bim − 1) tanh(φ) By comparing this relationship with the solution for the effectiveness factor in the absence of interphase concentration gradients [Eq. (51)], it becomes obvious that the overall effectiveness factor η can be expressed as the product of separate pore and external (film) effectiveness factors: η = ηpore ηext

(61)

where ηpore = re /rs and ηext = rs /rb . The term in the first square brackets in Eq. (60), together with the preceding factor 3/φ, is equivalent to Eq. (51); this represents the pore effectiveness factor ηpore , whereas the expression in the second square brackets denotes the external effectiveness factor ηext . For practical purposes however, Eq. (60) again suffers from the disadvantage that the Thiele modulus must be specified in order to calculate the catalyst efficiency. Hence the intrinsic rate constant must be known. In this situation, instead of directly plotting Eq. (60), it is more convenient to relate the effectiveness factor to the Weisz modulus, calculated from Eq. (58). For selected values of the Biot number Bim , such a diagram is given in Fig. 9. From this figure, it can be concluded that the reduction of the effectiveness factor at large values of ψ becomes more pronounced as the Biot number is decreased. This arises from the fact that the reactant concentration at the external pellet surface drops significantly at low Biot numbers. However, a clear effect of interphase diffusion is seen only at Biot numbers below 100. In practice, Bim typically ranges from 100 to 200. Hence the difference between the overall and the pore effectiveness factor is usually small. In other words, the influence of intraparticle diffusion normally is far more crucial than the influence of interphase diffusion. Thus, in many References see page 1781

1738

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

1

n=1 T = const.

reaction rate under bulk fluid phase conditions and dividing both sides of the resulting equation by 4πR 2 yields the relationship

Sphere

Ts = Tb + h

106 1

10

100

θs = 1 + 0.01 0.01

0.1

1

10

100

y Fig. 9 Effectiveness factor η as a function of the Weisz modulus ψ. Combined influence of intraparticle and interphase mass transfer on the effective reaction rate (isothermal, first-order irreversible reaction in a sphere, Biot number Bim as a parameter).

β β ηφ 2 = 1 + ψ 3Bih 3Bih

practical situations the overall catalyst efficiency may be replaced by the pore efficiency, as a good approximation. Film and Pore Diffusion Together with Interphase Heat Transfer When treating fast reactions which are accompanied by a substantial release or consumption of heat (e.g. oxidation or dehydrogenation reactions), one often has to take into account a possible limitation of the heat transfer between the catalyst pellet and the surrounding fluid phase, besides the combined effects of film and pore diffusion, since this may cause the catalyst to operate at a temperature significantly above or below the temperature of the bulk fluid phase. However, if the effective thermal conductivity of the catalyst material is high enough, intraparticle temperature gradients will be absent, irrespective of whether or not external temperature differences exist and the pellet itself may still be treated as isothermal. Assuming a constant temperature inside the pellet, it is sufficient to focus on the heat transfer across the external boundary layer instead of solving the general enthalpy balance given by Eq. (36). Therefore, we utilize the condition that the amount of heat generated (or consumed) inside the catalyst pellet per unit time, in a stationary mode of operation, must equal the amount of heat transported through the external fluid film per unit time. For a spherical catalyst, 6.3.3.3

(62)

Replacing the effective reaction rate re in Eq. (62) with the effectiveness factor, multiplied by the intrinsic

(64)

Supposing that heat and mass are transported by an identical convective mechanism, the catalyst temperature, given by Eq. (64), can also be expressed as a function of the Biot number for mass transport, Bim . For this purpose, the equivalence of the heat and mass transport factors (j -factors) is utilized [21]: jh = jm =

4 πR 3 re (−HR ) = 4πR 2 hf (Ts − Tb ) 3

(63)

which allows the unknown temperature at the external pellet surface to be determined. With the non-dimensional numbers already introduced, this equation can be rewritten in dimensionless form:

0.1

Bim = 0.1

R ηkcbn (−HR ) 3 hf

2 hf kf 2 P r 3 = Sc 3 ρucp u

where Pr is the Prandtl number: ν Pr = a

(65)

(66)

where a = λ/ρcp and Sc is the Schmidt number: Sc =

ν D

(67)

The ratio of the Schmidt number to the Prandtl number, which is known as the Lewis number, can be interpreted as the ratio of the thermal diffusivity to the molecular diffusivity of the fluid: Le =

Sc a = Pr D

(68)

According to Eqs. (65) and (68), the ratio hf /kf can be expressed using the Lewis number: 2 hf = ρcp Le 3 kf

(69)

By substituting hf from Eq. (69) into Eq. (64), the Biot number for heat transport can be replaced by the Biot number for mass transport, when additionally a modified Prater number β ∗ is introduced: β∗ =

cb (−HR ) 2

(70)

ρcp Tb Le 3 The unknown normalized pellet temperature θs is now given by θs = 1 +

β∗ ψ 3Bim

(71)

1739

6.3.3 Single Reactions (Conversion Problem)

According to Eq. (71), the temperature of the catalyst pellet can be calculated as a function of the Weisz modulus, for given values of the modified Prater number and the Biot number for mass transport. To derive an equation for determining the overall effectiveness factor, we first introduce a Thiele modulus φ ∗ , which is related to the unknown surface temperature Ts : (φ ∗ )2 = R 2 = R2

k(Ts )cbn−1 De k(Tb )cbn−1 k(Ts ) = φ 2 eγ (1−1/θs ) De k(Tb )

(72)

Replacing φ in Eq. (35) with φ ∗ transforms the mass balance into the same form as in the isothermal case. For a first-order irreversible reaction, the solution of the isothermal problem has already been derived (see Section 6.3.3.2). The reactant concentration is obtained analogously to Eq. (59) as a function of the radial coordinate, for given values of φ ∗ and Bim . We then have   Bim tanh(φ ∗ ) sinh(φ ∗ x) (73) f = x sinh(φ ∗ ) φ ∗ + (Bim − 1) tanh(φ ∗ ) The derivative of f , calculated at the point x = 1, is now substituted into Eq. (34), resulting in the following expression for the catalyst efficiency:    3 φ∗ Bim tanh(φ ∗ ) η= 2 −1 φ tanh(φ ∗ ) φ ∗ + (Bim − 1) tanh(φ ∗ ) (74)

which relates the Weisz modulus ψ to the modulus φ ∗ . Then, for a given value of φ ∗ , the corresponding value of ψ is calculated and from ψ the unknown catalyst temperature θs [Eq. (71)]. This temperature is substituted into Eq. (72) to obtain the corresponding value of the Thiele modulus φ. Dividing ψ by φ 2 finally yields the overall effectiveness factor, which is then plotted against ψ. This procedure yields the curves depicted in Fig. 10 for fixed values of Bim and γ and the modified Prater number β ∗ as a parameter. It is obvious that for exothermal reactions (β ∗ > 0) and large values of the Weisz modulus, effectiveness factors well above unity may be observed. The reason for this is that the decline of the reactant concentration over the external boundary layer and further towards the center of the catalyst pellet might be overcompensated by an increase in temperature. This happens when the heat flux transferred from the pellet surface across the boundary layer to the surrounding fluid per unit time is considerably smaller than the amount of heat generated inside the pellet per unit time. Then, the effective reaction rate is higher than would be expected under the conditions in the bulk fluid phase. Whether or not such an effect occurs in a practical situation and, if so, how pronounced it will be depends basically on the modified Prater number β ∗ [cf. Eq. (71)], that is, on the maximum amount of heat effectively produced inside the pellet, as compared with the

100

φ∗

The term in the first square brackets in Eq. (75), together with the preceding factor 3/φ ∗ , denotes the isothermal intraparticle effectiveness factor [cf. Eq. (51)]. The term in the second square brackets is identical with the isothermal interphase effectiveness factor [cf. Eq. (60)]. The exponential factor between the two terms describes the influence of the deviating catalyst temperature. Again, Eq. (75) cannot be used immediately to calculate the overall effectiveness factor, since the modulus φ ∗ , which is related to the unknown catalyst temperature, can only be determined when the overall efficiency has been specified [cf. Eqs. (71) and (72)]. Therefore, both sides of Eq. (74) are multiplied by φ 2 , resulting in an expression

10

n=1 Bim = 100 g = 20

b* = 1

Sphere

0.75 0.5

1

h

By substituting from Eq. (72) into Eq. (74), φ can be eliminated. The overall effectiveness factor can then be expressed in a more illustrative form:   1 3 1 − eγ (1−1/θs ) η= ∗ φ tanh(φ ∗ ) φ ∗   Bim tanh(φ ∗ ) × (75) φ ∗ + (Bim − 1) tanh(φ ∗ )

0.25 0

0.1

0.01 0.1

−1

1

10

100

y

Effectiveness factor η as a function of the Weisz modulus ψ. Combined influence of intraparticle and interphase mass transfer and interphase heat transfer on the effective reaction rate (first-order irreversible reaction in a sphere, Biot number Bim = 100, Arrhenius number γ = 20, modified Prater number β ∗ as a parameter). Fig. 10

References see page 1781

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

maximum amount of heat transported across the external boundary layer. Additionally, the Arrhenius number plays an important role, which, as a normalized form of the activation energy, is a measure of the increase in the reaction rate due to an increase in temperature. With the Thiele modulus, the Biot number for mass transport, the modified Prater number and the Arrhenius number, four dimensionless numbers are necessary to characterize this problem fully. Any of the curves in Fig. 10, which refer to different values of the modified Prater number β ∗ , tend to approach a certain limiting value of the Weisz modulus for which the overall effectiveness factor obviously becomes infinitely small. This limit can be easily determined, bearing in mind that the effective reaction rate can never exceed the maximum interphase mass transfer rate (the maximum rate of reactant supply) which is obtained when the surface concentration approaches zero. To show this, we formulate the following simple mass balance, analogous to Eq. (62): 4 πR 3 re = 4πR 2 kf (cb − cs ) 3

kf

(77)

b* = 1

n=1 Bim = 100 g = 30

100

0.75

Sphere 0.5

10

1

0.25 0 −1

0.01 0.001

0.01

0.1

1

hDaII

Fig. 11 Effectiveness factor η as a function of the observable variable ηDaII . Combined influence of intraparticle mass transfer and interphase heat and mass transfer on the effective reaction rate (first-order irreversible reaction in a sphere, Biot number Bim = 100, Arrhenius number γ = 30, modified Prater number β ∗ as a parameter).

n=1 Bim = 100 g = 10

b* = 2 1.5

1 1 0.5 0.1

cs = 1 − ηDaII cb

(79)

0 −1

The maximum effective reaction rate is obtained for the limiting value of cs = 0. This means that the product of the effectiveness factor and the second Damk¨ohler number can never exceed unity. A comparison of the definition of the Weisz modulus [Eq. (56)] with the definition of DaII [Eq. (78)] gives the following equivalence: R 2 re = D e cb

1000

Sphere

into Eq. (77), yields

ψ = ηφ 2 =

Hence, as an alternative to Fig. 10, the overall effectiveness factor can also be plotted against the product ηDaII , which again contains only measurable quantities and which, as already stated, can never exceed unity. This leads to the representation shown in Figs. 11 and 12 for two different values of the Arrhenius number.

(78)

3 cb R

(81)

h

DaII =

ψ 3Bim

0.1

Substituting the second Damk¨ohler number DaII , which is defined as the ratio of the intrinsic reaction rate to the maximum rate of reactant supply, k(Tb )cbn

ηDaII =

(76)

Dividing the above equation by the outer pellet surface area 4πR 2 and by the bulk concentration cb and solving for cs /cb gives ηk(Tb )cbn cs re =1− =1− 3 3 cb kf cb kf cb R R

or shorter:

h

1740

re 3kf R = ηDaII 3Bim 3 De kf cb R

(80)

0.01 0.001

0.01

0.1

1

hDaII

Fig. 12 Effectiveness factor η as a function of the observable variable ηDaII . Combined influence of intraparticle mass transfer and interphase heat and mass transfer on the effective reaction rate (first-order irreversible reaction in a sphere, Biot number Bim = 100, Arrhenius number γ = 10, modified Prater number β ∗ as a parameter).

6.3.3 Single Reactions (Conversion Problem)

Film and Pore Diffusion Together with Interphase and Intraparticle Heat Transfer When a fast reaction is highly exothermic or endothermic and, additionally, the effective thermal conductivity of the catalyst is poor, then significant temperature gradients across the pellet are likely to occur. In this case the mass balance [Eq. (35)] and the enthalpy balance [Eq. (36)] must be simultaneously solved using the corresponding boundary conditions [Eqs. (37)–(40)], to obtain the concentration profile of the reactant and the temperature profile inside the catalyst pellet. The exponential dependence of the reaction rate on the temperature thereby imposes a nonlinear character on the differential equations which rules out an exact analytical treatment. Approximate analytical solutions [34, 35] and numerical solutions [36–38] of Eqs. (35)–(40) have been reported. In the most general case, i.e. when intraparticle and interphase transport processes have to be included in the analysis, the effectiveness factor depends on five dimensionless numbers, namely the Thiele modulus φ, the Biot numbers for mass and heat transport Bim and Bih , the Prater number β (or the modified prater number β ∗ ) and the Arrhenius number γ . Once external transport effects can be neglected, the number of parameters reduces to three, because the Biot numbers then approach infinity and thus can be discarded. In Fig. 13, typical curves for the effectiveness factor as a function of the Thiele modulus φ are given for a first-order irreversible reaction in a spherical catalyst pellet. These curves were obtained numerically by Weisz and Hicks [38] for the case of negligible interphase concentration and temperature gradients. In the light of the previous results, it comes as no surprise that also under non-isothermal conditions inside the catalyst pellet effectiveness factors above unity are encountered (exothermic reactions). This happens when the heat of reaction, generated inside the porous pellet, accumulates due to poor thermal conductivity of the catalyst material. The temperature inside the pellet then increases until finally a steady state between heat generation and heat removal is reached, at a temperature level somewhat above the temperature at the external pellet surface. The corresponding rise of the reaction rate towards the center of the catalyst pellet may overcompensate the influence of the decreasing reactant concentration, thus producing a net increase in the reaction rate over the intrinsic rate which would be expected under the conditions in the bulk fluid phase. This effect will be particularly emphasized at small values of the Thiele modulus where the intrinsic rate of reaction and the effective rate of diffusion assume the same order of magnitude. At large values of φ, the effectiveness factor again becomes inversely proportional to the Thiele modulus, as observed under isothermal conditions (Section 6.3.3.1). Then the reaction takes place 6.3.3.4

1741

300.0 g=

100.0

EA RTS

= 20

b=

cs(−DHR)De leTs

h

10.0

1.0 b = 0.8 0.6 0.1

0.4 0.3 0.2 0.1 0 −0.2

−0.4 −0.8 0.01 0.1

1.0

10.0 f

100

1000

Effectiveness factor η as a function of the Thiele modulus φ. Combined effects of intraparticle heat and mass transfer on the effective reaction rate (first-order irreversible reaction in a sphere, Arrhenius number γ = 20, Prater number β as a parameter) (adapted from Weisz and Hicks [38]). Fig. 13

only within a thin shell close to the external pellet surface. Here, controlled by the Arrhenius and Prater numbers, the temperature may be distinctly higher than at the external pellet surface, but constant further towards the pellet center. However, whereas effectiveness factors above unity under non-isothermal conditions can be explained fairly easily, the observation of multiple steady states is a new and unexpected feature. These arise at small values of φ, provided that the reaction is substantially exothermic and, additionally, has a high activation energy. This means that, for a single value of the Thiele modulus, several possible solutions for the steady-state overall effectiveness factor may exist (operating points), usually up to three. The middle operating point is normally unstable. Whenever the temperature and/or the reactant concentration is increased by small-scale fluctuations of the operating conditions, the reaction is accelerated and the amount of heat additionally produced by the reaction inside the pellet per unit time surpasses the amount of heat that can be transported to the external pellet surface via effective heat conduction. The consequence is a further rise of the temperature inside the pellet. The system will reach a new stationary state between heat production References see page 1781

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

and heat removal when finally approaching the upper operating point. The opposite case is observed when a sudden drop of the temperature and/or the reactant concentration occurs. The temperature decreases steadily until the lower operating point is reached. Whether the upper or the lower operating point will prevail under stationary conditions depends on the direction from which the stationary state is approached. When the system reaches the stationary point from higher temperatures it will arrive at the upper operating point. If, however, the approach comes from lower temperatures, the system will operate at the lower stationary point. In the case that the operating points are fairly close together, a cyclic switching between the upper and lower operating points is also possible (ignition/quenching behavior). Whether or not multiple steady states will appear and how large the deviation of the effectiveness factors between both stable operating points will be are determined by the values of the Prater and Arrhenius numbers. Effectiveness factors above unity generally may occur when β > 0 (exothermic reactions). However, for the usual range of the Arrhenius number (γ = 10–30), multiple steady states are possible only at larger Prater numbers (Fig. 13). For further details on multiple steady states, the interested reader may consult the monograph by Aris [1] or the work of Luss [39, 40]. As stated above, instead of plotting the effectiveness factor against the Thiele modulus φ, which contains the unknown intrinsic rate constant, it is often more convenient to relate it to the observable Weisz modulus ψ. This leads to the representation given in Fig. 14 for the same situation as depicted in Fig. 13. The dashed portions of the curves indicate the regions in which a unique solution of the effectiveness factor does not exist, corresponding to the regions of multiple solutions in Fig. 13. Figures 13 and 14 refer to the situation where only intraparticle transport effects influence the observable reaction rate. However, a similar behavior is observed if, besides intraparticle heat and mass transport processes, the heat and mass transfer between the catalyst pellet and the bulk fluid phase is also considered. More information about this situation can be found, for example, in the work of Cresswell [41], McGreavy and Cresswell [42, 43], Kehoe and Butt [44], Hatfield and Aris [45], Morbidelli and Varma [46] and Butt [31, 47]. Although multiplicities of the effectiveness factor have also been detected experimentally, these are of minor importance, practically, since, for industrial processes and catalysts, Prater numbers above 0.1 are less common. In contrast, effectiveness factors above unity in real systems are frequently encountered, although the dominating part of the overall heat transfer resistance normally lies in the external boundary layer rather than

300.0 g= 100.0

EA RTS

= 20

b = 0.8

10.0

0.6 h

1742

1.0 0.4

0.2

0.1 b=

cs(−DHR)De

0.01 0.1

leTs 1.0

−0.8 10 y

−0.2 100

0 1000

Fig. 14 Effectiveness factor η as a function of the observable Weisz modulus ψ. Combined effects of intraparticle heat and mass transfer on the effective reaction rate (first-order irreversible reaction in a sphere, Arrhenius number γ = 20, Prater number β as a parameter) (adapted from Weisz and Hicks [38]).

inside the catalyst pellet. For mass transfer the opposite holds; the dominating diffusional resistance is normally located within the pellet, whereas the interphase mass transfer most frequently plays a minor role (high space velocity). The fact that the effective heat conduction within the catalyst pellet normally is not the crucial process in determining the excess temperature of the catalyst pellet can be illustrated by analyzing the contribution of the interphase temperature gradient to the total excess temperature. To demonstrate this, we first eliminate the common reaction term occurring in the mass and enthalpy balances [Eqs. (35) and (36)]: 
d2 θ b dθ b df d2 f (82) + + = −β dx 2 x dx dx 2 x dx Upon introducing the new variables u = fx and v = θx [see Eqs. (43) and (44)], this equation simplifies to d2 v d2 u = −β 2 2 dx dx

(83)

Integrating Eq. (83) twice and utilizing the boundary conditions at the pellet center,   df du  =x +f =f (84) dx x=0 dx x=0

6.3.3 Single Reactions (Conversion Problem)

  dv  dθ = x +θ =θ dx x=0 dx x=0

1

(85)

50 20

and, at the external pellet surface,

0.8

(86)

v(x = 1) = θs

(87)

gives the following simple relation between θ and f : (88)

From Eq. (88), we conclude that the maximum temperature difference inside the pellet is related to the point at which the maximum concentration difference is observed, i.e. where the reactant concentration drops to zero (f = 0). Therefore, we have a maximum internal temperature difference Tpore of Tpore = βTb

cs cb

(90)

The interphase temperature difference is given by Eq. (71) as Text = Tb

β∗ ψ = Tb β ∗ ηDaII 3Bim

rβ ηDaII Text = Text + Tpore 1 + ηDaII (rβ − 1)

β∗ Bim = β Bih

1 = rb = 0.4

Bim Bih

0.2

0.0 0.0

0.2

0.4

0.6

0.8

1

hDaII

Fig. 15 Ratio of interphase temperature difference to total intraparticle–interphase temperature difference rT as a function of the observable variable ηDaII (spherical catalyst, ratio of Biot numbers rβ = Bim /Bih as a parameter) (after Carberry [21]).

be treated as isothermal, at a temperature level which is controlled by the interphase heat transfer resistance. However, this leads to a reduction of the general problem to the case which has already been treated in the previous section. External Heat and Mass Transfer A special type of catalyzed fluid–solid reaction is obtained when either the reaction rate is so fast that the reactants are completely exhausted at the external catalyst surface (i.e. at very high reaction temperatures) or when the catalyst is non-porous. Then, pore diffusion and effective heat conduction inside the pellet need not be considered. Thus, the problem is reduced to a treatment of the coupled interphase heat and mass transport. The conservation equations for mass and enthalpy for this special situation have already been given with Eqs. (76) and (62). As there is no diffusional mass transport inside the pellet, the overall catalyst effectiveness factor is identical with the film effectiveness ηext , which is defined as the ratio of the effective reaction rate under surface conditions divided by the intrinsic chemical rate under bulk fluid phase conditions [cf. Eq. (61)]. For an nth-order irreversible reaction, we have the following expression:   n  k(Ts )csn cs k(Ts ) = (94) η = ηext = k(Tb )cbn k(Tb ) cb 6.3.3.5

(92)

where rβ is the ratio of the Prater numbers or the Biot numbers: rβ =

2

0.6

(91)

If we denote by rT the ratio of the interphase temperature difference Text versus the total excess temperature of the catalyst Text + Tpore , we finally obtain, after some rearrangement, rT =

5

(89)

The concentration ratio cs /cb in Eq. (89) can be expressed as a function of the observable variable ηDaII [Eq. (79)]. Hence, Tpore = βTb (1 − ηDaII )

10

rDT

u(x = 1) = fs

θ − θs = β(fs − f )

1743

(93)

The term rT in Eq. (92) is depicted in Fig. 15 as a function of the observable variable ηDaII for different values of rβ . When rβ = 1, a linear rise of the curve is obtained. However, with increasing rβ , the external temperature difference dominates already at low values of ηDaII . In reality, the Biot number for mass transport in most cases is considerably higher than the Biot number for heat transport, i.e. the ratio rβ is frequently larger than 40–50 [21]. This means that the catalyst pellet can usually

References see page 1781

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

According to this equation, the effectiveness factor is controlled by two terms, namely the ratio of the rate constants ks /kb , governed by the temperature difference over the external fluid film, and the ratio of the surface versus the bulk concentration cs /cb . Defining equations for both of these terms have already been given with Eqs. (71) and (79). Substituting these into Eq. (94) and using Eq. (81), we obtain the effectiveness factor for arbitrary reaction order as a function of the observable variable ηDaII : η = (1 − ηDaII )n eγ [1−1/(1+β

∗ ηDa )] II

100

n=1 g = 20

b* = 0.5

Sphere 0.2

h

0.1 1

0 −0.2 −0.5

0.1

0.01 0.001

0.01

T = const. −1

Sphere 1

1/2

n=2

1

0.1

(95)

Figure 16 shows an effectiveness factor diagram for a first-order irreversible reaction which has been calculated from Eq. (95) for various values of the modified Prater number β ∗ . It can be seen that for exothermal reactions (β ∗ > 0), effectiveness factors above unity may be observed when the catalyst operates at a temperature substantially above the bulk fluid phase temperature. This is caused by the limited heat transfer between the pellet and the surrounding fluid. The crucial parameters controlling occurrence and size of this effect are again the modified Prater number and the Arrhenius number. The reaction order also has a great influence on the effectiveness factor. This is illustrated in Fig. 17 for the isothermal case (β ∗ = 0). Clearly, the effectiveness factor drops with increasing reaction order at a given value of ηDaII . For reactions with a negative order, effectiveness factors above unity are possible even under isothermal conditions.

10

10

h

1744

0.1

1

hDaII Fig. 16 Effectiveness factor η as a function of the observable variable ηDaII . Combined influence of interphase heat and mass transfer on the effective reaction rate (first-order irreversible reaction in a sphere, modified Prater number β ∗ as a parameter).

0.01 0.001

0.01

0.1

1

hDaII Fig. 17 Effectiveness factor η as a function of the observable variable ηDaII . Influence of interphase mass transfer on the effective reaction rate (isothermal, nth-order irreversible reaction in a sphere, reaction order n as a parameter).

Use of Complex Rate Expressions In the previous sections, only simple, irreversible reactions have been considered, the kinetics of which were assumed to obey a simple power rate law of the type r = kcn . The reason for this assumption was to have analytical solutions for most of the important problems in order to demonstrate the key effects in a clear manner. Moreover, many heterogeneously catalyzed reactions, although not strictly obeying a power rate law, can nevertheless be described by this kind of rate expression for practical purposes, at least when the concentration range to be covered is not too wide. In reality, however, situations also exist where a more complex form of the rate expression has to be applied. Among the numerous possible types of kinetic expressions, two important cases will be discussed here in more detail, namely rate laws for reversible reactions and rate laws of the Langmuir–Hinshelwood type. Basically, the purpose of this is to point out additional effects concerning the dependence of the effectiveness factor upon the operating conditions which result from a more complex form of the rate expression. Moreover, without going too much into the details, it is intended at least to demonstrate to what extent the mathematical effort required for an analytical solution of the governing mass and enthalpy conservation equations is increased and how much a clear presentation of the results is hindered whenever complex kinetic expressions are necessary. To keep the discussion as simple as possible, the view is restricted to the isothermal case without interphase concentration and temperature gradients (i.e. intraparticle diffusion and reaction only). For a more detailed treatment 6.3.3.6

6.3.3 Single Reactions (Conversion Problem)

of complex kinetic expressions, the monographs by Satterfield [18], Aris [17] and Emig [2] may be consulted. Simple Reversible Reactions We start with the simplest reversible reaction, A1   A2 , where both the forward and the reverse reaction follow a first-order rate law. The net reaction rate of such a process is given by

1

n = 10 T = const. Sphere

h

6.3.3.6.1

r = k + c1 − k − c2

1745

0.1

Keq = 0.01

1 10 106

0.1

(96)

This type of reaction has been investigated by Smith and Amundsen [48] and Carberry [49]. Without derivation, it may be stated that under isothermal conditions the same solution for the effectiveness factor is obtained as in the case of an irreversible reaction if a modified Thiele modulus φrev is introduced [18]:  k− k+ + (97) φrev = R D1,e D2,e If the effective diffusivities D1,e and D2,e do not differ markedly, then Eq. (97) can be expressed in simplified form by introducing the equilibrium constant Keq = k + /k − :  k + (1 + 1/Keq ) (98) φrev = R De From this relation, it can be seen that the modulus φrev transforms to the standard Thiele modulus φ [Eq. (27)] when the equilibrium constant approaches infinity. Additionally, it is obvious that the effectiveness factor decreases when, at a given value of the forward rate constant k + , the reverse reaction becomes increasingly important (Fig. 18). This holds for all types of reversible reactions [2, 18]. Therefore, the effectiveness factor of a truly reversible reaction might be considerably overestimated if the reaction is treated as irreversible. This problem arises even at low conversion, because although the product concentration may then be negligible at the external pellet surface, in general this is not the case inside the pellet, unless diffusion effects upon the effective reaction rate are absent. Therefore, normally a difference exists between the effectiveness factor of irreversible and reversible reactions which becomes increasingly important as the equilibrium constant Keq of the reaction is shifted to smaller values. At large values of φrev (i.e. φrev > 10), the proportionality η ∼ 1/φrev [cf. Eq. (51)] for the given example of a firstorder reaction [Eq. (98)] dictates a decrease of the catalyst efficiency by the factor (1 + 1/Keq )1/2 . The mathematical effort increases considerably when the influence of a second reactant is considered. We then have a reaction of the type A1 + A2   A3 + A4 . According

0.01 0.01

0.1

1

10

100

y

Effectiveness factor η of a first-order reversible reaction versus the Weisz modulus ψ (related to the forward rate constant k+ ). Influence of intraparticle diffusion on the effective reaction rate (isothermal reaction in a sphere, equal diffusivities D1,e = D2,e , equilibrium constant Keq as a parameter).

Fig. 18

to Eq. (99), the net reaction rate for this general case of a bimolecular equilibrium reaction is a function of four different concentrations: r = k + c1 c2 − k − c3 c4

(99)

Assuming a flat plate geometry, Maymo and Cunningham [50] developed a relationship for calculating the effectiveness factor for reactions of the above type. This depends on the Thiele modulus and the equilibrium constant of the reaction, but additionally also on the stoichiometric ratio of the reactants and their effective diffusivities. In the most general case, eight parameters have to be specified to determine the effectiveness factor. Even if a single common effective diffusivity can be used to describe the diffusion of the various reactants inside the porous pellet with sufficient accuracy, there still remains a five-parameter equation. Therefore, a graphical representation of the results is not very illustrative and hence has been omitted here. Moreover, the dependence of the effectiveness factor on the Weisz modulus and on the equilibrium constant is very similar to the previously treated case of a first-order reversible reaction. The main contribution of the work of Maymo and Cunningham [50] therefore is a criterion allowing a decision on whether or not a reversible reaction obeying a truly second-order rate law may be represented in simplified form by an irreversible, pseudo-first-order expression. It turns out that the error in the effectiveness factor, associated with such a simplification, will be References see page 1781

1746

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

smaller than 10% whenever the product of the normalized concentration and the normalized diffusivity of the second reactant E2 D2 is ≥2, where E2 and D2 are defined by E2 = c2,s /c1,s and D2 = D2,e /D1,e . Finally, it is worth noting that, according to Carberry [21], heat transfer effects are of only minor importance in reversible reactions, since the effect of temperature on the equilibrium constant compensates at least partly for the effect of temperature on the rate constants. As an example, the temperature rise towards the pellet center, which is observed for exothermic equilibrium reactions, leads to a decrease in the equilibrium constant, but in addition also to an increase in the forward rate constant [18]. 6.3.3.6.2 Simple Irreversible Reactions with Langmuir– Hinshelwood-Type Kinetics Similarly to reversible reactions, the mathematical treatment becomes more complicated if a Langmuir–Hinshelwood (Hougen– Watson)-type kinetic expression is used. For the simplest case of an irreversible, monomolecular reaction A1 → products, the following rate law holds:

r=

kp1 1 + K1 p1 +

(100) Ki pi

r=

i =1

Among others, Roberts and Satterfield [51, 52] analyzed this type of reaction. On the basis of numerical calculations for a flat plate, these authors presented a solution in the form of effectiveness factor diagrams, from which the effectiveness factor can be determined as a function of the Weisz modulus as well as an additional parameter Kp1,s which considers the influence of the different adsorption constants and effective diffusivities of the various species [18]. The constant K involved in this parameter is defined as follows:

1 + K1 p1 + K2 p2 +

(103) Ki pi

Compared with the previous case of a monomolecular reaction, here an additional parameter E is required to take into account the influence of the stoichiometric ratio of the reactants on the effectiveness factor. This is defined as  −D2,e p2,s −1 (104) E= ν2 D1,e p1,s

Ki νi i =1

K=

kp1 p2

i =1,2

1.0

Di,e

ω

where ω is given by

 D1,e Ki pi,s + p1,s νi ω =1+ Di,e

Kp1,s ∞(zero order)

(101) 0.0 (first order)

h

K1 − D1,e

Kp1,s tends towards infinity, we obtain a simple zeroorder expression. If at least one of the products is strongly adsorbed at the surface, negative values of Kp1,s (i.e. of K) may result, which means that the reaction rate is inhibited. The smallest possible value of Kp1,s is −1. In Fig. 19, calculated curves of the effectiveness factor versus the Weisz modulus are shown for different values of Kp1,s [18]. For comparison, this diagram also contains the curves corresponding to the results which apply to simple, irreversible power rate laws of zeroth, first and second order. From this figure, it is obvious that a strong adsorption of at least one of the products leads to a decrease in the effectiveness factor similar to that observed in the case of a reversible reaction. In addition to the convenient representation of the effectiveness factor as a function of the observable Weisz modulus, this diagram has the advantage that the error arising from an approximation of the truly hyperbolic form of the rate expression by a simple power rate law with integer order can be estimated. Roberts and Satterfield [51, 52] also treated bimolecular reactions of the type A1 + bA2 → products, described by Langmuir–Hinshelwood-type models. In this case, we have the rate expression

0.1

(102)

i =1

p1,s is the partial pressure of reactant A1 at the external pellet surface and vi is the stoichiometric coefficient of species i which is positive for all products and negative for the reactants. When the adsorption constant K1 , and also all other adsorption constants Ki , tend towards zero, the parameter Kp1,s also approaches zero. We then have a simple first-order reaction. On the other hand, when

0.01 0.01

20.0 5.0

−0.30 −0.50 (second order) −0.70 −0.90 −0.95 −0.98 0.1

1.0

1.0

10.0

100.0

y

Fig. 19 Effectiveness factor η of an irreversible monomolecular reaction with Langmuir–Hinshelwood-type kinetics versus the Weisz modulus ψ. Influence of intraparticle diffusion on the effective reaction rate (isothermal reaction in a flat plate, Kp1,s as a parameter) (adapted from Satterfield [18]).

6.3.3 Single Reactions (Conversion Problem)

In practice, A1 is chosen to designate that reactant that permits E to be zero or positive. E will be large whenever reactant A2 is present in large excess and/or has an effective diffusivity substantially larger than reactant A1 . It is therefore called the modified stoichiometric excess. A special feature of bimolecular reactions is that ω, defined by Eq. (102) and normally positive, might become negative when reactant A2 is adsorbed much more strongly than reactant A1 and when, additionally, the value of D2,e p2,s /v2 is very small. For ω < 0, the method of Roberts and Satterfield cannot be applied [18]. Inspection of the curves of the effectiveness factor versus the Weisz modulus for different values of Kp1,s and E reveals two interesting phenomena when E > 0 (Fig. 20) [18, 51, 52]. At first, for large values of Kp1,s (10–100), effectiveness factors above unity may occur even though isothermal conditions prevail. This can be explained by the fact that the reaction rate given by Eq. (103) has a maximum for certain combinations of p1 and p2 . This maximum results from the assumption that the rate is proportional to the concentration of the adsorbed reactants A1 and A2 which compete for adsorption sites on the active (inner) surface. When, for example, A1 is adsorbed more strongly than A2 , then a raised partial pressure of A1 , at constant partial pressure of A2 , will lead to a displacement of A2 from the surface and hence to a lower reaction rate. By a quantitative analysis, it can be shown that effectiveness factors above unity will appear whenever Kp1,s is greater than (E + 2)/E [18]. The second effect concerns the phenomenon that, for large values of E and Kp1,s , a region is observed where the effectiveness factor is no longer a unique function of the Weisz modulus. This is in analogy to the multiplicities

Kp1,s = 100.0

E = 10

10.0 zero order 1.0 0

h

1.0

0.1 −0.40 −0.70 −0.90 0.01 0.001

observed in the non-isothermal case for simple integerorder power rate laws (Figs. 13 and 14). To give an example, Fig. 20 shows a diagram for E = 10 and various selected values of Kp1,s . The dashed lines indicate the range over which multiple steady states of η(ψ) might occur. Here, by means of numerical methods it is not possible to determine a unique solution for the effectiveness factor of the pellet for given conditions at the external pellet surface [18]. Which operating point is observed in a real situation again depends on the direction from which the stationary state is approached [18]. Simple Reversible Reactions with Langmuir– Hinshelwood-Type Kinetics The same approach as for irreversible Langmuir–Hinshelwood-type models can be extended to reversible reactions. Kao and Satterfield [53] developed a graphical method for monomolecular reversible reactions of the type A1   A2 , which is presented here as our last example. The method is based on the following formulation of the net reaction rate:   k p1 − p2 (p1,eq /p2,eq ) (105) r=

1 + K1 p1 + K2 p2 + Ki pi 6.3.3.6.3

i =1,2

where p1,eq is  the equilibrium partial pressure of reactant A1 and Ki pi represents the inhibition of the reaction rate by adsorption of inert species. Compared with the irreversible case [Eqs. (100)–(102)], which is unambiguously characterized by the Weisz modulus and the parameter Kp1,s , two new parameters B and C are now required. Of these, B is an extension of the parameter Kp1,s to the reversible case, whereas C is truly an additional parameter. They are defined as follows: K(p1,s − p1,eq ) 1 + Kp1,eq p1,eq C= p1,s B=

10.0

0.01

0.1

1.0

10.0

100.0

y Fig. 20 Effectiveness factor η for a bimolecular irreversible reaction with Langmuir–Hinshelwood-type kinetics versus the Weisz modulus ψ. Influence of intraparticle diffusion on the effective reaction rate (isothermal reaction in a flat plate, modified stoichiometric excess E = 10, Kp1,s as a parameter) (adapted from Satterfield [18]).

1747

(106) (107)

For p1,eq = 0 we have an irreversible reaction. B is then equivalent to the parameter Kp1,s . Consequently, B ranges from a minimum value of −1 to infinity (as well as Kp1,s ), where B can never exceed Kp1,s . Possible values of C range from 0 for an irreversible reaction to 1 when the reaction reaches the equilibrium and the net reaction rate approaches zero. The effectiveness factor versus the Weisz modulus according to Kao and Satterfield [53] is shown in Fig. 21 for C = 0.5 and different values of B. From this diagram, a similar behavior is seen as in the case of a simple, first-order reversible reaction (Fig. 18): with decreasing References see page 1781

1748

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

1.0 B = 50.0 5.0 1.0 0.0

C = 0.5 B = −0.50 −0.70 −0.90 −0.95

h

0.1

−0.98

0.01 0.001

0.01

0.1 y

1.0

10.0

Effectiveness factor η of a monomolecular reversible reaction with Langmuir–Hinshelwood-type kinetics versus the Weisz modulus ψ. Influence of intraparticle diffusion on the effective reaction rate (isothermal reaction in a flat plate, equilibrium parameter C = 0.5, B as a parameter) (adapted from Satterfield [18]). Fig. 21

values of B, the effectiveness factor is reduced. A decline of the effectiveness factor is also observed for an increase in the parameter C, which corresponds to a shift towards the chemical equilibrium and hence to a reduction of the net reaction rate [18]. 6.3.4

Temperature Dependence and Reaction Order of Transport-Limited Reactions

As stated in the Introduction, the observable activation energy and normally also the observable order of a heterogeneously catalyzed fluid–solid reaction, carried out on a porous catalyst, may differ from the respective properties of the intrinsic chemical reaction, whenever notable heat and mass transfer resistances exist (Fig. 3). Therefore, the observable quantities are frequently termed effective or apparent properties in order to provide a clear distinction from the true or intrinsic parameters of the pure chemical reaction. This aspect is now discussed in more detail, since it is of great importance for practical situations. Two important restrictions must be introduced to allow a general representation of the temperature and concentration dependence of the effective reaction rate in the diffusion-controlled regime. The first concerns the restriction to simple reactions, i.e. which can be described by only one stoichiometric equation. Whenever several reactions occur simultaneously, it is obvious that the individual activation energies and reaction orders may be influenced differently by transport effects. Therefore, how the coupled system in such a case finally will respond to a change of temperature or concentration cannot be specified in a generally valid form. The second necessary condition is isothermal operation. This is apparent from the results in Sections 6.3.3.3

and 6.3.3.4, where it was shown that heat and mass transport may drive the effective reaction rate in opposite directions. Normally, mass transfer control of a reaction means a drop of the effective reaction rate (for positive reaction order), whereas limited heat transfer in the case of an exothermal reaction will cause the temperature inside the catalyst pellet to rise and will thus increase the effective reaction rate. When both effects occur simultaneously, either an increase or a decrease in the effective rate may be observed, indicating either a lower or a higher apparent activation energy (or reaction order). Again, this rules out a generalized treatment. However, when the view is restricted to simple, irreversible reactions obeying an nth-order power rate law and if, additionally, isothermal conditions are supposed, then – together with the results in Section 6.3.3 – it can be easily understood how the effective activation energy and the effective reaction order will change during the transition from the kinetic regime to the diffusioncontrolled regime of the reaction. Intraparticle Diffusion When the effective rate is controlled by pore diffusion, then the asymptotic solution of the catalyst effectiveness factor as a function of the generalized Thiele modulus can be utilized [Eq. (108)]. This (approximate) relation was derived in Section 6.3.3.1. It is valid for arbitrary order of reaction and arbitrary pellet shape. 6.3.4.1

η=

  φpn > 3

1 φpn

(108)

Combining Eq. (108) with the definition of the effectiveness factor [Eq. (32)], the following expression for the effective reaction rate is obtained: re = ηrb = ηkcbn =

1 kcn φpn b

(109)

Substituting the generalized Thiele modulus from Eq. (55) yields re =

 Vp Sp

kcbn n−1

(110)

n + 1 kcb 2 De

The temperature dependence of the rate constant k is normally expressed by an Arrhenius law with intrinsic activation energy EA . In contrast, the temperature dependence of the effective diffusivity De is much weaker. Normally, De is obtained from De =

εp D τ

(111)

6.3.5 Diagnostic Criteria and Experimental Methods for Estimating the Influence of Heat and Mass Transfer on the Effective Reaction Rate

1749

where εp denotes the void fraction (porosity) of the catalyst and τ is the empirical tortuosity factor. In practice, the diffusivity D is most often approximated either by the bulk diffusivity DM or the Knudsen diffusivity DK , depending on whether the diffusion process is governed by bulk or Knudsen diffusion, i.e. whether the characteristic pore size of the catalyst is substantially larger than or close to the mean free path of the molecules (cf. Chapter 6.2). The dependence of DM and DK on temperature is usually described by

(115)

3

DM ∼ T 2 ;

DK ∼ T

1 2

(112)

As both εp and τ are virtually uninfluenced by the temperature, the temperature dependence of De equals that of DM or DK , respectively. If we then formally introduce, instead of Eq. (112), an Arrhenius law with ED representing the activation energy of the diffusion process, after some rearrangement, we obtain from Eq. (110) the following solution: re = Ce

−(EA +ED ) (n+1)/2 2RT cb

(113)

where  k0 D0,e  C= Vp n + 1 Sp 2

Interphase Mass Transfer If the effective rate is controlled by interphase mass transfer, then we may utilize the equivalence between the effective rate and the mass transfer rate: 6.3.4.2

re = ηkcbn = kf a(cb − cs )

Provided that the interphase mass transfer resistance 1/kf is sufficiently large, the reactant concentration at the external pellet surface will drop almost to zero, hence we may neglect the surface concentration cs compared with the bulk concentration cb . With cs → 0 in Eq. (115), it is obvious that in this case the reaction will effectively follow a first-order rate law. Moreover, it is also clear that the temperature dependence of the effective reaction rate is controlled by the mass transfer coefficient kf . This exhibits basically the same temperature dependence as the bulk diffusivity DM , since the boundary layer thickness δ is virtually not affected by temperature (kf = DM /δ). Therefore, we have the rule of thumb that the effective activation energy of an isothermal, simple, nth-order irreversible reaction will be less than 5–10 kJ mol−1 when the overall reaction rate is controlled by interphase diffusion. 6.3.5

(114)

From this relationship, it is obvious that the effective activation energy is given by the mean of the intrinsic activation energy of the catalyzed reaction EA and the activation energy of the effective diffusion ED . Additionally, it can be noticed that the effective reaction order (n + 1)/2 also tends to the average of the intrinsic reaction order n and the order of the diffusion process nD = 1. Because of the weak dependence of the effective diffusivity upon temperature [Eq. (112)], which corresponds to an activation energy of less than 5–10 kJ mol−1 , ED can normally be neglected compared with EA . Therefore, we obtain, as a rule of thumb, that the observable activation energy of an isothermal, simple, nth-order irreversible reaction will drop to roughly half of the true value when the reaction is carried out under intraparticle diffusion control. As can be easily derived from Fig. 3, in combination with Eqs. (113) and (114), the range of the diffusioninfluenced region may be extended or restricted by changing the characteristic length Vp /Sp of the porous catalyst pellet. By application of small enough particles, one can usually manage to stay in the kinetic regime. This will be treated in more detail in Section 6.3.5.

Diagnostic Criteria and Experimental Methods for Estimating the Influence of Heat and Mass Transfer on the Effective Reaction Rate

In real situations, the question frequently arises as to whether or not a marked influence of heat and mass transfer on the observable reaction rate may be expected under certain reaction conditions. Fairly often, then, one has to deal with reactions obeying complex kinetics where either no or only a very cumbersome analytical solution is possible based on the methods described in Section 6.3.3. For such cases a number of useful diagnostic criteria have been developed in the past, derived either from asymptotic solutions of the governing differential equations or from perturbation methods [54]. Most of these criteria have been explained in a detailed review by Mears [55]. More recent surveys of diagnostic transport criteria have been given by, for example, Butt [31] and Madon and Boudart [56]. Typically, a criterion is derived on the premise that the net transport effect should not alter the true chemical rate by more than some arbitrarily specified amount, normally 5%. Because of the uncertainty involved in knowing some of the necessary parameters and since they are based on approximate rather than exact solutions, the philosophy of using the criteria should be conservative. References see page 1781

1750

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

As a general rule, a clear decision on whether a reaction takes place under kinetic or diffusion control is possible only when the calculated value of a criterion is significantly above or below the respective limiting value (i.e. an order of magnitude), otherwise a more detailed analysis is recommended. Corresponding to the different use of the criteria, a subdivision into two groups appears to be useful. Experimental criteria are needed when the kinetics of the reaction under consideration are still unknown, i.e. neither the type of rate law nor the intrinsic values of the kinetic parameters have yet been identified. This may be the case during an early stage of a laboratory kinetic study when a new reaction is analyzed for the first time. Experimental criteria in general contain only directly observable quantities, i.e. the measured effective rate of reaction and some (effective) physical properties of the catalyst and the reaction mixture (R, De , λe , etc.). Therefore, these can be easily applied. However, experimental criteria suffer from the disadvantage of sometimes being less conservative when more complex kinetics prevail. Theoretical criteria normally contain an explicit expression of the intrinsic chemical rate and optionally also a measured value of the observed reaction rate. Hence these criteria are useful only when the intrinsic kinetics are available and one is interested in, for example, whether or not transport effects are likely to influence the performance of the catalyst as the operating conditions are changed. If it is not possible to generate a numerical solution of the governing differential equations, due either to a lack of time or to other reasons, then the use of theoretical criteria will not only save experimental effort, but also provide a more reliable estimation of the net transport influence on the observable reaction rate than simple experimental criteria can give, which do not contain any explicit information about the true concentration and temperature dependence of the intrinsic rate. Experimental Criteria Table 2 lists most of the available experimental criteria for intraparticle heat and mass transfer. These criteria apply to single reactions only, where it is additionally supposed that the kinetics may be described by a simple nth-order power rate law. The most general of the criteria, 5 and 8 in Table 2, ensure the absence of any net effects (combined) of intraparticle temperature and concentration gradients on the observable reaction rate. However, these criteria do not guarantee that this may not be due to a compensation of heat and mass transfer effects (this point has been discussed in the previous section). In fact, this happens when γβ ≈ n [31]. 6.3.5.1

As the most conservative policy, it may therefore be recommended to analyze separately whether the necessary conditions of isothermicity inside the pellet are met, for example by criterion 4 or 7. Then, when isothermal conditions can be ensured, criterion 3 may be used to check for concentration gradients. A similar situation occurs when interphase transport effects are considered. Table 3 gives a survey of experimental criteria for the estimation of interphase transport effects. The most general relationship here is criterion 4. However, again it may be suggested that the separate isothermicity criterion 5 be used first and then, if isothermal conditions prevail, be looked for concentration gradients using relation 1 or 2 in Table 3 or, if necessary, the respective criteria in Table 2. The criteria given in Tables 2 and 3 apply to spherical catalyst pellets. If other geometries have to be treated, then the sphere radius may be replaced by the ratio of the pellet volume divided by the external pellet surface area, Vp /Sp [31]. At this point, it should be mentioned that there may be some doubt about how successful the non-isothermal criteria are in involving observable quantities only. This concerns the fact that, in the non-isothermal case, one has to specify the Arrhenius number, which contains the true activation energy of the catalyzed reaction. The above statement would obviously define the true activation energy as a directly observable quantity in the nonisothermal criteria. However, this would presumably be an experimental value derived from studies in which intraparticle transport effects were absent, which is precisely what one is attempting to define [31]. Theoretical Criteria Table 4 summarizes a number of well-known theoretical diagnostic criteria for the estimation of intraparticle transport effects on the observable reaction rate. Table 5 gives a survey of the respective criteria for interphase transport effects. It is obvious that these are more difficult to use than the simple experimental relations given in Tables 2 and 3. In general, the intrinsic rate expression has to be fully specified and, additionally, either an integration of the rate expression must be performed or the first derivative of the intrinsic rate with respect to concentration at surface conditions is required. In the non-isothermal case, i.e. for estimating the combined effect of heat and mass transport limitations, the rate expression to be integrated or differentiated must be uniquely related to the concentration, which is generally achieved by substituting the heat balance into the reactant material balance. In contrast to the experimental criteria in Tables 2 and 3, which apply to power law kinetics only, the criteria in Tables 4 and 5 can be used for arbitrary forms of rate 6.3.5.2

6.3.5 Diagnostic Criteria and Experimental Methods for Estimating the Influence of Heat and Mass Transfer on the Effective Reaction Rate Tab. 2 Experimental diagnostic criteria for the absence of intraparticle transport effects in simple, irreversible reactions (power law kinetics only)

Application Pore diffusion (T = constant)

No.

Criterion R2 re < 1; n = 1 De cs n=0 R2 re 20. Conversely, for Ke2,1 < 0.05 an approximation of the effective diffusivity valid in the viscous friction dominated range would be  1 κ1 1 = + (178) D1,e − D12 B0 The quantity κ1 in Eqs. (175), (177) and (178) denotes the fractional viscous contribution to the overall friction between component 1 and the pore wall. For the general case of component i, this is defined as κi =

ηi0 1 m p xj ςij

(179)

j =1

where ηi0 represents the pure-component viscosities and ςij the binary Wilke parameters for calculating the mixture viscosity [79]:

1   1 1 + ηi0 /ηj0 2 Mj /Mi 4

ςij =



8(1 + Mi /Mj )

1

2 (180)

2

For a porous catalyst, DiM and DiK in the above equations are replaced by the effective properties analogous to Eq. (111) using the porosity and the tortuosity to correct for the pore structure. However, in situations that depart significantly from the ideal case of equimolar counter-diffusion, which are frequently found References see page 1781

1762

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

in real applications, Eqs. (174), (175) and (178) can only be used for approximations. This is because the effective diffusivity according to Fick’s law then becomes a function of the concentration and eventually also of the magnitude and the direction of the diffusional fluxes.

The overall flux density of pressure-driven flow through a porous catalyst is normally described by Darcy’s law, which holds for Newtonian fluids:

Pressure-Driven Viscous Flow and Surface Diffusion In addition to bulk and Knudsen diffusion, pressure-driven viscous flow and surface diffusion occur in porous catalysts and may contribute to the flux in the pores in certain situations. In the following it is of interest to what extent these effects can be treated by definition of a combined effective diffusivity according to Fick’s law in order to maintain a simplified treatment. It is common to assume, as a first approximation, that the various transport mechanisms occur in parallel, i.e. the fluxes assigned to the different phenomena are independent. This is often illustrated by the analogue of an electric circuit shown in Fig. 25. The overall flux density of component i then is

with η denoting the viscosity and ct the total molar concentration of the mixture and B0 is the permeability coefficient. For a single capillary, B0 has been given in Eq. (176). For a porous catalyst, the correction factor ε/τ [cf. Eq. (111)] must be included. Therefore,

6.3.7.1.2

Ni = NiD + NiC + NiS

(181)

where NiD denotes the diffusional flux density in mol m−2 s−1 (combined effect of bulk and Knudsen diffusion), NiC denotes the pressure-driven viscous flux density and NiS stands for the surface flux density. All fluxes are defined relative to the fixed coordinate system of the catalyst. Viscous flow is caused by a pressure gradient inward or outward of the catalyst. This can be generated in gas-phase reactions inside the catalyst if the reaction stoichiometry leads to a change in the overall number of moles. It can also be imposed by a pressure difference between two compartments if the catalyst is placed as a barrier in between, for example in the case of a catalytic contactor or a catalytic membrane.

Ni = Ni D + Ni C + Ni S Surface diffusion

Bulk diffusion

Ni D

dpi dz

Ni C

Ni S Convective viscous flow

Knudsen diffusion

Ni Fig. 25 Circuit diagram of the diffusional resistances in the transport of an ideal gas mixture at constant T according to the dusty gas model (adapted from Keil [29]).

NC = −

B0 =

B0 ct ∇p η

(182)

2 ε dpore τ 32

(183)

By multiplying Eq. (182) with the mole fraction, the molar flux density of component i attributed to pressuredriven viscous flow is obtained: NiC = −xi NC = −xi

B0 ct B0 ci ∇p = − ∇p η η

(184)

From Eq. (183), we may expect that viscous flow has a significant influence only in large pores; due to the quadratic dependence of B0 on the pore diameter the permeability coefficient becomes very small in pores only a few nanometers in diameter. Notwithstanding, there are examples of reactions where pressure-driven viscous flow should be taken into account. These include cracking and dehydrogenation of hydrocarbons (outward flow) and hydrogenations (inward flow). It has been shown that pressure changes inside the catalyst, depending on the catalyst activity and the reaction stoichiometry, may rise to the level of 40–50% and sometimes even cause concern about the mechanical stability of the catalyst [80]. Surface diffusion is important mainly in small mesopores and in micropores, because whenever the molecules of a diffusing species are adsorbed on the solid surface, then in pores of very small dimensions surface forces become dominant and the adsorbed molecules never completely escape from the force field of the pore wall even when at the center of the pore. The diffusion in this regime is an activated process. A general perception of adsorption is that on the surface distinct adsorption sites exist where the interaction potential between the adsorbent and the adsorbed molecules has a minimum. Consequently, between two sites a potential barrier exists. Depending on the adsorption strength (i.e. the adsorption energy), one can distinguish two types of movement. In physical adsorption, the molecules are weakly bound to the surface (the adsorption energy is rather low) and hence their average thermal energy is sufficient to overcome the potential barrier. What results is called mobile diffusion, characterized by a low activation energy of diffusion. Conversely, in chemisorption the molecules are more

6.3.7 Mathematical Description of the Diffusional Transport

strongly bound to the surface (the adsorption energy is higher); their average thermal energy is not sufficient to overcome the potential barrier and to move freely over the surface. The result is a hopping mechanism where the molecules jump from one adsorption site to another. The principles and the theory of chemisorption are addressed in detail in Chapter 5.1. In addition, Chapter 5.5.2 is devoted to the molecular simulation of adsorption. The molar flux density of component i due to surface diffusion can be described by Fick’s law, with the gradient of the surface concentration of this species as the driving force: NiS = −DiS,e SV ∇ciS

(185)

Equation (185) introduces an effective surface diffusivity DiS,e , which is defined analogously to Eq. (111) by DiS,e =

ε DiS τs

(186)

to account for the porosity and for the surface tortuosity. The quantity SV in Eq. (185) denotes the internal surface area per unit volume of the catalyst. This is required to convert the molar flux along a line on the surface in direction of diffusion to the molar flux density through a plane orthogonal to the direction of diffusion, as ciS denotes the concentration related to the surface (mol m−2 ) and not to the volume. Surface diffusion coefficients defined by Eq. (185) are available for a number of systems in the literature. The paper by Sladek et al. [81] may be cited here as an example; the authors have also presented a useful general correlation that allows to estimate the surface diffusivity DiS as a function of the adsorption enthalpy. In logarithmic form, this reads log DiS = 1.8 − 0.20

Hads,i mRT

(187)

where m is an empirical, integer-value parameter characteristic of the bonding type (cf. Table 6). Equation (187) is based on a variety of data from different systems covering a wide range of adsorption enthalpies (1.3–837 kJ mol−1 ) and temperatures (–230 to 600 ◦ C). It represents 11 orders of magnitude change in DiS to within 1.5 orders of magnitude over a range of Hads /mRT from 1 to 60. The combination of the different transport mechanisms according to Eq. (181) is demonstrated here for a binary ideal gas mixture at constant T in a flat plate geometry with the z-coordinate as the direction of transport. The molar flux density due to combined bulk and Knudsen diffusion is described according to Fick’s law [Eq. (171)] using the effective diffusivity from the Bosanquet equation [Eq. (174)] multiplied by ε/τ . The reason is that in the alternative Eq. (175) viscous friction forces are already considered by the term B0 /κi , which leads to an increased diffusional resistance. Through the viscous flow term the influence of the permeability coefficient B0 would be introduced a second time, which would be confusing. Substitution of the expressions for pressure-driven viscous flow [Eq. (184)] and surface diffusion [Eq. (185)] into Eq. (181) yields, for component 1, N1 = −D1,e

dc1 dc1S B0 c1 dp − − D1S,e SV dz η dz dz

References see page 1781

Classification of adsorption bonds to define the empirical parameter m in Eq. (187) (adapted from Sladek et al. [81])

van der Waals polar adsorbate van der Waals non-polar adsorbate Ionic Covalent

(188)

This equation can be recast into a form according to Fick’s law only if p and c1S can be expressed as unique functions of c1 . We may expect that adsorption and desorption are fast compared with diffusion and reaction, i.e. the adsorption–desorption equilibrium is reached, and further, we assume that the adsorption can be described by a single-site Langmuir model. For the

Tab. 6

Adsorption bond type

1763

Solid

m

Conductor Insulator Conductor Insulator Conductor Insulator Conductor Insulator

2 1 1 1 2 1 3 1

Examples of available surface diffusion data SO2 −carbon SO2 −glass, NH3 −glass Ar−W, N2 −carbon Kr−glass, C2 H4 −glass Cs−W, Ba−W None H−metals, O−W None

1764

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

fractional occupancy of component 1 we then have θ1 =

c1S K1 c1 = cSat 1 + K1 c1 + K2 c2

(189)

We restrict ourselves to low surface coverage (θi  1) where the denominator tends to unity and Eq. (189) simplifies to c1S = cSat K1 c1

(190)

This is the linear region of the Langmuir adsorption isotherm, where c1S is not influenced by the concentration of component 2. Equation (190) can be used to express the gradient of the surface concentration c1S in Eq. (188) in terms of the gradient of the gas phase concentration c1 . This allows one to combine the two diffusional flux terms, and Eq. (188) becomes   dc1 B0 c1 dp − N1 = − D1,e + D1S,e SV cSat K1 dz η dz ∗ = −D1,e

dc1 B0 c1 dp − dz η dz

(191)

The term in parentheses is constant and represents ∗ for combined bulk, Knudsen an effective diffusivity D1,e and surface diffusion. An analogous equation holds for component 2. For treatment of the pressure-driven viscous flow term, more information on the system must be given. As an instructive example, we discuss the case of a simple irreversible first-order reaction A1 → bA2 in a catalyst pellet. To express the pressure gradient in Eq. (191) in terms of dc1 /dx requires some effort. First, we make use of the continuity equation, which, in the steady state, forces the sum of the ingoing (component 1) and outgoing (component 2) mass fluxes to be zero: N1 M1 + N2 M2 = 0

(192)

For the chosen simple stoichiometry, the ratio of the molecular weights M1 /M2 is identical with the stoichiometric factor b. Substituting the molar flux densities from Eq. (191) into Eq. (192) and utilizing the relationship  dc1 dc2 dp = RT + dz dz dz after some rearrangement leads to the desired expression for dp/dz as a function of the concentration gradient dc1 /dz: ∗ − D∗ bD1,e dc1 dp 2,e =− ∗ dz B0 (bc1 + c2 ) D2,e dz + η RT

(193)

This can be substituted into Eq. (191), which finally takes the form of Fick’s law:
   ∗ D2,e c1 b − ∗   D1,e  dc1  ∗  N1 = −  D −  1,e η  bc1 + c2  dz  + ∗ D1,e B0 RT # (c) = −D1,e

dc1 dz

(194)

The term in square brackets in Eq. (194) turns out to # for combined transport be the effective diffusivity D1,e by bulk and Knudsen diffusion, surface diffusion and # obviously is pressure-driven viscous flow. However, D1,e not independent of concentration. Which of the various transport mechanisms have to be considered in a given situation can be judged best by analyzing the contributions of the individual terms to the effective diffusivity, i.e. in the above example, bulk diffusion vs. Knudsen diffusion in Eq. (174), surface diffusion vs. combined bulk and Knudsen diffusion in Eq. (191) and pressure-driven viscous flow vs. combined bulk, Knudsen and surface diffusion in Eq. (194). Less important contributions may be neglected to permit a simplified treatment. In cases where this is not permissible, the concentration # must be taken into account when dependence of D1,e solving the material balance. Its general form reads ∂ci = −∇Ni + Ri ; ∂t

i = 1, . . . , m

(195)

where Ri is the rate of formation or consumption of component i. In the stationary state the accumulation term vanishes and so the gradient of the molar flux density must equal Ri . The boundary conditions in the case of a catalyst pellet are the same as discussed earlier (cf. Table 1), i.e. symmetry of the concentration and pressure gradients at the pellet center and, for example, known concentrations at the external pellet surface. The stationary-state solution of Eq. (195) for a concentration-dependent effective diffusivity [cf. Eq. (194)] or different driving forces for the individual contributions to Ni analogous to Eq. (188) must be achieved by numerical integration. From the resulting flux of component i at the external pellet surface, the effective rate of production or consumption is obtained. Maxwell–Stefan Equations and Related Models Whenever the assumption of parallel fluxes assigned to the relevant transport mechanisms and effective diffusivities independent of the fluxes of the other components is expected to introduce an error exceeding the usual precision in determining these parameters, i.e. roughly 5–25% 6.3.7.2

6.3.7 Mathematical Description of the Diffusional Transport

depending on the particular situation (liquid phase or gas phase, complexity of the molecules/mixtures), a description of the diffusional transport closer to reality is necessary. The most versatile approach available to date goes back to Maxwell [82] and Stefan [83]. Their formulation of the diffusional flux is based on the kinetic theory of gases. Two important characteristics, as summarized by Krishna [77], will be emphasized here: 1. The force effecting the diffusional transport of species i is balanced by friction between the diffusing molecules i and j , where j = 1, . . . , m, i  = j denotes all other species in the mixture. As a consequence, the diffusional fluxes of the various species mutually influence each other. For example, if species i has a large diffusional flux compared with all other species, the friction exerted by these molecules on the molecules of a second component j will be large. Depending on the direction of movement of species j , this will then accelerate or delay the movement of the molecules of species j . 2. The driving force for diffusion of component i is the gradient of its chemical potential rather than the concentration gradient. This permits one to take into account explicitly the non-ideal behavior of mixtures, that is, it allows one to separate kinetic effects from thermodynamic effects, whereas the diffusivity according to Fick’s law is influenced by both. As a consequence, the diffusivities defined in the context of the Maxwell–Stefan equations are much less dependent on the mixture composition than the diffusivities defined by Fick’s law. These characteristics allow us to account for phenomena observed in multicomponent diffusion that could not be explained by Fick’s law, i.e. the occurrence of a diffusional flux when there is no concentration gradient (osmotic diffusion), the diffusion against a concentration gradient (reverse diffusion) and the absence of a diffusional flux in presence of a significant concentration gradient (diffusion barrier) [84]. For a more comprehensive explanation of the nature and advantages of the Maxwell–Stefan approach to multicomponent diffusion, the review by Krishna and Wesselingh [85] is recommended. Bulk Diffusion, Knudsen Diffusion and PressureDriven Viscous Flow For the case of a mixture of m chemical components at constant T and p, the Maxwell–Stefan equations read 6.3.7.2.1

xj NiD − xi Nj D xi ; j = 1, . . . , m ∇T ,p µi = RT ct − D ij m

−

j =1 j  =i

(196)

1765

where ∇T ,p µi is the gradient of the chemical potential of component i in direction of diffusion, ct is the total concentration, − Dij denotes the Maxwell–Stefan diffusion coefficient of species i and j , which has the meaning of an inverse drag coefficient for the interaction between the molecules of these two species, xi is the mole fraction of component i and NiD stands for the molar flux density of the diffusional flux of species i (relative to a fixed coordinate frame). Note that in this general formulation only m − 1 of Eq. (196) are independent because of the Gibbs–Duhem relation, which requires that the sum of all chemical potential changes, at constant T and p, adds up to zero. m

xi ∇T ,p µi = 0

(197)

i=1

Equation (197) provides the missing relationship to determine the flux densities of all m components. It is often useful to express the chemical potential gradients in Eq. (196) in terms of the mole fraction gradients. This is achieved by introducing the activity (µi = µ0i + RT ln ai ) and, further, the activity coefficients (ai = xi γi ). By differentiation and rearrangement of the resulting expressions, the left-hand side of Eq. (196) can be rewritten in the form [77] −

m−1

x1 ij ∇xj ; ∇T ,p µi = RT j =1 j  =i

ij = δij + xi

∂ ln γi ; ∂xj

i, j = 1, . . . , m − 1

(198)

where ij is an (m − 1) × (m − 1) matrix of thermodynamic factors portraying the non-ideal behavior and δij is the Kronecker delta, which is equal to 1 for i = j and 0 for i  = j . By substituting Eqs. (197) and (198) into Eq. (196), one finally obtains, in matrix notation, − c[][∇x] = [B][J ]

(199)

The elements of the (m − 1) × (m − 1) matrix [B] in Eq. (199) are related to the Maxwell–Stefan diffusivities − Dij as follows: Bii =

m

xi xk + ; − Dim − Dik k=1 k =i

 Bij,i=j = xi

1 1 − − Dim − Dij



(200) The (m − 1) column vector [J ] comprises the diffusional flux densities of the different species relative to a References see page 1781

1766

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

coordinate frame moving with the mean velocity of the fluid: Ji = NiD − xi Nt ;

i = 1, . . . , m − 1

(201)

where Nt denotes the net molar flux density which determines the mean fluid velocity: Nt =

m

NiD

(202)

i=1

Note that in Eq. (199), the Maxwell–Stefan equations are written in terms of [J ] rather than with the fluxes in the fixed coordinate frame, [ND ] [Eq. (196)]. By substituting NiD from Eq. (201) into Eq. (196), it can be easily shown that both forms are equivalent, as the two terms containing Nt cancel each other out and leave Eq. (196) in the same form, but with [J ] instead of [ND ]. For information on the prediction of the Maxwell–Stefan diffusivity for gaseous and liquid mixtures, Taylor and Krishna [86] and Wesselingh and Krishna [87] can be consulted. In the simplest case of an ideal gas mixture, the matrix of thermodynamic factors [] reduces to the identity matrix and hence Eq. (199) simplifies to −

m

xj NiD − xi Nj D 1 ; j = 1, . . . , m ∇p,T pi = RT − Dij j =1 j  =i

(203)

Equation (203) holds for bulk diffusion, that is, for absence of wall effects. However, combined bulk and Knudsen diffusion has also been treated using the Maxwell–Stefan approach. For this, the solid wall is interpreted as fictitious particles with zero diffusional flux. In addition to the friction between the species in the mixture caused by molecule–molecule collisions, the particles as an extra species m + 1 also exert a frictional force on the diffusing molecules. In this way, the interaction between the diffusing molecules and the pore wall during molecule–wall collisions is accounted for. The Knudsen diffusivity DiK [Eq. (173)] is adopted as the relevant parameter to quantify this interaction. The concept of solid particles (‘‘dust’’) has also donated the name ‘‘dusty gas model’’ to the resulting equations, which are shown here without derivation in Eq. (204) for the case of an ideal gas mixture. There, once more, effective diffusivities − Dij,e based on the Maxwell–Stefan diffusivities − Dij and the empirical correction factor ε/τ according to Eq. (111) have been introduced to take into account the geometric aspects of the pore

system: −

m

xj NiD − xi Nj D 1 NiD + ∇p,T pi = RT − Dij,e DiK,e j =1 j  =i

(204)

Note that also an effective Knudsen diffusivity appears in Eq. (204). More generic than by Eq. (173), this is defined as  8R T 4 (205) DiK,e = K0 3 πMi where K0 is a structure factor characteristic of the pore system. For cylindrical pores, K0 =

ε dpore τ 4

(206)

A detailed explanation of the dusty gas model including its derivation has been given by Mason and Malinauskas [88]. The effect of pressure-driven viscous flow on the mass transport has also been built into the dusty gas model [77, 85, 88]. The perception is, in accordance with Fig. 25, that the convective transport is superimposed on the diffusional transport. The approach again is illustrated here by example of an ideal gas mixture; constant T is assumed. To account for the influence of viscous flow on the transport of component i, Eq. (204) is expanded by an additional term giving its share of the viscous flux density multiplied by a friction factor. This factor is the inverse of the effective Knudsen diffusivity, which means that the friction forces between the wall and the diffusing molecules are described in the same way as in absence of a pressure gradient. With Eq. (184) for the viscous flux density of component i, this leads to −

m

xj NiD − xi Nj D NiD 1 + ∇T pi = RT − Dij,e DiK,e i=1 i =j

+

xi B0 p ∇T p RT ηDiK,e

(207)

Also, instead of Eq. (202), the total molar flux density becomes Nt =

m

NiD + NC

(208)

j =1

where the viscous flux density NC is given by Eq. (182). It is often more convenient to rewrite Eq. (207) in terms of the mole fraction gradient. With the relation ∇T pi = p∇T xi + xi ∇T p and after rearrangement, the

6.3.7 Mathematical Description of the Diffusional Transport

final form of the dusty gas model is obtained:  p xi B0 p ∇T xi − − + 1 ∇T p RT RT ηDiK,e =

m

xj NiD − xi Nj D NiD + − Dij,e DiK,e

(209)

i=1 i =j

Note that the determination of the molar flux densities of the m components in Eq. (209), as compared with Eq. (204), requires one additional relationship, because with the pressure the system has one more unknown variable. Moreover, it should be pointed out that, contrary to the basic Maxwell–Stefan equations [Eq. (196) or (203)], in the generalized formulations considering Knudsen diffusion [Eq. (204)] and pressure-driven viscous flow [Eq. (207) or (209)] it makes a difference whether the flux densities relative to the fixed coordinate frame, NiD , or those relative to the mean fluid velocity, Ji , are used. This is due to the Knudsen diffusion term, where, after substitution of NiD through Ji from Eq. (201), the additive element xi Nt remains. The dusty gas model has been validated for the description of mass transport in porous catalysts and membranes in numerous publications. In addition to the main reference by Mason and Malinauskas [88], further examples can be found in, amongst others, the reviews by Krishna [77] and Krishna and Wesselingh [85], and, for the field of porous catalytic membranes (cf. Chapter 10.7), in Sloot et al. [89] and Tuchlenski et al. [90]. The extension of the dusty gas model to non-ideal fluids has brought forth the so-called ‘‘dusty fluid model’’. Through a generalization of the driving force term, external body forces such as an electrostatic potential gradient or a centrifugal force field can be included in the formulation of the Maxwell–Stefan equations. By properly considering the thermodynamic non-idealities of the respective systems, such models have been successfully applied, for example, to describe the transport of ions in liquid electrolytes (for more information, see Ref. [85]). Other fields of application concern heterogeneous reactive distillation [91] and proton exchange membrane fuel cells [92, 93]. Despite the success of the dusty gas model, there is an ongoing debate in the literature about conceptual errors in its basic derivation. Kerkhof [78, 94, 95] and others (e.g. [96]) have demonstrated that the picture of a parallel viscous flux added to the diffusional flux in the dusty gas model is conceptually wrong because it is the friction forces between the species and the wall that should be added, just as the friction forces between the components are added in the Maxwell–Stefan equations.

1767

Kerkhof showed [94] that the appearance of the separate viscous flux term in the dusty gas model is due to derivational errors and leads to double counting of the viscous terms. In addition, the description of the friction forces between the diffusing components and the pore wall by the term NiD /DiK,e [cf. Eq. (209)], which contains the diffusional flux and the effective Knudsen diffusivity of component i, is inconsistent in that it lacks the viscous nature of the fluid–wall interaction. Fortunately, both effects partly compensate each other, but at low Knudsen number incorrect results are expected. The alternative ‘‘binary friction model’’ developed by Kerkhof [78, 94, 95] on the foundation of the Lightfoot friction model [97] relates the pore-averaged fluxes to driving forces by introducing friction coefficients associated with individual viscosity contributions. Friction forces between the diffusing molecules and friction forces between the molecules and the wall are included. The model is able to account for the mass transport over the whole range of bulk and Knudsen diffusion including viscous and diffusive slip effects. For a mixture of gases at constant T the equations read −

p xi ∇T ,p xi − ∇T p RT RT m

ij (xj Ni − xi Nj ) = + fim Ni − Dij,e

(210)

j =1

where fim

 B0 = DiK,e + κi

(211)

is a friction factor between the diffusing components and the pore wall which takes into account both the viscous wall-friction (B0 /κi ) and the interaction due to molecule–wall collisions (DiK,e ). Equation (210) can be directly compared with Eq. (209). The coefficients κi representing the fractional viscous contributions to the overall friction between component i and the wall were already given in Eq. (179). The factor ij has been formally introduced to allow for disappearance of the interspecies friction at the zero-pressure limit. However, for many practical cases good agreement between theory and experiment was obtained with ij = 1. Thus, Eq. (210) can be simplified by neglecting ij [78]. The binary friction model has also been adopted for liquid-phase systems; one of the applications described in Kerkhof’s first paper [94] was ultrafiltration. In a recent study, it was selected as a basis for the development of a transport model for polymer electrolyte membranes [96]. Due to the consistent derivation (as opposed to the dusty gas model), it is expected that the binary friction model References see page 1781

1768

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

will increasingly replace the dusty gas model in the future. In respect thereof, attention is called also to a recent book by Struchtrup [98] on the theoretical background and the mathematical methods to develop macroscopic transport equations for gas flow in confined structures from kinetic theory. A third model based on the Maxwell–Stefan equations is the ‘‘mean pore transport model’’ proposed by Schneider and co-workers [99, 100]. This also takes into account bulk diffusion, Knudsen diffusion and pressure-driven viscous flow. It assumes that the decisive part of the flux is through cylindrical transport pores with radii distributed around the mean value rp . An effective porosity is assigned to these pores through the correction factor εp /τ [cf. Eq. (111)] for the Maxwell–Stefan diffusivities, just as in the dusty gas model. The width of the pore radius distribution is characterized by the mean value of the squared radii rp2 . For the diffusional flux densities the generalized Maxwell–Stefan equations are used without modification, whereas the description of the viscous flux density differs in that it includes contributions of Knudsen diffusion and viscous slip on the pore wall. This leads to component-specific permeability coefficients Bi . For an ideal gas mixture at constant T , the mean pore transport model takes the following form [101], which can be directly compared with Eq. (209) (dusty gas model) and Eq. (210) (binary friction model):   −

m  B

xj (Bi − Bj )  p  i  xi + ∇T xi −  ∇T p   DiK,e  RT RT − Dij,e j =1 j  =i

=

m

xj N i − xi N j Ni + − Dij,e DiK,e

(212)

j =1 j  =i

The term in square brackets in Eq. (212) describes the viscous flux contribution. The permeability coefficients Bi have been calculated slightly differently in various ˇ ˇ publications. Capek et al. [102] and Solcov´ a et al. [103] used the following definition, which is a simplified version of the equation originally proposed by Schneider [104]: Bi = DiK,e

ω + Kni p + B0 1 + Kni η

(213)

where DiK,e is the effective Knudsen diffusivity, defined according to Eqs. (205) and (206) with dpore = 2rp . B0 denotes the permeability coefficient obtained from Eq. (183) 2 = 4rp2 and ω is a factor characterizing the slip for dpore at the pore wall (typical values of 0.6–1.4 [102]). Kni stands for the Knudsen number [Eq. (172)] of component i and p and η denote the total pressure and the

ˇ viscosity of the mixture, respectively. Later Capek and Seidel-Morgenstern [101] dropped the term involving the slip parameter ω. In that study, the mean transport pore model, the dusty gas model and the binary friction model were compared for their ability to reproduce transient measurements of the pressure change in a modified Wicke–Kallenbach cell during counter-diffusion of binary and ternary gas mixtures. The conclusion was that all three models were able to reproduce the experimental results with acceptable accuracy for engineering purposes. However, the mean transport pore model gave the best fit of the experimental data, the binary friction model showed the largest deviation and the dusty gas model ranged between the other two. The advantage of the mean pore transport model was more pronounced in ternary mixtures than for binary systems and it showed a dependence on the average pore diameter of the catalyst studied: for a mesoporous catalyst used for methanol synthesis the mean transport pore model was clearly superior whereas for a macroporous catalyst for butane oxidation it gave almost identical results as the dusty gas model. Other applications of the mean transport pore model have been described, for example, by Ackmann et al. [105], who studied the heat and mass transport in a planar substrate solid oxide fuel cell, and by Divisek et al. [106], who implemented it into a two-phase flow model of a direct methanol fuel cell. 6.3.7.2.2 Surface Diffusion Multicomponent surface diffusion can also be described adopting the Maxwell–Stefan equations. This concept has been developed mainly by Krishna [77, 107, 108] on the foundation of the dusty gas model. It keeps the idea of parallel transport mechanisms in the pore and on the surface, which do not directly interfere with each other. To illustrate the concept, Krishna used the picture of ‘‘crated dust golf balls’’ interacting with the diffusing molecules through friction forces. The golf balls stand for the dust particles representing the porous medium in the dusty gas model and the craters are linked to the adsorption sites [77]. The mass transport in the free pore volume occurs according to the mechanisms constituting the dusty gas model. On the surface, it is assumed that the diffusional transport takes place by adsorbed molecules hopping from one adsorption site to another. There, the vacant adsorption sites are considered to be the (m + 1)th pseudospecies of the surface system, analogous to the dust particles in the bulk system. Hence the diffusional velocity of the vacant sites is assumed to be zero. The driving force acting on species i to move it along the surface is the gradient of the surface chemical potential. With the correction factor ε/τs analogous to Eq. (186), this leads to the following formulation for the molar surface flux

6.3.7 Mathematical Description of the Diffusional Transport

Depending on the isotherm employed to describe the adsorption equilibrium, different expressions apply to calculate the elements of []. A typical model is that of a Langmuir isotherm:

density of component i: −SV csat

m

θj NiS − θi Nj S θi ∇µi = RT − Dij S,e j =1 j  =i

+

NiS ; − DiS,e

θi = i = 1, . . . , m

(214)

θi ij ∇θj ; ∇µi = RT m

j =1

θi ∂pi ; pi ∂θj

i, j = 1, . . . , m

ciS = cSat

bi p i ; m

1+ bj p j

bi p i =

θi θVac

(216)

j =1

Equation (214) contains two types of diffusivities: − DiS,e denotes the effective Maxwell–Stefan surface diffusivity of component i which, apart from the geometric correction factor ε/τS , is equivalent to the thermodynamically corrected diffusivity as defined by Ruthven [109]. In terms of the assumed hopping mechanism, this is related to the displacement, to the number of nearest neighbor sites and to the jump frequency of the adsorbed molecular species. In general, the last factor is expected to be dependent on the surface coverage, which complicates the treatment. The second type of coefficient, − Dij S,e , accounts for the interaction between adsorbate i and adsorbate j . It is interpreted as a measure of the facility of replacement of an adsorbed species j at an adsorption site by species i and is therefore expected to be dependent on the jump frequencies of these two species (i.e. also on the surface coverage). The net effect of this counter-exchange is a slowing of the faster moving species due to interactions with a species of lower mobility and vice versa. The Onsager reciprocity Dj iS [107]. The quantities relation holds, i.e. − Dij S = − NiS , θi , SV and csat have already been introduced in Eqs. (188)–(190) in Section 6.3.7.1. Note that Eq. (214) holds for the case that a common saturation loading, given by the saturation concentration cSat , applies to all species. This implies the assumption of just one type of adsorption site, the concentration of which on the surface determines the adsorption capacity. It also implies that the different species show a similar adsorption behavior, for example the formation of a monolayer. Following the derivation of Eq. (199), the surface chemical potential gradient in Eq. (214) can be related to the gradient of the fractional occupancy by introducing a matrix of thermodynamic factors []. For a gaseous phase in equilibrium with the surrounding fluid, the surface chemical potential of component i is given by µi = µ0i + RT ln fi . At low to moderate pressure, the component fugacity fi can be replaced by the partial pressure pi . Differentiation of µi and rearrangement of the resulting equations yields

ij =

1769

(215)

for which ij = δij +

θi

(217)

θVac

where θVac denotes the fraction of vacant adsorption sites, bi is the adsorption equilibrium constant of component i in Pa−1 , ciS denotes its concentration on the surface in equilibrium with the surrounding fluid (mol m−2 ) and cSat the surface concentration at saturation. By substituting Eq. (215) into Eq. (214) we have, in matrix notation, −SV cSat [][∇θ ] = [BS ][NS ]

(218)

where the elements of matrix [Bs ] are related to the two types of Maxwell–Stefan surface diffusivities as follows: BiiS =

m

θj 1 + ; − DiS,e − Dij S,e j =1 j  =i

Bij S = −

θi ; − Dij S,e

i, j = 1, . . . , m

(219)

Of these two, the component Maxwell–Stefan surface diffusivity − DiS,e can be determined from pure-component experiments (Wicke–Kallenbach cell or similar methods). The counter-exchange Maxwell–Stefan surface diffusivity − Dij S,e , however, is difficult to obtain experimentally. To estimate − Dij S as a function of − DiS and − Dj S , Krishna [107] proposed a generalized form of an empirical correlation originally developed for diffusion in bulk liquid mixtures [110]: θi

θj

DiS ] θi +θj [− Dj S ] θi +θj ; − D ij S = [−

i, j = 1, . . . , m (220)

This equation was adopted, for example, by van den Graaf et al. [111] to describe the transport of binary mixtures of light alkanes through a silicalite-1 membrane. Their results clearly showed that adsorbate–adsorbate interactions must be included in a proper description of these systems. Neglecting the counter-exchange Maxwell–Stefan diffusivities − Dij S,e caused even qualitative failure in predicting the observed behavior, which References see page 1781

1770

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

was characterized by a pronounced delay of the faster diffusing species (weakly adsorbed) by the slower moving component (more strongly adsorbed) while leaving the flux of the slower component nearly unchanged. The Maxwell–Stefan model for surface diffusion, as represented by Eqs. (214)–(220), has been further developed and validated mainly for the diffusion in zeolites in a significant number of studies over the last 5 years. One important aspect concerns an extension to different saturation loadings assigned to various species. This is observed, for example, in certain zeolites where distinctly different adsorption sites exist at which molecules of different size or chemical nature are preferentially adsorbed [112]. If this applies, the saturation loading of such species may differ greatly. The fractional occupancy is then defined by relating the actual loading to the saturation loading of the respective species: θi =

ciS ci,Sat

(221)

This changes Eq. (214) to the form −SV ci,Sat

m

cj S NiS − ciS Nj S θi ∇µi = RT cj,Sat −Dij S,e j =1 j  =i

+

NiS ; − DiS,e

i = 1, . . . , m

(222)

where ciS and cj S denote the surface concentrations of the components i and j in mol m−2 and ci,Sat and cj,Sat represent the concentrations at saturation. The equations are given here in terms of the surface concentrations ciS and the specific internal surface area SV for notational consistency, although in most of the publications cited the loading per kilogram, qi (mol kg−1 ), and the density, ρ (kg m−3 ), or the loading per unit cell (mol) and the number of unit cells per volume (m−3 ) are used. The latter definition is useful only for zeolites or other materials with regular pore structure. The definition of the matrix of thermodynamic factors given by Eq. (215) applies unchanged, but not the calculation of its elements, which becomes  cj,Sat ciS ∂pi ij = (223) ci,Sat pi ∂cj S Equation (223) includes the ratio of the saturation loadings of the species j and i, which in general differs from unity for i  = j , i.e. for the off-diagonal terms. With these changes, Eq. (218) transforms to − SV [][∇θ ] = [BS ][csat ]−1 [NS ]

(224)

where [csat ] is a diagonal matrix of saturation loadings and [Bs ] is still given by Eq. (219).

This extension improved the predictive capabilities of the Maxwell–Stefan surface diffusion model as it allows more complex adsorption isotherms to be applied, e.g. determined based on the theory of ideal adsorbed solutions (IAS) [113] from the individual component adsorption isotherms (cf. Kapteijn et al. [112]). For more information on the theory of mixture adsorption, in particular on refined methods such as non-ideal adsorbed solutions (NIAS), Myers [114] and Siperstein and Myers [115] and the references therein may be consulted. Another important aspect concerns the development of a generalized strategy for estimating the Maxwell–Stefan counter-exchange surface diffusivities − Dij S from pure-component parameters and its validation by molecular dynamics simulations and configurational bias Monte Carlo simulations for different zeolite structures and various gases. Skoulidas et al. [116], Chempath et al. [117] and Krishna and co-workers [118, 119] and the references therein may be consulted for derivational details and application examples. This approach relates − Dij S no longer to the two purecomponent Maxwell–Stefan surface diffusivities − DiS and − Dj S , as in Eq. (220), but to the pure-component Maxwell–Stefan self-exchange surface diffusivities − DiiS and − Djj S . These naturally arise when Eq. (214) is applied to tracer diffusion, that is, j becomes i ∗ , a tagged variant of species i with unchanged properties. Following Ref. [120], from Eqs. (214)–(217) it is easily derived by letting Ni + Ni∗ = 0 and ∇θi + ∇θi∗ = 0 and by considering − DiS = − Di∗S , that the tracer (self) diffusivity is Dself ,iS =

1 θi 1 + − DiS − DiiS

(225)

Rearrangement for − DiiS gives − D iiS =

θi 1 Dself ,iS

−

1 − DiS

(226)

which allows − DiiS to be related to Dself ,iS and − DiS , which both in general are dependent on the occupancy. Hence − DiiS also depends on the occupancy. The Maxwell–Stefan counter-exchange surface diffusivity − Dij S is now estimated from − DiiS and − Djj S according to the interpolation equation [119]   θi   θj cj,Sat − D ij S = cj,Sat − D iiS θi +θj ci,Sat − D jj S θi +θj i, j = 1, . . . , m

(227)

which is an extension of Eq. (220). Note that, in line with the Onsager reciprocity condition, cj,Sat − D ij S = ci,Sat − D j iS [118].

6.3.8 Control of Selectivity in the Disproportionation of Ethylbenzene to Benzene and Diethylbenzenes

Equation (225) has been generalized for mixtures with different saturation loadings of the components [cf. Eqs. (222) and (223)] [121]. The generalized form reads Dself ,iS =

1 m

θj θi 1 + + − DiS − DiiS − Dij S

(228)

j =1 j  =i

By substituting Eq. (219) into Eq. (228), the selfexchange coefficient DiiS in a mixture of m components then becomes [119] − D iiS =

θi 1 Dself ,iS

(229) − Bii

DiS ]−1 in the case of a pure which, due to Bii = [− component, simplifies to Eq. (226). Krishna and van Baten [119] showed that the self-exchange coefficient − DiiS at a given occupancy is the same whether determined from pure-component, binary or ternary mixture data. This allows us to estimate all necessary diffusivities from pure-component data. What remains crucial is a proper description of the occupancy dependence of the pure-component Maxwell–Stefan surface diffusivity − DiS and the self diffusivity Dself ,iS . Different models and correlations have been proposed based on experimental data and computer simulations. More information can be found, for example, in the papers by Krishna and co-workers [118, 119]. In the treatment of the Maxwell–Stefan equations for surface diffusion, here it was so far assumed that surface diffusion is the only relevant transport mechanism. Formally, based on Fig. 25, a second set of Maxwell–Stefan equations (e.g. the dusty gas model) could be added to describe the parallel mass transport in the free pore space. However, in large pores surface diffusion usually does not contribute greatly to the overall fluxes, just because the effective surface diffusivity is much smaller than the effective diffusivity due to bulk or Knudsen diffusion. Conversely, in small micropores, as a matter of the confinement, there is no free pore volume because the molecules never escape the influence of the pore wall. In large micropores or small mesopores there could be contributions by both mechanisms, but due to the possibility of desorption strong coupling of the fluxes of the various species as observed in small micropores is normally not expected (unless capillary condensation occurs). This allows a simplified treatment of surface diffusion. Two examples have already been mentioned earlier [89, 90]. Industrial solid catalysts often consist of a microporous, amorphous or crystalline, powdery active material formed into larger particles or pellets or applied as a washcoat on a

1771

support, e.g. a monolith or a structured packing, together with a binder. These are systems with a bimodal pore size distribution, i.e. macropores or mesopores between the active particles and micropores inside. To account for the influence of mass transport on the reaction rate in such catalysts, concentration gradients in both types of pores may have to be considered at sufficiently high reaction rates. This can be handled best by a hierarchical modeling approach. For mass transport through the macropores, one of the three variants, i.e. dusty gas model, binary friction model or mean transport pore model, is recommended, while Fick’s law with constant effective diffusivity should be used only for highly diluted reaction systems or for approximation purposes. To account for the diffusional transport in the micropores, the Maxwell–Stefan equations for surface diffusion represent a good choice if enough information on the component Maxwell–Stefan surface diffusivities and the self-diffusivities as a function of the occupancy is available. If not, or if the system lacks strong adsorbate–adsorbate interactions, a simpler version ignoring the counterexchange terms and also the occupancy dependence of the diffusivities might be sufficient. In setting up such a model, the geometric situation, i.e. size and shape of the catalyst and the crystallites, fraction of active material, porosity and connectivity of the pores, etc., also have to be considered. The simplest approach would be to connect the diffusion in the micropores to the modeling of the transport in the macropores by adding a second dimension to the problem, but more complicated three-dimensional pore-network models are also conceivable. 6.3.8

Control of Selectivity in the Disproportionation of Ethylbenzene to Benzene and Diethylbenzenes on an HZSM-5 Catalyst by Utilizing Diffusion Effects

The discussion of the Maxwell–Stefan equations for surface diffusion in micropores in Section 6.3.7.2 has already highlighted that the mass transport in this domain may follow very complex relationships effectuated by the interaction of the diffusing molecules with the pore wall and by adsorbate–adsorbate interactions under extremely confined conditions. This diffusional regime is known as ‘‘configurational diffusion’’ after Weisz [122]. It has been, and still is, a very active and important research field in catalysis and adsorption science and technology [123–125]. Chapter 6.2 covers this topic in detail. On balance, in this regime, 1. The diffusivity is much lower than in the bulk and Knudsen diffusion regimes (cf. Fig. 26). References see page 1781

1772

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

Bulk 1

1 bar 10 bar

10−2

D / cm2s−1

10−4

CH3OH + Gases

Knudsen Liquids

10−6 10−8 Configurational

10−10 10−12 10−14

H2O + Reactant shape selectivity

0.1

1

10

100 1000 10 000 r / nm

Fig. 26 Diffusivity versus pore size. Classical regions of bulk and Knudsen diffusion and regime of configurational diffusion (adapted from Weisz [122]).

2. The diffusivity of species with slightly different kinetic diameter may vary by several orders of magnitude. 3. In mixtures, faster diffusing species may be significantly delayed by more slowly moving components and vice versa. 4. The diffusivity of certain species might be dependent on the direction (anisotropic diffusion). Shape-Selective Catalysis Configurational diffusion effects can be utilized on purpose to influence the activity and selectivity in the catalysis of multiple reactions, for example when classical attempts at tailoring catalyst properties come up against severe difficulties. This is possible by using microporous catalysts with regular, highly ordered pore systems, such as the crystalline framework of zeolites or other comparable structures of related materials (aluminophosphates, pillared clays, etc.). They can be modified concerning the chemical properties and also with respect to the pore geometry by different techniques during and after synthesis over a wide range as to match the requirements of the desired reaction. This field, fascinating to catalysis scientists and reaction engineers, is identified with the term ‘‘shape-selective catalysis’’. It offers a multitude of new opportunities to develop materials with tailor-made catalytic properties. Shape-selectivity effects may occur whenever the pore size of a microporous catalyst is in the same range as the diameter of the molecules or transition states involved in the reacting system. According to Weisz et al. [126] and Csicsery [127], shape-selectivity effects may be classified into the following three types (Fig. 27). 6.3.8.1

+

Restricted transition state shape selectivity

Product shape selectivity

Fig. 27 Classification of shape-selective effects according to Weisz et al. [126] and Csicsery [127].

• Reactant shape selectivity This is found when only the tallest of several reactants is able to diffuse into the pore structure of the catalyst and is reacted. • Product shape selectivity This is given when several products are formed from the reactants inside the crystallite, but those with a rather bulky structure cannot diffuse out. These molecules may be reacted further on the active surface to give either the desired products (i.e. by isomerization) or other, smaller molecules (i.e. by cracking) which then can readily leave the pore system. In the worst case, the bulky molecules will be trapped inside the cavities or channels, block the pores and prevent other reactant molecules from entering (deactivation by coking). • Restricted transition state shape selectivity This means that in a network of multiple reactions those products are suppressed which require the intermediate formation of a bulky transition state, whenever the size of the cavities or channels is too narrow for this to be established. Basically, reactant and product shape selectivities are mass transfer effects, where the diffusivities of the various species in practice frequently do not differ that extremely as indicated above. Instead, in most cases only a preferred diffusion of certain species is observed, a fact which often hinders a clear understanding of product shape selectivity. This is because the various products, during their way through the pore system, may be reacted when contacting the catalytically active surface of the wall. This combined effect of diffusion and reaction will be discussed in detail in the following, as it is of great importance for the product distribution in zeolitecatalyzed reactions.

6.3.8 Control of Selectivity in the Disproportionation of Ethylbenzene to Benzene and Diethylbenzenes

The disproportionation of ethylbenzene (EB) to benzene (B) and diethylbenzenes (DEB) serves as a case study. This reaction was introduced by Karge et al. and Weitkamp et al. to compare the activity of different zeolite catalysts [128, 129] and to distinguish between medium- and large-pore zeolites [130, 131]. Moreover, it has been further developed as a standardized test reaction by the International Zeolite Association (IZA) [132], see Chapter 2.6.2 of this Handbook. A general representation of the individual transformations which occur in this type of reaction, is given in the following scheme: kF,1

2EB

− −   −− − − o − DEB + B

2EB

−−   −− − − m − DEB + B

2EB

−−   −− − − p − DEB + B

o − DEB

− −  −−   −− − − m − DEB  −− − − p − DEB

kF,2 kF,3 kt,1

kt,2

where the prefixes o−, m− and p− indicate the ortho-, meta- and para-isomers, respectively. It may be obvious to note that the effects to be discussed here will basically apply to similar reacting systems, such as the alkylation of other monoalkylbenzenes using various alkenes or alcohols or the disproportionation of toluene. Olson and Haag [133] showed that the yield of p-xylene observed during the disproportionation of toluene on various modified and unmodified ZSM-5 catalysts is actually influenced by product shape selectivity. They attributed the observed effects to an interaction of diffusion and reaction, characterized by means of a dimensionless modulus similar to the classical Thiele modulus φ. The mathematical treatment of shape selectivity in zeolite catalysts, which will be applied in this section, is largely based on the theory of Olson and Haag [133], although some modifications and extensions to this are given. 6.3.8.1.1 Microkinetics As an essential requirement for the fundamental effects of shape selectivity to be developed in an instructive manner, we first simplify the above general reaction scheme. In this sense, we restrict our considerations to a sufficiently small reactant conversion, where we may replace the separate formation reactions of the three isomers with a single pseudoirreversible reaction, described by a common rate constant kF . This is justified by the fact that all three reactions obey the same kinetic order and the backward reactions, due to the small isomer concentration at low level of conversion, may be neglected. Moreover, according to Wei [134] we treat the isomerization, which actually involves two reversible reactions, also as a single pseudoirreversible reaction. This may be interpreted in a sense

1773

that the reacting system globally approaches the chemical equilibrium with a rate characterized by a common rate constant kI . Thus, we have two consecutive reactions: disproportionation:
o−  kf

2EB −−−→

m− p−

DEB + B

(230)

pri

followed by isomerization:
o− 
o−  m− p−

kI

DEB −−−→ pri

m− p−

DEB

(231)

sec

where the subscripts pri and sec denote the primary and secondary isomer distribution, respectively. The primary isomer distribution, which is the result of the disproportionation reaction, may deviate significantly from the thermodynamic equilibrium composition, for two reasons. First, the reaction may be controlled by the kinetics rather than the thermodynamics, i.e. mechanistic reasons may exist which cause the reaction to proceed along a certain path. Second, in the case that the reaction obeys a bimolecular mechanism, it may pass through a transition state which would presumably favor the (taller) para-isomer. Hence it is possible that the primary product contains an enhanced fraction of the para-isomer. The departure from the equilibrium composition then gives the driving force for the subsequent (monomolecular) isomerization reaction. This will reduce the fraction of the para-isomer, provided that the formation of the bulky ortho- and meta-isomers is not inhibited by steric effects, i.e. when the micropore diameter is sufficiently large or there is a chance for the isomerization reaction to take place at the outer surface of the crystallites. Thus, the secondary isomer distribution may approach the thermodynamic equilibrium composition, as a limiting case. 6.3.8.1.2 Macrokinetics As discussed in Section 6.3.6, the apparent selectivity of the catalyst in a network of multiple reactions may be affected by interphase and intraparticle heat and mass transfer. In zeolites or other microporous materials, the intraparticle diffusion process normally is governed by two different mechanisms. To illustrate this, the zeolite catalyst pellet may be visualized as consisting of a large number of individual crystallites, typically around 1–100 µm in diameter. The void space between these crystallites constitutes the macropores of the pellet, whereas the micropores are given by the cavities or channels inside the zeolite crystals. Hence these materials show a bimodal pore size distribution. The mass References see page 1781

1774

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

transport inside the macropores is governed by bulk (and/or Knudsen) diffusion, whereas the mass transport inside the micropores is controlled by configurational diffusion. However, as we are concerned here primarily with the investigation of product shape selectivity, which is basically something that happens inside the zeolite crystals, for the subsequent discussion we may neglect any concentration and temperature gradients along the macropores of the pellet and also across the external boundary layer surrounding the pellet. Moreover, to keep the discussion as simple as possible, temperature gradients inside the crystallites will also be ignored. Hence we have to consider the following steps, which are similar to steps 1–7 described in the Introduction, for the case of a macroporous pellet: 1. adsorption of the reactant on the outer surface of the crystallite 2. disproportionation of the reactant at the catalytically active sites on the outer surface 3. diffusion of the reactant into the microporous structure 4. disproportionation of the reactant and isomerization of the products at the catalytically active sites inside the pore system 5. diffusion of the products out of the microporous structure 6. isomerization of the products at the catalytically active sites on the outer surface 7. desorption of the products from the outer surface. At this point, it should be emphasized that steps 2–6 are not linked in a sequential manner. Instead, diffusion and reaction inside the pore system, steps 3–5, occur simultaneously, as also observed in the macroporous pellet. Moreover, the opportunities for a reactant molecule either to be converted at the outer surface of the crystallite or to enter the porous structure also represent parallel reaction paths. Although both the microporous (crystallite) and the macroporous (pellet) problem share instructive similarities, they nonetheless differ in some aspects. These differences have important implications with respect to the control of selectivity. These include the fact that the boundary conditions at the outer surface of the crystallite are controlled by the laws of adsorption, whereas the boundary conditions at the external surface of the pellet are governed by interphase diffusion. Hence a different adsorption behavior of the various species may have a significant effect on selectivity. In contrast, the interphase mass transfer coefficients of the various species are not likely to differ substantially and hence in the case of an isothermal, macroporous pellet interphase diffusion is not expected to have a marked influence on selectivity.

Moreover, in the case of a crystallite, there will be a difference in selectivity depending on whether the reaction takes place at the outer surface or inside the pore system. This is because in the former case the reaction will proceed unaffected by geometrical restrictions, whereas in the latter shape-selective effects may control the selectivity. It is obvious that such a difference will not be observed in the macroporous pellet. Modeling of Shape Selectivity Effects In the light of the previous discussion, it is apparent that a detailed mathematical simulation of the combined chemical reaction and transport processes, which occur in microporous catalysts, would be highly desirable to support the exploration of the crucial parameters determining conversion and selectivity. Moreover, from the treatment of the basic types of catalyst selectivity in multiple reactions, given in Section 6.3.6, it may be clear also that an analytical solution to this problem, if at all possible, will presumably not favor a convenient and efficient treatment of real-world problems. This is because of the various assumptions and restrictions which usually have to be introduced in order to meet a complete or even an approximate solution. Hence numerical methods are required. Concerning these, one basically has to distinguish between three fundamentally different types, namely molecular dynamic models, stochastic models and continuous models. Molecular dynamic simulations are characterized by a solution of Newton’s laws of motion for the molecules traveling through the zeolite pore system under control of the force field given by the properties of the host frameworks, by interactions between the host and the molecules and by interactions among the molecules. To date this has been possible only for the diffusion of simple molecules (e.g. methane or benzene) inside a zeolite framework of limited dimensions [135–137]. To take into account also the effects of a chemical reaction would require quantum-mechanical considerations; however, such simulations are in their infancy. In stochastic models, the regular zeolite framework is projected onto an adequate rectangular grid, where each grid element corresponds to the adsorption sites located either inside a certain cavity or crossing of channels, depending on the type of the zeolite. The siteto-site movement (jumping) of the adsorbed molecules is described in these models by the probability of jumping. In Monte Carlo methods, which may be considered as a special class of stochastic models, a random generator serves to initialize the activated adsorption sites and the directions of movement of the activated molecules. By additionally introducing the probability of reaction, chemical transformations may be included in such 6.3.8.2

6.3.8 Control of Selectivity in the Disproportionation of Ethylbenzene to Benzene and Diethylbenzenes

simulations. Theodorou and Wei [138] were the first to report about the application of a Monte Carlo method to the modeling of combined diffusion and reaction in zeolite catalysts. Since then, several related papers on this subject have appeared in the literature, for example the work of Tsikoyiannis and Wei [139], Frank et al. [140] and Wang et al. [141]. However, most of these simulations use rather simple model reactions, for example systems of the type A   C, or A → B → C. B  B, A  The principles and basic equations of continuous models have been introduced already in Section 6.3.2. These are based on the well-known conservation laws for mass and energy. The diffusion inside the pores is usually described in these models by the Fickian laws or by the theory of multicomponent diffusion (cf. Section 6.3.7.2). However, these approaches basically apply to the mass transport inside the macropores, where the necessary assumption of a continuous fluid phase essentially holds. In contrast, in the microporous case, where the pore size is close to the range of molecular dimensions, only a few molecules will be present within the cross-section of a pore, a fact which poses some doubt as to whether the assumption of a continuous phase will be valid. Wei [134] again was the first to come up with a continuous pseudo-homogeneous model which allowed the simulation of shape-selective effects observed during the alkylation of toluene using methanol to yield xylene isomers on an HZSM-5 catalyst. He treated diffusion and reaction of the xylene isomers inside the pores in a one-dimensional model. The isomer concentration at the pore mouth was set to zero, as a boundary condition. This allowed the model equations to be solved analytically, but it also limited the application of the results to small conversions. Do [142] used a method based on Wei’s approach [134] to simulate again the alkylation of toluene on an HZSM-5 catalyst but, in contrast to Wei, for a fixed-bed reactor where it was assumed that concentration gradients exist only inside the crystallites and along the axial coordinate of the catalyst bed. Liang et al. [143] later extended the original model of Wei [134] by taking into account the adsorption at the outer surface of the crystallite. They assumed that the adsorption takes place at the pore mouth and that the reaction and the diffusion into the microporous structure proceed in an adsorbed state. A further improvement of Wei’s approach [134] was reported by Hashimoto et al. [144], who considered not only adsorption effects, but also the non-selective reactions occurring at the outer surface of the crystallites. The non-selective influence of these reactions was also recognized by Fraenkel [145], who studied the formation of xylene from toluene on an HZSM-5 catalyst. Fraenkel assumed that inside the crystallite only p-xylene is

1775

formed, whereas the ortho- and meta-isomers are sterically inhibited there. Hence he concluded that the amount of o− and m-xylene observed during his experiments must be due to the isomerization of p-xylene at the outer surface of the crystallites. This two-step mechanism was first suggested by Paparetto et al. [146] for the ethylation of toluene. It may also be worth noting that Fraenkel’s model took into account not only the non-selective isomerization but also the non-selective alkylation at the outer crystallite surface. Finally, Emig et al. [147] presented a continuous model, which uses a versatile numerical solver and thus allows the simulation of a variety of situations occurring when different steps of the overall process are significant (see the steps 1–7 above). In the following, the mathematics and the use of this model are explained. Simplifying Assumptions 1. The model equations are based on the assumption that the disproportionation of EB, the present case study, is carried out in a gradientless recycle reactor (Fig. 28), where neither concentration gradients nor temperature gradients exist throughout the entire reacting volume. Moreover, such gradients will be absent inside the macropores of the zeolite pellet and also across the external boundary layer. Finally, the crystallites will be isothermal. Hence we have to treat only the coupled adsorption, diffusion and reaction inside the micropores of the crystallite. 2. The disproportionation of EB is treated as a first-order irreversible reaction, making use of the simplifications described earlier. Hence the rate of formation of the three DEB isomers is given by 6.3.8.2.1

RF,i = yi,pri kf cEB

(232)

where yi,pri denotes the primary isomer distribution which, precisely, is the ratio of the mole fractions of the various isomers. This may be chosen arbitrarily. Hence, by varying the primary isomer distribution,

in

qN

out

q out (N2)

2

in c EB

out c EB

in c DEB =0

Fig. 28

mcat, Sm

c iout , y out i i = o −, m −, p − DEB

Schematics of a gradientless recycle reactor.

References see page 1781

1776

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

restricted transition state shape selectivity and mechanistic aspects of the disproportionation reaction may be included in the simulations. 3. The isomerization reaction is described in analogy to Wei’s approach [134]: RI = K I c

(233)

where c denotes the isomer concentration vector, RI is the vector of the rates of production or disappearance of the isomers due to the isomerization reaction and KI is the isomerization rate matrix, which is defined as follows:
κ κ κ  KI = kI

1,1

1,2

1,3

κ2,1 κ3,1

κ2,2 κ3,2

κ2,3 κ3,3

(234)

The elements κi,j are obtained by letting RI (ceq ) = 0, which essentially means the condition that the isomerization rate must tend to zero as the isomer concentration approaches the thermodynamic equilibrium concentration ceq . At this point, it should be noted that a direct transformation of o− into p-DEB is not possible, hence the elements κ1,3 and κ3,1 are zero. 4. Reactions at the outer surface of the crystallites are neglected. 5. The diffusion of the various species inside the micropores is described according to the Fickian relations, i.e. neither multicomponent diffusion nor a concentration dependence of the diffusivities are taken into account. 6. An inert gas (N2 ) is assumed to be present in large excess. Concerning the small partial pressures of the reacting species under these conditions, the adsorption isotherms may be approximated by linear relationships [Eq. (235)]; the volume change due to reaction is neglected: ci,s = Ki ciout

(235)

7. The individual crystallites are assumed to be of spherical shape and of the same size (radius R). 6.3.8.2.2 Model Equations With these assumptions, the mass conservation laws for the reactant EB and the three DEB isomers inside the crystallite may be written as follows:

d2 cEB 2 dcEB kf cEB + = dr 2 r dr DEB,e

2 dci d2 ci + =− dr 2 r dr

(236)

yi,pri kf cEB + kI

3

This is a system of four coupled, second order ordinary differential equations. The solution of each equation must satisfy two boundary conditions. The first is dictated by the symmetry of the crystallite (sphere), which requires the concentration gradients of all species to disappear at the center of the crystallite:  dci  = 0; i = EB, o−, m−, p−DEB (238) dr r=0 The second condition can be derived from the overall mass balance (reactor) for the reacting species. As we have assumed a gradientless recycle reactor (Fig. 28), the fluid inside the reactor is supposed to be perfectly mixed. Thus, we have an ideal CSTR for which the mass balance of species i takes the form    dci  = 0; q ciin − ciout − mcat Sm Di,e dr r=R i = EB, o−, m−, p−DEB

where q denotes the volumetric flow rate to the reactor, mcat the catalyst mass and Sm the outer surface area of the crystallites per unit mass. Substituting the linear adsorption law of Eq. (235) into Eq. (239) and rearranging yields the following expression for the concentration gradient at the outer surface of the crystallite:  ci,s in  q ci − dci  Ki = ; i = EB, o−, m−, p−DEB  dr r=R mcat Sm Di,e (240) Inspection of the above equations shows that the EB mass balance [Eq. (236)] together with the corresponding boundary conditions [Eqs. (238) and (240)] is independent of the respective equations for the DEB isomers [Eqs. (237), (238) and (240)]. Hence Eq. (236) may be solved separately. The solution of the non-dimensional form of Eq. (236) has already been given in Section 6.3.3 [Eq. (49)]. Therefore, for convenience, we introduce the generalized Thiele modulus φF related to the disproportionation reaction [see Eq. (54)], which gives sinh(3φF r/R) cEB (r) = cEB,s r/R sinh(3φF )

Di,e i = o−, m−, p−DEB

cEB,s

;

in cEB

(237)

(241)

The unknown surface concentration cEB,s can be determined by differentiating Eq. (241) and substituting the result into Eq. (240). After rearranging, this yields

κi,k ck

k=1

(239)



1 DEB,e mcat Sm = + KEB qR



3φF −1 tanh(3φF )

−1 (242)

6.3.8 Control of Selectivity in the Disproportionation of Ethylbenzene to Benzene and Diethylbenzenes

Hence, if we substitute Eqs. (241) and (242) into Eq. (237), we have reduced the system to three secondorder ordinary differential equations. This boundary value problem is then solved numerically by applying a (mass-conservative) finite volume method. In this, the crystallite is divided into a number of small volumes, i.e. thin spherical shells of thickness h, over which the isomer mass balances are separately invoked. The thickness of the shells may either be constant throughout the entire crystallite (evenly spaced nodes) or may vary with the steepness of the actual concentration gradients (dynamic spacing). A dynamic node spacing, depending on the absolute values of the concentration gradients, provides an efficient control of the approximation error, i.e. the solution becomes virtually independent of the node spacing already at a smaller number of nodal points, which leads to significant savings in computation time. Figure 29 illustrates the subdivision of the crystallite into finite volumes: volume j is limited by coordinate rj to the east (e) and by coordinate rj −1 to the west (w). The concentration of species i inside volume j is assigned to the geometric center of the volume (P ). Towards the edges of the crystallite, i.e. the outer surface and the center, the node spacing is cut to half. The mass balance for an arbitrary volume element inside the crystallite can be formulated as Ni,w Aw − Ni,e Ae + Ri δV = 0

(243)

where Ni,w denotes the molar flux density of isomer i at the western boundary and Ni,e the same at the eastern boundary, Aw and Ae denote the respective boundary areas, Ri is the rate of production or disappearance of species i per unit volume and time and δV denotes the capacity of the volume element. The molar flux densities Ni,w and Ni,e can be expressed according to Fick’s first law:   dci  dci  ; N = −D (244) Ni,w = −Di,e i,e i,e dr w dr e e R

w

W

r1 c0 c1

rj − 2 cj − 1

P

rj − 1

The rate of production or disappearance of species i is given by Ri = yi,pri kf cEB + kI

3

κi,k ck

(245)

k=1

The boundary areas and the capacity of the volume element are calculated as follows: Aw = 4πrj2−1 ; δV =

Ae = 4πrj2 ;

4 π(rj3 − rj3−1 ) 3

(246)

Substituting Eqs. (244)–(246) into Eq. (243) and rearranging finally yields     3 dci  2 dci  2 − r r j j −1 dr e dr w rj3 − rj3−1 3 kI kf + κi,k ck = yi,pri cEB Di,e Di,e

(247)

k=1

The derivatives in Eq. (247) are approximated by central differences. This leads to a discrete representation of the problem in the form of linear equations. For j = 1 and j = N we have slightly modified equations, since here the boundary conditions have to be considered. Because the concentration of three isomers has to be calculated for each of the N volume elements as a whole, we have a system of N × 3 linear equations. This can be expressed in matrix notation as Ac = b

(248)

where the coefficient matrix A has a heptadiagonal structure. Such problems are solved very efficiently by means of LU decomposition [148]. 6.3.8.2.3 Simulation Results As already stated, the model was applied to simulate the effects of product shape selectivity during the disproportionation of EB on an HY zeolite. The various parameters required for the simulations, i.e. the adsorption constants and diffusivities, were taken from the literature [149]. The remaining simulation conditions were chosen similar to the operating conditions of related experimental studies on this reaction [150]. Table 7 gives the range of the respective variables covered in the simulations.

E

rj cj

1777

rN

r

cj + 1 cN + 1

Fig. 29 Subdivision of the crystallite into finite volumes and indexing of the volumes.

A Product Shape Selectivity (Unaffected by Other Factors) For these simulations, the primary isomer distribution is chosen according to the thermodynamic equilibrium References see page 1781

1778

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

Tab. 7

Parameter range used for the simulation of the disproportionation of EB on an HY zeolite

Parameter

Symbols

Primary isomer distribution

yo,pri , ym,pri , yp,pri KEB , Ko , Km , Kp DEB,e , Do,e, Dm,e, Dp,e

10−14 m2 s−1

Thiele modulus

φF

–

0.01–100

Isomerization rate constant

kI

s−1

10−8 –104

Reactant concentration

cEB

mol m−3

Flow rate

Q

mL s−1 (STP)

Catalyst weight

mcat

Crystallite radius

R

g µm

mcat −3 q = 43.8 kg sm

100

50

RD = Dp,e / Do,e

10 5

40

RD = 2 thermodyn. equilibr.

10−6

9, 14, 14, 14–1400 0.5, 1, 1, 1–1000

1.5 8 0.35 20

gas per m3 zeolite.

1000

y pout / %

0.067, 0.638, 0.295

m3

m−3a

Adsorption constants

60

10−4

10−2

k1 /

1

102

104

s−1

Simulated para-selectivity yout p at the reactor outlet as a function of the isomerization rate constant kI . Dependence on the ratio of the effective diffusivities of p- versus o-DEB, RD = Dp,e /Do,e .

Fig. 30

–

Effective diffusivities

(cf. Table 7). Such a situation would be encountered in practice when neither the reaction mechanism kinetically favors a particular isomer nor restricted transition state shape-selectivity effects occur. The disproportionation reaction is assumed to be unaffected by diffusion (i.e. φF < 0.01). The effective diffusivities of the ortho- and meta-isomers are fixed and assumed to be equal, but by a factor of RD smaller than the effective diffusivity of the para-isomer. Figure 30 shows the simulated, normalized fraction of the para-isomer in the DEB product stream at the reactor outlet. In the subsequent discussion, this will be called the para-selectivity for notational simplicity. From Fig. 30, it is clear that the observable para-selectivity passes through a maximum as the isomerization rate constant is increased. This may be explained as follows:

20 10−8

Numerical values

103

a m3

30

Units

on the left side of the maximum, where the isomerization reaction is slow compared with the rate of diffusion of the para-isomer, an increased isomerization rate will cause the fast-diffusing para-isomer to be more readily supplied. However, as the isomerization rate is increased beyond a certain limit, the para-isomer will be converted faster than it can diffuse out of the pore structure. Hence, at a certain value of the effective diffusivity of the para-isomer, the para-selectivity will begin to drop upon further increasing the isomerization rate constant. The maximum gain of the para-selectivity, which gives the size of the product shape-selective effect, naturally increases as the ratio of the diffusivities RD becomes larger, although not to 100%, but instead to a limiting value which is approached at RD > 100. This is controlled by several other factors, namely the absolute values of the diffusivities. The simulation results also show that the product shape-selective effect does not depend either on the inlet concentration of EB or on the rate of the disproportionation reaction (i.e. the rate constant kF ). Instead, it is affected by the ratio of the catalyst weight, divided by the total flow rate, mcat /q, which can be interpreted as a modified residence time. This is illustrated in Fig. 31, where the para-selectivity is plotted against the isomerization rate constant for different values of mcat /q. A reduced residence time leads to an increase in the para-selectivity. However, at the same time the conversion drops and thus the actual para-concentration at the reactor outlet is reduced. In addition, the maxima of the curves shown in Fig. 31 become less distinct. In the limiting case of zero conversion, the para-selectivity increases steadily with increasing value of the isomerization rate constant. Then a maximum is no longer observed. This is exactly the result obtained by Wei [134]. However, it is true only for

6.3.8 Control of Selectivity in the Disproportionation of Ethylbenzene to Benzene and Diethylbenzenes

60

100

mcat −3 q = 0.01 kg sm

RD = 10

mcat −3 q = 43.8 kg sm

80

RD = 10

50

43.8 40

1000 thermodyn. equilibr.

10−6

10−4

10

40

fF = 100 30

20

0 10−8

0.1 5

0.1

5.0

60

ypout / %

ypout / %

1779

10−2

1

102

104

20 10−8

k1 / s–1

small conversions. As soon as the conversion increases, the assumption of a negligible isomer concentration at the pore mouth, which was introduced as a boundary condition in Wei’s approach [134], is no longer valid. Figure 31 shows that in this case the para-selectivity begins to drop whenever the isomerization rate exceeds some limit. This has the important consequence that an increased activity of the zeolite does not automatically produce an enhanced yield of the para-isomer. Instead, an optimum range of kI is observed where the para-yield is at a maximum. B Product Shape Selectivity and Reactant Diffusion When the diffusion of Limitation (Combined Effect) the reactant begins to influence the effective rate of the disproportionation reaction, the observable product-shape selective effect is reduced. This is obvious from Fig. 32, which shows the para-selectivity as a function of the isomerization rate constant kI for various values of the Thiele modulus φF (by variation of the rate constant kF ). In the limiting case of φF approaching infinity, no shapeselective effect at all is encountered. This comes as no surprise, since the penetration depth of the reactant into the porous structure of the crystallite in this case approaches zero. Hence the distance for the isomer molecules to pass from the point where they are formed until they reach the outer surface of the crystallite tends to zero, which is precisely the range where differences in the diffusivities of the isomers – together with the isomerization reaction taking place – generate product shape selectivity. It may be obvious from this explanation that the paraselectivity will become virtually independent of the size

10−6

10−4

10−2

102

1

104

k1 / s−1

Simulated para-selectivity yout p at the reactor outlet as a function of the isomerization rate constant kI . Dependence on the Thiele modulus of the disproportionation reaction φF .

Fig. 32

of the crystallites when the disproportionation rate is severely affected by a limited diffusivity of the reactant EB. C Product Shape Selectivity and Selective Adsorption The effect of different adsorption (Combined Effect) constants of the isomers is illustrated in Fig. 33, where the para-selectivity is plotted against the isomerization rate constant for different ratios of the adsorption constants RK = Kp /Ko (Ko = Km = constant). From this we note that the maximum para-selectivity increases as the

100

80

ypout / %

Simulated para-selectivity yout p at the reactor outlet as a function of the isomerization rate constant kI . Dependence on the modified residence time mcat /q(cEB = constant).

Fig. 31

Thermodyn. equilibr.

mcat −3 q = 43.8 kg sm

0.01

RD = 10

0.1 = Kp / Ko

60 0.5

40

1

Thermodyn. equilibr.

20

0 10−8

2 3 10−6

10−4

10−2

k1 /

1

102

104

s−1

Simulated para-selectivity yout at the reactor outlet as p a function of the isomerization rate constant kI . Dependence on the ratio of the adsorption constants of p- versus o-DEB, RK = Kp /Ko (Ko = Km = constant). Fig. 33

References see page 1781

1780

6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

adsorption constant of the para-isomer is reduced below the respective constants of the other isomers. This effect is most pronounced at large values of kI . Moreover, when the ratio of the adsorption constants RK tends to zero, the paraselectivity increases steadily with increasing value of the isomerization rate constant kI . A maximum is no longer observed. For RK above unity, where the para-isomer is more strongly adsorbed than the others, the paraselectivity drops below the thermodynamic equilibrium composition once the isomerization rate constant has become sufficiently large. This can be explained by the fact that, if the isomerization reaction is very fast, then the isomer concentration will approach the thermodynamic equilibrium at any point inside the crystallite. Hence the isomer distribution in the surrounding fluid is controlled only by the adsorption of the various species. D Product Shape Selectivity and Restricted Transition In Fig. 34, State Shape Selectivity (Combined Effect) the influence of the primary isomer distribution on the observable para-selectivity is shown by plotting the para-selectivity versus the isomerization rate constant for different primary isomer distributions. It is apparent that an increase in the fraction of the para-isomer produced by the disproportionation will only lead to a net increase of the observable para-selectivity if the isomerization rate is not too high. Otherwise, most of the para-isomer, which is selectively formed by the primary disproportionation reaction, will be converted to the other isomers by the secondary isomerization reaction. As kI approaches infinity, in any case the thermodynamic equilibrium distribution of the isomers must be obtained. 100 0/0/100 80

mcat = 43.8 kg sm−3 q RD = 10

10/10/80

ypout / %

10/20/70 60

yo, pri / ym, pri / yp, pri

10/30/60 10/40/50

40

10/50/40 7/63/30 Thermodyn. equilibr.

20 10−8

10−6

10−4

10−2

1

102

104

kI / s–1

Simulated para-selectivity yout p at the reactor outlet as a function of the isomerization rate constant kI . Dependence on the primary isomer distribution yi,pri .

Fig. 34

Controlled Modification of the Pore Structure The observation of shape-selective effects in zeolite catalysis suggests the possibility of tailoring catalyst properties, in the sense that the selectivity to certain products could be maximized. A variety of methods have been developed to modify the pore structure and the catalytic activity of zeolites and related materials. These may be divided into two groups, namely methods applied during the synthesis and actions to be taken after the zeolite material has been synthesized (post-synthesis modification). In the following brief discussion of such methods, we will focus on the case of zeolites applied to acid-catalyzed reactions, where the activity of the material is due to Brønsted acid sites, stoichiometrically linked to the Al atoms of the host framework and to Lewis acid sites. The choice among the variety of different types of zeolites and related materials in a practical situation will depend on the characteristics of the reacting system and the types of selectivity effects to be expected. The pore size, the deactivation behavior and the chemical and thermal stability of the zeolite material determine whether or not a particular catalyst is attractive. The necessary condition for shape-selectivity effects to occur is that the pore size has to meet the dimensions of the reacting molecules. The radius of the crystallites and the strength and number of the acid sites may then be adapted to the actual requirements during synthesis. Post-synthesis methods (pore-size engineering) allow an existing shape-selectivity effect to be intensified and also a new one to be established. However, normally not only the pore size will be influenced by most of these methods, but also the catalytic activity. Vansant [151] gave a classification of post-synthesis modification methods which covers the entire range of zeolite applications (gas separation, gas purification, encapsulation of gases and catalysis). Of particular interest for shape-selective catalysis are the modification of the zeolite by means of cation exchange and the modification of the inner and/or outer crystallite structure by treatment with chemically reacting agents which leads to a deposition of additional functional groups or compounds. This can be done either in the gas phase [i.e. by chemical vapor deposition (CVD)] or in the liquid phase. By means of ion exchange using metal cations of different size and specific charge, the geometric restrictions, the number and strength of the Brønsted acid sites and the adsorption properties of the zeolite material can be influenced. Investigations of this kind have been reported, for example, for ZSM-5 and mordenite catalysts [152, 153]. 6.3.8.3

References

During a treatment with chemically reacting agents, these may penetrate into the pore system, where they react to form deposits, which may lead to a narrowing of the pores and to a change in the catalytic activity. Boric acid, trimethyl borate, phosphoric acid and triphenylphosphine may serve as examples of suitable reacting agents. Such methods have been applied successfully to the selective production of p-xylene by disproportionation of toluene on a ZSM-5 zeolite, where para-selectivities in excess of 90% were claimed [154, 155]. CVD of silanes, along with subsequent calcination using steam, can be utilized to deposit silica (SiO2 ) inside the pore system. By variation of the temperature, the partial pressure of the silane and the duration of the treatment, location and amount of the deposited material can be controlled [151]. When, for example, tetraethoxyor tetramethoxysilane is used as a reacting agent on a mordenite, ZSM-5 or β-zeolite, then a controlled deactivation of only the external crystallite surface is possible [156, 157]. This is because these are rather bulky molecules which are not able to diffuse into the pore system of the crystallite. Alternatively, an irreversible adsorption of bulky bases may serve to destroy the undesired external acidity. Suitable basic compounds are 4-methylquinoline for ZSM-5 [158] and tributyl phosphite for mordenites [159]. Finally, a controlled precoking, and also the undesired coking during reaction, may cause a change in selectivity. Coke in this general sense here is meant to include any type of undesired higher hydrocarbons which may be formed from an organic reactant [160]; however, polycondensed aromatics are of particular importance. The blocking of the catalytically active surface by coking may be traced back to basically two reasons. Either the hydrocarbon, due to its size, is trapped inside the zeolite cage or channel crossing or its vapor pressure at the reaction conditions is too low. The latter may also cause deactivation of the external surface of the crystallite. During the deposition of coke inside the pores, either pore blockage or site coverage [160] may occur. In the case of pore blockage, the deactivation, at the same level of coke deposited, is more pronounced than in the case of site coverage. Analogous to CVD, an appropriate choice of the precoking agent along with the precoking conditions (temperature) allows control of whether the deposits will be placed at the outer surface of the crystallites or inside the microporous structure. Investigations concerning the effects of precoking have been carried out, for example, on ZSM-5 catalysts for xylene production by disproportionation or alkylation of toluene [161]. During these experiments, an increase in the p-xylene selectivity could be observed. By a

1781

defined precoking, it is also possible to generate product selectivity effects initially not present, as has been shown for the disproportionation of ethylbenzene on HY zeolites [150]. 6.3.9

Concluding Remarks

For chemists and chemical engineers, this chapter is meant to give insight into the problems of transport phenomena combined with simultaneous chemical reactions in porous media. The intention was to facilitate a quick and effective understanding of this important field and also to offer the most important references for extended studies. A further goal was the presentation of a consistent nomenclature for the most important quantities in this area. Finally, a recommendation for engineering applications is given: when analyzing a new system in the field of heterogeneous catalysis, first the criteria in Section 6.3.5 should be applied in order to decide whether or not transport phenomena play a role in a particular case. In the next step, a decision has to be made as to which of the different models presented above should be used. The main guideline should always be to use models that are as simple as possible with a minimum number of parameters. However, in special cases, e.g. if the reaction is hindered by strong adsorption effects in micropores or if non-stationary operation must be described, it is worthwhile to accept a larger mathematical effort in using more complicated models, for example based on the Maxwell–Stefan formulation of diffusion. References 1. R. Aris, The Mathematical Theory of Diffusion and Reaction in Permeable Catalysts. Vol. 2: Questions of Uniqueness, Stability and Transient Behavior, Clarendon Press, Oxford, 1975, pp. 232. 2. G. Emig, Top. Curr. Chem. 1970, 13, 451. 3. V. Hlav´acˇ zek, M. Marek, J. Catal. 1969, 15, 17. 4. V. Hlav´acˇ zek, M. Marek, J. Catal. 1969, 15, 31. 5. V. Hlav´acˇ zek, Chem. Eng. Sci. 1970, 25, 1517. 6. V. Hlav´acˇ zek, J. Catal. 1971, 22, 364. 7. G. Damk¨ohler, Chemieingenieur 1937, 3, 359. 8. Y. B. Zeldovich, Acta Phys. Chim. USSR 1939, 10, 583. 9. E. W. Thiele, Ind. Eng. Chem. 1939, 31, 916. 10. A. Wheeler, in Advances in Catalysis, W. G. Frankenburg, E. K. Rideal, V. I. Komarewsky (Eds.), Vol. 3, Academic Press, New York, 1951, p. 249. 11. P. B. Weisz, C. D. Prater, in Advances in Catalysis, W. G. Frankenburg, E. K. Rideal, V. I. Komarewsky (Eds.), Vol. 6, Academic Press, New York, 1954, p. 143. 12. G. F. Froment, K. B. Bischoff, Chemical Reactor Analysis and Design, Wiley, New York, 1979, Chapter 3. 13. E. L. Cussler, Diffusion–Mass Transfer in Fluid Systems, Cambridge University Press, Cambridge, 1984, 536 pp.

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6.3 Simultaneous Heat and Mass Transfer and Chemical Reaction

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154. W. W. Kaeding, C. Chu, L. B. Young, B. Weinstein, S. A. Butter, J. Catal. 1981, 67, 159. 155. M. Sayed, J. Vedrine, J. Catal. 1995, 101, 43. 156. Y. Chun, X. Chen, A.-Z. Yan, Q.-H. Xu, in Zeolites and Related Microporous Materials: State of the Art 1994, J. Weitkamp, H. G. Karge, H. Pfeifer, W. H¨olderich (Eds.), Studies in Surface Science and Catalysis, Vol. 84, Elsevier, Amsterdam, 1994, p. 1035. 157. T. Hibino, M. Niwa, Y. Murakami, Zeolites 1993, 13, 518. 158. J. R. Anderson, K. Foger, T. Mole, R. A. Rajadhyaksha, J. V. Sanders, J. Catal. 1979, 58, 114. 159. T. Matsuda, E. Kikuchi, in Zeolites and Microporous Crystals, T. Hattori, T. Yashima (Eds.), Studies in Surface Science and Catalysis, Vol. 83, Elsevier, Amsterdam, 1993, p. 295. 160. M. Guisnet, P. Magnoux, in Catalyst Deactivation 1994, B. Delmon, G. F. Froment (Eds.), Studies in Surface Science and Catalysis, Vol. 88, Elsevier, Amsterdam, 1994, p. 53. 161. M. A. Uguina, D. P. Serrano, R. van Grieken, S. Venes, Appl. Catal. A 1993, 99, 97.

6.4

Magnetic Resonance Imaging Lynn F. Gladden∗ , Michael D. Mantle, and Andrew J. Sederman

6.4.1

Introduction

In the last 10 years, magnetic resonance imaging (MRI) has found increasing application in the study of catalysts and catalytic reactors. Early work focused on imaging single-phase liquid hydrodynamics in reactors and this continues to be an area of considerable interest. More recently, these studies have been extended to investigations of multi-phase flows, single-phase gas flows and also to chemical mapping inside reactors [1]. The principal attribute of MRI in the context of heterogeneous catalysis is that it is possible, in principle, to resolve spatially measurements of flow velocity, molecular diffusion and chemical composition within a reactor. These measurements are non-invasive and do not require any chemical or physical tracer to be introduced into the system. As with tomographic methods, MRI offers the opportunity to see what is happening inside the reactor. Although MRI may not compete with optical and capacitance tomographies in terms of data acquisition times, it does have the ability to study optically opaque systems with sufficiently fast data acquisition times (1–20 ms) that unsteady-state phenomena can be studied. Magnetic resonance (MR) techniques do have their limitations. Ferromagnetic materials cannot be studied and any process environment must be designed to fit ∗

Corresponding author.

inside the bore of the superconducting magnet. In a vertical geometry, such magnets usually require that a sample does not exceed 6 cm in diameter. However, magnets capable of studying reactors with external dimensions up to 30 cm can be obtained. Overall, the role of MRI as currently used is to provide images that challenge our current understanding of the way catalytic reactors operate. In particular, MR methods are particularly well suited to investigating both transport (i.e. diffusion, dispersion and flow) and reaction processes within the same sample environment. In many cases MRI offers the opportunity to study 3-D systems directly for the first time. The data obtained are used in two main areas of application: (i) gaining insight into the effect of catalyst properties, reactor design and process operating conditions on the hydrodynamics and chemical conversion occurring with the reactor and (ii) critically evaluating the predictive power of numerical codes for simulating reactor performance. MRI also plays an important role in the direct measurement of parameters to be used in numerical modeling schemes and design correlations. For example, MRI has already been used to provide the first direct measurements of catalyst wetting in a trickle-bed reactor. In situ local determinations of molecular diffusion within catalyst pellets and molecular dispersion within the inter-particle (void) space of fixedbed reactors have also been measured. Ongoing activities are addressing the in situ measurement of mass transfer coefficients within reactors. The key contribution that MRI can make is that it provides measurements of quantities on both a local and global scale, i.e. we can study the ‘‘global’’ flow field across the entire reactor cross-section and we can resolve the flow field between individual catalyst pellets within the bed. MRI provides images that clearly demonstrate that the ‘‘local’’ and ‘‘global’’ properties of a catalytic reactor can differ widely. Indeed it is the heterogeneity in local characteristics of the reactor that may often determine its overall performance. Section 6.4.2 outlines the basic principles of MR imaging and transport measurement techniques and highlights recent developments in the field. In Section 6.4.3, several examples are presented with the aim of giving the reader an overview of the state-of-the-art capabilities of MRI applied to catalysts and catalytic reactors. 6.4.2

Basic Principles

There are two main families of MRI methods used in catalysis: microimaging and flow imaging. Microimaging usually refers to the imaging of the internal structure of a sample, perhaps with spatial mapping of chemical composition, distribution of gas and liquid and even transport

6.4.2 Basic Principles

properties such as molecular diffusivity; such experiments typically give a spatial resolution of ∼30–50 µm. Flow imaging is usually performed at slightly poorer spatial resolution of ∼100–500 µm and it gives images of the flow field within the system of interest. In this section, the basic principles of spatially resolved measurements are presented and an overview of how data acquisition times are decreased is given. Increasing the rate of image acquisition has been an important step forward in applying MR techniques to the study of catalyst and reactors. The principles of transport measurements are then introduced. Transport measurements include quantification of diffusion, dispersion and flow. All transport measurements can be made with or without spatial resolution. To achieve spatial resolution the transport measurement pulse sequences (Section 6.4.2.3) are integrated into an appropriate imaging sequence (Section 6.4.2.1). Magnetic Resonance Imaging Only the very basic concepts of MRI are presented here; the interested reader should refer to excellent texts by Callaghan [2] and Kimmich [3] for more detailed discussion. The principles of MRI develop directly from those of NMR spectroscopy. Indeed, the power of MRI is readily appreciated once we identify that, in principle, all the techniques of NMR spectroscopy can be integrated into imaging experiments – we are simply adding spatial resolution to the spectroscopic measurement. The challenge in doing this is that since NMR is a relatively insensitive spectroscopic tool it is often difficult to achieve required levels of signal-to-noise ratio in the acquired spatially resolved data. It is for this reason that most MRI studies are reported for 1 H observation. The 1 H nucleus is the most ‘‘NMR friendly’’ nucleus with ∼99.99% natural abundance and the highest NMR sensitivity of all nuclear spins (except 3 H). As follows from the principles of NMR spectroscopy, when a nucleus of non-zero nuclear spin quantum number is placed in an external magnetic field (typically a superconducting magnetic field of 2–10 T), its nuclear spin energy levels become non-degenerate. By exposing the system to electromagnetic energy of appropriate frequency [radiofrequency (r.f.)], a resonant absorption occurs between these nuclear spin energy levels. The NMR spectrum is obtained by Fourier transformation of the time domain response or NMR signal, S(t), of the nuclear spin system following the r.f. excitation. This time domain signal is recorded as a decaying voltage in a receiver coil placed around the sample. The frequency, ω0 , at which the resonant absorption occurs is called the resonance (or Larmor) frequency and is proportional to the strength of the external magnetic field, B0 , used in the 6.4.2.1

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experiment and the specific nuclear isotope under study: ω0 = γ B0

(1)

where γ is the gyromagnetic ratio, which is an isotopespecific property. The precise energy-level splitting is slightly modified by the electronic environment of the nucleus under study; thus ω0 is also modified and becomes specific to individual molecules containing the element of interest. Hence we can take a conventional NMR spectrum of a mixture of chemical species and identify the presence of particular molecular species in that mixture. To obtain spatial resolution, the basic spatially unresolved experiment is still performed but, by applying a spatially varying magnetic field in addition to the large static field B0 , the resonance frequency of species within the sample becomes a function of position and strength of the applied gradient. Thus, for a magnetic field gradient applied along the z-direction, Gz : ωz = γ (B0 + Gz z)

(2)

The measurement is calibrated such that the relationship between resonance frequency and spatial position is known. Upon application of the field gradient, the nuclear spins at every spatial location along the direction of that gradient have a different resonance frequency and therefore we acquire a time domain signal that represents, after Fourier transformation, a 1-D projection of that nuclear spin density along the direction of the applied gradient. 2-D and 3-D images are acquired by applying gradients in two and three orthogonal directions, respectively. As with NMR spectroscopy, after appropriate calibration, the acquired NMR signal is a quantitative measure of the number of nuclear spins present, i.e. a quantitative measure of chemical concentration. To understand imaging pulse sequences the concept of the k-space raster is used [4, 5]. Rewriting Eq. (2) for the general case of the variation of resonance frequency with spatial position r, we obtain ω(r) = γ (B0 + G · r)

(3)

Further, neglecting the effect of nuclear spin relaxation, the acquired signal dS in an element of volume dr at position r with spin density ρ(r) is given by dS(G, t) = ρ(r) exp[iω(r)t] dr

(4)

Inserting Eq. (3) into Eq. (4) gives dS(G, t) = ρ(r) exp[i(γ B0 + γ G · r)t] dr References see page 1800

(5)

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A transformation into the rotating frame of reference of the nuclear spin system then allows us to rewrite Eq. (5) as  S(t) = ρ(r) exp[iγ G · rt] dr (6) By defining a k-space vector as k = (γ Gt/2π), it follows that Eq. (6) can now be written as  S(k) = ρ(r) exp[i2πk · r] dr (7) and that the spatial distribution of spins is then given by the inverse 3-D Fourier transform:  ρ(r) = S(k) exp[−i2πk · r] dk (8) The simple imaging experiment described above is therefore achieved by sampling all relevant k-space. From the definition of k-space, it is seen that this can be done by collecting data at different times with a fixed field gradient. Alternatively, k-space can be sampled by changing the magnetic field gradient. Modern imaging pulse sequences acquire data as a function of both magnetic field gradient and time to produce a k-space data array, which, following Fourier transformation, gives a real space spin density image. Figure 1 shows a schematic of a simple 2-D imaging sequence. In this case, let us assume that the sample is cylindrical and oriented along the z-axis and an xy image is to be recorded. The first component of the pulse sequence is the so-called ‘‘slice selection’’ phase. The procedure comprises the application of a narrowband r.f. excitation simultaneously with a magnetic field gradient

imposed along the direction in which the 2-D image is to be taken. The effect of this procedure is that the only spins that will be excited will be those that resonate within the bandwidth ω of the r.f. pulse – and therefore only those spins that lie within a certain ‘‘image slice thickness’’ z. The rest of the sequence acquires data along different rows of the k-space raster for successive r.f. excitations. With reference to Fig. 1, a field gradient is applied in the x-direction, simultaneously with the maximum magnitude negative field gradient in the ydirection. A slice-selective π ‘‘refocusing’’ pulse is then applied; this is represented on the k-space raster as a move from kx,max , −ky,max to −kx,max , ky,max . A second gradient is then applied along the x-direction while data, typically 128 or 256 complex data points, are acquired at a specified digitization rate. The digitization rate will define the spacing of the points acquired in k-space. The signal, S(t), that is acquired during application of the second x-gradient is said to be frequency-encoded, because the signal is acquired in the presence of a magnetic field gradient and spatial position is related directly to signal frequency. This gradient along the x-direction is therefore referred to as the frequencyencoding gradient, also being termed the ‘‘read’’ gradient. The acquisition of complex data points in the presence of a constant linear ‘‘read’’ gradient yields a straightline k-space data trajectory the direction of which is defined by the Cartesian orientation of the gradient. A straight, equally spaced k-space trajectory will always result, as long as the read amplitude gradient is kept constant and the digitization (acquisition) rate of the complex data is fixed. The spin system is then allowed to

TE

p/2

p

S(t )

ky (phase)

r.f.

Gz Gx

kx (read)

Gy

repeat N times (a)

time

(b)

(a) Schematic representation of a slice-selective two-dimensional spin-echo pulse sequence. In this pulse sequence the magnetic field gradient (Gy ) is varied for successive acquisitions of different rows of the k-space raster. (b) The corresponding k-space raster showing the action of one pass through the pulse sequence. Following a sufficient T1 -relaxation period, the sequence is repeated to acquire a second row of the k-space raster. Acquisition of each row of k-space requires a separate r.f. excitation and application of a Gy -gradient of different magnitude. The signal or ‘‘echo’’ is acquired after an echo time TE.

Fig. 1

6.4.2 Basic Principles

return to equilibrium and the pulse sequence is repeated, this time with the second largest negative y-gradient being applied – hence ‘‘reading’’ the next row of k-space. This process is repeated until the entire raster has been sampled. A 2-D Fourier transformation of these data followed by modulus correction gives a 2-D spin density image. 2-D images are typically acquired in a few minutes using this approach. As in NMR spectroscopy, nuclear spin relaxation properties play a key role in implementing individual pulse sequences. Each pulse sequence is initiated using an r.f. pulse. Following this excitation pulse, the nuclear spin system has excess energy. To return to thermal equilibrium, the spin system has to redistribute this excess energy – a process known as ‘‘relaxation.’’ A number of different relaxation times characterize different mechanisms for this redistribution of energy. The most important are the spin–lattice relaxation (T1 ) and spin–spin relaxation (T2 ) time constants. Albeit a generalization, T1 influences the choice of the recycle time between application of successive r.f. pulses during acquisition, while T2 influences selection of the echo time, TE, employed in the r.f pulse sequence. Since each chemical species will have its own T1 and T2 characteristics and these will vary depending on the physical state in which that species exists, careful selection of the recycle and echo times employed in the pulse sequence of interest allows us to ‘‘weight’’ preferentially the contrast obtained in the image to a particular subset of nuclear spins (that is, chemical species in a given physical state) within a multicomponent, multiphase mixture. Fast Data Acquisition At the heart of recent developments in applying MRI in reaction engineering research has been the development and implementation of fast spatially resolved MR measurements. Fast imaging is considered here to describe the acquisition of say a 128 × 128 2-D image in less than 1 s. Such pulse sequences are also referred to as ‘‘ultra-fast’’ and ‘‘rapid’’ imaging sequences. In using a fast sequence, it is often necessary to relax a desire for high spatial resolution (∼30–50 µm) and take great care, if quantitative data are required, to account for relaxation contrast effects in the final image. In general, the acquisition speed of an MR image may be improved by two basic methods [6]: 6.4.2.2

(i) the sampling of more than one line of k-space for each r.f excitation of the spin system (ii) the use of rapid multiple r.f excitations (and subsequent acquisitions). Three general ultra-fast sampling strategies are used: echo planar imaging (EPI) [4, 7–11], rapid acquisition

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with relaxation enhancement (RARE) [12] and low excitation angle imaging (also known as fast low-angle shot imaging, FLASH or SNAPSHOT imaging) [13]. The first two methods are based on the sampling of more than one line of k-space for each r.f. excitation and the third uses rapid multiple r.f. excitations. All of these methods have been used in application to catalysis and examples are given in Section 6.4.3. In reality, the quality of the data obtained from application of a particular pulse sequence is very system dependent. However, in general, one can say that although EPI ultimately provides the fastest data acquisition speeds, it is also the least robust in application and can often produce significant artifacts when studying multiphase systems. This is because the performance of the EPI pulse sequence is particularly adversely affected by the local heterogeneities in magnetic susceptibility that occur across gas/liquid/solid boundaries within a sample. Measurement of Diffusion, Transport and Flow The most widely applicable and robust strategy for measuring transport processes, which include molecular diffusion, dispersion and bulk flow phenomena, exploits the principle of phase-encoding. Other methodologies do exist, most notably time-of-flight methods [14–16] and rapid image acquisition techniques such as the single excitation multiple image RARE (SEMI-RARE) pulse sequence [17]. Both of these approaches have been used in specific applications, most notably in studies of flow in ceramic monoliths [18–20]. Although details of the implementation may vary, the basic principles of transport measurement by phase encoding are common to all measurements and are illustrated in Fig. 2. The same methodology is used to quantify incoherent (for example, diffusion and dispersion) and coherent (flow) processes occurring within the same system. Typically, two magnetic field gradients (g) of equal but opposite direction are applied for short time periods, δ (∼1–2 ms). When considering the application of pulsed magnetic field gradients to measure transport processes, we use the lower-case symbol g, as opposed to G, which is reserved for use in describing the imaging gradients [Eq. (2)]. Considering Fig. 2, the action of the first pulsed magnetic field gradient (applied along z) 6.4.2.3

∆ gz d

time

The principle of transport measurements using the ‘‘phase shift’’ approach. Two pulsed magnetic field gradients (of magnitude g and duration δ) are applied a time  apart.

Fig. 2

References see page 1800

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6.4 Magnetic Resonance Imaging

is to impose a phase offset (say φ1 ) to a spin characteristic of the spatial position of that spin in the z-direction when the pulse was applied. As follows from Eq. (2), in the rotating frame of the spin system φ1 = γ gδz1 , where δ is the duration of the applied gradient, z1 is the position of the spin and g is the magnitude of the magnetic field gradient along the z-direction. The spin system is then left to evolve for an observation time, , after which an equal but opposite polarity magnetic field gradient pulse is applied which gives the spins a further phase offset, φ2 , such that the total phase offset is φ1 + φ2 = γ gδ(z1 –z2 ). Clearly, if the molecule (spin) has not moved during the time , it will experience a net phase shift φ1 + φ2 = 0 at the end of the pulse. However, if the molecule has moved during the time  and therefore z1  = z2 , then φ1 + φ2  = 0 and observation of the magnetization associated with the spin will show a phase shift that is proportional to the distance moved (z1 –z2 ). Since γ , g and δ are known, the displacement or average velocity over the time-scale  is obtained. A typical transport measurement would proceed by making several measurements at differing values of δ or g and recording the resulting phase shift and amplitude of the signal. Coherent and incoherent transport processes are distinguished because of the different effects that they have on the overall signal. In the case of coherent motion all spins will have moved the same distance in a given observation time and therefore have the same phase offset, while the signal amplitude remains constant. However, in the case of incoherent transport, where there is a distribution of displacements, there will be a corresponding distribution of phase shifts such that the signal amplitude will decay. When applied to the measurement of molecular diffusion in catalysts, this transport measurement strategy is often referred to as the pulsed field gradient (PFG) or pulsed gradient spin-echo (PGSE) NMR method [21, 22]. 6.4.3

Applications Imaging Individual Catalyst Pellets Microimaging has been used to characterize heterogeneity in both structure and transport within catalyst pellets. The in-plane spatial resolution achieved in these investigations is typically 30–50 µm for pellets of dimensions 1–5 mm. In the majority of cases, investigations have addressed the pure (usually oxide) support so that the quantitative nature of the data obtained is not lost because of the presence of metal (which introduces an unknown degree of nuclear spin-relaxation time contrast into the images). Clearly, by working with typical spatial resolutions of ∼30–50 µm, individual pores within the material are not resolved. However, a wealth of information is 6.4.3.1

obtained even at this lower resolution [23–26]. Typically, three characteristics of catalyst pellets are spatially resolved in these experiments: porosity, locally averaged pore size and liquid diffusivity within the pore structure. The measurements are made by imbibing a liquid within the porous catalyst. The liquid used may be water or a liquid of particular relevance to the reaction of interest; maps or images of 1 H spin density are most commonly acquired. Assuming that relaxation contrast effects are made negligible, the spin density map is a quantitative measure of the amount of water present within the porous catalyst pellet, that is, it is a spatially resolved map of porosity (also known as voidage or void volume). Estimates of overall pellet porosity obtained from the MR data agree to within 5% with those obtained by gravimetric analysis. At the same level of spatial resolution it is also possible to map the spin–lattice relaxation time (T1 ) and the diffusion coefficient of the imbibed liquid. The T1 map yields information about the spatial distribution of mean pore size within a given image pixel. Longer and shorter T1 values identify locations at which the liquid is confined within larger and smaller diameter pores, respectively. Relaxation time imaging of catalyst pore size is a useful tool to complement the 1-D determinations of pore-size distribution obtained from nitrogen adsorption or mercury porosimetry. The nitrogen- and mercury-based methods provide a more accurate measure of the individual pore sizes whereas the MR data reveal any gross variations in the spatial distribution of those pore sizes within the catalyst. A useful application of this MR technique has been to characterize the spatial heterogeneity in porosity within a catalyst pellet introduced during the manufacturing process. The influence of spatial variations in porosity within the catalyst pellet on the transport processes occurring within it are then studied directly by acquiring maps of molecular diffusion coefficient [23]. Images of pellet structure and the transport processes occurring within that structure have been used in developing strategies for predicting transport in porous catalyst support pellets [27–31]. MRI has also been used to gain insight into catalyst preparation. For example, Khitrina et al. [32] have mapped the hexachloroplatinate dianion distribution during various impregnation procedures. This was achieved by imbibing cyclohexane within the catalyst. Interaction of cyclohexane molecules with hexachloroplatinate dianions increases the spin–lattice relaxation time of the cyclohexane. Therefore, by spatially resolving the T1 of cyclohexane within the catalyst, the distribution of the hexachloroplatinate dianion is determined. More recently, changes in liquid distribution within a catalyst pellet during reaction have been studied [33]. In particular, 1 H images of liquid distribution during α-methylstyrene evaporation accompanied by its hydrogenation (with vapor-phase reactants) within a cylindrical Pt/γ -Al2 O3 catalyst pellet

a) 2 min 11 s 109 °C

b) 6 min 33 s 109 °C

d) 19 min 39 s 167 °C

g) 31 min 45 s 100 °C

e) 23 min 01 s 148 °C

c) 10 min 55 s 141 °C

f) 27 min 23 s 113 °C

h) 36 min 07 s, 64 °C

Low

Liquid content

High

6.4.3 Applications

5 mm Spatial maps of the liquid phase in a catalyst pellet under reactive conditions. For each image, the temperature and time of detection are indicated. The intensity scale is shown on the left-hand side. Reprinted with permission from Ref. [33]. Copyright (2002) American Chemical Society.

Fig. 3

have been reported. Figure 3 shows the spatial maps of the liquid phase within the catalyst pellet under catalytic reaction conditions. The in-plane spatial resolution was 230 µm × 140 µm, with an image slice thickness of 2 mm; the pellet was of diameter and length 2 mm. The images show that under conditions of simultaneous evaporation of the reactant and hydrogenation of its vapor, two regions with very different compositions can form within the catalyst, with one region having a large liquid content and the other filled with the reacting gas mixture. Microimaging has also been used at the single-pellet level [34] and in small beds of pellets to map coke distribution [35, 36]. The effect of coke deposition on pore structure and molecular diffusion within supported metal catalysts has also been investigated [37]. A range of strategies have been employed to study coked catalysts. The most straightforward is to image a liquid imbibed into the coked pellet. Regions of coke laydown are associated with reduced signal intensity compared with the uncoked state [34]. Alternative approaches exploit the effect of the interaction of a probe molecule with any coke present to provide T2 contrast within an image. For example, Bonardet et al. [35] imaged the 1 H spin density associated with 2,3-dimethylpentane probe molecules adsorbed on pellets of HY zeolites coked to levels of 7.5 and 10% w/w. By varying the echo time used in the pulse sequence it was found that the T2 of the probe molecule, which is a function of the aromaticity of the coke, varies within

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the pellet. In particular, the more heavily coked zones were associated with coke characterized by the shortest T2 values, indicative of coke of the most graphitic content. The 1-D image profiles of a 7.5% w/w coked sample showed a heterogeneous coke distribution in the pellet at the macroscopic level. The most heavily coked region was that first exposed to reaction and the heterogeneity in coking was found to be associated with structural heterogeneities arising from the pellet manufacturing process. The more highly coked sample had a more homogeneous coke distribution, with the coke being of a homogeneous graphitic content throughout. Bar et al. [36] investigated the location of coke deposits in industrial HZSM-5 pellets, of diameter 5 mm, contained in a small fixed-bed reactor by imaging the coke directly by use of the SPRITE technique, a special MR imaging sequence for detecting materials with short T2 relaxation times. In the example given, SPRITE was used to produce a 1-D coke profile along the axis of the model fixed bed of inner diameter 3 cm, containing two layers of coked pellets (20.5 wt% coke) separated by a 3.3-cm layer of fresh pellets. A spatial resolution of 0.5 cm was obtained, this being limited by the rapid nuclear spin relaxation times of the sample. 6.4.3.2

Hydrodynamics in Reactors

6.4.3.2.1 Single-Phase Flow High-resolution MRI investigations of fluid flow in packed beds with columnto-particle diameter ratios in the range 10–20 are now routine. Interest has focused on these narrow packed beds because of the diameter of the superconducting magnets available. If high spatial resolution of the flow field within the inter-particle space of the bed is not required, reactors of far greater column-to-particle ratio can be studied. Figure 4 shows 2-D sections through 3-D volume images of the x, y and z components of flow within a fixed bed of non-porous spherical particles. The map of the z-component of the flow velocity is the most interesting; the +z direction is the direction of superficial flow in the reactor. In this example, the superficial flow velocity was 0.56 mm s−1 , corresponding to a Reynolds number of 2.8; hence flow in much of the bed is dominated by viscous forces, associated with flow velocities less than, or of order of, the superficial velocity. The most striking characteristic of these images is the extent of heterogeneity in the flow field; a relatively small fraction of the inter-particle space carries a high percentage of the liquid flow [38–41]. On the basis of these images, it is clear that any theoretical analysis of the flow within such a reactor must account for distinct populations of fast- and References see page 1800

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6.4 Magnetic Resonance Imaging

yz

yz

yz

xy

xy

xy

xz

xz

xz

y x z (a)

–7.2 mm s–1

Vx , Vy

7.2 mm s–1

–2.7 mm s–1

Vz

9.0 mm s–1

(b)

(c)

MR visualization of water flowing within a fixed bed of non-porous glass spheres; the spheres have no MR signal intensity associated with them and are identified as black voxels. Flow velocities in the (a) z-, (b) x- and (c) y-directions are shown with slices taken in the xy, yz and xz planes for each of the velocity components. For each image the positions at which the slices in the other two directions were taken are identified. Voxel resolution is 195 µm × 195 µm × 195 µm. The glass spheres were of diameter 5 mm and were packed within a column of internal diameter 46 mm. Reprinted from Ref. [38], Copyright (2001), with permission from Elsevier.

Fig. 4

slow-moving liquid – channeling does not occur just at the walls of the bed. The heterogeneity in the flow field within the reactor will also cause local variations in the flow profile between individual catalyst pellets. In turn, this will cause significant variations in the contact time and hence mass transfer between reactants and products between the inter- and intra-particle void space throughout the reactor. MR data of this type have direct use in identifying how catalyst particle size, shape and method of loading into the reactor influence heterogeneities in hydrodynamics [42–44]. The ability to map the flow field and the internal structure of the bed also allows the study of the structural characteristics of the bed that promote or reduce fines production and deposition [45]. An area of increasing interest in reactor design is the understanding of unsteady-state operation, and this is a topic that will be addressed later in more detail in the context of two-phase flow in fixed-bed (or trickle-bed) reactors. In the case of single-phase flow in fixed beds, there remains a significant challenge in designing, for example, narrow fixed beds typical of those required for processes requiring high rates of heat or mass transfer [46], exemplified by the steam-cracking process for conversion of long-chain hydrocarbons. Greater rates of heat and mass transfer can be attained by using turbulent flows, but the associated flow fields are difficult to predict. Computational fluid dynamics (CFD) codes are being developed [47], but a technique for imaging these flow fields is required to validate the CFD results. To use MRI to study unsteady-state phenomena, new ultra-fast techniques have been developed. All the early

imaging studies of flow in fixed-bed reactors discussed previously focused on reactors operating under conditions of relatively low Reynolds number (Re < 200), where Re is defined as (ρdv/µ) where ρ, v and µ are the density, superficial velocity and dynamic viscosity, respectively, of the liquid and d is a characteristic length scale in the system. This was done because under these conditions the flow field was independent of time. The study of steady-state systems was essential because at this time the only pulse sequences used to map flow in reactors were those based on the type of pulse sequence shown in Fig. 1 integrated with the transport measurement sequence shown in Fig. 2. Such pulse sequences took several minutes to acquire a 2-D image of a particular component of a flow vector (vx , vy or vz ) within the reactor. To study unsteady-state flow, the imaging time must be substantially decreased from the time-scale of minutes to tens of milliseconds. This has been achieved in the development of the gradient echo rapid velocity and acceleration imaging sequence (GERVAIS) [48], which is based on a variant of the EPI sampling strategy and allows acquisition of images of three orthogonal velocity components from a single excitation over a time-scale of 60 ms, with each velocity component itself being acquired in 99 β + 100 β + 99.9 β + 97 β − 100 β − 100 β − 100 β − 100

Max. energy/MeV

Max. linear range in Al/mm

Max. specific activity/ Bq mol−1

0.96 1.19 1.72 0.635 0.0186 0.155 1.71 0.167

1.4 1.9 3.0 0.9 0.002 0.1 30 0.1

3.4 × 1020 7.0 × 1020 3.4 × 1021 6.3 × 1019 1.1 × 1015 2.3 × 1012 3.4 × 1017 5.5 × 1016

6.5.3 Production of Labeled Material

chemical purification. As the shortest-lived of these four radionuclides lives for more than 2 weeks, there will be ample time to carry out the synthesis and to deliver the labeled compound for experiments at a radionuclide laboratory. As the range of β − -particles is limited, protection against radiation hazards is easily achieved by shielding the radioactivity with layers of Perspex. In general 11 C, 13 N, 15 O and 18 F are produced by deuteron (8 MeV) or proton (16 MeV) bombardment of suitable target materials, in which the labeled precursor molecules will be formed instantaneously. The labeled precursor compound [5] is utilized either for ‘‘direct’’ conversion into the required chemical compound or for conversion into another labeled precursor compound to be used as a starting compound for chemical (micro-)synthesis (Table 2). With respect to the half-life of the 11 C, 13 N, 15 O and 18 F radionuclides, synthesis and purification steps should be fast and experiments with labeled compounds have to be carried out near the site where the radionuclides were produced. Next to lead shielding (hot cells and dedicated shielding) when handling β + -emitters, the synthesis and experimental procedures should as far as possible be remotely controlled for protection against radiation hazards. Through the emergence of positron emission computed tomography (PET) during recent decades as a medical imaging tool [6], the general reputation of 11 C, 13 N, 15 O and 18 F has become more widespread. For medical PET applications even dedicated self-shielding cyclotrons have been developed to allow medical research and diagnostic studies [6, 7]; however, an existing nuclear accelerator facility at a university or nuclear research institute generally may be used for 11 C, 13 N, 15 O and/or 18 F production. The successful application of 11 C, 13 N and/or 15 O to heterogeneous catalysis requires close interdisciplinary cooperation between radionuclide production experts, radiochemical synthesis experts, radioanalytical experts and experts in the discipline of application. The production and synthesis methods of some molecules of interest to catalysis is described below. Tab. 2

13 N 15 O 18 F

11 CO, 11 CO , 11 CH C H 2 3 5 11

Carbon-11 is synthesized by irradiating a target of gaseous nitrogen with 12-MeV protons, which results in the 14 N(p,α)11 C nuclear reaction. Oxygen impurities in the target gas are sufficient to oxidize completely all of the 11 C, and 11 CO2 results. This labeled carbon dioxide can be reduced to 11 CO over zinc at 650 K. Labeled alkanes have been produced in a two-step procedure [8]. Labeled CO is used in a two-step alkene homologation reaction previously developed at Eindhoven University of Technology [9]. 11 CO is pulsed over a vanadium-promoted Ru/SiO2 catalyst at 620 K. After quenching the temperature to 380 K, 1-pentene is fed to the catalyst followed by desorptive hydrogenation. As homologation of shorter alkanes, resulting from cracking of 1-pentene over Ru, also occurs, a range of alkanes from C1 to C6 are produced. The hydrocarbons are separated by a process of freezing, flash-heating and gas chromatography. The desired product (∼1–3 MBq or 10−15 mol of 11 CH3 C5 H11 ) is then frozen out ready for use in the experiments. When desired, one of the other labeled alkane fractions can also be isolated. The radiolabeled fraction is minute compared with the amount of non-labeled n-hexane that is produced as the total amount of the injected pulse is typically on the order of 10−6 mol. The production process (that is, the adsorption of 11 CO on Ru/SiO2 and the separation) is monitored by NaI scintillation detectors. The process has been optimized such that pulses of radiolabeled n-hexane can be prepared every 45 min. In addition to labeled n-hexane, labeled n-pentane and methylpentanes have also been produced by this method. 6.5.3.2

13 NO, 13 NH

3

Nitrogen-13 is produced by irradiating a water target with a beam of 16-MeV protons. During irradiation the 16 O(p,α)13 N nuclear reaction occurs. The 13 N-labeled species typically exist as nitrates (more than 85%) and nitrites [10] under the conditions in the aqueous target. References see page 1810

Precursor molecules available immediately or with limited synthesis ( 1985: autocat. Europe

Price/$ g−1

25 -> 1975 autocat. US

20 15 10 5 0 1900

1910

-> 1920s Pt-catalyst nitric acid

1920

-> mid1990s Pt-jewellery China 1988: high Ptinvestment Japan

1930

1940

-> 1985 Pd in MLCC

Mid1950s Pt-reforming cat.

1950

1960

-> 2004 Pdjewellery China -> mid1980s: Pt for LCD glass 1970

1980

1990

2000

Long-term price development for Pt and Pd and milestones in their applications.

Fig. 1

Net demand

Net-demand

300

250 Palladium

250

200

150

Autocat.-net Others Jewellery Electronics Dental

100

Platinum

100

0

0 1980 1982 1984 1986 1988 1990 1992 1994 1996 1998 2000 2002 2004 2006

1980 1982 1984 1986 1988 1990 1992 1994 1996 1998 2000 2002 2004 2006

Development of net demand for Pt and Pd by application (data based on Ref. [1]).

far less vulnerable to such price changes, since only for the small lifecycle losses (and demand for new capacities) does additional Pt have to be purchased at market prices. Recycling Technologies The different types of catalysts require individual treatment and refining processes. The greatest impact comes from the catalyst support, but the combinations and absolute loadings of PMs and other metals involved (Re, Sn, Ge, Pb, Co, Ni, etc.) are also important. Moreover, the spent catalyst often is contaminated with carbon and hydrocarbons from the catalytic process (e.g. highly 7.2.1.3

150

Autocat.-net Others Petroleum Jewellery Investment Glass Electronics Chemical

50

50

Fig. 2

Amount / t yr−1

200

Amount / t yr−1

-> 2000 Europe diesel boom

-> late1960s Pt-jewellery Japan

1900: Pd /C-cat. (fine chemistry)

1900–1930 Pt-cat. H2SO4

1849

coked ‘‘heel catalyst’’ from CCR reactors), it can contain hazardous elements from the crude oil (As, Hg, etc.), halogens (Cl, F, etc.), such as found in isomerization processes, or Fe, Ni and Cr from in process corrosion of reactor walls and tubes. All these factors can significantly impact the refining process [11, 12]. Before the actual PM recovery process can take place, conditioning of the recycling-material is necessary in many cases. Examples for such a pretreatment are the decanning of car catalysts (extracting the monolith from the steel case) and the burning off of oil refining References see page 1863

1850

7.2 Recycling of Spent Catalysts

Pt/Pd: London fixing, Rh:JM-base 26.01.01: $1094 troy oz−1

1000 Palladium 800 600

Price/ $ troy oz−1

400 200

19.05.06: $1322 troy oz−1

1200 Platinum

21.11.06: $1390 troy oz−1

1000 800 600 400 200

81 82 83 84 85 86 87 88 89 90 91 92 93 94 95 96 97 98 99 00 01 02 03 04 05 06 6000 22.5.06: 5000 $6275 troy oz−1 Rhodium 1 troy oz = 31.1035 g 4000 3000 $100/troy oz−1 = $3215 kg−1 2000 1000 0 90 91 92 93 94 95 96 97 98 99 00 01 02 03 04 05 06 JM Base Price(monthly average)

Price/$ troy oz−1

600 500 Iridium 400 300 200

Ruthenium

100 0 96 96 97 97 98 9899 99 00 00 01 01 02 02 03 03 0404 05 05 06 06

Fig. 3

Price development for Pt, Pd and Rh (1981–2006) and Ir and Ru (1996–2006).

catalysts contaminated with carbon. This first step is often not carried out by the refiner itself, but by other companies. However, for some pyrometallurgical refining processes (see below) the preburning of carbon can be waived. The metals recovery process itself can be broken down into four basic steps: homogenization/sampling, preconcentration, dissolution and PM isolation and finally PM purification (refining). First the spent catalyst is prepared and homogenized, in order to obtain a representative sample for the purpose of incoming analysis and settling accounts with suppliers.

With supported catalysts on extrudates or pellets this is mostly done through screening and/or blending of the material and separation of a raw sample via a rotary divider or similar equipment from the continuous material stream. In the case of ceramic car catalysts or agglomerated oil-refining catalysts, the material first has to be milled, and sometimes also drying is required. High-grade PGM alloys such as catalytic gauzes are often dissolved in aqua regia or HCl–Cl2 and the sample is drawn from the stirred solution. The accuracy of sampling and assaying of the spent PM-catalyst is a decisive factor for the monetary results that a supplier

7.2.1 Recycling of Spent Catalysts Containing Precious Metals

Low & mediumgrade Upgrading of PM content (“furnace” or wet-chemical)

Separation of non-precious metals Non-PM (+PM)

Dissolving and removal of non-PM elements

Separation of PM from one another

Purification of PM to high grade sponge/powder

Separation of PGM (raw salts) Rh Ir Ru

Aftertreatment of mother liquids with PGM residues

(a)

Fig. 4

Aftertreatment of sidestreams out of purification

Separation of Ag, Au Ag/Au

(b)

Multistep purification of raw salts

Pt Pd Pt R

s aw

alt

s

Pd Fin e m Rh eta ls

Processes applied - Precipitation - Reduction / oxidation - Hydrolysis

PM containing sidestreams

PM-concentrate refining (wet-chemical)

Highgrade

Backfeed into preliminary stage

Dissolving Residue cont.PM

Preparation/homogenization sampling

1851

Ir Ru

- (solvent)-extraction - Ion exchange - Calcination

(a) Main steps in PM refining and (b) sequence of PM concentrate processing [5].

achieves when supplying material for precious metals refining. Here, some basic requirements have to be fulfilled: • Only well homogenized material gives a good sample. • The exact determination of weights, moisture content and volatile elements is as important as the PM analysis itself. • Sampling should be conducted as early as possible in the refining chain; any treatment before sampling bears the risk of undetected PM losses. It needs sound technical know-how and experience and high ethical standards of everyone involved, which unfortunately is not the case everywhere [13]. Where the PM content is above about 30%, preconcentration is generally not necessary and the material can be directly channeled into dissolution and PM isolation. In most cases, however, the content is lower, so that preconcentration is necessary. This can be done by incineration, pyrometallurgical smelting or hydrometallurgical dissolution of the matrix or selective leaching of the precious metals. The choice of the optimum method depends on the composition of the material and on the specific capabilities of the refiner and has a significant effect on recycling yields and environmental performance. After preconcentration, concentrates generally have a >30% content of PMs. These are dissolved and the individual PMs are then separated from each other and refined. The isolation of PGMs is complicated and expensive and reference is made in this particular respect to specialist publications [14]. To achieve a high PM yield it is necessary to re-treat carefully the side-streams that occur in these processes and to recover any traces of

PMs present. Figure 4 displays the principal steps in PM refining and the basic sequence of concentrate processing. The following are examples of recovery technologies for important solid catalysts. PGM catalysts on activated carbon supports contain typically 5% Pd (referring to dry weight), but also Pt, Pd/Pt and Ru catalysts are in use, with PM loading ranging from 0.5 to 10%. The spent catalyst usually contains >50% moisture and significant amounts of pollutants can also be present. Spent catalysts are carefully incinerated in special furnaces; the resulting ash contains PGMs in high concentrations [5, 15]. Originally, the incineration process was used at the same time for homogenization and the sample was drawn from the ash, thus accepting a small PM loss in the offgas. In some new processes, however, the sample is taken from the suspended catalyst before incineration and the suspension is then incinerated in a continuously operating, newly developed furnace or added to versatile pyrometallurgical processes. For PTA catalysts (Pd on coked coconut shell granules) or graphite support, both of which incinerate very badly, a sampling from the well prepared original material bears clear advantages and usually delivers much better results than after incineration (a 2005 blind test, where the same batch of several tons of spent PTA catalyst was weighed and sampled twice, demonstrated 11% more Pd returns from the sampling before incineration compared with weighing and sampling after incineration at another refiner for the same material) [11]. Ceramic monoliths of car catalysts and support catalysts based on α-Al2 O3 , zeolite or SiO2 are treated References see page 1863

1852

7.2 Recycling of Spent Catalysts

pyrometallurgically by reducingly melting in special furnaces. The ceramic support is slagged, the PMs are collected, according to the process, in Cu, Fe or Ni, and the collector metal containing PMs is then subject to further wet-chemical processing. Several types of furnace processes are used: some specialized refiners operate electric arc or plasma furnaces with the addition of CaO at temperatures of around 1600 ◦ C. With proper design and control of the furnace, good yields of Rh, Ir and Ru can be achieved. With high SiO2 shares in the matrix or corresponding blending with materials that melt at lower temperatures, such material can also be treated in gasheated rotary furnaces or converters, where the precious metals are mostly collected in Cu. The disadvantage with these processes is that large proportions of Ru are lost through volatilization, while Ir and Rh become trapped in slag, copper matte and the furnace lining. Finally, some copper and nickel smelters channel such sweeps into existing systems for the processing of Cu/Ni primary and secondary concentrates, with PMs accumulating as residue in copper and nickel refining. The same problems basically occur here with Rh, Ir and Ru, PMs also dispersing to a certain extent in the different metal paths (Cu, Ni, Pb, etc.). This can have an adverse effect on throughput times and PM yields. Modern ‘‘integrated smelters and refineries’’, such as the process newly installed at the end of the 1990s by Umicore Precious Metals Refining in Hoboken, Belgium, combine the advantages of base metal smelters with those of specialized precious metal refiners. Through skilled process management and aftertreatment of primary slag and other side-streams (Fig. 5) as well as high throughput, favorable production costs, enhanced flexibility and relative insensitivity to contamination can be achieved together with PM yields that are comparable to those from special processes. In addition, non-ferrous metals such as Cu, Ni, Pb, Sn, As, Se, Te, Bi and Sb can be recovered. In all pyrometallurgical processes Re is lost (volatilized) in the offgas stream. Typical catalysts treated in pyrometallurgical operations are car catalysts, catalytic hydrocrackers (in two-step hydrocracking processes a 0.2–0.7% Pd on zeolite catalyst can be used in the second step), lube oil catalysts (typically 0.5% Pd on zeolite), VAM catalysts (Pd/Au or Pd/Cd on α-alumina), hydrogen peroxide (sintered Al2 O3 tubes with Pt coating from the BMA/Degussa process), nitric acid catalysts (Russian process with 0.5–2% Pd on α-alumina support) or contaminated reforming (mostly CCR) or isomerization catalysts. The silicate slag from pyrometallurgical operations can usually be used as a construction material, e.g. for harbor or dyke fortification or for roadworks. Several attempts have been made to treat ceramicsupported autocatalysts and catalysts on an α-alumina support hydrometallurgically, mainly by leaching the

Sulfuric acid plant Matte Process Lead blast furnace

Gas SO2

ISA-Smelter Leaching & electrowinning

Cu bullion

Lead bullion

Slag

Aggregate for concrete

Umicore plant in olen Speiss

Ni, As

Lead refinery

Precious metals residues

Special metals refinery

Cupellation

PM-refinery

H2SO4

Cu

Ag, Au, Pt, Pd, Rh, Ru, Ir

Pb, Sn, In, Se, Te Sb, Bi

Basic procedures in an integrated metals smelter and refinery (example Umicore Precious Metals, Hoboken/Antwerp, Belgium).

Fig. 5

washcoat with the PGM from the substrate. On a production scale, none of these processes has been commercially successful so far. The main problems are the insufficient yields that could be achieved for recovering the PGMs, the treatment/disposal of the stripped leaching solution and the high abrasiveness of finely milled ceramic – necessary to create a high surface area for the leaching reagent – in piping systems and sealings. In some car segments, automotive catalysts with a metallic support are also used; here a thin stainlesssteel foil is coated with the PGM-containing washcoat. Due to the high iron content of such catalysts, a direct pyrometallurgical treatment is not feasible. Also for sampling purposes, an initial separation between steel foil and washcoat is required. This can be achieved using a special shredding process that mechanically removes the washcoat from the support, wherein special attention is required to prevent encapsulation of washcoat in the shredded foil. After this separation, the washcoat is blended and sampled and then treated pyrometallurgically like ceramic-supported autocatalysts. Reforming or isomerization catalysts containing Pt, Pt/Ir and Pt/Re or other catalysts on a γ -alumina support usually are treated in hydrometallurgical processes, since the support can be dissolved in special reactors by means

7.2.1 Recycling of Spent Catalysts Containing Precious Metals

of NaOH or H2 SO4 . The leach residue contains the PGMs (together with other insolubles) and is further treated. The support is transferred into a sodium aluminate or aluminum sulfate solution, which can be used in wastewater treatment. If the catalyst contains Re, this dissolves as perrhenic acid in the aluminate solution from where it can be removed using anion exchange [15]. Such a process, e.g. as run by Heraeus in Germany, is quick and very efficient, provided that the catalyst is sufficiently clean, and it also offers the possibility of recovering Re with yields up to 95%. However, it is only applicable for soluble supports (γ -alumina), and carbon, coke or hydrocarbons from the catalytic process have to be burnt off before, if a threshold value (usually 3–5%) is reached. This regeneration process often has to be subcontracted to third parties, which implies additional costs and process time and the risk of metal losses. Elements such as Pb, Ni, As and Hg can prevent the further use of aluminate solution and halogens can negatively impact the process. Thus, from a certain level of insolubles (e.g. phase transfer from γ - into α-alumina due to overheating in the reactor) or contamination, a pyrometallurgical process can become a viable alternative. In the case of VAM, two types of catalysts exist. In the case of Pd–Au on an α-alumina/silica support, the Pd and Au are leached from the support with aqua regia for further PM refining. In the case of Pd–Cd on an α-alumina/silica support, the Pd is chemically reduced to the metal state and Cd is washed from the surface. Subsequently, Pd is leached with aqua regia. For both processes the support often still carries residues of Pd and Au after the process and has to be fed into a scavenger pyrometallurgical process [15]. Alternatively, VAM catalysts are also treated directly in appropriate pyrometallurgical operations. Ethylene oxide catalysts (14% Ag on α-alumina) are also treated hydrometallurgically by leaching the Ag from the support with nitric acid. The silver nitrate solution is further processed into Ag; the support has to be washed intensely to mobilize all the Ag [15]. Gauze catalysts from nitric acid and hydrocyanic acid production consist of pure PGM alloys of Pt, Pt/Rh (90 : 10 or 95 : 5) or Pd for catchment gauzes. These are usually dissolved in aqua regia or HCl–Cl2 . After homogenization and sampling, the PGMs are then recovered from the solution. For new catalyst developments, the end of life recyclability should be one point of investigation, especially when new metal combinations or substrates are used. There are some recent examples of ‘‘exotic’’ metal–metal and metal–substrate combinations that do not fit into the existing routes and require the development of (in part) new refining processes. Challenges are the combined recovery of precious and non-ferrous metals

1853

(e.g. in GTL catalysts), new substrate types such as SiC (with a high melting point) for diesel particulate filters (DPFs) and process or environmental constraints deriving from other components/impurities. The catalytic layers of the membrane–electrode-assembly (MEA) in low-temperature fuel cells such as PEMFC, DMFC or PAFC contain, in addition to Pt (and Ru), high contents of fluoropolymer in a carbon matrix. This prohibits conventional thermal treatment in recycling because of toxic fluorine and hydrogen fluoride gas emissions. Spent GTL catalysts typically have a high wax content, which impacts sampling and smelting operations. Spent diesel particulate filters contain soot and mineral ashes which have to be handled carefully, especially during collection, milling and sampling. It is therefore recommended to contact an experienced precious metals refiner at an early stage of a new product development. Table 2 summarizes the main treatment and refining methods for PM catalysts. Responsible Care and Trans-Border Shipments For the selection of the optimum recovery process for a specific spent catalyst, in addition to technological and economical considerations, the environmental issues also cannot be neglected: 7.2.1.4

• What environmental, health and safety (EHS) risks arise in a certain recycling chain? • To what liabilities is the spent generator exposed? • How secure can the spent generator be about the technical, commercial and environmental performance of the selected catalyst refiner? The right balance between economical and ecological performance is crucial when recovering metals from spent catalysts. Recycling costs have to be considered with regard to their interdependencies with environmental and recycling efficiencies, where recycling efficiency can be used as a synonym for metal yields in the recovery process. For PM catalysts this has a direct monetary benefit for the catalyst generator. In addition to their valuable content of PMs, spent catalysts often contain a complex mix of different substances, some of them with toxic or hazardous characteristics. The recycling chain has to cope with those complex mixtures. The generator of a spent catalyst has to assure that all operators involved along the entire recycling chain are acting in an EHS-compliant way when treating the spent catalyst. This responsibility is related to individual regulations in specific countries, to the Basel liability protocol and to the principles of Responsible Care , to which the chemical and oil refining References see page 1863

1854 Tab. 2

7.2 Recycling of Spent Catalysts Overview of treatment and refining methods for precious metals catalysts

Substrate/catalyst type Activated carbon

Al2 O3 (γ -type)

Main sampling method Out of ash after incineration or from suspension of original catalyst Screening or blending

Zeolite, SiO2 , Al2 O3 (all types), ZrO2 , BaSO4 , CaCO3 Gauzes

Screening, blending, (milling)

Homogeneous

Out of mixed pulp or after incineration

Preconcentration

PGM yield/ % (after sampling step)

Incineration or, ‘‘smelting’’

PGM concentrate or collector metal with PGM

96–99

Dissolution of substrate by NaOH or H2 SO4 Pyrometallurgy (smelting)

PGM concentrate

98–99

Dissolution or melting

– Chemical collection or incineration or ‘‘smelting’’

industries have committed themselves. In the context of catalyst recycling, especially the following Responsible Care guidelines are of relevance: • to make health, safety, the environment and resource conservation critical considerations for all new and existing products and processes • to provide information on health or environmental risks and pursue protective measures for employees, the public and other key stakeholders • to work with customers, carriers, suppliers, distributors and concentrators to foster the safe use, transport and disposal of chemicals • to operate facilities in a manner that protects the environment and the health and safety of employees and the public. 7.2.1.4.1 Shipment of Spent Catalysts – What Has to Be Considered? Shipping of spent catalysts for recovery across borders falls under the trans-frontier movement of waste regulations, also often referred to as the ‘‘Basel Convention’’. The following information is based on Council Regulation 259/93 of 1 February 1993, which describes trans-frontier movements of waste within, into and out of the European Community [16, 17]. Similar procedures are regulated by the OECD [revision of decision C(92)39/final on the control of transboundary movements of wastes destined for recovery operations] [18].

A Classification as Green or Amber Listed Waste In principle, spent catalysts are classified as green listed waste under the entry GC 060: GC 060 Spent metal-bearing catalysts containing any of:

• precious metals: gold, silver

PGM separation and purification out of

Collector metal with PGM PGM solution or PGM alloy PGM concentrate or collector metal with PGM

94–98.5

98–99 85–98

• platinum-group metals: ruthenium, rhodium, palladium, osmium, iridium, platinum • .... However, the trans-frontier movement of green listed wastes may be amber-list controlled if they are contaminated by other materials to an extent which (a) increases the risks associated with the waste sufficiently to render it appropriate for inclusion in the amber or red lists or (b) prevents the recovery of the waste in an environmentally sound manner. More information should be obtained from the local competent authorities. Some of these hazardous characteristics are listed in Annex III of Council Directive 91/689/EEC [16, 17]. When considered as amber listed waste, spent catalysts should be classified under entry AB 080: AB 080 Spent catalysts not on the green list This results in two possible EURAL codes (European Waste List): • 16 08 01 Spent catalysts containing gold, silver, rhenium, rhodium, palladium, iridium or platinum (except 16 08 07) • 16 08 07* Spent catalysts contaminated with dangerous substances. Combining the entries on green and amber list on the one hand and Eural codes on the other results in: • 16 08 01 GC 020 green list procedure (see below) • 16 08 07* AB 080 amber list procedure (see below). B Green List Procedure Transport of green listed catalysts does not require a notification procedure to the competent authorities of the countries of despatch,

7.2.1 Recycling of Spent Catalysts Containing Precious Metals

transit and destination. Some documents signed by the holder should, however, accompany the shipment of the green listed waste and contain the following information: • the name and address of the holder • the usual commercial description of the waste + entry on the green list + Eural-code • the name and the address of the consignee • the operations involved in recovery, as listed in Annex II B to Directive 75/442/EEC (in most of the cases this will be code R4) • the anticipated date of shipment. C Amber List Procedure The amber list procedure implies a notification of the intended shipment to the competent authorities of countries of dispatch, transit and destination. These authorities must provide consent prior to commencement of shipment. Notification is done by means of a notification form. Each shipment should be accompanied by a movement tracking form and by a copy of the consent. The procedure here is therefore more complex. In the case that the amber listing applies for a recycling material, an experienced catalyst recycler usually can advise and assist their customers to do it in the appropriate way. For documentation within the amber list procedure, standardized notification and movement tracking forms are available. Among others, decisive factors for the decision of the authorities are the source, composition and quantity of the waste for recovery, a description of the process for recovery and the amount of recycled material in relation to the residual waste. Summarizing, the main difference between green and amber listed materials is that in the green category, only information to the relevant authorities is needed, whereas for trans-border movement of wastes regarded as amber also the consent of the authority is required. For trans-border shipping from outside the OECD, in principal similar procedures apply, but other document forms might be needed. Details must be checked on an individual case basis with the relevant authorities [19]. Commercial Execution of a Refining Job Spent PM catalysts are usually recycled under a so-called ‘‘toll refining’’ contract. The PM refiner offers the service of recovering the metals from the catalyst under agreed contract terms and credits them back to the supplier of the spent catalyst. In the case of chemical and oil refining catalysts, where the supplier/owner usually replaces the spent by a fresh catalyst to keep the process running, this metal credit is in most cases given on a weight base; this means either physically as fine metals returned to the user or an appointed catalyst manufacturer or in the 7.2.1.5

1855

form of crediting the metals on a specific customer weight account (comparable to a currency bank account), from where it can be withdrawn or transferred on demand. The advantage of such a weight-based procedure is that the PMs remain the property of the user at any time and that fluctuating PM prices have no direct impact on the catalyst user except for the small lifecycle losses that need to be filled up with purchases of additional metals. Volatile metal prices only impact the asset value of the user’s metals inventory. A time span of several months is needed between discharging and shipping the spent catalyst, receiving the credit from refining and having the fresh catalyst manufactured and installed in the reactor again. This span can be either covered by a corresponding stock of additional PMs at the user (which can mean a significant amount of capital employed) or by a ‘‘bridge lease’’ for the metals, which the refiner or another source provides to the catalyst user against payment of a lease rate. The most optimum way of such ‘‘precious metals management’’ has to be determined on a case-by-case basis between the user and the refiner and/or the manufacturer of the catalyst. A toll-refining contract has three main parameters: • A metal credit in percent of the analytical content of the spent catalyst, which is determined in the sampling process. This metal credit reflects the technical process losses during refining. • A metals return time (until the metal credit takes place), usually in weeks after the end of the sampling process. This reflects the process throughput time to recover and purify the precious metals. • A refining price, usually comprising treatment charges (per kilogram of spent catalyst received), refining charges (per kilogram of metal credited) and handling and assay charges (per lot sampled and analyzed). In some cases, additional costs can be charged for handling impurities (‘‘penalties’’). The refining price reflects the total recovery costs. All these parameters can be transferred into monetary numbers and summed up in a ‘‘bottom-line’’ calculation to make different refining offers with different parameters comparable. For such a cost comparison, all costs involved have to be considered, including transport and conditioning (e.g. burning off carbon at a specialized subcontractor, if not avoidable). It also becomes obvious why the accuracy of sampling is so important: any mistake here has a direct monetary impact (more or less metals credited back to the catalyst owner), and the same applies for any undetected metal/catalyst losses before the sampling process (any deviation from the true value in References see page 1863

1856

7.2 Recycling of Spent Catalysts

weight, moisture content or PM analysis can easily turn a virtual ‘‘best offer on the paper’’ into a real bottom-line disadvantage of several tens of thousands of dollars) [13]. Evaluating the true bottom-line profitability of a spent catalyst recycling job requires experience and a complex calculation. Compared with the intrinsic metal value of a spent PM catalyst, especially for Pt spent, the reclamation costs in most cases are only a small fraction. Taking risks or uncertainties in the performance of the recycling chain could easily offset the recycling cost. Hence a bottomline cost calculation has to be balanced against general expertise, professionalism and financial soundness of the catalyst refiner and also against transparency and EHS compliance along the entire recycling chain. Modern integrated PM smelter and refinery plants can benefit from economies of scale. In the total feed of such a plant, spent catalysts make up only a small proportion and other feed materials are more complex in their composition. The processes are designed to cope with very challenging recycling materials and comprise all the necessary environmental installations for materials that are much more ‘‘nasty’’ in their composition than spent catalysts. In total, this provides a very protective recycling environment. Although, as shown before, total EHS-related investments are very high, spent catalyst generators can benefit from the cost distribution over a large quantity of input recycling materials. Thus, cost-efficient recycling solutions can be offered which simultaneously do not expose the generator of the spent catalyst to any environmental risk. It is advisable that the generator of a spent catalyst at an early stage enters into commercial and technical discussions with the catalyst refiner to be sure about the overall performance and to develop mutually the optimum approach to how a specific refining material can be treated. In Table 3, a checklist is given to facilitate the execution of a refining job, and Fig. 6 illustrates the parameters needed for a bottom-line calculation. In the case of automotive catalyst recycling, in principal the same applies as said above. For car manufacturers with a permanent need for PMs for new catalysts, weight-based loop business makes sense and is in many cases current practice. In the case of catalysts from end-of-life cars collected by specialized companies or scrap dealers, lot sizes from more than 1 t catalyst weight (after decanning) are usually also refined under a toll refining contract; however, here the supplier often sells the credited metals to the refiner at current market prices (this is possible also for chemical and oil refining catalysts). Smaller quantities and undecanned catalytic converters are often sold on a ‘‘telquel price’’. Such an ‘‘as-is’’ price per piece depends on the size and type of the catalyst and on the PGM price development. Taking into account today’s large variety of

catalyst types and loadings, telquel prices bear significant risks for the parties involved (see Section 7.2.1.6). Catalyst Lifecycles and Future Recycling Potential An in-depth analysis of the lifecycle efficiencies for PM catalysts was performed in a research project conducted jointly by Umicore Precious Metals Refining and the ¨ German environmental research institute Oko-Institut between 2001 and 2004. The project covers all application segments for PGMs, among which catalysts play an important role. Focus and system boundaries of the project were Germany but the global market environment for PGM products was considered. The project was cofunded by the German Federal Ministry of Education and Research (BMBF). Entitled ‘‘Materials Flow of Platinum Group Metals – System Analysis and Measures for a Sustainable Optimization’’, an English-language version of the report was also produced in collaboration with the London-based metals analyst GFMS, which also put the German study into an international context [4, 5]. Most of the findings achieved in Germany should, in principle, be valid also for other industrialized countries. For Germany, the following results were found (for 2001) with respect to catalysts [4, 5]: 7.2.1.6

• Autocatalyst production accounted for 40% of total PGM gross demand, industrial catalysts [all PM catalysts (including homogeneous) except automotive catalysts] 22%, glass industry products 13% and jewellery 9%. • Solid industrial catalysts accounted for a PGM inventory of 10 t and an annual gross demand of 6.4 t, of which 6 t could be covered from recycling of spent catalysts. Only the gap of 400 kg (6%) had to be supplied by ‘‘new’’ metal from the international markets (= net demand) (Fig. 7). • PGMs recycled from scrapped autocatalysts represented only 12% (static) of the gross autocatalyst demand of 15.4 t yr−1 . The total PGM inventory in autocatalysts was 102 t, about 25% of the global annual PGM mine production. • In 2001, 10 t of PGMs escaped recycling and were considered ‘‘lost’’. About 50% of this was Pt, with Pd comprising 40% and rhodium 10%. Autocatalysts make up for over 30% of these PGM losses. The analysis showed substantial differences between application fields, ranging from PGM lifecycle efficiencies (dynamic recycling ratios) of >90% on the one hand (e.g. chemical and oil refining catalysis) to icrit electrode B is more suitable. Fig. 6

8.1.1 Fundamentals of Electrocatalysis

Second, for small values of η, the exponential can be expanded in a Taylor series up to the linear term, which yields a linear dependence of current on overpotential: i=

i0 F η RT

Rct =

RT i0 F

(28)

Hence i0 can be in principle accessed from the slope of a measured i –η curve in the range |η| ≤ 5 mV, provided that the reaction proceeds in a single step. Third, for large values of η, i.e., |η| ≥ 50 mV, one of the terms in brackets can be neglected and an exponential current–potential characteristic is obtained, yielding for large negative η i = i0 e−αF η/RT

(29a)

and for large positive η i = i0 e(1−α)F η/RT

(29b)

Thus, a plot of log i vs. η, known as a Tafel plot, yields a straight line: η = a + b log i

becomes important the i − η relation obeys

   i i e(1−α)F η/RT − 1 − e−αF η/RT i = i0 1 − il,a il,c (31)

(27)

The resulting equation has the same form as Ohm’s law, which led to the definition of the charge transfer resistance Rct :

(30)

Evaluation of the slope b allows the determination of α and interpolation of the linear segments to the equilibrium potential, i.e. η = 0, yields i0 . Tafel plots are a very useful diagnostic tool to determine kinetic parameters. They are also helpful for reactions that do not proceed in a single step; then, however, the interpretation of the slope and the intercept changes – see Section 8.1.1.3.2. On increasing the overpotential in the positive or negative direction, a corresponding exponential increase in current density as predicted by the Butler–Volmer equation can only take place as long as the transport of educts to and products from the electrode is sufficiently fast such that the surface concentrations are identical to the bulk concentrations. It is obvious that sooner or later in any system the mass transport rate will come in the range of the particle fluxes owing to the electrochemical reaction when increasing |η|. Then, we have a mixed control of mass transport and electrode kinetics. At very large |η|, finally, the current is totally determined by the mass transfer rate and a limiting current independent of η adjusts. It can be shown that when mass transport

1881

where il,a and il,c are the anodic and cathodic limiting current densities [40]. Again, Eq. (31) can be linearized for small η, which leads to 
1 1 RT 1 − + η= i (32) F i0 il,c il,a Corresponding to the definition of the charge transfer resistance above, we can define an anodic and cathodic mass transfer resistance and rewrite Eq. (32) as follows: η = i(Rct + Rmt,a + Rmt,c )

(33)

Clearly, if Rct  Rmt,a(c) , that is, i0 il,a(c) , then the current is dominated by the mass transfer resistance and the corresponding overpotential is determined by the difference between the surface and the bulk concentrations and termed concentration overpotential ηconc . In the opposite case, i.e. when Rct Rmt , the overpotential is dominated by the activation of charge transfer and is referred to as activation or charge transfer overpotential ηct [see also Eq. (17)]. For large overpotentials, one of the two terms in brackets in Eq. (31) can again be neglected and plots of log[(il,c − i)/i] (which is usually referred to as mass transport-corrected current) vs. η yield straight lines. From their slopes and intercepts the kinetic parameters α and i0 can be evaluated even in the presence of mass transfer. 8.1.1.3.2 Multistep Reactions So far, we have considered reactions that involve just one electron that is transferred in the overall reaction and assumed that the mechanism constitutes of only a single step. Many electrochemical reactions require the transfer of more than one electron for the overall reaction to occur once. In addition, the reaction mechanism often contains also purely chemical steps. Thus, a general electrochemical reaction in which n electrons are transferred will proceed in at least n electrochemical steps, which might be coupled to m chemical steps. The succession of electrochemical and chemical elementary steps determines the reaction mechanism. Especially electrocatalytic reactions that involve the adsorption of an intermediate at the electrode surface are multistep reactions. A general theory of current–potential relations for multistep reactions would need to take into account all potential dependences and References see page 1902

1882

8.1 Electrocatalysis

surface concentrations of the intermediates in addition to the bulk concentrations of educts and products, many of which will not be experimentally available. Fortunately, in most experimental situations the problem can be significantly simplified because one of the steps is much more sluggish than all the others. The current is then determined by the rate of this slowest step, which we shall abbreviate as RDS (rate-determining step). As a rule of thumb, a step qualifies as RDS if it is at least 10 times slower than all the other reactions. In the following we assume that an RDS exists. Let us consider the overall reaction −−  O + ne −− − −R

(34)

A general mechanism will contain n electron transfer reactions preceding the RDS, possibly intermixed with m chemical steps, one RDS that might be of electrochemical or chemical nature and n (with 0 ≤ n ≤ n − n ) charge transfer steps in addition to m chemical steps following the RDS. It is thus of the general character  −−  O + n e −− − −O 

kbRDS

(35)

where r = 0 if the RDS is of chemical nature and 1 if it is a charge transfer step. Obviously, n = n + n + r. The existence of an RDS indicates that all the other steps are virtually at equilibrium. In the case of an electrochemical RDS, its current density is thus given by s −αF η/RT − constant b iRDS = constant f × cO e

n − n n − → + rα and ← γ− = − rα γ = ν ν

(39)

(40)

and ν is the stoichiometric number of the RDS. This determines how many times the RDS takes place for the overall reaction to occur once. To elucidate this definition, consider the overall reaction −  2A+ + 2e − −− − − A2

(41)

Let us assume that it takes place in two consecutive steps, the first of which is the RDS: −−  A+ + e −− − − Aad

(36)

where α is the transfer coefficient of the RDS and the constants depend on the respective rate constants k RDS , the Faraday constant and the standard electrode potential of the RDS. At a steady state, all sequential electron transfer steps proceed at the same rate and therefore the total current is given by i = niRDS

(38)

The so-called observable charge transfer coefficients − → γ and ← γ− of the forward and backward reactions, respectively, now depend, however, on the specific reaction mechanism. This makes generalized Tafel plots a useful tool for the validation of postulated reaction mechanism of multistep reactions in general and of electrocatalytic reactions in particular. Note that also the exchange current density i0 in Eq. (38) is not given by Eq. (25) any longer, but depends on the equilibrium constants of the steps in equilibrium, the bulk b or cb , the rate constants of the RDS concentrations cO  R and its transfer coefficient and the number of electrons transferred in the steps preceding the RDS, n . The above derivation of Eq. (38) leads also to explicit expressions for the observable transfer coefficients [56]:

n − → γ +← γ− = ν

 −−  O + re −− − −R

× cRs  e(1−α)F η/RT

− − →  ← i = i0 e γ F η/RT − e− γ F η/RT

where

kfRDS

−−  R + n e −− − −R

current–potential relation:

(37)

Since steps preceding and following the RDS are at s and cs can equilibrium, the surface concentrations cO  R b and be expressed in terms of the bulk concentration cO  cRb  , respectively, and the equilibrium constants of the individual reactions [56]. On inserting the corresponding expressions into Eq. (36), combining it with Eq. (37) and rearranging terms, one arrives at a Butler–Volmer-type

−−  2Aad −− − − A2

(42)

Then, clearly, the first step must take place twice for the overall reaction to occur once. Hence ν = 2. The diagnostic value of Tafel plots for the determination of the reaction mechanism of a multistep electron transfer reaction can now be appreciated. For any postulated reaction mechanism, the values of the observable charge transfer coefficients can be calculated (provided that α is known) and they can be extracted from the experimental data by analyzing the slopes of the linear portions of log i vs. η curves. Hence considerable experimental evidence for or against a certain reaction mechanism and in particular the RDS can be obtained. The determination of the latter, in turn, is among the most important tasks in electrocatalytic research, since any attempt to accelerate an electrocatalytic reaction is tantamount to making the RDS faster.

8.1.1 Fundamentals of Electrocatalysis

The above discussion should not give the impression that the reaction mechanism can always be obtained in the above-mentioned way. Often there are experimental constraints that make such a complete analysis impossible. For example, it might not be possible to eliminate the mass transport limitation (note that so far we have assumed that the current is entirely charge transfer controlled), forward and backward reactions can follow different mechanisms in the Tafel regimes, the slowest reaction might not be sufficiently separated in rate such that equilibrium conditions cannot be assumed for the remaining reactions, etc. Also, the same Tafel slopes may occur for different RDS (see Section 8.1.1.5.1). These aspects render multistep electron transfer kinetics a complicated topic. How one can deal with more complicated situations, in which the above analysis cannot be applied, is discussed in Ref. [40], where references to more specific elaborations are also given. Fundamentals of Electrocatalysis Electrocatalysis is characterized by the fact that widely different reaction rates are observed for different electrodes at the same electrode potential (i.e. equal electron energies) and also otherwise identical parameters (concentration, temperature, etc.). The primary origin of electrocatalysis is the adsorption of educts, intermediates or products on the electrode, the heats of adsorption depending on the chemical nature or the structure of the electrode. The signature of electrocatalysis is a ‘‘volcano-shaped’’ dependence of the reaction rate (or current) on the adsorption energy of key reaction intermediates, a so-called volcano curve. The rationalization of volcano curves in catalysis was first given by Sabatier in the early 1900s. The Sabatier principle states that an active catalyst should adsorb a key intermediate neither too weakly nor too strongly. Studies in heterogeneous catalysis and electrocatalysis demonstrated innumerous examples of volcano curves. In Section 8.1.1.4.1 we will discuss the origin of volcano curves from an electrochemical perspective. In addition to the bonding strength of intermediates, the structure of the double layer, which varies at identical electrode potential for different materials, affects the overall reaction rate. These effects are referred to as ‘‘secondary effects’’ and are reviewed in Section 8.1.1.4.2. 8.1.1.4

proceeds through an intermediate Y: k1

Primary Effects: Heat of Adsorption and Volcano The qualitative dependence of electrocatalytic Plots activity on the adsorption energy, which, independent of the reaction considered, follows a volcano curve, can be rationalized most easily when analyzing a simple example. Consider a one-electron reduction of a species O to R that

k2

−  O+e− −− − − Y −−−→ R

(43)

k−1

If Y adsorbs at the electrode, its standard Gibbs energy of formation and hence also the activation energy of the formation will be lowered. Evidently, as long as the formation of Y is the rate-determining step, the overall reaction rate increases as the adsorption energy increases since the adsorbed intermediate is stabilized by the electrode surface. However, the bond strength of Y to the electrode necessarily affects also the rate with which R is formed from Y: the latter reaction becomes slower as the bond strength becomes stronger. These contrary responses of formation and removal of Y with increasing heat of adsorption lead to an increase in the coverage of Y with increasing heat of adsorption. A high coverage, in turn, slows down the rate of formation of Y, which is proportional to the free surface sites, but increases the reaction rate of formation of R. These simple considerations already indicate that the production rate of R, and thus the electrocatalytic activity of the electrode, will be maximum at an intermediate value of the heat of adsorption. For the reaction scheme above, a quantitative relationship can be easily deduced, giving further insight into how the three partial reaction rates are affected by a change in the adsorption energy and how their interaction yields a volcano curve. Let us first consider the steady-state coverage that will establish for different ratios of the three rate constants. The rates of formation of Y and its removal due to oxidative desorption or consecutive reaction are given by νad = cO k1 (1 − θ) νdes = k−1 θ and

νreac = k2 θ

(44)

respectively, where θ is the fraction of surface sites which are covered by Y and cO , as before, the concentration of educt in the reaction plane. At the steady state dθ = νad − νdes − νreac = 0 dt

(45)

which yields θss =

8.1.1.4.1

1883

1 k−1 k2 1+ + k1 cO k1 cO

(46)

Equation (46) tells us that as long as the formation of Y remains much slower than both its reductive References see page 1902

1884

8.1 Electrocatalysis

desorption and consecutive reaction to R, i.e. as long as k1 cO  k−1 + k2 , the coverage remains negligible. In the opposite limit, i.e. if the adsorption is much faster than either desorption or reaction, i.e. k1 cO k−1 + k2 , the electrode surface is nearly fully covered with Y. As the adsorption strength increases, k1 increases and both k2 and k−1 decrease. Hence θ increases monotonically with increasing adsorption strength. To obtain a quantitative relation between the electrocatalytic activity, which depends on θss and the adsorption strength, we have to express the rate constants ki in terms of the Gibbs free energy of adsorption of Y, Gad , i.e. the change in the standard Gibbs energy due to the adsorption of Y on the electrode surface. Figure 7 illustrates how the activation energies of the three partial reactions depend on Gad . The dashed curve indicates the energetic situation if Y has no interaction with the surface. In this case the Gibbs activation energy is denoted by G#1 . According to the Brønsted–Evans–Polanyi rule, there is a linear relationship between the activation energy and the change in the free energy for an elementary reaction, i.e. G#1,ad

=

G#1

+ αGad

and G#2,ad = G#2 − (1 − β)Gad

The change in the activation energy upon adsorption is also illustrated in Fig. 7. Equations (49) and (50) allow us to choose the rate constants for the case when Y does not interact with the surface, ki , as convenient reference constants: −G#1,ad /RT

k1 = A1 e

G#−1,ad = G#1 − (1 − α)Gad

(48)

= A1 e−G1 /RT e−αGad /RT

k−1 =

 k−1 e(1−α)Gad /RT

k2 = k2 e(1−β)Gad /RT

(50) (51) (52)

Note that Gad < 0 and thus k1 increases and k−1 and k2 decrease with increasing adsorption strength, as required. To a first approximation, we can assume that α = β = 0.5. Inserting Eqs. (50)–(52) into Eq. (46) θss can be calculated as a function of Gad : θss =


1+

1  k2 eGad /RT + k1 cO k1 cO  k−1

(53)

The stationary current density follows directly from the rate laws Eq. (44):

i = iox − ired = F k1 eαGad /RT [1 − θss (Gad )]   e−αGad /RT θss (Gad ) (54) −k−1 Plots of ln i and of θss vs. Gad as calculated from Eqs. (53) and (54) are shown in Fig. 8. As expected,

a∆Gad Y ∆Gad

∆G1#

#

= k1 e−αGad /RT

(47)

Equally, the activation energies of desorption and reaction can be written as

(49)

1.0 0.8 0.6

q

Yad

In i

# ∆G1,ad

0.4 0.2

O+e

−∆Gad

Catalytic activity (solid curve) and coverage of intermediate Y (dashed curve) as a function of the standard Gibbs energy of adsorption of Y for reaction scheme (44). The curves were calculated  /k c + k /k c = 10 000, i.e. a with Eqs. (54) and (53) for k−1 1 O 2 1 O situation in which without adsorption the formation of Y is the rate-limiting step. (Note that shapes and relative positions of the curves do not depend on the quantitative values of the rate constants as long as without adsorption the formation of Y is rate determining.)

Fig. 8

R

Standard Gibbs energies for reaction scheme (44) for the case that the intermediate Y does not adsorb at the electrode (dashed curves) and for the case that Y adsorbs at the electrode (solid curves).

Fig. 7

0.0

8.1.1 Fundamentals of Electrocatalysis

the catalytic activity goes through a maximum with increasing adsorption strength whereas the coverage increases monotonically with |Gad | and exhibits a sigmoidal shape. Equations (53) and (54) imply that as long as θss is low, adsorption is rate determining whereas a high θss is found only if the reaction is rate determining. Hence the RDS necessarily changes with increasing adsorption energy, i.e. as we go through the maximum of the volcano curve. In this view, discussions of volcano curves that assume the same RDS over the entire range of Gad , as found fairly often in the electrochemical literature (e.g. in Chapter 10 in Ref. [1]) have to be treated with caution. At maximum current, θss = 0.5 and thus k1 cO − k−1 = k2

(55)

Note that for more complex mechanisms with more than one type of adsorbate, optimum coverage of key reaction intermediates may differ significantly from 0.5. Hence the best catalyst is the one for which the rate constants of ‘‘net adsorption’’ (i.e. the difference in the adsorption rate at zero coverage and the rate constant of desorption) is equal to the rate constant of the reaction. This suggests that the standard Gibbs free energies of formation of Y from O and formation of R from Y should be similar. So far, our discussion has neglected the influence of the electrode potential on the reaction rate. This dependence  that is implicitly included in the rate constants k1 or k−1 can be written as 



k1 = k10 e−αF (E−E and





0

)/RT

0 (1−α)F (E−E k−1 = k−1 e

0 )/RT

(56)

respectively. Hence the catalytic activity of different electrode materials should always be compared at equal electrode potentials. Owing to the different chemical potentials of electrons in different materials, the magnitude of surface charge and possibly even the sign of the surface charge will differ at equal E for different electrocatalysts. As a consequence, the double layer structure also differs, which in turn affects the reaction rate. Changes of the reaction rate that originate from a changed double-layer structure are called secondary effects and are discussed in the next section. One consequence of the different double-layer structures for different materials that is usually not considered when secondary effects are discussed is that the adsorption strength of an adsorbate is influenced by the surface charge, a fact that further complicates an analysis of catalyzed electrode reactions. Experimental evidence thereof is the Stark tuning effect observed in

1885

IR spectra, that is, variation of vibrational frequencies of adsorbates with the electrode potential [57]. Differences in catalytic activity of the same material in an electrochemical environment and in heterogeneous catalysis are therefore not only due to the fact that the educts and products are solved in an electrolyte, but also arise because of different surface charge densities. In electrocatalysis, the surface charge at the metal/electrolyte interface is to a large extent controlled by the electrode potential under the operating conditions and influenced by the adsorbed ions/molecules. In heterogeneous catalysis, the surface charge is entirely determined by the nature of the adsorbates. The connection between solid/liquid electrified and solid/gas interfaces has been discussed in relation to the so-called double layer modeling in ultra-high vacuum, which was pioneered by Sass and coworkers in the 1980s [58, 59]. Secondary Effects: Double Layer Structure At equal driving force for the reaction, different electrode materials develop different potential drops across the double layer (see also Section 8.1.1.2.2) and, furthermore, they exhibit a different tendency to chemisorb ions, which again, as discussed above, influences the doublelayer structure. An influence of the electrode material on the reaction rate may therefore even be observed in the absence of any interaction of a species involved in the overall reaction with the electrode surface. Such electrocatalytic effects are called ‘‘secondary effects’’ or, in recognition of Frumkin’s pioneering work, ‘‘Frumkin effects’’. Their origin lies in the different double-layer structures that establish at electrodes of different nature. Above we derived that the current density is proportional to the concentration of the educt in the reaction plane and depends exponentially on the potential: 8.1.1.4.2

|i| ∼ ce|E−E

0

|

(57)

In interpreting Eq. (57), we assumed (a) that if c is not identical with the bulk concentration, mass transfer comes into play, and (b) that the entire potential drop across the double layer occurs between the electrode and the reaction plane. Both assumptions are idealizations which are met best in electrolytes with high ionic strength at potentials far away from the pzc and in the absence of specific adsorption at the electrode surface. If conditions (a) and (b) are not met, the following effects come into play that alter the current density: 1. For charged species and electrolytes with medium or low conductivity, Coulomb forces exerted by the surface charge of the electrode on the charged reactants cause References see page 1902

1886

8.1 Electrocatalysis

a static concentration profile across the double layer with an increased or decreased concentration in the reaction plane with respect to the bulk concentration, depending on the signs of the charges on the electrode and the reacting ions. Let us denote the distance of closest approach from which the electron transfer occurs from the electrode surface by x2 , the potential in the reaction plane, i.e. at x2 by φ2 , and let us set the potential in the bulk electrolyte to zero, φ0 = 0. Then, the concentration of the reacting species at x2 is c(x2 ) = cb e−zF φ2 /RT

(58)

where z is the charge on the reacting species, for anions z < 0, for cations z > 0. 2. The lower the electrolyte conductivity, the greater is the extension of the double layer into the electrolyte. Hence only the fraction φ − φ2 of the entire potential drop across the double layer φ drops between electrode and x2 and thus also only a fraction of the change in electric energy of an electron between electrode and electrolyte, Fφ, is operating in the reaction plane. In the case of a reduction reaction the current is thus proportional to i ∼ e−(E−φ2 −E

0 )αF /RT

= e−φ2 αF /RT e−(E−E

0 )αF /RT

(59)

A corresponding expression holds for oxidation currents. Hence, compared with the ideal situation, the current density differs by a correction factor, which is called the Frumkin correction: i = iideal e−(α−z)F φ2 /RT

(60)

The Frumkin correction can be calculated if φ2 is known. In the absence of specific adsorption, φ2 can be determined readily from measurements of the charge density and the Gouy–Chapman–Stern theory [40]. However, a theory that would allow one to extract φ2 in the presence of specific adsorption, either of ions of the supporting electrolyte or of reacting species, is not available. This makes it difficult to predict the influence of double-layer effects in electrocatalytic reactions. Whenever possible, in an electrocatalytic investigations one will avoid them by using a large excess of supporting electrolyte. The Hydrogen Oxidation/Evolution Reaction The hydrogen evolution reaction (HER): 8.1.1.5

2H2 O + 2e −−−→ H2 + 2OH−

(61)

2H+ + 2e −−−→ H2

(62)

[Eq. (61) in alkaline and Eq. (62) in acidic medium] was the first electrochemical reaction ever investigated. It was mentioned in 1800 by Nicholson and Carlisle [60], who used a Volta pile to perform water electrolysis. For many years, the HER was studied as a model reaction in electrochemistry. For example, it provided a basis for establishing the famous phenomenological Tafel equation [61] relating the current to the electrode potential [see Section 8.1.1.3.1, Eq. (30)]. In 1930, Kobosev and Nekrasova first recognized the influence of the hydrogen adsorption energy on the rate of the HER [62]. It was again the HER which inspired Kobosev and Monblanova to coin the term ‘‘electrocatalysis’’ [63]. The HER is important for water electrolysis and it also occurs as a side-reaction for many cathodic processes (hydrogenation, etc.). The reverse reaction – the oxidation of molecular hydrogen (HOR) – has been studied less extensively due to the interference of mass transport. However, during recent decades the HOR has attracted much attention in connection with the development of hydrogen fuel cells, in particular polymer electrolyte membrane fuel cells (PEMFCs) [1]. The HER and HOR and also related processes of hydrogen adsorption have been reviewed comprehensively (see, e.g., Refs. [13, 64–72] and references therein). Below, we first discuss possible mechanisms of the HER/HOR (Section 8.1.1.5.1) and then review how the reaction rates and mechanisms are related to different electrode materials (Section 8.1.1.5.2). In the last part (Section 8.1.1.5.3), we present some details on hydrogen adsorption on electrode surfaces. 8.1.1.5.1 Reaction Kinetics and Mechanisms The HER/HOR may occur by two different reaction sequences, the Tafel (63) – Volmer (64) or the Heyrovsky (65) – Volmer (64) reactions, which are given below for the case of an acidic electrolyte. Here ∗ and Had stand for free adsorption sites and hydrogen atoms adsorbed on the electrode surface, respectively.

−−  H 2 + 2∗ −− − − 2Had ∗

(63) +

−

−  Had − −− − − +H +e

+ − −−  H2 +∗ −− − − H + Had + e

(64) (65)

Hence the HER and the HOR are examples of complex multielectron–multistep electrochemical reactions, comprising charge transfer steps (64) and (65), chemical step (63) and also mass transport of H2 and H+ from and to the electrode surface. The kinetic equation for the HER/HOR and hence the dependence of the overall current on the electrode potential vary markedly depending on which

8.1.1 Fundamentals of Electrocatalysis

reaction is the RDS. Let us discuss some typical cases, which are summarized in Table 1 (see Refs. [66] and [73] for details). Case A in Table 1 assumes that the charge transfer reaction (64) is the RDS and the Butler–Volmer equation Eq. (26) holds. Equation (i) in Table 1 is obtained from Eq. (26) when considering the coverage of the electrode surface with hydrogen (θH ) and the fact that the evolution 0 is the of one H2 molecule involves two electrons. θH hydrogen coverage at equilibrium, i.e. at η = 0. Note that θH is a function of the electrode potential, which in general leads to a potential-dependent Tafel parameter. However, for metals poorly adsorbing hydrogen θH is low in a wide potential window and Tafel slopes close to 118.3 mV decade−1 at 298 K (corresponding to αV = 0.5) are observed. Case B corresponds to the case when the chemical reaction (63) is the RDS. This implies that the rate constant for the Tafel reaction is much smaller than that for the Volmer reaction and the latter may be considered to be in quasi-equilibrium. Under these simplifying conditions, the relation between the current and overpotential is usually expressed as
 RT i η= ln 1 + (66) 2F iT Note that Eq. (66) takes into consideration only the cathodic reaction. Here the rate of the heterogeneous Tafel reaction is formally converted into electric current: iT = 2F kT (θH )2

(67)

It is often assumed that θH does not differ markedly 0 and hence i ≈ i from the equilibrium coverage θH T 0T = 0 2 2F kT (θH ) . However, when writing θH as a function of the overpotential, it is also possible to obtain an analytical solution for the current without resorting to this simplification. Assuming a Langmuir isotherm, θH is given by θH =

0 θH 0 + (1 − θ 0 ) exp(ηF /RT ) θH H

(68)

which leads to Eq. (ii) in Table 1. A more general description is obtained when using a Temkin [74] or a Frumkin [75] isotherm. According to Eqs. (66) and (ii) in Table 1, a Tafel slope of 29.6 mV decade−1 is expected and has indeed been observed for the platinum metals. It should be noted, however, that although Eqs. (66) and (ii) give useful guidelines for the analysis of the HER/HOR on metals that strongly adsorb hydrogen, the treatment presented is oversimplified and in fact θH established is the result of the interplay of reactions (63) and (64). Analysis of Eq. (ii) in Table 1 leads to the conclusion

1887

that at high cathodic and anodic overpotentials limiting reaction currents must be observed independent of mass transfer limitations: il,c = −

i0T i0T ; il,a = 0 0 )2 2 (θH ) (1 − θH

(69)

This prediction has been widely discussed in the literature (see, e.g., Refs. [73, 76]), but has not been unambiguously supported experimentally. For the HOR the limiting anodic current is determined by the mass transport rather than by the slow chemical reaction (see case E, Table 1). For cases C and D in Table 1, i.e. when reactions (64) and (65), respectively, are rate determining in the case of the Volmer–Heyrovsky sequence, kinetic equations similar to case A are obtained. More complex situations arise when the exchange current densities for different reaction steps have comparable values [66, 68, 73]. 8.1.1.5.2 Influence of the Electrode Material on the HER/HOR The HER has been studied on various metal and alloy electrodes and the exchange current density has been shown to vary by up to 10 orders of magnitude. Many authors have attempted to establish correlations between the rate of the HER and physical and chemical properties of electrode materials. The dependence of the HER on the metal work function  was first discussed by Bockris in 1947 [77]. In his seminal papers, Trasatti presented critical analysis of experimental data on metal work functions and potentials of zero charge [78] and demonstrated that log i0 varies linearly with , irrespective of the detailed nature of the mechanism involved in the rate-determining step [79]. Moreover, transition metals and sp metals with positively charged surfaces and sp metals with negatively charged surfaces (Fig. 9) fall on two different straight lines that are shifted in parallel along the  axis. The shift has been interpreted in terms of differing orientations of water dipoles at the interface of the two groups of metals. An issue which has attracted much attention and which has been approached by a number of authors is the relationship between the catalytic activity of materials towards the HER and hydrogen adsorption energies. Conway and Bockris [80] demonstrated that the linear dependence of log i0 on  originated from a correlation between  and M–H adsorption energy. Parsons [81] showed that volcano-type curves (see Fig. 8 and Section 8.1.1.4.1) should arise when log i0 values for a series of metals are plotted against the standard Gibbs energy of hydrogen adsorption. It took quite a References see page 1902

Volmer (A)

Tafel (B)

Volmer (C)

Heyrovsky (D)

Diffusion (E)

Volmer–Tafel

Volmer–Tafel

Volmer–Heyrovsky

Volmer–Heyrovsky

Volmer–Tafel or

a Concentration



θH

exp

Kinetic equation

(1 − αV )ηF −αV ηF 1 − θH exp − (i) RT RT θH0 1 − θH0 exp(2ηF/RT) − 1 i = i0T 0 (ii) [θH + (1 − θH0 ) exp(ηF/RT)]2 See case A

1 − θH (1 − αH )ηF −αH ηF θH i = 2i0,H exp − 0 exp (iii) RT RT 1 − θH0  θH

RT RT i i η= − (iv) ln 1 − ln 1 − 2F il,c 2F il,a i = 2i0,V

polarization is neglected in all cases except for (E).

Volmer–Heyrovsky

Rate-determining step

Kinetic Equations and Tafel Parameter b for HER/HORa

Mechanism

Tab. 1

2.3RT (for θH ≈ constant) αH F 2.3RT − 2F −

See case A

2.3RT (for θH ≈ constant) (1 − αH )F 2.3RT 2F

2.3RT (for θH ≈ constant) αV F 2.3RT − (for θH ≈ constant) 2F −

Cathodic reaction (HER)

See case A

2.3RT (for θH ≈ constant) (1 − αV )F 2.3RT (for θH ≈ constant) 2F

Anodic reaction (HOR)

Tafel parameter b

1888 8.1 Electrocatalysis

8.1.1 Fundamentals of Electrocatalysis

13

Hydrogen evolution exchange current, – log i0 / A cm−2

Hg Tl

11

Cd In

Mn

Pb Ag Zn Ga

Tl

Bi b

Sn

ln

9 Ta Ti Nb

Sb Ga Ag

Al

Cu

Sn Bi Mo

7

Cr

Au W Fe

a

5

Sb

Co Ni Ru Os

Ir Rh Pd Re Pt

3

4.0

4.5 Work function, f / eV

5.0

Exchange current densities for the HER vs. work functions of metals: (a) corresponds to transition metals and sp metals with positively charged surfaces and (b) refers to sp metals with negatively charged surfaces. Reproduced with permission from Ref. [79].

Fig. 9

while, however, until the theoretically predicted volcano relationship was confirmed experimentally. The critical aspect is the source for M−H bond energies. Most often reproduced (although not always acknowledged as such) is the volcano relationship between log i0 and the heat of adsorption of hydrogen on metals as derived by Krishtalik [82] from experimental data for electrochemical hydrogen evolution [79]. However, as pointed out by Trasatti [79], the significance of a volcano plot would have been much greater if the electrochemical data could be correlated with independently measured heats of hydrogen adsorption in the gas phase. This has been done for transition metals [79]. However, for sp metals, which adsorb hydrogen rather poorly, data for hydrogen adsorption in the gas phase are generally not available. Hence the heat of formation of metal hydrides was taken as a measure. Such a volcano plot is presented in Fig. 10. Data for sp and transition metals are shifted with respect to each other along the x-axis. The shift is attributed to the fact that the values for M−H bond

1889

energies derived from heats of bulk hydride formation are generally about 20 kcal mol−1 (1 kcal = 4.184 kJ) higher than those obtained from adsorption heats. For further details the reader is referred to Refs. [79, 83]. Although the qualitative validity of volcano plots for the HOR/HER is beyond any doubt, quantitative correlations between the exchange current densities and the M−H bond energies must be verified in view of limited experimental data for well-characterized metal surfaces (see the discussion in Ref. [72]). In agreement with the discussion in Section 8.1.1.4.1, it has been found that for the metals on the ascending branch of the volcano plot, the RDS is associated with the slow discharge reaction (64), whereas for those on the descending branch slow hydrogen atom removal in reaction (63) or (65) is usually the RDS. More precisely, mechanism C (see Table 1) is usually proposed for Pb, Tl, Hg, Cd, Ag, Au and Cu, mechanism D for W, Mo and Nb and mechanism B for platinum metals at low overpotentials; at high overpotentials, some authors proposed for Pt the Volmer–Heyrovsky mechanism (see the discussion in Ref. [68]). It should also be noted that the mechanism of the HER is strongly influenced by the electrode pretreatment and the reaction conditions (potential, pH, temperature, electrolyte composition, etc. [72, 73]). Pt possesses the highest catalytic activity and both HER and HOR on Pt and other platinum metals exhibit high current densities in the vicinity of the equilibrium potential. The activities of Pt electrodes in acidic electrolyte considerably exceed those observed in alkaline solutions. Different explanations have been offered to account for this fact [67, 83], a very likely one being competitive adsorption of OH and H. In practical applications, for water electrolysis massive noble metal electrodes are too expensive to justify their utilization. In alkaline electrolytes, nickel is a reasonable alternative due to its high electrocatalytic activity and stability. Even more suited for industrial electrolysis is steel [68]. The HOR proceeds through the same series of reaction steps as the HER, but in the reverse direction. As already mentioned, the interest in the HOR has greatly increased recently due to the research in the area of low-temperature fuel cells, in particular PEMFCs. Experimental investigations of the HOR in aqueous electrolytes are complicated by mass transport limitation of H2 due to its low solubility in water (∼10−3 mol L−1 ). This problem can be attenuated either by application of a high-speed rotating disk electrode (RDE) [84, 85] or impedance spectroscopy (IS) ([83] and references therein). The HOR has been investigated on noble metal electrodes, most extensively on Pt, and has proven to be structure sensitive and to differ markedly depending on the crystallographic orientation of Pt single-crystal References see page 1902

8.1 Electrocatalysis

Exchange current for H2 evolution, –log i0 / A cm−2

1890

∆g ° > 0 qH − 0

∆g ° < 0 qH − 1

Pt Pd

3 Rh

Cu*

5

Au*

a

Ni

Cu

Fe

w Mo

b

7

Cr Ag* Zn Al

Ti

Ta

Ga

9

Cd In Tl

50

Tl*

70

70

80

M-H Bond strength / kcal mol−1 Exchange current densities for the HER vs. strength of M−H bond derived from the heat of hydride formation in the case of sp metals and from the heat of hydrogen adsorption from the gas phase in the case of transition metals. Reproduced with permission from Ref. [79].

Fig. 10

Hydrogen Adsorption on Electrode Surfaces As already pointed out, the formation of chemisorbed hydrogen is of paramount importance in the overall mechanism of both the HER and the HOR. Electrochemical adsorption 8.1.1.5.3

1.2 H2 10 ppm CO in H2 40 ppm CO in H2 100 ppm CO in H2 Solid Pt Dashed PtRu

1.0

Cell voltage / V

surfaces [83, 84]. No agreement has been achieved so far concerning the particle size effects in the HOR/HER (see discussion in Ref. [83]). The oxidation of H2 is strongly suppressed when CO impurities are contained in the hydrogen stream. This is illustrated in Fig. 11, which shows the current–voltage characteristic of a fuel cell fed by pure and COcontaminated hydrogen. The decrease in the cell performance is due to the fact that CO is strongly adsorbed and blocks surface sites, since COad is not oxidized on Pt in the potential interval relevant for the operation of PEMFC anodes. Much attention has been drawn to the search for CO-tolerant anodes for the hydrogen oxidation in view of the development of low-temperature fuel cells fed by reformate gas. Figure 11 shows considerable improvement of the cell performance when PtRu alloys are used at the anode instead of Pt. However, despite substantial improvements, the cell performance in the presence of CO (even with a PtRu anode) does not reach the level achieved with pure hydrogen. The development of CO-tolerant anodes continues to be an active research area; for more information the reader is referred to Refs. [67, 86].

0.8 0.6 0.4 0.2

200

400

600

800

1000

1200

Current density / mA cm−2 Fig. 11 Influence of CO poisoning in a PEMFC with pure Pt (solid lines) and Pt0.5 Ru0.5 alloy anodes (dashed lines). The anodes were prepared from 20 wt.% Pt/Vulcan XC72R or 20 wt.% Pt+10 wt.% Ru/Vulcan XC72R at a loading of 0.25 mg Pt cm−2 . The cathode uses 40 wt.% Pt/Vulcan XC72R at a loading of 0.6 mg Pt cm−2 . The membrane–electrode assemblies (MEAs) are based on catalyzed substrates bonded to Nafion NE-115 membrane. The Ballard Mark 5E single cell is operated at 80 ◦ C with full internal membrane humidification. Reproduced with permission from Ref. [86].

of hydrogen has been extensively studied on noble metal electrodes (for references see, e.g., Ref. [71]). It occurs via the discharge of protons in acidic or the reduction of water molecules in basic medium. The phenomenon of the so-called ‘‘underpotential deposition’’ (UPD) of hydrogen

8.1.1 Fundamentals of Electrocatalysis

was first recognized (although not yet designated as such) by Frumkin and Slygin [87] when they analyzed charging curves on Pt. The term comes from the fact that hydrogen adsorption occurs positive of the reversible potential of the hydrogen electrode. The reaction is strongly sensitive to the electrode material and surface crystallography. Cyclic voltammograms (CVs) for well-ordered low index Pt(hkl) single crystals were first reported by Clavilier et al. [88, 89] and are presented in Fig. 12 for the so-called ‘‘HUPD region’’, i.e. the potential region in which UPD of hydrogen occurs. The negative currents in the CVs stem from the discharge of hydronium ions and formation of Had , while the positive currents are due to the reverse reaction of Had oxidation.

Current density / µA cm−2

200

100

0

−100

−200

0.0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

E / V vs. RHE

Cyclic voltammograms for Pt(hkl) in H2 SO4 : (111), solid lines; (100), dashed lines; (110), dotted lines. Reproduced with permission from Ref. [83].

Fig. 12

Partial pressure of H2 / arb. units

40

Current / µA

Numerous attempts have been made to correlate hydrogen UPD at solid/electrolyte to hydrogen adsorption at solid/gas interfaces. Figure 13 compares HUPD formation and removal on the high index plane Pt(533) from 0.5 M H2 SO4 with temperature-programmed desorption (TPD) data for the same surface. This surface consists of four-atom-wide terraces of (111) structure separated by (100) steps. One may note striking similarities between the electrochemical and gas-phase data. The high-temperature peak (corresponding to so-called strongly adsorbed hydrogen) in TPD is attributed to hydrogen recombinative desorption from the (100) step sites whereas the low-temperature peak (corresponding to weakly adsorbed hydrogen) stems from H2 desorption from the (111) terraces [90]. For Pt(111), the high-temperature TPD peak is absent and only the low-temperature peak remains [91]. Analogously, in the electrochemical environment strongly adsorbed hydrogen (above 0.25 V) is attributed to HUPD on (100) sites (cf. Figs. 12 and 13) and the peaks in the interval from 0.1 to 0.2 V to HUPD from (110) and (111) sites. Hence both HUPD and gas-phase hydrogen adsorption are very sensitive to the surface crystallography. The similarities between hydrogen adsorption at the solid/gas and the solid/electrolyte interface are not only qualitative. Also the energy of M−H bonds calculated, e.g., for Pt(111) from the gas phase [92] and from the electrochemical data are in very good agreement [71]. The energetics of hydrogen electrochemisorption has been accessed by a few research groups [93–95] studying the temperature dependence of HUPD . The determination of Had (HUPD ), Gad (HUPD ) and Sad (HUPD ) on Pt(111) in different electrolytes revealed that the thermodynamic quantities depend only very weakly on the nature of the electrolyte [95]. The bond energy of References see page 1902

18000

60

20 0 −20 −40 −60

15000 12000 9000 6000 3000 0

0.0 0.1 0.2 0.3 0.4 0.5 (a)

1891

E / V vs. RHE

100 150 200 250 300 350 400 450 500 550 (b)

Temperature / K

Comparison of (a) HUPD on Pt(533) in 0.5 M H2 SO4 measured at 100 mV s−1 [171] and (b) thermal desorption spectra of H2 from the Pt(533) surface. The coverages increase from 0.09 to 0.9. H2 was dosed at Ts = 120 K and the heating rate was 1 K s−1 . Reproduced with permission from Ref. [90]. Fig. 13

8.1 Electrocatalysis

150 PdML/PtRu(111) PdML/Pt(111)

j /µA cm−2

Pt(111)–HUPD , EPt(111)−HUPD , is equal to 262 kJ mol−1 at zero coverage [94] (∼240 kJ mol−1 [95]) and was found to depend only slightly on the hydrogen coverage. The experimental values for EPt(111)−HUPD are in very good agreement with gas-phase adsorption data (EPt(111)−HUPD = 255 kJ mol−1 ) [92] and also with the values calculated using DFT (EPt−Had = 252 kJ mol−1 [96]; for more references the reader is referred to Ref. [97]). Along with the similarities between hydrogen electrochemisorption at solid/liquid and hydrogen chemisorption at solid/gas interfaces, one should consider also the differences which exist between these processes. First, gas-phase hydrogen adsorption involves the dissociation of hydrogen molecules and hydrogen desorption involves the association of hydrogen atoms, whereas HUPD formation or oxidation does not. Instead, the latter processes include highly energetic H+ solvation or desolvation. The enthalpy of H3 O+ formation is −754 kJ mol−1 ; hydration of H3 O+ to H9 O+ 4 releases an additional enthalpy of −356 kJ mol−1 [71]. Second, in contrast to hydrogen chemisorption, HUPD formation/oxidation is associated with a charge transfer and thus depends on the electrode potential. Also, the presence of ions may influence HUPD . Indeed, for Pt single-crystal surfaces in sulfuric acid solutions, (bi)sulfate adsorption is superimposed on HUPD . The two adsorption processes are clearly resolved in the CV of Pt(111) in H2 SO4 electrolyte: the low-potential contribution is related to HUPD , whereas the high-potential component is due to chemisorption of (bi)sulfate (see hatched part in Fig. 12). On Pt(111) in H2 SO4 the maximum HUPD coverage amounts to 0.66, whereas that of sulfate/bisulfate is 0.21 [98]. In alkaline electrolytes, H competes for the adsorption sites with OH. One of the main goals in electrocatalysis is to establish relationships between the structure and the electronic properties of an electrode material and its chemisorptive and catalytic properties. It has recently been shown by DFT calculations that the electronic properties of a surface can be modified considerably by changing nearest-neighbor separations through the formation of pseudomorphic overlayers on foreign substrates [99, 100]. This theoretical prediction has been verified experimentally by a number of research groups. For example, Kibler et al. [101] observed a systematic shift of the HUPD peak on pseudomorphic Pd monolayers on different substrates of (111) orientation (Fig. 14a). The position of the HUPD peak correlated with the d-band center δεd calculated using DFT (Fig. 14b) [100]. The formation of UPD hydrogen positive of the hydrogen reversible potential is a unique property of noble metal electrodes. However, adsorbed hydrogen is formed on all metal electrodes evolving hydrogen as an intermediate in reactions (63)–(65). In order

100

PdML/Ir(111) Pd(111) PdML/Rh(111)

50

PdML/Ru(0001) PdML/Au(111)

PdML/Re(0001)

0 −0.4

−0.3

−0.2

−0.1

0.0

0.1

0.2

ESCE / V

(a) 0.05

PdML/Au(111)

0.00

Pd(111)

−0.05

ESCE / V

1892

PdML/Pt(111)

−0.10 −0.15 −0.20 −0.25

PdML/Rh(111)

PdML/PtRu(0001)

PdML/Ir(111) PdML/Ru(0001) PdML/Re(0001)

−0.30 −1.0 −0.8 −0.6 −0.4 −0.2 0.0 0.2 0.4 0.6 (b)

ded /eV

Fig. 14 (a) Positive sweeps of the CVs for Pd(111) and pseudomorphic palladium monolayers (PdML ) on seven different single-crystal substrates in 0.1 M H2 SO4 , revealing a spectrum for hydrogen desorption. Scan rate, 10 mV s−1 . (b) Plot of the hydrogen desorption potentials versus the shift of the d-band center δεd (calculated using DFT [100]). Reproduced with permission from Ref. [101].

to make a distinction between these two types of chemisorbed hydrogen at the solid/electrolyte interface, it was suggested (see, e.g., Refs. [102, 103]) to name the latter overpotentially deposited (OPD) hydrogen. We would like to stress that chemically HUPD and HOPD are identical, but they are believed to differ in terms of adsorption sites and bonding energies. On noble metal electrodes HOPD is assumed to form on top of a HUPD monolayer, but may also form when HUPD is fully suppressed, e.g. by a monolayer of chemisorbed sulfur ([71] and references therein). In contrast to this widely accepted view, Breiter argues ([66] and references therein) that an assumption on the existence of two types of adsorbed hydrogen (HOPD and HUPD ) is excessive and that it is HUPD which becomes mobile in the vicinity of a monolayer coverage and participates in the HER/HOR. It has been proposed that in the case of the HOR the HUPD is not a reactive intermediate either. Some authors considered this reactive intermediate as HOPD , which by definition is incorrect since the HOR occurs above the reversible potential of the hydrogen electrode. This semantic problem, however, does not resolve the issue

8.1.1 Fundamentals of Electrocatalysis

of the reactive intermediate in the HOR or the HER. Given the experimental data on the RDE curves for H2 oxidation and recent IR spectroscopic evidence [104], it is indeed likely that the reactive intermediate both in HER and HOR is different from HUPD . An important issue which has not been fully resolved yet concerns the adsorption sites for HUPD and for the reactive intermediate in HOR/HER. It is usually speculated that HUPD occupies multi-coordinated sites [three-fold sites on (111) and four-fold sites on (100) surfaces], whereas HOPD resides in an atop position (see discussion in Ref. [71] and references therein). However, somewhat conflicting evidence comes from IR/VIS SFG (sum frequency generation) investigations by Tadjeddine et al., which favor mono-coordinated HUPD [105, 106]. Interesting data on the adsorption of hydrogen on platinum have recently been obtained with X-ray absorption spectroscopy (XAS) [107]. Highly dispersed 1.5–2.0-nm Pt particles supported on carbon were studied in situ in 0.1 M HClO4 electrolyte with Pt L2,3 XAS. The experimental results combined with real-space full multiple scattering calculations on model clusters led the authors to the following conclusions: (i) at low coverage a chemisorbed hydrogen atom is highly mobile and possibly delocalized on the surface, (ii) at higher coverage it localizes into fcc sites and (iii) at very high coverage H is also found in atop sites presumably at or near edges. Further research is needed to sort out the issue of adsorption sites of HUPD and of the reactive intermediate(s) in the HOR/HER [71]. The Oxygen Reduction and the Oxygen Evolution Reaction The oxygen reduction reaction (ORR) and the oxygen evolution reaction (OER) have been studied extensively because of their importance for many practical systems. The OER occurs in water electrolyzers and the ORR in metal–air batteries and fuel cells. The overall reactions involve four electrons: 8.1.1.6

−−  O2 + 4e− + 4H+ −− − − 2H2 O

at an appreciable rate. Because of this, cathodic and anodic reactions occur on essentially different electrode surfaces: The ORR takes place on pure or oxygencovered metal surfaces (depending on the type of metal) whereas the OER proceeds on oxide phases. The exchange current density of the ORR/OER is typically in the range 10−9 –10−11 A cm−2 . Because of these small values, the activities of different electrode materials towards the ORR/OER are usually characterized by current densities at specified electrode potentials rather than by the exchange current densities. The ORR and the OER have been subject of many reviews, e.g. Refs. [108–112]. In Sections 8.1.1.6.1–8.1.1.6.4 we summarize important aspects of the ORR and in Section 8.1.1.6.5 those of the OER, both in aqueous electrolyte solutions. We start our discussion with reaction mechanisms and their relation to the electrode material (Section 8.1.1.6.1). Since the formation of various oxygen-containing species plays a key role in the ORR, Section 8.1.1.6.2 is devoted to water electrochemisorption and molecular oxygen adsorption on noble metal electrodes. In Section 8.1.1.6.3, we briefly discuss structural and particle size effects in the ORR, and Section 8.1.1.6.4 deals with catalyst development for the ORR in polymer electrolyte fuel cells. High-temperature studies of the ORR are not treated below; interested readers are referred to a review paper [113] and references therein. In Section 8.1.1.6.5, the most important aspects of the OER are summarized. 8.1.1.6.1 Reaction Mechanisms: Relation to the ElecThe ORR is a multielectron electrode Material trochemical reaction that, depending on the electrode material and the reaction conditions, may involve various reaction intermediates, in particular , HO2,ad , Oad , OHad , H2 O2,ad . The latter may O2,ad , O− 2,ad desorb from the electrode surface and can be detected in the electrolyte. A scheme showing different pathways of the ORR, including the intermediate formation of H2 O2 was proposed by Wroblowa et al. [114] and Bagotskii et al. [115] and is represented in Fig. 15. The 4 e− path (k1 )

(70)

with E 0 = 1.23 V vs. SHE at 298 K (in acidic electrolytes), and − −−  O2 + 4e− + 2H2 O −− − − 4OH

1893

k1 k2

(71)

with E 0 = 0.40 V vs. SHE at 298 K (in alkaline electrolytes). In contrast to the reactions at the hydrogen electrode, those occurring at the oxygen electrode are characterized by a very sluggish reaction kinetics even at the most catalytically active Pt electrodes, hence a high negative (positive) overpotential is required to reduce (form) O2

O2

O2,ad

k3 H2O2,ad

k4

H2O

k5 H2O2

Fig. 15

Simplified scheme of the ORR from Refs. [108, 114].

References see page 1902

1894

8.1 Electrocatalysis

is usually termed the direct or parallel pathway and that involving intermediate H2 O2 formation the series pathway (k2 and k3 ). H2 O2 may also catalytically decompose on the electrode surface (k4 ). Isotope measurements showed that hydrogen peroxide formation occurs without O−O bond splitting (see Ref. [108] and references therein), while water formation obviously requires it. Much research has been directed towards the understanding of the ORR mechanism on various electrode materials. In 1959, Frumkin et al. proposed to use a rotating ring disk electrode (RRDE) [116] for detecting stable intermediates formed at the working electrode in the course of complex electrochemical reactions. Detection of H2 O2 at the ring electrode gives evidence that the ORR follows the series rather than the direct mechanism. Since 1960, many research groups have put considerable effort into investigating the ORR using RRDE with various disk electrode materials. Details of these studies can be found in review articles [108–111]. According to the mechanism of the ORR, electrode materials can be divided into two groups. The first comprises metals which catalyze the ORR predominantly via the 2e− peroxide mechanism. These include mercury, graphite, gold {except for Au(100) in alkaline solution [117]}, the majority of metal oxides and oxide-covered metals. The ORR was studied most thoroughly on Hg and graphite electrodes in alkaline electrolytes. Two reduction waves were observed and could be assigned unambiguously: the first wave corresponds to the reduction of oxygen to peroxide and the second to the reduction of peroxide to water. Investigations of the reaction orders with respect to molecular oxygen and H+ suggest that the RDS on Hg and graphite electrodes is the first electron transfer both in the first wave, i.e. O2 reduction, and in the second wave, i.e. the H2 O2 reduction reaction. Predominantly 4e− oxygen reduction occurs on Pt and other platinum metals, and also on Pt-based alloys, Ag and, as already mentioned, Au(100) in alkaline electrolytes. Despite the fact that the ORR has been extensively investigated on Pt metals (especially on Pt), first on polycrystalline and later on single-crystal electrodes [13], there is as yet no consensus regarding the detailed reaction mechanism. For polycrystalline Pt in alkaline electrolytes some H2 O2 was detected (although in small quantities) at the ring of an RRDE. Hence the series mechanism was proposed. In acidic solutions, the amount of hydrogen peroxide formed is usually lower. The absence of H2 O2 in the electrolyte, however, does not imply that H2 O2,ad is not formed as a reaction intermediate. It just means that if it is produced it does not leave the working electrode due to either a high k3 or a high k4 value. The Tafel slopes for polycrystalline and for supported Pt nanoparticles in acidic solutions have been found to change from ca. −120 mV decade−1 in the potential interval below ca. 0.8 V vs. RHE

to ca. −60 mV decade−1 at higher electrode potentials. The change in the Tafel slope has been attributed to the fact that below ca. 0.8 V vs. RHE the Pt surface is free from oxides, whereas above 0.8 V it is covered by oxides. Hence understanding the ORR on surfaces adsorbing oxygen and forming surface oxides requires understanding the surface state under relevant conditions. 8.1.1.6.2 Oxygen Adsorption and Formation of Surface Oxides and Their Relation to the ORR Formation of adsorbed oxygen species and surface oxides may occur on electrode surfaces either by electrochemical adsorption and decomposition of water or via the interaction with molecular oxygen. The electrochemical oxidation of noble and IB metal surfaces has been investigated by numerous electrochemical methods and also by surface-sensitive techniques performed either in situ using IR spectroscopy, ellipsometry, electrochemical quartz-crystal microbalance (EQCM), XAS and X-ray scattering, electrochemical scanning tunneling microscopy (STM) and Raman spectroscopy, or ex situ with the electrodes removed from the electrochemical cell and transferred to ultra-high vacuum (UHV) using, e.g., X-ray photoelectron spectroscopy (XPS), Auger electron spectroscopy (AES) or STM (see, e.g., review articles [118, 119] and references therein). Whereas the formation of oxide phases on the electrode surfaces has been studied fairly extensively, much less is known about the initial steps of metal oxidation. The first step of surface oxidation is usually assumed to be OHad formation, which in acidic and alkaline electrolytes reads + − −−  M + H2 O −− − − M − OHad + H + e

(72)

− −  M + OH− − −− − − M − OHad + e

(73)

respectively. This step has a clear voltammetric signature for Pt(111) in alkaline and acidic aqueous electrolytes of weakly adsorbing anions (e.g. HClO4 ) [13], Au(111) [120] in neutral and alkaline electrolytes and Ag(hkl) in alkaline electrolytes [121, 122]. The reversible nature of OH adsorption on Pt(111) is reflected by mirror-like anodic and cathodic peaks. Using density functional theory, Anderson ([123] and references therein) calculated the reversible potential of reaction (72) on Pt to be 0.62 V vs. RHE. The Pt(111)–OHad bond energy has been estimated as ∼136 kJ mol−1 in alkaline electrolyte, which is much smaller than the Pt−Oad bond energy (∼350 kJ mol−1 ) at a gas/solid interface (see Refs. [13, 95] and references therein for details). In solutions containing strongly adsorbing anions (e.g. SO2− 4 ), OH adsorption is inhibited by anion chemisorption. Despite the importance of OHad and Oad for the ORR and the oxidation of CO and other organic molecules, papers reporting

8.1.1 Fundamentals of Electrocatalysis

placeexchange

M − OHad −−−−−−−−−→ (OH − M)quasi 3D lattice (74) It was proposed that the incorporation of OH into the metal lattice is initiated via place exchange. After the OH species have been transformed into O, the subsequent growth of thick metal oxides presumably takes place by field-assisted transfer of metal cations through this film, into the 3D-type layer via the Mott–Cabrera field-assisted growth mechanism [128]. In more recent investigations, Birss et al. [129] and then Jerkiewicz et al. [130] questioned the above mechanism for Pt electrodes. Through combined cyclic voltammetry, EQCM and Auger electron spectroscopy measurements, Jerkiewicz et al. proposed that the oxidation of Pt surfaces occurs via the formation and subsequent place exchange of Oad rather than OHad [130]. An essential question is whether the electrochemically formed oxides are equivalent to those formed through gasphase metal oxidation. This issue has been addressed by Weaver’s group in a series of publications ([131] and references therein). Five Pt-group metals, namely platinum, palladium, iridium, rhodium and ruthenium, were examined by means of surface-enhanced Raman spectroscopy (SERS) in aqueous electrochemical and gaseous dioxygen environments as a function of electrode potential and temperature, respectively, with the objective of comparing systematically the conditions required for surface oxide formation and of elucidating the reaction mechanisms involved. In order to obtain surface enhancement, noble metal films were deposited on roughened Au templates. Figure 16 compares SERS spectra of 1–3 monolayers of oxides grown on the surface of a Pt film. Essentially the same band at 575 cm−1 develops, in the electrochemical environment above 0.6 V SCE and in the gas-phase environment at 200 ◦ C. It coincides with the wavenumber characteristic of Pt−O stretch previously observed for bulk amorphous PtO2 [132]. Remarkable is the fact that electrode potential reversal recovers the pristine Pt surface, whereas the gas-phase Pt oxide reduction is kinetically hindered. It was therefore concluded that the oxides formed

V (SCE) 0.0 25 °C 0.4

Intensity / arb. units

their direct spectroscopic observation are scanty. Among spectroscopic methods applied, Raman spectroscopy [124, 125], IR spectroscopy [120], XPS [121, 126] and XAS [127] should be mentioned. Whereas the onset of OHad formation is believed to be reversible, at more positive electrode potentials it becomes irreversible. In order to account for the irreversibility observed for Pt and Au electrodes, Angerstein-Kozlowska and Conway ([118] and references therein) proposed the so-called place-exchange mechanism between metal atoms and OH moieties:

1895

400 °C 350 °C

0.6

300 °C 0.8 250 °C 1.0

575 200 °C

0.8

575 150 °C

0.6 100 °C 0.0 25 °C

600

200

Raman (a)

600

200

shift / cm−1 (b)

Fig. 16 (a) Potential-dependent SER spectra acquired for electrochemical oxidation and subsequent reduction of a platinum film deposited on an Au substrate in 0.1 M HClO4 . The initial potential was 0 V vs. SCE (bottom spectrum), followed by potential increments of 0.2 V up to 1.0 V and finally returning to 0 V. (b) Temperature-dependent SER spectra acquired for thermal oxidation of a platinum film in 1 atm of flowing O2 . The temperature was increased in 50 ◦ C increments before decreasing back to 25 ◦ C. Reproduced with permission from Ref. [131].

on noble metal surfaces at solid/liquid and solid/gas interfaces were largely similar (although distinct differences were observed for Pd and Ru). However, the kinetics of their formation and reduction are significantly different: they occur through direct oxide formation at the gas/metal interface and a metal–oxygen place-exchange mechanism, expedited by interfacial solvation in the electrochemical environment. Thus, oxide formation in the anhydrous gas-phase environment is energetically unfavorable. Among other techniques utilized for the investigation of metal oxidation, XAS, which allows in situ monitoring of structural transformations at the surfaces of both smooth and dispersed metals, can be mentioned. Figure 17 References see page 1902

8.1 Electrocatalysis

Fourier transform modulus, k 3c

1896

0

5.0 10 0 Tim e/

4.0

s

2.0

20 0 30 0

0

Fourier transform modulus, k 3c

(a)

1.0

3.0 Å R/

O2 + H+ + e− −−−→ O2 Had

0

5.0 10 0 Tim e/ s

(b)

accompanied by surface restructuring. Interestingly, in their studies of oxide formation and reduction kinetics, the authors did not find a signature of the place-exchange mechanism proposed for smooth surfaces. Overall we would like to stress that the structure and dynamics of the interface under conditions relevant for the ORR and other electrocatalytic reactions is a key issue of modern electrocatalysis and further studies are necessary to achieve a better understanding. As already mentioned, surface oxides formed on Pt and other metal electrode surfaces strongly influence the kinetics and mechanism of the ORR. The change in the Tafel slope observed for polycrystalline Pt electrodes during the ORR has been attributed to the build-up of Pt oxides on the electrode surface above 0.8 V vs. RHE [134]. The apparent coverage of electrodes by oxygen-containing species determined from the charge has been found to increase linearly with the electrode potential. For polycrystalline Pt, the reaction order with respect to oxygen is 1 in both potential intervals (below and above 0.8 V vs. RHE), whereas the reaction order with respect to protons changes from 1 in the low-potential region to ca. 1.67 in the high-potential region [134]. These results were interpreted in terms of the addition of the first electron to the adsorbed oxygen molecule as the RDS:

4.0 3.0 20 0

2.0 30 0

1.0

R/

Å

0

Fig. 17 Fourier transform moduli of the Pt EXAFS −1 (k = 2.2 − 9.7 A˚ ) acquired during (a) the oxidation and (b) the reduction of Pt/C catalyst (particle diameter ca. 2 nm) as a function of time after jumping the electrode from 0.1 to 1.2 V. The peak at ˚ 2.24 A˚ corresponds to the first shell of Pt near neighbors at 2.76 A. The peak at 1.50 A˚ is a combination of the side-lobe from the Pt shell and a shell of O near neighbors at 2.01 A˚ [133].

shows Fourier transform moduli (describing the radial distribution function around Pt atoms) extracted from Pt L3 edge EXAFS spectra of carbon-supported Pt nanoparticles measured in the dispersive mode [133]. The catalyst was incorporated in a polymer electrolyte membrane fuel cell operating at 80 ◦ C with 1 atm (101 kPa) of water-saturated N2 over the working electrode and H2 on the counter electrode. The appearance and the concomitant growth of the peak corresponding to the Pt−O bond occurred simultaneously with the decrease in the intensity of the peak corresponding to Pt−Pt bond. It was concluded that Pt surface oxidation is

(75)

For the oxide-free surface this agrees well with a −120 mV decade−1 Tafel slope for α = 0.5. For partially oxide-covered surfaces, assuming that the free energy of activation for the ORR depends on the coverage of the surface oxide, a change in the Tafel slope to −60 mV decade−1 can be predicted, which is in agreement with the experimental data. The relevance of water electrochemisorption to the change in the Tafel slope has recently been confirmed by varying the content of water in the electrolyte [135]. The experiments were performed in H2 O–TFMSA (trifluoromethanesulfonic acid) mixtures with the water:acid mole ratio varied from 50 : 1 to 4 : 1. Whereas at high water contents the Tafel slope changed from −112 to −59 mV decade−1 in agreement with what has been observed previously in aqueous solutions of H2 SO4 and HClO4 , at low water contents no change in the Tafel slope was observed (Fig. 18). This substantiates the involvement of water in the formation of oxides on the Pt surface. One of the critical issues is related to whether the adsorbates formed through dissociative adsorption of molecular oxygen are equivalent to those formed via water electrochemisorption. Yeager [136] suggested that the mechanism of the ORR depends on the type of molecular oxygen bonding to the electrode surface. Three models for molecular oxygen bonding have been

8.1.1 Fundamentals of Electrocatalysis

1.00

−1

1M TFMSA l.o.r. (59 mV dec ) −1

1M TFMSA h.o.r. (112 mV dec )

0.95

−1

Potential / V

6M TFMSA (110 mV dec )

0.90 0.85 0.80 0.75 0.70 0.65 −3.0

−2.5

−2.0

−1.5

Log i k /

mAcm−2

−1.0

−0.5

0.0

Fig. 18 Tafel plots for the ORR at room temperature on a polycrystalline Pt bulk electrode at 1225 rpm based on the potential sweep from 1.2 to 0.3 V at 25 mV s−1 in 1 and 6 M CF3 SO3 H. Reproduced with permission from Ref. [135].

proposed: (a) according to the Griffith model [137], an O2 molecule interacts with a single substrate atom by forming a bond between its π orbitals and the empty d2z orbitals of the metal surface atom; (b) in the end-on Pauling model [138], the σ orbital of an O2 molecule donates electron density to an acceptor d2z orbital of the metal; (c) the bridge model holds if an oxygen molecule binds to two surface atoms. Yeager [136] suggested that type (a) and (c) adsorption of oxygen favors O−O bond splitting and thus the direct 4e path of the ORR, whereas type (b) adsorption results in the 2e pathway of the ORR and the formation of H2 O2 . To gain further information on the type of O2 adsorption during the ORR, Adzic and Wang [139] used the adsorption of foreign (Ag) adatoms to probe the oxygen adsorption reaction at a Pt(111) surface. The inhibition of the ORR on Pt(111) by sub- and monolayer coverages of Ag was studied using electrochemical and in situ surface X-ray scattering techniques. The analysis of the extent of the inhibition of the ORR as a function of the Ag coverage showed that the data are best interpreted with O2 adsorbed at a bridge site. Oxygen adsorption on Pt surfaces from the gas phase has been studied with many spectroscopic techniques with the aim of identifying and characterizing different adsorption states. O2 is physisorbed on Pt(111) below 90 K; at higher temperatures two chemisorbed molecular states have been identified, superoxo (O− 2 ) and peroxo ), with different extents of electron donation from (O2− 2 the Pt surface to the oxygen molecule. It was demonstrated that the two chemisorbed molecular states have O2 lying on bridging di-σ and µ − π sites. Above 150 K O2 dissociates to form atomically adsorbed oxygen in 3-fold hollow surface sites. For references the reader is referred to Ref. [140].

1897

Recently, the ORR has attracted the close attention of theoreticians who performed quantum chemical calculations of reaction intermediates using different computational approaches [29, 140–146]. Potential energy surface profiles for the 2e and 4e O2 reduction on Pt have been considered, the 4e path being identified as the dominant one [144, 145, 147]. Anderson et al. considered both end-on adsorption of O2 to a single Pt atom [148] and a bridge-bonded O2 molecule adsorbed to a Pt2 cluster [140]. In agreement with the earlier hypothesis of Yeager, it was concluded that oxygen bonding to a single Pt atom would result predominantly in H2 O2 formation, whereas bridge-bonding O2 favors 4e reduction. Comparison of the activation barriers for different reaction intermediates suggests that O2 dissociation does not occur before electron and proton transfer, which greatly diminish the activation barrier for O−O bond splitting. The first electron transfer step was identified as the RDS and its activation energy at the reversible electrode potential was estimated as 0.60 [140] to 0.5 eV [149], which is close to the experimental value of 0.44 eV on Pt(111) in H2 SO4 [150]. Electric field dependences of adsorbates on Pt(111) have been calculated [146]. Lowering the field , causes an increase in the O−O bond length of O− 2ad attracting the molecule to the Pt surface and increasing the charge transfer from Pt to 2π ∗ orbitals of the oxygen molecule. Structural and Particle Size Effects in the ORR An influence of surface crystallography on the kinetics and the mechanism of the ORR has been observed for different materials. For example, the basal plane of graphite was found to be significantly less active than the edge plane. This has been explained by the lack − of adsorption sites for O2 , O− 2 and HO2 on the basal plane. In contrast, on edge planes oxygen can adsorb at the edges of graphene layers [108, 110, 151, 152]. Structural effects for Pt(hkl) have been investigated by Markovic et al. in both acidic and alkaline electrolytes ([13] and references therein, [153]). In Fig. 19, the significant influence of the crystallographic orientation of Pt on the rate of the ORR (which is proportional to the disk current shown) and the amount of H2 O2 formed (for which the ring current is a measure) can be seen. The order of activity of Pt(hkl) in 0.1 M KOH increased in the sequence (100) < (110) < (111) for both oxygen and peroxide reduction. These differences were attributed to the structure sensitivity of hydroxyl anion (OH− ) adsorption on Pt(hkl) and its inhibiting (site-blocking) effect on oxygen kinetics. In recent studies, Wang et al. [154] proposed that along with site-blocking effects, 8.1.1.6.3

References see page 1902

1898

8.1 Electrocatalysis

10 000 20

ID /mA

10 0

0.1 M KOH 50 mV s−1 1600 rpm

−10 −20 0.2

ID / mA

(a)

IR /mA

40

−

HO2

−0.2

100

−0.6

0.0

0.2

−ID / mA

0

0.4 0.6 E/V

0.8

0

30

60

90

120

APt,cat / m2 gPt−1 Fig. 20 ORR specific activities for Pt/C catalysts, polycrystalline Pt (shown at 0 m2 g−1 ) and Pt black (at 5 m2 g−1 ) at 0.9 V vs. RHE in 0.1 M HClO4 at 60 ◦ C as a function of their specific surface areas. Reproduced with permission from Ref. [156].

0.5 1.0 1.5 0.0

(b)

1000

0.0

−0.4

20

iS(0.9V) / µA cm−2Pt

Pt(111) Pt(110) Pt(100)

0.2

0.4

0.6

0.8

1.0

E/V

(a) Cyclic voltammetry of Pt(hkl) in oxygen-free 0.1 M KOH electrolyte in the RRDE assembly. (b) Disk (ID ) and ring (IR ) currents during ORR on Pt(hkl) (ring potential 1.15 V). Insert: reduction of HO− 2 on Pt(hkl) mounted in the RRDE assembly; 0.1 M KOH, 50 mV s−1 , 1600 rpm. Reproduced with permission from Ref. [172]. Fig. 19

anions also exert electronic effects on the ORR. The mechanism of O2 reduction is also affected by adsorbed hydrogen, with the increased formation of peroxide ions in the HUPD region. Adsorbed hydrogen has an inhibiting effect on peroxide reduction, the effect decreasing in the order (111) > (100) (110). Particle size effects were discovered for the ORR on Pt nanoparticles in phosphoric acid many years ago in connection with the development of cathode catalysts for phosphoric acid fuel cells (PAFCs) ([110, 155] and references therein). Similar dependences have been reported more recently also in electrolytes containing weakly adsorbing anions, such as in HClO4 [156]. The specific electrocatalytic activities were found to decrease substantially with increase in the Pt specific surface area (corresponding to a decrease in the particle size) (Fig. 20). Different hypotheses have been proposed to account for this negative particle size effect. According to Kinoshita [157, 158] and Mukerjee [159], the size effect ensues from different fractions of (100), (111) and (110) sites on particles with different sizes. Gasteiger et al. [156]

attributed the negative particle size effect to an increase in the strength of OH adsorption (and thus surface blocking) with decreasing particle size. An alternative viewpoint was presented by Stonehart and Watanabe (see discussion in Ref. [159]), who proposed that the apparent correlation between the catalytic activity of nanoparticles and the particle sizes stems from interparticle diffusive interference between the platinum crystallites. This interpretation is based on the observation that the specific activities of platinum particles of different sizes became identical on different carbon supports when the crystallite separations on the carbon supports were similar. Further studies are needed to understand better the influence of particle sizes on the electrocatalysis of the ORR. 8.1.1.6.4 ORR Electrocatalysis for Low-Temperature Fuel Cells The ORR occurs at the cathode of both low- and high-temperature fuel cells and its sluggish kinetics is greatly responsible for the fact that the cell voltage Vcell is much below the equilibrium value of the H2 /O2 cell Eeq (which depends on the temperature and the gas partial pressures). This is illustrated in Fig. 21. The cell voltage can be expressed as follows (see also Section 8.1.1.2.3): ∗ Vcell = Eeq − |ηORR | − |ηHOR | − Eohmic

(76)

∗ and ηHOR represent the overpotentials at the where ηORR cathode and the anode, respectively. In a fuel cell fed by pure hydrogen ηHOR is small and is often neglected [156]. ∗ ηORR can be separated into the reaction overpotential ηORR and concentration overpotential ηconc . The former is the consequence of the intrinsically sluggish ORR

8.1.1 Fundamentals of Electrocatalysis

1.15

1899

Eeq = 1.169 V

1.10

∆Ecell / V

1.05

IR-free and hconc-free ∆Ecell

1.00

hconc-free ∆Ecell

0.95

Measured ∆Ecell

0.90

50% hconc-free ∆Ecell

hORR

0.85 0.80 0.75

∆Eohmic

0.70 0.65

hconc

0.60 0.55 0.0

0.2

0.4

0.6

0.8

1.0

1.2

1.4

i / A cm−2 Current–voltage characteristic obtained in a 50 cm2 hydrogen–air fuel cell at 80 ◦ C at a total pressure of 150 kPa. The anode and cathode layers consist of ca. 50 wt.% Pt/carbon [0.4/0.4 mg Pt cm−2 (anode/cathode)] and ionomer (ca. 900 EW; ionomer:carbon ratio = 0.8 : 1). Circles, experimental data; triangles, mass transport-free Ecell ; squares, mass transport-free and ohmically corrected Ecell . The current density is referred to the geometric surface area of the MEA. Reproduced with permission from Ref. [156].

Fig. 21

kinetics discussed above, whereas the latter arises from the oxygen mass transport limitations in the catalytic layer. Eohmic comprises the contributions of (i) the contact resistance between the flowfield plates, (ii) the membrane resistance and (iii) the resistance of the catalytic layers and the diffusion layers. Eohmic can be measured directly via either current-interrupt or high-frequency resistance measurements. Since the electrochemical fuel cell efficiency is given by the ratio between Vcell and Eeq , Eq. (76) and Fig. 21 give a guidance to which contributions must be minimized in order to improve the fuel cell efficiency. As is apparent from Fig. 21, at low current densities the difference between the actual cell voltage and the thermodynamic equilibrium value is fully determined by the cathode overpotential (ηORR ). At high current densities, mass transport of oxygen to the active centers and ohmic losses also make significant contributions to the overall losses. Eohmic can be decreased by utilizing polymer membranes with higher ionic conductivities and optimization of the operation conditions. However, ohmic losses in other cell components such as bipolar plates and catalytic layers must also be minimized. ηconc depends hugely on the architecture of the catalytic layers and porosities of catalytic supports utilized. Since the main losses come from ηORR , much effort has been directed towards the development of novel electrode materials for PEMFC applications. Here, the goal is to maximize the electrocatalytic activity while decreasing the amount of precious metals.

Pt alloys (such as PtCr, PtNi, PtCo and PtFe and also ternary alloys) show noticeable enhancement of the ORR activity both in PAFC and PEMFC [156, 159]. Different hypotheses have been offered to account for the enhanced electrocatalytic activity of Pt alloys with transition metals [156]. However, most authors nowadays agree that the enhanced ORR kinetics can be explained by an electronic effect of the transition metals, which results in the inhibition of OH adsorption on Pt [135, 160]. This is supported by in situ EXAFS results [135] and by quantum chemical calculations [161]. Another approach towards the improvement of cathode catalysts for PEMFCs, first proposed by Brankovic et al. [162], utilizes Pt monolayers on foreign metal nanoparticles. This allows (i) the electrocatalytic activity of Pt to be tuned by inducing lattice strain and (ii) the amount of Pt per unit current density to be decreased significantly. Electrocatalysis of the ORR on Pt monolayers supported on nanoparticulate and single-crystal electrodes has been considered extensively by several research groups (see, e.g., Ref. [163] and references therein). For example, in Fig. 22 the electrocatalytic activities of platinum monolayers on Ru(0001), Ir(111), Rh(111), Au(111) and Pd(111) towards the ORR are compared. A strong influence of the substrate on both the reduction current and the amount of H2 O2 produced can be observed. In this context, it should be mentioned that References see page 1902

1900

8.1 Electrocatalysis

18

−4

0.2

0.4

0.6

0.8

1.0

1.2

−3.5

10 8

PtML/Rh(111)

PtML/Au(111) −4.0

6

−3.0

−2.5

Current

−4.5

B.E. of O

−2.0

−5.0 −1.5

ed -eF /eV

10

Kinetic currents (jK ; squares) at 0.8 V for O2 reduction on platinum monolayers supported on different single-crystal surfaces in 0.1 M HClO4 solution and calculated binding energies of atomic oxygen (BEO; filled circles) as functions of calculated d-band center relative to the Fermi level of the respective clean platinum monolayers. The current data for Pt(111) are included for comparison. Reproduced with permission from Ref. [143].

Fig. 23 PtML/Ru(0001) PtML/Ir(111) PtML/Rh(111) PtML/Au(111) PtML/Pd(111)

8

Iring / µA

12

0 −3.5

E / V vs. RHE

(a)

6 4 2 0 0.0

(b)

Pt(111)

4 PtML/Ir(111) PtML/Ru(0001) 2

−6 0.0

−3.0

14

Binding energy of O/eV/O

−2

−2.5

PtML/Pd(111)

16

PtML/Ru(0001) PtML/Ir(111) PtML/Rh(111) PtML/Au(111) Pt(111) PtML/Pd(111)

−j k /mA cm−2

j / mA cm−2

0

0.2

0.4

0.6

0.8

1.0

1.2

E / V vs. RHE

Fig. 22 Disc (a) and ring (b) currents for the ORR at the RRDE on platinum monolayers (PtML) on Ru(0001), Ir(111), Rh(111), Au(111), Pd(111) and for Pt(111) electrode in 0.1 M HClO4 solution. The rotation rate is 1600 rpm and the sweep rate is 20 mV s−1 [50 mV s−1 for Pt(111)]. The ring potential is 1.1 V and the collection efficiency is 0.20. The electrode labelling: from left to right (a) and from top to bottom (b). Reproduced with permission from Ref. [143].

in terms of the development of PEM fuel cells it is necessary to obtain cathode materials catalyzing 4e oxygen reduction, since hydrogen peroxide production not only decreases the efficiency of a cell, but also leads to the degradation of the cell components, in particular polymer membranes. The platinum monolayers on Ru(0001), Rh(111) and Ir(111) are compressed compared with Pt(111), whereas on Au(111) it is stretched by more than 4%. DFT studies have shown that compressive strain tends to downshift the weighted center of the d-band whereas tensile strain has the opposite effect [164]. The experimentally determined kinetic currents (corrected for the mass transport effects) show a volcano-type dependence on the

center of their d-bands as determined by DFT calculations (Fig. 23) [143]. The platinum monolayer supported on Pd(111) is at the top of the volcano curve and shows improved ORR activity over pure Pt(111). The oxygen binding energy to the surface has also been computed and is plotted in Fig. 23. It has been suggested that the volcano-type behavior is determined by two opposite trends: whereas a higher lying d-band center tends to facilitate O−O bond breaking, a lower lying one tends to facilitate bond formation (hydrogen addition). The results suggest that an improvement in the overall fuelcell efficiency can be combined with substantial cost savings that result from using less platinum at the cathode. Similar logics can be applied when discussing the behavior of PtM alloys. Along with Pt and Pt alloys, other materials have been investigated in the ORR, in particular metal oxides [165], macrocyclic N4transition metal chelates [166] and Ru chalcogenide materials [167]. The interested reader should consult the above references. 8.1.1.6.5 Oxygen Evolution Reaction The oxygen evolution reaction [the reverse reaction of (70) in acidic and of (71) in alkaline electrolytes] has been investigated on various electrode materials ([73] and references therein). It occurs in the potential interval where metal surfaces are covered by phase oxides. Participation of Pt oxides in the OER was proposed in early isotope experiments of Rosenthal and Veselovski [168]. They covered the Pt

8.1.1 Fundamentals of Electrocatalysis

electrode surface with surface oxide enriched with 18 O. Mass spectrometric analysis of the oxygen evolved on the anode proved its enrichment with 18 O. Although Pt and its alloys are the best catalysts for the ORR, they are not the most active catalysts for the OER. Their poor catalytic activity towards the OER stems in part from the insulating properties of Pt oxides. The best catalysts for the OER are materials which are not only catalytically active but are also conductive and stable. The field of the OER has been reviewed by Kinoshita [110] and will be discussed only briefly here. The ability of an oxide material to change its valence state readily is considered a prerequisite for their high catalytic activity in the OER. Hence it is postulated [110, 112] that the catalytic activity towards the OER is related to the change in enthalpy when the oxide undergoes a transition from a lower to a higher oxidation state [169]. Those oxides, which are characterized by a low transition enthalpy, form stable oxides and O2 is not readily released. On the other hand, oxides with a very high transition enthalpy will not stabilize the OER intermediates. This explains the volcano-type dependence of the overpotential of the OER vs. the enthalpy change from lower to higher oxide (Fig. 24). Along with the recent development of the regenerative fuel cells, interest in active catalysts for the OER has increased [170]. A regenerative fuel cell is a hydrogen–oxygen cell, which can operate both as a fuel cell and as an electrolyzer. When the cell works as an electrolyzer splitting water into H2 and O2 , the hydrogen gas (and in some applications also the oxygen gas) is stored; on demand, electricity can then be generated from

RuO2

h/ v

0.2

PtO2

0.4

IrO2 MnO2

Co3O4

NiO

Fe3O4

0.6 PbO2 0

−100

−200

−300

∆H °t / kJ/mol−2 Fig. 24 Volcano plot of the overpotential for the OER versus the enthalpy of the lower to higher oxide transition. Open circles correspond to alkaline and closed circles to acidic electrolytes. Reproduced with permission from Ref. [169].

1901

the stored H2 through the fuel cell process. A regenerative fuel cell has a distinct advantage over state-of-the-art battery systems, because power and energy are separated: the energy is directly related to the fuel storage, whereas the rated power depends on the electrode area. This means that by simply increasing reactant storage, without changing the reactor stack(s), it is possible to increase the stored energy. Therefore, the mass advantage of typical advanced batteries loses out to that of regenerative fuel cells when the discharge time is increased beyond a few tens of minutes. Catalysts for the oxygen electrode of a regenerative fuel cell must therefore be active towards both oxygen evolution and oxygen reduction. One of the promising approaches in catalyst development relies on the combinatorial discovery of novel materials. Using this approach, Chen et al. [170] discovered a ternary Pt4.5 Ru4 Ir0.5 catalyst with superior catalytic activity and stability. It was suggested that the addition of Ru to PtIr alloys increases the reaction rate by stabilizing M−O bonds and accelerating the oxidative deprotonation of M−OH groups. Outlook This chapter has introduced some fundamental principles of electrocatalysis and discussed the electrocatalytic reactions that occur at hydrogen and oxygen electrodes in detail. In the first part, emphasis was put on those properties of the electrochemical environment that introduce a qualitative difference to heterogeneous reactions and also on the theoretical foundation of phenomenological approaches that are used in practice to judge the quality of an electrocatalytic material. The second part demonstrated the application of these methods and exemplified at the same time the large variety of experimental methods that are available to study the solid/liquid interface and the diversity of catalyst forms (ranging from single crystals to supported nanoparticles) and materials that may be used. Furthermore, emphasis was put to demonstrate parallels between heterogeneous catalysis and electrocatalysis. The material covered here is just the tip of the iceberg of the vast field of electrocatalysis which has developed rapidly during recent decades. Along with the further development of in situ spectroscopic techniques that allow the electrode of the solid/liquid interface to be monitored in real time and the evolution of ab initio quantum mechanical calculations of the structure of the interface and the kinetics of complex interfacial reactions including adsorption, bond breaking and bond formation, the rapid progress is likely to continue in the following decades. Improvements in the temporal and 8.1.1.7

References see page 1902

1902

8.1 Electrocatalysis

spatial resolution of traditional methods (e.g. vibrational spectroscopic techniques) will go in parallel with the application of new techniques, in particular those based on synchrotron and neutron facilities, which might allow atomic and electronic structures to be monitored under real conditions. Hence an understanding of some electrocatalytic processes on the atomic and molecular level seems to be within reach. During recent decades, electrocatalysis has moved from the investigation of polycrystalline electrodes to welldefined single crystals, thin films and supported metal particles. Until now the interest in Pt and its alloys has been overwhelming; however, emphasis nowadays is shifting towards complex materials, but also to Ib group metals and other non-noble metals. Inducing lattice strain by fabricating metal overlayers on foreign substrates has proven to be a powerful tool to produce materials with desired properties. Further progress in this direction is expected through concerted experimental and theoretical approaches. A better understanding of supported electrocatalysts requires the utilization of well-defined systems allowing control of the particle size, shape and interparticle separations. In this context, of particular interest is the development of templating and nanopatterning approaches. Further progress in practical applications requires the development of complex multifunctional materials which will comprise both metallic and molecular blocks. The catalyst support is nowadays attracting increased attention and it is widely recognized that it affects all vital properties of electrocatalytic materials spanning from the intrinsic catalytic activity to macrokinetics. The range of materials which are utilized as electrocatalytic supports has expanded greatly recently and includes novel carbon materials, e.g. carbon nanotubes, nanofibers, mesoporous carbons with well-defined structures and non-carbonaceous supports, in particular metal oxides with semiconducting properties and electron-conducting polymer materials. A very promising direction is the preparation of catalytic supports with predefined properties in order to enhance metal utilization and catalyst performance in real devices, e.g. fuel cells. Thanks to the mentioned recent advancements, the development of novel catalysts for practical applications is at the edge of moving towards rational design rather than trial and error approaches. The growing technological interest in fuel cells has already brought the catalytic and electrochemical communities closer together and this tendency will gather momentum. Although the evolution of fuel cells, both SOFC and PEMFC, has been really spectacular during recent decades, other applications of heterogeneous electrocatalysis, e.g. in electrosynthesis and the ‘‘electrochemical incineration’’ of organic and inorganic hazards,

has lagged behind but will pick up speed in the coming years. Acknowledgement

The authors would like to express their gratitude to G. Ertl, H. Gasteiger and R. Adzic for careful reading the manuscript and to R. Adzic, H. Gasteiger, B. Hayden, L. A. Kibler, N. M. Markovic and S. Trasatti for kindly offering figures for this chapter. References 1. W. Vielstich, A. Lamm, H. A. Gasteiger (Eds.), Handbook of Fuel Cells. Fundamentals, Technology and Applications, Vol. 2: Electrocatalysis, Wiley, Chichester, 2003, 783 pp. 2. S. McIntosh, R. J. Gorte, Chem. Rev. 2004, 104, 4845. 3. C. Lamy, J.-M. L´eger, S. Srinivasan, in Modern Aspects of Electrochemistry, Vol. 34, B. E. Conway, J. O. ’M. Bockris, R. E. White (Eds.), Kluwer/Plenum Press, New York, 2002, p. 53. 4. S. C. Thomas, X. M. Ren, S. Gottesfeld, P. Zelenay, Electrochim. Acta 2002, 47, 3741. 5. S. Wasmus, A. Kuver, J. Electroanal. Chem. 1999, 461, 14. 6. C. Lamy, A. Lima, V. LeRhun, F. Delime, C. Coutanceau, J. M. L´eger, J. Power Sources 2002, 105, 283. 7. S. Trasatti, in Interfacial Electrochemistry: Theory, Experiment and Applications, A. Wieckowski (Ed.), Marcel Dekker, New York, 1999, p. 769. 8. A. T. Kuhn, (Ed.), Industrial Electrochemical Processes, Elsevier, 1971, pp. 632. 9. J. G. daSilva, M. O. F. Goulart, M. Navarro, Tetrahedron 1999, 55, 7405. 10. S. Kaneco, K. Iiba, N. Hiei, K. Otha, T. Mizuno, T. Suzuki, Electrochim. Acta 1999, 44, 4701. 11. D. M. Kolb, Angew. Chem. Int. Ed. 2001, 40, 1162. 12. A. Wieckowski (Ed.), Interfacial Electrochemistry: Theory, Experiment and Applications, Marcel Dekker, New York, 1999, 996 pp. 13. N. M. Markovic, P. N. Ross, Surf. Sci. Rep. 2002, 45, 121. 14. H. E. Hoster, H. A. Gasteiger, in Handbook of Fuel Cells. Fundamentals, Technology and Applications, Vol. 2. Electrocatalysis, W. Vielstich, A. Lamm, H. A. Gasteiger (Eds.), Wiley, Chichester, 2003, p. 236. 15. A. J. Bard, M. Stratmann, P. R. Unwin (Eds.), Instrumentation and Electroanalytical Chemistry. Encyclopedia of Electrochemistry, Vol. 3, Wiley-VCH, New York, 2003, 689 pp. 16. M. J. Weaver, Top. Catal. 1999, 8, 65. 17. C. Korzeniewski, Crit. Rev. Anal. Chem. 1997, 27, 81. 18. A. Rodes, J. M. Perez, A. Aldaz, in Handbook of Fuel Cells. Fundamentals, Technology and Applications, Vol. 2: Electrocatalysis, W. Vielstich, A. Lamm, H. A. Gasteiger (Eds.), Wiley, Chichester, 2003, p. 191. 19. A. E. Russell, A. Rose, Chem. Rev. 2004, 104, 4613. 20. R. R. Adzic, X. Wang, B. M. Ocko, J. McBreen, in Handbook of Fuel Cells. Fundamentals, Technology and Applications, Vol. 2: Electrocatalysis, W. Vielstich, A. Lamm, H. A. Gasteiger (Eds.), Wiley, Chichester, 2003, p. 279. 21. N. J. Tao, C. Z. Li, H. X. He, J. Electroanal. Chem. 2000, 492, 81. 22. K. Itaya, Prog. Surf. Sci. 1998, 58, 121.

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1903

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8.1.2

Electrochemical Modification of Catalytic Activity Constantinos G. Vayenas∗ , Alexandros Katsaounis, Susanne Brosda, and Ahmad Hammad

8.1.2.1

Introduction

8.1.2.1.1 Catalytic and Electrocatalytic Kinetics Heterogeneous catalysis and aqueous or solid electrochemistry have been treated traditionally as different branches of physical chemistry, yet similar concepts are used to model their kinetics [1–4] and similar surface science techniques are used to investigate their fundamental aspects at the molecular level [1–8]. The growing technological interest in fuel cells, both high-temperature solid oxide fuel cells (SOFCs) and low-temperature polymeric electrolyte membrane (PEM) fuel cells, has brought the catalytic and electrochemical communities closer, as the merits of catalysis in designing and operating efficient anodes and cathodes is being more widely recognized [9–11]. An important additional operating parameter in electrochemical (electrocatalytic, i.e. net charge transfer) vs. catalytic (no net charge transfer) kinetics is the electrical potential dependence of the electrochemical rate, yet in recent years its has been shown that for electrochemically promoted catalysts (i.e. catalysts in contact with a solid electrolyte [12–18]) the catalytic rate also depends dramatically on catalyst potential, similarly to the electrochemical rate. It has been known for some years that electrochemistry can be used to activate and tune precisely heterogeneous catalytic processes when the catalyst is in contact with ionic or mixed ionic–electronic conductors as supports (e.g. YSZ, TiO2 , CeO2 ) [12–39]. These materials act as specific anionic or cationic conductors and, depending on their composition, have catalytically useful electrical conductivities at temperatures of 25–1000 ◦ C. Within this temperature range, which covers practically all heterogeneous catalytic reactions, solid electrolytes can be used as reversible in situ promoter donors and/or poison acceptors to affect the catalytic activity and selectivity of metals deposited on solid electrolytes in a very pronounced, reversible and, to some extent, predictable manner. This is accomplished by applying a potential (±1–2 V) between the conductive catalyst film and a second metal film (counter electrode) also interfaced with the solid electrolyte and thus causing a controlled migration (backspillover) of promoting ions from the solid electrolyte on to the catalyst surface. References see page 1934 ∗ Corresponding author.

1906

8.1 Electrocatalysis

The pronounced reversible promotional phenomena observed on varying the electrical potential of metal catalysts interfaced with solid electrolytes are known as ‘‘non-Faradaic electrochemical modification of catalytic activity’’ (NEMCA effect) [15, 16, 34] or ‘‘electrochemical promotion’’ (EP) [16, 23, 24] or ‘‘in situ controlled promotion’’ (ICP) [24]. These three terms are used interchangeably in this chapter as they refer to the same phenomena. Detailed [15–18] and shorter [40–43] reviews on electrochemical promotion have been published previously. The underlying electrochemical and catalytic principles are discussed in detail in Ref. [16] and the importance of NEMCA in catalysis and electrochemistry in Refs. [44–46]. In addition to the potential technological applications of electrochemical modification of catalytic activity, the ability of solid electrolytes to dose reversibly, precisely and in situ catalyst surfaces with promoters, by ‘‘knob-turn’’ variation of the catalyst potential, provides a unique opportunity for the systematic study of the role of promoters and poisons in heterogeneous catalysis. 8.1.2.1.2 Solid Electrolytes Michael Faraday was first to observe in 1834 that solid PbF2 , when heated at 500 ◦ C, becomes an electrical conductor. It took almost a century to explain this observation and establish that PbF2 is a F− ion conductor. In the meantime, other solid electrolytes such as AgI, an Ag+ conductor, had been discovered by Tubandt and Strock and it soon became apparent that ions can diffuse as rapidly in some solids as in aqueous salt solutions. The atomistic interpretation of ionic conduction in solids was largely established by the pioneering work of Joff´e, Frenkel, Wagner and Schottky in the 1920s and early 1930s [16]. These studies established that ion conduction can take place either by hopping of ions through a series of interstitial sites (Frenkel disorder) or by hopping of vacancies among lattice positions (Schottky disorder). Today, the term solid electrolyte or fast ionic conductor or, sometimes, superionic conductor is used to describe solid materials whose conductivity is wholly due to ionic displacement. Mixed conductors exhibit both ionic and electronic conductivity. Solid electrolytes range from hard, refractory materials, such as 8 mol% Y2 O3 -stabilized ZrO2 (YSZ) or sodium β-Al2 O3 (Na1+x Al11 O17+x/2 ), to soft proton exchange polymeric membranes such as Nafion and include compounds that are stoichiometric (AgI), non-stoichiometric (sodium β-Al2 O3 ) or doped (YSZ). The preparation, properties and some applications of solid electrolytes have been discussed in a number of books and reviews. The main commercial application

of solid electrolytes is in gas sensors. Another emerging application is in solid oxide fuel cells. The classification of solid electrolytes is usually based on the ion mainly responsible for the conductivity. There exist: (i) Oxygen ion conductors: these are solid solutions of divalent and trivalent metal oxides (e.g. Y2 O3 , Yb2 O3 , CaO) in quadrivalent metal oxides (e.g. ZrO2 , ThO2 , CeO2 ). Calcia- or yttria-stabilized zirconia (YSZ), containing 5–15 mol% CaO or 6–10 mol% Y2 O3 in ZrO2 , is widely used in oxygen sensors, normally in the temperature range 400–1200 ◦ C. (ii) H+ and Li+ conductors: several polymeric solid electrolytes belong here. Of particular importance are the proton exchange membranes (PEMs), such as Nafion, which is a copolymer of polytetrafluoroethylene and poly(sulfonyl fluoride) containing pendant sulfonic acid groups, which exhibit substantial conductivity at room temperature [46]. High cationic conductivity is also exhibited by several alkali metal salt solutions in poly(ethylene oxide). Proton conduction is also exhibited by CsHSO4 [20], by H+ -substituted β  -Al2 O3 and by SrCeO3 -based compounds. (iii) Na+ conductors: these are β- and β  -aluminas which are non-stoichiometric compounds corresponding to Na1+x Al11 O17+x/2 (0.15 ≤ x ≤ 0.3) and Na1+x Mx Al11−x O17 , respectively, where M is a divalent metal (e.g. Mg2+ , Ni2+ , Zn2+ ). They exhibit high conductivity in the range 150–300 ◦ C. (iv) K+ , Cs+ , Rb+ , Tl+ conductors: they are substituted β and β  -Al2 O3 and are conductive in the range 200–400 ◦ C. (v) Ag+ conductors, e.g. α-AgI, RbAg4 I5 and Ag2 HgI4 are conductive in the range 150–350 ◦ C. (vi) Cu+ conductors, e.g. Cu2 Se and KCu4 I5 are conductive in the range 250–400 ◦ C. (vii) F− conductors, e.g. PbF2 and CaF2 are conductive above 500 and 600 ◦ C, respectively. Detailed information about the conductivity of solid electrolytes can be found in Refs. [16, 47]. In general the temperature dependence of the ionic conductivity σ can be described by the semiempirical equation:
 σ  −EA 0 σ = exp (1) T kb T where σ0 is a function of the ionic charge, the concentration of the mobile ions and the frequency with which these ions attempt to move to a neighboring site (attempt frequency); EA is the activation energy for defect motion and kb is the Boltzmann constant. The activation energy EA is usually on the order of

8.1.2 Electrochemical Modification of Catalytic Activity

0.5–2 eV. The minimum necessary ionic conductivity value of a solid electrolyte for practical fuel cell applications [46] is 0.1–1 ohm−1 cm−1 . This places very stringent restrictions on the choice of material and operating temperature. For catalytic (promotional) and sensor applications, however, much lower conductivity values (∼10−4 ohm−1 cm−1 ) are usually sufficient. This permits the use of a large variety of solid electrolytes over a very wide temperature range. Electrochemical promotion (EP) studies have so far utilized [16]: O2−

conductor, (i) Yttria-stabilized-zirconia (YSZ), an at temperatures of 280–650 ◦ C. (ii) β  -Al2 O3 , an Na+ conductor at temperatures 130–400 ◦ C. (iii) CsHSO4 and Nafion, which are proton conductors at temperatures of 150 and 25 ◦ C, respectively. (iv) CaF2 , an F− conductor, at temperatures of 550–700 ◦ C. (v) Aqueous KOH solutions (0.01–0.2 M) at temperatures of 25–60 ◦ C.

pO2,W

Working electrode (W)

O

Solid electrolyte

U WR

pO2,R

(a)

Reference electrode (R)

Reactants Catalyst film (W) G/P

Solid electrolyte

(b)

Counter electrode (C)

Reactants

Products

W Reactants C

8.1.2.1.3

where F is Faraday’s constant (96 460 C mol−1 ) and pO2 ,W and pO2 ,R are the oxygen partial pressures over the two electrodes. The superscript 0 designates hereafter opencircuit conditions, i.e. there is no current (I = 0) flowing between the two electrodes. The Nernst equation, Eq. (2), is valid provided that there is equilibrium between gaseous oxygen and oxygen, O(tpb), adsorbed at the three-phase boundaries (tpb), solid electrolyte–metal–gas. It is also necessary that the net charge transfer (electrocatalytic) reaction at the tpb is 2− −−  O(a) + 2e− −− − − O (YSZ)

(3)

i.e. that there is no interference from other gases, e.g. H2 , CO, which may react with O2− (YSZ) at the tpb. Carl Wagner first proposed the use of such galvanic cells in heterogeneous catalysis, to measure in situ the thermodynamic activity of oxygen O(a) adsorbed on metal electrodes during catalytic reactions [48]. This led to the

UWR

Reference electrode (R)

W

Solid Electrolyte Potentiometry (SEP) When a solid electrolyte component is interfaced with two electronically conducting (e.g. metal) films (electrodes), a solid electrolyte galvanic cell is formed (Fig. 1a). Cells of this type with YSZ solid electrolyte are used as oxygen 0 which develops sensors. The potential difference UWR spontaneously between the two electrodes (W and R designate working and reference electrodes, respectively) is given by

 RT pO2 ,W 0 ln (2) UWR = 4F pO2 ,R

1907

C

R

R

Auxiliary gas

(c)

Fuel-cell type configuration

(d)

Single-pellet type configuration

Electrode configurations for SEP (a) and for PPR or NEMCA (b) studies. The latter can be carried out using the fuel-cell type configuration (c) or the single pellet-type configuration (d).

Fig. 1

technique of solid electrolyte potentiometry (SEP) [16, 49]. In this technique, the working electrode W (e.g. Pt) is exposed to the reactive gas mixture (e.g. C2 H4 plus O2 ) and also serves as the catalyst for a catalytic reaction, e.g. −−  C2 H4 + 3O2 −− − − 2CO2 + H2 O

(4)

0 is related to the The measured potential difference UWR oxygen activity, aO , on the catalyst surface via [4, 13–16, 33, 34]
  2  aO RT 0 UWR = (5) ln 4F pO2 ,R

which is again derived on the basis of the equilibrium (3). The SEP technique, used in conjunction with kinetic studies, is a useful tool for mechanistic investigations, and is particularly suitable for the study of oscillatory reactions [16]. The limitations of Eq. (5) References see page 1934

1908

8.1 Electrocatalysis

together with detailed reviews of the SEP literature can be found elsewhere [16]. Today, it is well established both theoretically [16] and experimentally [14, 16] that SEP with metal catalyst electrodes is a work function () measuring technique: 0 = W − R eUWR

(6)

where W and R are the (average [15, 16]) work functions  of the gas-exposed surfaces of the working (W) and reference (R) electrodes. Equation (6) is more general than Eq. (5) as it does not depend on the nature of the solid electrolyte and does not require the establishment of any specific charge-transfer equilibrium [e.g. Eq. (3)] at the tpb [16]. It shows that solid electrolyte galvanic cells are work function probes for their gas-exposed, i.e. catalytically active, electrode surfaces. 8.1.2.1.4 Electrocatalytic Operation of Solid Electrolyte Cells Solid electrolyte cells based on YSZ can be used as fuel cells for electrical power generation [16, 46]. One porous electrode (cathode) is exposed to air and acts as an electrocatalyst for the reduction of O2 :

1 2− −  O2 (g) + 2e− − −− − − O (YSZ) 2

(7)

The other electrode (anode) is exposed to the fuel (e.g. H2 or CH4 ) and acts as an electrocatalyst for the anodic fuel oxidation, e.g. − −−  H2 (g) + O2− (YSZ) −− − − H2 O + 2e

(8)

State-of-the art solid oxide fuel cells operate at 900–1000 ◦ C and utilize mixed conductor perovskite cathodes (La1−x Srx MnO3 ) and Ni-YSZ cermet anodes. Similar cells with appropriate electrocatalytic anodes can be used for ‘‘chemical cogeneration’’ [50, 51], i.e. for the simultaneous production of electrical power and industrial chemicals. This mode of operation, which combines the concepts of a fuel cell and of a catalytic reactor, was first demonstrated for the case of NH3 oxidation to NO using Pt and Pt−Rh anodes [50, 51]: Ammonia is supplied via the gas phase and oxygen via the solid electrolyte as O2− . Several other reactions have been investigated, including the oxidation of H2 S to SO2 [52], of CH3 OH to H2 CO [53] and of methane to ethylene [54]. In the last case, it was found that ethylene yields of up to 85% can be obtained in a gas-recycle solid electrolyte cell reactor–separator using a Ag−Sm2 O3 anode and a molecular sieve adsorbent [54]. A simple rule which has emerged from chemical cogeneration studies [50–54] is that suitable electrocatalysts for the electrocatalytic (i.e. net charge transfer) anodic reaction can be chosen on the basis of proven catalysts for

the corresponding catalytic reaction (e.g. Pt−Rh for NH3 oxidation to NO, Ag for CH3 OH oxidation to H2 CO, etc.). In addition to chemical cogeneration studies, where the anodic and cathodic reactions are driven spontaneously by the cell-generated voltage, several other electrocatalytic reactions have been investigated in solid electrolyte cells via external voltage application. These include the dissociation of H2 O and NO, the hydrogeneration of CO and the partial oxidation of CH4 , as reviewed elsewhere [16]. 8.1.2.1.5 Electrochemical Promotion of Catalysis The idea of using an electronically conductive metal or metal oxide porous film simultaneously as a catalyst and as an electrode can be traced to the last work of Wagner [48]. This led not only to the passive technique of solid electrolyte potentiometry (SEP) for measuring in situ the chemical potential of oxygen on catalyst electrodes [49], but also, much more importantly, to the discovery of the effect of the electrochemical promotion of catalysis (EPOC) or non-Faradaic electrochemical modification of catalytic activity (NEMCA effect) [12–39]. The basic phenomenology of this effect when using O2− -, Na+ - and H+ -conducting solid electrolytes is shown in Figs. 2–5. The (usually porous) metal catalyst electrode, typically 2–5 µm thick, is deposited on the solid electrolyte and, under open-circuit conditions (I = 0, no electrochemical rate), produces a catalytic rate r0 for, e.g., C2 H4 oxidation [34] (Figs. 2 and 4) or CO oxidation [36] (Fig. 3). Application of an electrical current, I , or potential (±2 V) between the catalyst and a counter electrode causes very pronounced and non-Faradaic (i.e. r I /2F ) alterations to the catalytic rate, r, and, fairly often, to the product selectivity (e.g. Fig. 5) [15, 16, 28]. The rate of the catalytic reaction, r, can become up to 200 times larger than the open-circuit rate, r0 , and up to 3 × 105 times larger than the Faradaic rate (I /2F for O2− , −I /F for Na+ and H+ ) of ion supply (or removal) to (or from) the catalyst electrode [15, 16]. The Faradaic efficiency, , defined by

≡

r(catalytic) (I /nF )

(9)

where n is the ion charge, can thus reach values up to 3 × 105 or down to −104 [16]. Electrocatalysis is limited to || ≤ 1 and this is the main distinguishing feature of electrocatalysis and electrochemical promotion. Up to 2001 [16], more than 70 different catalytic reactions (oxidations, hydrogenations, dehydrogenations, isomerizations, decompositions) had been electrochemically promoted on Pt, Pd, Rh, Ag, Au, Ni, IrO2 , RuO2 catalysts deposited on O2− (YSZ), Na+ (β  -Al2 O3 ), H+ (CaZr0.9 In0.1 O3−α , Nafion), F− (CaF2 ), aqueous [32, 37],

8.1.2 Electrochemical Modification of Catalytic Activity

Electrochemical oxidation rate re = I /nF

Catalytic oxidation rate r current induced rate change ∆r >>I /nF

C2H4 + 6O2−→

C2H4 + 3O2→ 2CO2 + 2H2O

−

12e + 2CO2 + 2H2O I

1909

e−

Catalyst-electrode

U WR U

e−

Solid electrolyte

O 2− ↑ O 2− ↑ O 2− ↑

Counter and reference electrodes

O2(g)

(a) 50

100

I = + 1µA

I=0

800

40

TOF/ s−1

60

40

20

0

r / 10−8 mol O s−1

600 80

30

400

20

200

10

0

r0 + (I /2F ) × 104

Catalyst potential U WR / mV

I=0

−200

r0 0 0

τ

2FNG /I

(b)

20

40

110

130

Time / min

(a) Basic experimental setup and operating principle of electrochemical promotion with O2− conducting supports. (b) Catalytic rate, r, and turnover frequency, TOF, response of C2 H4 oxidation on Pt deposited on YSZ, an O2− conductor, upon step changes in applied current [13, 14, 16, 34]. T = 370 ◦ C, pO2 = 4.6 kPa, pC2 H4 = 0.36 kPa. Also shown (dashed line) is the catalyst electrode potential, UWR , response with respect to the reference (R) electrode. The catalytic rate increase, r, is 25 times larger than the rate, r0 , before current application and 74 000 times larger than the rate, I/2F, of O2− supply to the catalyst electrode. NG is the Pt/gas interface surface area, in mol Pt, and TOF is the catalytic turnover frequency (mol O reacting per surface Pt mol per s). (Reprinted with permission from Kluwer/Plenum Press [16].) Fig. 2

molten salt [30] and mixed ionic–electronic (TiO2 [38], CeO2 [39]) conductors. Clearly, EPOC is not limited to any particular class of conductive catalyst, catalytic reaction or ionic support. The first commercial electrochemically promoted soot combustion units were produced by Dinex in Denmark [16, 55]. Selectivity Modification and Promotional Effects on Chemisorption

8.1.2.2

8.1.2.2.1 Selectivity Modification One of the most promising applications of in situ-controlled promotion is in product selectivity modification. Two examples regarding the epoxidation of ethylene on Ag are discussed in Ref. [16]. In the former case, the Ag film is supported

on YSZ. For UWR > 0 V, ethylene oxide and CO2 are the only products and the selectivity to ethylene oxide is 55%. Decreasing the catalyst potential to UWR = −0.6 V causes a dramatic shift in selectivity. The selectivity to ethylene oxide vanishes and acetaldehyde becomes the dominant product with a selectivity of 55% [16]. In the latter case [16], the Ag catalyst is supported on β  -Al2 O3 and traces of C2 H4 Cl2 ‘‘moderator’’ are also added to the gas phase [16]. Ethylene oxide and CO2 are the only products. It was found that the combined effect of the partial pressure of C2 H4 Cl2 and of the catalyst potential leads to a selectivity to ethylene oxide of 88% [16]. This is one of the highest values reported for the epoxidation of ethylene. References see page 1934

1910

8.1 Electrocatalysis

Electrocatalytic reaction

Catalytic reaction

Na+(b″− Al2O3) + e− → Na(Pt)

CO + ½ O2 → CO2

e−

Pt

I

U

Na+ ↑ Na+ ↑ Na+ ↑

e−

Au

(a) 0.5

UWR / mV

0

0.02

U WR Au

I = −20µA 0.04

I=0 0.06

qNa

0

−0.5

−1

7.5

5

r / 10−7 mol s−1

10

2.5 −1

0

1

2

3

4

5

6

t / min

(b)

(a) Basic experimental setup and operating principle of electrochemical promotion with Na+ conducting supports. (b) Catalytic rate, r, response of CO oxidation on Pt deposited on β  -Al2 O3 , an Na+ conductor upon step changes in applied current [36]. Also shown is catalyst potential, UWR , response. T = 350 ◦ C, pCO = 2 kPa, pO2 = 2 kPa. Note that the rate passes through a maximum at θNa = 0.015, as the reaction rate of CO oxidation on Pt exhibits volcano-type behavior with respect to the catalyst potential and work function [36]. Upon current interruption (I = 0) the rate, r, and potential, UWR , do not return to their initial values. This is accomplished only by imposing potentiostatically the initial UWR value [36]. In this experiment the potentiostat, previously used to control UWR , is disconnected at t = −1 min and then at t = 0 the galvanostat is used to apply a constant current. Dashed curves correspond to rate and UWR transients obtained with different previously imposed UWR values. Note that the Na coverage (inset axis) always determines the r and UWR values during the transients [36].

Fig. 3

Promotional Effects on Chemisorption As noted previously, the strength of chemisorptive bonds can be varied in situ via controlled promotion. Figure 6 shows the effect of the catalyst potential UWR and work function  on the TPD peak desorption temperature Tp and on the binding strength Ed of oxygen dissociatively chemisorbed on Pt supported on YSZ [60–63]. Increasing eUWR and  by 1 eV causes a 150 ◦ C decrease in Tp and a 1 eV decrease in Ed . The latter is computed by varying the heating rate β via the modified Redhead equation of Falconer and Madix (p. 231 in [16]): 8.1.2.2.2



β ln Tp2





Rνn C0−1 = ln Ed

 −

Ed R




1 Tp

 (10)

where β is the heating rate, νn is the pre-exponential factor and C0 is the initial coverage. It is important to note that Ed decreases linearly with  with a slope of −1, in excellent agreement with the observed decrease in activation energy E with  in the Pt-catalyzed oxidation of C2 H4 and CH4 , where cleavage of the metal–oxygen bond is rate limiting [16]. The effect of UWR and  on the kinetics of oxygen adsorption and desorption on Ag or Au deposited on YSZ has also been investigated [16]. It was found (Fig. 6) that increasing  again causes a linear decrease in Ed with slopes of −1 and −4, respectively (Fig. 6), in very good semiquantitative agreement with rigorous ab initio quantum mechanical calculations [70, 71]. These results establish that increasing/decreasing the work function  causes a decrease/increase in

8.1.2 Electrochemical Modification of Catalytic Activity

Electrochemical reaction H+(CaZr0.9In0.1O3-α) + O(Pt) + e− → OH(Pt)

U

e−

(a)

Catalytic reaction C2H4 + 3O2 → 2CO2 + 2H2O Pt

e−

I

H+ ↑ H+ ↑ H+ ↑ Au

U WR Au

50

1

I=0

I=0

I = −3 µA

T = 405°C po2 = 5.05 kPa pC2H4 = 1.05 kPa

12 40

30 −0.5

20

−1

4

ro = 14.7 × 10−8mol s−1 ∆r = 27.5 × 10−8mol s−1

10

−1.5

(−I /2F ) = 1.55 × 10−11mol s−1 = 17700 r = 2.87 −2

2

0

U WR /V

6

0.5

0

r / 10−8 mol O s−1

TOF/s−1

10

8

1911

0

0t=2

(b)

20

40

t /min

(a) Basic experimental setup and operating principle of electrochemical promotion using an H+ conductor during C2 H4 oxidation on Pt deposited on CaZr0.9 In0.1 O3−α [35]. (b) Catalytic rate, r, catalytic turnover frequency, TOF, and catalyst potential response to step changes in applied current. The increase in O consumption, r, is 17 700 times larger than that anticipated from Faraday’s law and the corresponding rate, −I/F, of proton transfer to the Pt catalyst.

Fig. 4

the chemisorptive bond strength of electron acceptor adsorbates such as chemisorbed atomic oxygen and an increase/decrease in the chemisorptive bond strength of electron donor adsorbates such as benzene [16]. This is a key observation in the interpretation of in situ-controlled promotional phenomena. Basic Questions After it became apparent in the early 1990s [56] that electrochemical promotion is a general effect at the interface of catalysis and electrochemistry, several important questions were raised, which can be summarized as follows [75]: 8.1.2.2.3

What is the molecular origin of electrochemical promotion and how does it relate to (i) electrocatalysis (ii) classical (or chemical, conventional) promotion (where the promoter is added to the catalyst ex situ, i.e. during catalyst preparation), and (iii) metal–support interactions?

The last part of the question became relevant because of the following two discoveries: 1. That mixed electronic–ionic conducting supports, such as TiO2 [38] or CeO2 [39], could replace YSZ in inducing NEMCA, a very noteworthy observation in view of the fact that TiO2 and CeO2 are relatively common conventional catalyst supports and their ionic conductivity is, at best, only 3% of their electronic n-type conductivity [38, 39]. 2. That just short-circuiting the catalyst and counter electrode and taking advantage of the potential difference spontaneously generated, under opencircuit, between the catalyst and counter electrode due to their different activities for the catalytic reaction, just as in the case of single-chamber fuel cells [16], was sufficient to induce NEMCA at least References see page 1934

1912

8.1 Electrocatalysis

1.4

125 30 wt% Pd/C rcis rtrans rbutane I/ mA

1.2

100

75

0.8

I / mA

r /10−6 mol/s

1

0.6

50

0.4

(i) that a power source (galvanostat or potentiostat) is not necessary to induce NEMCA and (ii) that if the support has both electronic and ionic conductivity (e.g. TiO2 ), then NEMCA is induced even without external (via a wire) short-circuiting of the catalyst and counter electrode. The shortcircuiting is internal (Fig. 7b) and the Pt/TiO2 catalyst is electrochemically promoted without any wire attached to it, thus giving macroscopically all the symptoms of a metal–support interaction.

25

0

0

0.1

0.2

0.3

0

0.4

Cell voltage / V

(a) 100

80

Cis-butene Trans-butene Butene

50

60 30 40

20

20

(b)

Faradaic efficiency,

% Selectivity

40

0

Origin of Electrochemical Promotion It took several years and the use of (i) numerous surface science techniques, including XPS [16, 23, 55–59], UPS [59], TPD [60–63] (Figs. 8 and 9), PEEM [64] and STM [65, 67] (Figs. 10 and 11); (ii) several more conventional catalytic techniques, including rate transient analysis [16] and work function measurements [13, 16, 66]; (iii) electrochemical techniques, including cyclic voltammetry [16, 60] (Fig. 8) and AC impedance spectroscopy [68, 69]; and (iv) theoretical ab initio quantum mechanical calculations [70, 71], to understand fully the origin of electrochemical promotion. All these techniques have provided a unanimous answer to the problem: 8.1.2.3

0.2

10

0

0.05

0.1

0.15

0.2

0

Cell voltage / V

(a) Electrochemical promotion of an isomerization reaction [28]. Steady-state effect of cell potential on the cell current, I, and on the rates of formation of cis-2-butene, trans-2-butene and butane produced from 1-butene supplied over a dispersed Pd/C catalyst electrode deposited on Nafion, an H+ conductor at room temperature [28]. Reprinted with permission from the American Chemical Society. (b) Corresponding effect of cell potential on the selectivities to cis-2-butene, trans-2-butene and butane and on the apparent Faradaic efficiency, , defined as rtotal /(I/F). Thus, each proton catalyzes the isomerization of roughly 50 molecules of 1-butene to cis- and trans–2-butene [28]. Fig. 5

for electrophobic reactions, i.e. catalytic reactions where the rate increases with increasing potential or with O2− supply to the catalyst (Fig. 7a). This key experiment of Cavalca et al. [25] was the first ‘‘wireless’’ or, more precisely, ‘‘self-driven’’ NEMCA experiment. It showed:

Electrochemical promotion is due to the current or potential-controlled electrocatalytic (Faradaic) introduction of promoting species (e.g. O δ− , Na δ+ ) from the solid electrolyte to the catalyst/gas interface where an, overall neutral, double layer is formed. The density of this double layer (and the field strength in it) varies as the applied potential is varied and this affects both the work function of the surface and the chemisorptive bond strength of reactants and intermediates, thus causing dramatic and reversible alterations in catalytic rate.

In case the promoting species can also participate in the catalytic reaction (e.g. Oδ− originating from YSZ [58, 76] or TiO2 [38], which is distinct from chemisorbed oxygen originating from gas phase O2 ), then it acts as a sacrificial promoter, i.e. it promotes the catalytic reaction (via repulsive or attractive lateral interactions), but it also becomes consumed at a rate which is  times smaller than the rate of consumption of the catalytic reactant, e.g. atomic O originating from the gas phase [16] (Fig. 8). The concept of the sacrificial promoter, also discussed below, is the key to understanding electrochemical promotion with O2− conductors. This molecular mechanism is unanimously supported by all the above surface science, catalytic and electrochemical techniques. A combination of the results of any two or three of them would have sufficed to put together the puzzle. However, each of them had something new to offer, some new facet of the surface chemistry to reveal.

8.1.2 Electrochemical Modification of Catalytic Activity

1913

0.27V 0.4V 0.66V 0.58V 0.49V

∆Φ/eV 0.19V

0.14V

3

0.4

0.8

1.2

2.5

Ed /eV

In(b /T 2p)

−8.5

0.06V

−8

0

−9

2

−4

−1

Pt Au

−9.5

1.5 Ag

−1

−10 1.2 (a)

1.3

1.4

1.5

1

1.6

(1/Tp) /10−3K

0

0.4

0.8

1.2

eU *WR /eV

(b)

4

Ed, De (Pt− O) / eV

Stark SCF

3.5

−0.5

3

Pt25 − O 2.5 −2 (c)

−1

0

1

2

Work function change/eV

Oxygen chemisorption on polycrystalline Pt, Ag and Au films deposited on YSZ at low (θO < 0.1) coverages [16, 60–63]. (a) Redhead plots obtained at various imposed  values for O chemisorption on Pt. (b) Effect of  on the adsorption enthalpy of O on Pt, Ag and Au. (c) Effect of  on the adsorption enthalpy of O on Pt from the cluster quantum mechanical calculations of Pacchioni et al. [70]. SCF, self-consistent field. Note that (i) upon considering only electrostatic (Stark) effects a straight line is obtained [70] and (ii) the electrostatic (Stark) effect is dominant, as only minor deviations are obtained upon solving the full SCF equations.

Fig. 6

As Pritchard had correctly predicted in his 1990 Nature Editorial on NEMCA [72]: ‘‘The strong long-range effect implied by the correlation of work function change with activation energy change found by Vayenas et al. [14] in the presence of electrochemically induced promotion is particularly intriguing. So too is the nature of the electrochemically induced oxygen species that is believed to cause the increase in work function and catalytic promotion, yet which is less reactive than the adsorbed oxygen reactant that covers most of the surface. There is clearly much surface chemistry to be explored and it will be interesting to see how general the work function effect proves to be. In any case, the ability to vary the concentration of promoters by electrochemical control while under reaction conditions is a valuable development in catalytic research and one can expect it to

be rapidly exploited in conjunction with other in situ techniques of surface analysis.’’

Fifteen years and hundreds of publications later [16], one might only wish to change ‘‘that is believed to cause’’ to ‘‘which causes’’ in the above eloquent, concise and almost prophetic description of the origin of electrochemical promotion by Pritchard [72]. Several questions, of course, still remain open, such as the exact nature of the ‘‘permanent NEMCA’’ effect discovered by Nicole and Comninellis [73], which could have important applications in the scientific preparation References see page 1934

1914

8.1 Electrocatalysis

e− Pt

Ag CH3OH + 3/2O2

r °Ag

O2− O2− O2− YSZ

CO2 + 2H2O

Ag

CH3OH + 3/2O2 CH3OH + 3/2O2

rAg

r °Pt

O2− O2−

Pt CH3OH + 3/2O2

rPt

YSZ CO2 + 2H2O

O2− O2−

O2−

CO2 + 2HO2

O2−

CO2 + 2H2O

O2−

r Pt ° >> r Ag °

rPt >> r Pt °

UCW UCW

0.4 V mO2

mO2 mO2−

mO2− (a) e− CH3OH + 3/2O2

rAg

O2− e−

CH3OH + 3/2O2

rPt

O2− CO2 + 2H2O

e− O

(b)

Ag

CO2 + 2H2O 2−

TiO2

Pt

(a) Self-driven electrochemical promotion of a Pt catalyst for CH3 OH oxidation to CO2 , an electrophobic reaction [25] observed upon short-circuiting (UCW = 0) the Pt catalyst electrode with the Ag counter electrode deposited on YSZ at 380 ◦ C [25]. See text for discussion. (b) Principle of self-driven electrochemical promotion of the same catalytic reaction on a Pt catalyst deposited on a mixed electronic–ionic conductor such as TiO2 [16]. The Pt catalyst and Ag counter electrode are internally short-circuited via the electronic conduction of the support. Fig. 7

of supported catalysts [74], but the basic phenomenology of NEMCA is today not only well understood, but to a large extent predictable [16]. 8.1.2.3.1 Magnitude of Electrochemical Promotion and The magnitude the Sacrificial Promoter Mechanism of electrochemical promotion can be assessed from Figs. 2–5 and 12, which provide examples of two electrophobic (∂r/∂UWR > 0) catalysis systems (Figs. 2 and 12), of two electrophilic (∂r/∂UWR < 0) catalytic systems (Figs. 4 and 5) and one volcano (Fig. 3) catalytic system (Table 1). This is discussed further in Section 8.1.2.4. It has been shown recently [16, 77–79] that the catalytic rate dependence on catalyst potential or work

function [14, 16] can be predicted on the basis of the rate dependence on the partial pressures of the electron donor and electron acceptor reactant. These rigorous electrochemical and classical promotional rules appear to have no exceptions in the published electropromotion [16, 77–79] and classical promotion literature [80]. The molecular origin of electrochemical promotion has for years been rationalized on the basis of the sacrificial promoter mechanism [15, 16]. According to this mechanism, NEMCA results from the Faradaic (i.e. at a rate I /nF , where n is the ion charges) introduction of promoting species (Oδ− in the case of O2− conductors, H+ in the case of H+ conductors) on the catalyst surface. This electrochemically introduced O2− species acts as a promoter for the catalytic reaction

8.1.2 Electrochemical Modification of Catalytic Activity

1915

50 100

60 40

I=0

30

max

0

600

= 74000

NG = 4.2 × 10−9 mol Pt

400

2FNG /I = 800 s

20

200

Strongly bonded O backspillover state consumed over a period TOFmax /

Strongly bonded backspillover state gets populated at a rate I/2F catalytic rate is 99% due to weakly bonded state

10

20

800

r0 = 1.5 × 10−8 mol O s−1 −8 −1 ∆r = 38.5 × 10 mol O s I/2F = 5.2 × 10−12 mol O s−1 rmax = 26

40

r / 10−8 mol O s−1

TOF / s−1

80

I = +1 µA

0

Catalyst potential UWR / mV

I=0

−200

0 0 t 2FNG / I 1200

2400 6600 78000

Time / s

T/K 600

20

700

800

900

3900 s Weakly bonded highly reactive state

(dN /dT ) / 10−11 mol s−1

16

2030 s

Strongly bonded O backspillover state

12 700 s 2FNG/I = 2500 s 340 s

8 t 125 s 0s

4

0

300

400

500

600

T / °C 40 30

2FNG /I = 1200 s

I / µA

20 10 0 0s

−10 −20 −30 −1

20 s

t Strongly bonded O backspillover state

−0.7

−0.4

60 s 40 s 160 s 100 s 400 s 260 s 600 s

−0.1

0.2

Weakly bonded highly reactive O state

0.5

0.8

U WR / V NEMCA and its origin on Pt/YSZ catalyst electrodes [16]. Transient effect of the application of a constant current (a, b) or constant potential UWR (c) on (a) the rate, r, of C2 H4 oxidation on Pt/YSZ (also showing the corresponding UWR transient), (b) the O2 TPD spectrum on Pt/YSZ after current (I = 15 µA) application for various times t and (c) the cyclic voltammogram of Pt/YSZ after holding the potential at UWR = 0.8 V for various times t. Reprinted with permission from Kluwer/Plenum Press [16].

Fig. 8

1916

8.1 Electrocatalysis

Temperature/ °C

Temperature/ °C

15

200

300

400

500

600

High T adsorption

700

2 Pt/YSZ powder

b3

Electropromoted Pt film/YSZ

16O

0 0.3 0.2

2

Electropromoted b2

0.1

Pt film/YSZ

18

O2

500

600

700

300

400

500

600

700

b3 b2

0.8 16O18O

800

Pt/YSZ powder

0.6

b3

0.4 b2 16

0.2

0 300

900 1000

Temperature / K

O18O

Electropromoted Pt film/YSZ

Pt/YSZ powder

400

200

High T adsorption

16O

5

100

1

b3

b2

10

0 300 (a)

100

Desorption rate / 10−12 mol Os−1

Desorption rate / 10−12 mol Os−1

Desorption rate / 10−12 mol Os−1

20

400

500

600

700

800

900

1000

Temperature / K

(b)

Temperature / °C

Desorption rate/10−12 mol Os−1

30

450

650

25

Low T adsorption 16O

20

b3

15

5

b2

18

O2

2

60 45

16

O2

b2 + b3

30

×100

×100 ×100

16

O2

400

90 75

2

b1

10

850

b3 Pt/YSZ powder (UHV TPD) Pt/YSZ powder (atmospheric TPD) Electropromoted Pt film/YSZ b

0 200 (c)

250

600

800

1000

15

Desorption rate/10−10 mol Os−1

50

0

Temperature / K

Comparison of oxygen thermal desorption spectra from nanodispersed Pt/YSZ catalyst (1% Pt/YSZ) and from Pt/YSZ film taken under UHV conditions [81]. (a) Comparison of desorption of 18 O2 and 16 O2 , high Tad (200 ◦ C) adsorption. (b) Comparison of 16 O18 O desorption, high Tad (200 ◦ C) adsorption; Desorption was performed with linear heating rate β = 0.5 ◦ Cs−1 . (c) Comparison of O2 TPD spectra of the nanodispersed Pt/YSZ catalyst (UHV TPD and atmospheric pressure TPD) and of the electrochemically promoted Pt/YSZ film for low adsorption temperature, Tad = 70 ◦ C. Heating rate β = 0.5 ◦ C s−1 for the film and the UHV powder TPD and 1.5 ◦ C s−1 up to 750 ◦ C followed by 0.35 ◦ C above 750 ◦ C for the atmospheric pressure power TPD [81].

Fig. 9

(by changing the catalyst work function and affecting the chemisorptive bond strengths of coadsorbed reactants and intermediates) and is eventually consumed at a rate equal, at the steady state, to its rate of supply (I /2F ) which is  times smaller than the rate of consumption of the catalytic reactant, e.g. atomic O originating from the gas phase [15, 16]. Figure 8 shows the validity of the sacrificial promoter concept for the galvanostatic transient of Fig. 2, by presenting O2 TPD (Fig. 8b) and cyclic voltammetric (Fig. 8c) spectra obtained at times corresponding to those

of the NEMCA galvanostatic transient of C2 H4 oxidation (Fig. 8a) [16], under high-vacuum conditions. The sample consists of a polycrystalline Pt catalyst film deposited in the form of a half-ring on the outer surface of a tubular YSZ specimen with Au counter and reference electrodes deposited on the inside wall of the tubular YSZ element opposite to the Pt film. The sample is first exposed to gaseous oxygen followed by electrochemically supplied oxygen from the solid electrolyte. A linear temperature increase (Fig. 8b) or cyclic potential variation (Fig. 8c) is used to obtain the corresponding TPD spectra or cyclic

8.1.2 Electrochemical Modification of Catalytic Activity

Tab. 1

Classification of electrochemical promotion studies on the basis of global r vs.  behavior [16]

Reactants D

1917

Catalyst

Solid electrolyte

pA /pD

T/ ◦ C

Kinetics in D: (∂r/∂pD )

Kinetics in A: (∂r/∂pA )

Rule [16]

0 0 0 0 0 0 ≤0 0 ? 0 ≤0 0 ∼0 ? 0 0

G1 G1 G1 G1 G1 G1 G1 G1 ? G1 G1 G1 G1 ? G1 G1

≤0

? ? G1 G1 G1

+ + ? + + + + + ? ? ? + + ? + + + ?

G2 G2 ? G2 G2 G2 G2 G2 ? ? ? G2 G2 ? G2 G2 G2 G2

+ – + – – ? + +

G3 G3 G3 G3 G3 ? G3 G3

A

A. Purely electrophobic reactions C2 H4 C2 H4 C2 H4 C2 H4 C2 H4 C2 H4 C2 H4 CO CO CH4 C3 H6 CH4 C6 H6 C2 H2 H2 H2 H2 S CH4 NH3 NH3 CH4

O2 O2 O2 O2 O2 O2 O2 O2 O2 O2 O2 O2 H2 H2 CO2 C 2 H2 , C 2 H4 − − − − H2 O

Pt Pt Pt Rh Ag IrO2 RuO2 Pt Pd Pd Ag Ag Pt Pt Rh Pd

ZrO2 (Y2 O3 ) β  –Al2 O3 TiO2 ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) CaF2 ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) β  − Al2 O3 β  − Al2 O3 ZrO2 (Y2 O3 ) β  − Al2 O3

12–16 238 3.5–12 0.05–2.6 0.2–1.1 300 155 11–17 500 0.2–4.8 20–120 0.02–2 0.02–0.12 1.7–9 0.03–0.7 0.1–5.9a

260–450 180–300 450–600 250–400 320–470 350–400 240–500 500–700 400–550 380–440 320–420 650–850 100–150 100–300 300–450 70–100

+ + + + + + + + ?b + + + ≥0 ? + ≥0

Pt Ag Fe Fe Ni

ZrO2 (Y2 O3 ) SrCe0.95 Yb0.05 O3 CaZr0.9 In0.1 O3−α K2 YZr(PO4 )3 ZrO2 (Y2 O3 )

– – 4–12 kPae 4–12 kPa 0.05–3.5

600–750 750 530–600 500–700 600–900

? ? + + +

Pt Pt Pt Ag Ag Pt Ag Au Fe Ni Pt Ag Pt Pt Pt Pd Pd Pd

CaZr0.9 In0.1 O3−α CeO2 YZTi10 β  -Al2 O3 β  -Al2 O3 ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) CaZr0.9 In0.1 O3−α CsHSO4 ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) β  -Al2 O3 β  -Al2 O3 ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) Nafion

4.8 1.6–3.7 3 0.3–0.4 0.1–10 0.9–55 0–2 0.1–0.7 0–3 1 – 0–6 kPae 0.2–10 0.1–1.1 0.3–5 0.5–6.5 2–50 –

385–470 500 400–475 240–280 360–420 350–480 500 700–750 440 150–170 400–500 550–750 380–500 280–400 320–400 320–480 440 70

– – ?b – 0 ≤0 ? 0 ? ?

Pt Pt Pt Pt Pt Pt Pt Pt

Na3 Zr2 Si2 PO12 ZrO2 (Y2 O3 ) β  -Al2 O3 H2 O–0.1 M KOH Nafion V2 O5 − K2 S2 O7 β  -Al2 O3 β  -Al2 O3

1.3–3.8 0.2–55 0.5–20 0.3–3 0.2–5 1.8 2–70 0.3–6

430 468–558 300–450 25–50 25 350–450 375 360–400

B. Purely electrophilic reactions C2 H4 C2 H4 C2 H4 C2 H4 CO C3 H6 CH3 OH CH4 H2 H2 C2 H4 C2 H4 CO CO CO

O2 O2 O2 O2 O2 O2 O2 O2 N2 C 2 H4 CH3 OH CH3 OH NO NO NO NO N2 O 1-C4 H8

0 ? ≤0 ∼0 –

C. Volcano-type reactions C2 H4 CO CO H2 H2 SO2 C3 H6 H2

O2 O2 O2 O2 O2 O2 NO NO

– + – + + ?b – –

(continued overleaf )

1918 Tab. 1

8.1 Electrocatalysis (continued)

Reactants D

Catalyst

Solid electrolyte

pA /pD

T/ ◦ C

Kinetics in D: (∂r/∂pD )

Kinetics in A: (∂r/∂pA )

Rule [16]

0.2–0.3c 5 0.6–14 3.5–12.5 3–53 0.06–7 0.02–7 3–45 0.2–1.1 0.08–8d 0.33d

450–600 400–500 350–450 450–500 450–600 270–500 600–750 300–500 500–590 250–450 250–450

+ ?b + + + + + + + + +

+ ? + + ≥0 + + ? + NO: + O2 : 0 NO: + O2 : 0

G4 ? G4 G4 G4 G4 G4 ? G4 G4 G4

A

D. Inverted volcano-type reactions C2 H4 C3 H6 CO CO CO C2 H6 CH4 CH3 OH H2 C3 H6 CO

O2 O2 O2 O2 O2 O2 O2 O2 CO2 NO, O2 NO, O2

Pt Pt Ag Ag−Pd alloy Au Pt Pt Pt Pd Rh Rh

TiO2 YZTi10 ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 ) ZrO2 (Y2 O3 )

ap = p D C2 H2 + pC2 H4 . b ? = no data available. c Low

pA , pD region. d p /p is the ratio p NO /pC3 H6 A D e Monomolecular reaction.

and pNO /pCO . The pO2 range is between 0 and 6 kPa.

voltammograms. One clearly observes, both with TPD and with cyclic voltammetry, the Faradaic introduction, over a time period 2F NG /I (where NG is the catalyst surface area expressed in moles and thus 2F NG /I is the time required to form a monolayer of O2− on the catalyst surface), of a second (backspillover) strongly bonded oxygen species on the Pt catalyst surface which displaces the normally chemisorbed oxygen state to lower desorption temperatures. This displacement, which results from strong repulsive lateral interactions, between O2− and more covalently bonded atomic oxygen [16], causes the observed dramatic enhancement in the catalytic rate. The backspillover O2− state acts as a sacrificial promoter. It is termed ‘‘sacrificial’’ because it also becomes consumed at the steady state at a rate I /2F . This molecular picture has been recently confirmed by the use of 18 O2 TPD [81] (Fig. 9), which confirms that for the electropromoted Pt/YSZ catalyst the displaced weakly bonded state (β2 ) is populated by gaseous oxygen (18 O) whereas the strongly bonded promoting state β3 is populated by backspillover lattice oxygen 16 O. This fully confirms the sacrificial promoter mechanism of electrochemical promotion [16, 81, 82]. As also shown in Fig. 9, there is a strong and noteworthy similarity in the O2 TPD spectra of electropromoted Pt films deposited on YSZ and Pt/YSZ powder catalysts. The only difference is that, in the case of the Pt/YSZ powder, lattice oxygen also occupies state β2 (Fig. 9a) and gaseous oxygen 18 O2 is limited to the low-temperature desorption peak β1 [81, 82]. The latter appears both with the dispersed Pt/YSZ powder catalysts and with the electropromoted

Pt film/YSZ catalysts. These results corroborate the functional identity and only operational difference between electrochemical promotion and metal–support interactions with O2− conducting supports [15, 16]. The O2− backspillover mechanism of electrochemical promotion and metal–support interaction has been also confirmed recently via STM [16], as shown in Figs. 10 and 11. Anodic polarization causes the decoration of the Pt(111) surface of a Pt single crystal which is interfaced with the YSZ support, with backspillover O2− species which form a Pt(111)−(12 × 12)−O adlattice which coexists with the well-known underlying Pt(111)−(2 × 2)−O adlattice formed by gaseous O2 [16] (Fig. 11). the situation is similar (Fig. 10) when using β  -Al2 O3 , an Na+ conductor, as the solid electrolyte, i.e. active Na-donor catalyst support (Figs. 3 and 10). Mechanism of Metal–Support Interactions Metal–support interactions play an important role in the performance of many industrial supported catalysts. Since the time of Schwab [83, 84], understanding the mechanism of metal–support interactions has been one of the central and most challenging problems in heterogeneous catalysis. The effect can be pronounced, as shown in Fig. 12 for the case of C2 H4 oxidation on Rh dispersed on four different supports of increasing work function [85]. The sharp rate transition is due to surface Rh oxide formation, which poisons the oxidation rate [85, 86]. Figure 12 also shows one of the key experiments which proved the mechanistic equivalence of electrochemical 8.1.2.3.2

8.1.2 Electrochemical Modification of Catalytic Activity

1919

STM tip Pt film Pt(111)

Au connector

G/P b ″-Al2O3 I Pt counter electrode

Pt reference electrode

U WR (a)

25 Å

25 Å

(b)

(c)

500 Å

Fig. 10 (a) Schematic of the experimental setup for using STM to investigate the Pt(111) surface of a Pt single crystal interfaced with β  -Al2 O3 . (b) Low scanning-area STM images (unfiltered) of the (left) sodium-cleaned and (right) sodium-dosed Pt(111)−(2 × 2)−O adlattice. Total scan size 159 A˚ [16, 65]. Reprinted with permission from Elsevier Science. (c) Larger scanning area STM image (unfiltered) of a Pt single crystal surface consisting mainly of Pt(111) terraces and covered by a Pt(111)−(12 × 12)−Na adlattice formed via electrochemical Na+ supply on the Pt(111)−(2 × 2)−O adlattice. Each sphere on the image corresponds to an Na atom [16, 65]. Reprinted with permission from the Electrochemical Society.

promotion and metal–support interactions [85]. The r vs. pO2 behavior obtained on the finely dispersed Rh catalyst on the four supports can be reproduced by varying the potential (and thus the work function [16]) of a polycrystalline Rh film deposited on YSZ (Fig. 12, inset). Hence one can assign to each support an equivalent potential (Fig. 13a) which correlates linearly and with a slope of unity to the independently measured

absolute potential or work function [16] of the supports (Fig. 13b). Supports with higher absolute potential [66] or work function [66, 85] have enhanced propensity for O2− backspillover on the catalyst surface; therefore, they enhance C2 H4 oxidation, which is an electrophobic References see page 1934

1920

8.1 Electrocatalysis

STM tip Pt film Pt(111) G/P Y2O3-ZrO2

I

Pt counter electrode

Pt reference electrode

U WR

50 Å

(a)

250 Å (b)

50 Å (c)

Fig. 11 (a) Schematic of the experimental setup used to investigate via STM the Pt(111) catalyst electrode surface in contact with YSZ [67]. G/P is the galvanostat–potentiostat and UWR is the catalyst potential with respect to the reference electrode. (b) Unfiltered STM image of a 1275 × 1275 A˚ area of the Pt(111) surface following positive potential application (UWR = 1 V) (bias Ut = 450 mV, tunneling current It = 7 nA). (c) Unfiltered (top) and filtered (bottom) STM image of the Pt(111) surface following positive potential application (UWR = 1 V) showing the backspillover O2− -formed Pt(111)−(12 × 12)−O adlattice overlapping with the underlying Pt(111)−(2 × 2)−O adlattice. Inset: Fourier transform spectrum (bias Ut = 400 mV, tunneling current It = 8 nA) [67].

reaction, i.e. a reaction where the rate increases with increasing potential and work function (Fig. 12, inset). Another key experiment which has shown the equivalence of electrochemical promotion and metal– support interactions is shown in Fig. 14: IrO2 , a metallic oxide, and TiO2 exhibit a strong metal–support interaction, as shown by the volcano-type rate dependence of C2 H4 oxidation on catalyst composition [85, 87]. As shown in this ingenious experiment by Nicole and Comninellis [88], pure IrO2 can be electrochemically promoted by a factor of 11 (ρ = r/r0 = 11) (Fig. 14), but IrO2 −TiO2 catalysts are only marginally affected by electrochemical promotion. This is because they are already in a self-driven electrochemically promoted state via contact with TiO2 [85, 87] (Fig. 14). These two experiments, together with the self-driven electrochemical promotion experiments of Cavalca et al. [25] (Fig. 7), have shown conclusively that electrochemical promotion is an

electrically controlled metal–support interaction (Fig. 15) and that metal–support interactions on ZrO2 -, CeO2 - or TiO2 based supports are induced by reverse spillover of oxygen anions from the carrier on to the surface of the metal crystallites. Thus the same mathematical models describing diffusion and consumption of the sacrificial Oδ− promoter on electrochemically promoted NEMCA catalysts are applicable for dispersed catalysts deposited on ZrO2 , CeO2 and TiO2 supports [16, 89]. It is worth noting that the ‘‘classical’’ approach to interpreting metal–support interactions is based on the view of ‘‘electron transfer’’ between the catalyst and the support [83, 84, 86]. It neglects the ionic conductivity and thus O2− donating capacity of ZrO2 -, CeO2 - and TiO2 -based supports, which pins the Fermi level of supported nanoparticles to that of the support. Hence the ‘‘classical’’ approach focuses on the semiconducting and not the ionic properties of the support. The ‘‘electron

8.1.2 Electrochemical Modification of Catalytic Activity

4

120

3

1

0

1

2

3

pO / k Pa 2

T = 320°C p C2H4 = 3kPa

TiO2-4%WO3 3

0.05% Rh/ TiO2 (4% WO3) ZrO2 (8% Y2O3) SiO2 TiO2

40

0

0

2

pC2H4 = 3 kPa

2

0

80

4

Open circuit +1340 mV +950 mV +570 mV −1000 mV

eU * WR / eV

CO2 TOF / s−1

r CO2 / 10−6 mol g-cat−1s−1

160

1921

4

6

YSZ (ZrO2-8% Y2O3) SiO2 1

TiO2

0

−1 0.1

8

p O2 / kPa

2

0.2

0.4

0.6 0.8 1

2

(pO2 / pC2H4)

(a) 3.5

Effect of pO2 on the rate of C2 H4 oxidation on Rh supported on four supports of increasing . Catalyst loading 0.05 wt.% [16, 85, 87]. Inset: electrochemical promotion of an Rh catalyst film deposited on YSZ: Effect of potentiostatically imposed catalyst potential UWR on the rate and TOF dependence on pO2 at fixed pC2 H4 [16, 85, 87].

Fig. 12

8.1.2.3.3 Interrelation of Promotion, Electrochemical Promotion and Metal–Support Interactions: the Double-Layer Promotion, electrochemical proModel of Catalysis motion and metal–support interactions are three, at first glance, independent phenomena which can affect

2.5

eU * WR / eV

transfer’’ view is correct, but only as the first step for inducing O2− reverse spillover. Thus a support with high absolute potential or work function, such as YSZ (0 = 5.14 eV) will initially receive some electrons from a supported metal of initially lower , e.g. supported polycrystalline Rh, but the positive charge on Rh will induce O2− reverse spillover to neutralize it [16]. The higher the work function and absolute potential of the support, the higher is its ability to donate O2− [16]. Consequently, the ‘‘electrochemical promotion’’ view of metal–support interactions correctly predicts that the rate of electrophobic reactions, such as C2 H4 oxidation (Fig. 12), increases with increasing work function of the support (Figs. 12 and 13) while the ‘‘electron transfer’’ approach leads to the opposite conclusion: increasing support work function enhances the positive charge on the Rh nanoparticles, thus enhancing oxygen binding to the surface, which is the opposite of what Fig. 12 shows. One may therefore conclude that there is compelling evidence (Figs. 7 and 9–14) for the metal–support interaction mechanism shown in Fig. 15.

TiO2-4%WO3 3

2

1.5

1

0.5

YSZ (ZrO2-8% Y2O3)

SiO2 TiO2

4.8

(b)

5

5.2

5.4

5.6

Φo / eV

(a) Effect of (pO2 /pC2 H4 )∗ ratio at rate transition on the potential U∗WR where the rate transition occurs during C2 H4 oxidation on Rh films deposited on YSZ (Fig. 12 inset, circles) and on the equivalent potential U∗WR where the same rate break occurs on different supports (Fig. 12). (b) Correlation between the equivalent potentials of the supports defined in (a) and of the work function or absolute potential [16, 60] of the supports measured via the Kelvin probe technique in pO2 = 1 atm at 400 ◦ C [16, 60]. Fig. 13

catalyst activity and selectivity in a dramatic manner. We have already discussed the (functional) similarities and (operational) differences of promotion and electrochemical promotion. We have also discussed the functional similarities and only operational differences of electrochemical promotion and metal–support interactions References see page 1934

1922

8.1 Electrocatalysis

0.25

12

l = 0, Metal support interaction l = 200 µA, Electrochemical promotion r

10

0.1

r

6

Me tal –s up po rt

0.15

8

int era cti on

Ele ctr oc he mi ca lp rom oti on

r / 10−6 mol Os−1

0.2

4

0.05 2

0

0

25

75

100

0

XlrO2 (lrO2 / lrO2 + TiO2) / mol%

(a) O2

O2− O2−

2−

2− O2− O

TiO2

O O2−

O2−

IrO2 O

O2− O2−

Donor phase

O2−

C 2H 4

C2 H 4 + O 2

2−

C 2 H4 + O 2

CO2 + H2O

Acceptor phase

CO2 + H2O O

O2−

H4 C2

O2− O2−

O2−

O

O

(b)

50

O

2−

IrO2 2−

O2−

O2−

O2− O2− 2− O2− O2− O IrO2 IrO2 O2− O2−

O2− O2−

O2− O2− YSZ Au

(a) Effect of IrO2 mole fraction in the IrO2 −TiO2 catalyst [16, 85, 88] on the open-circuit catalytic rate, r0 , of C2 H4 oxidation (◦), on the electrochemically promoted (I = 200 µA) catalytic rate, r (•) and on the corresponding rate enhancement ratio, ρ (
). (T = 380 ◦ C, pO2 = 20 kPa, pC2 H4 = 0.15 kPa). (b) Mechanism of metal (IrO2 )–support (TiO2 ) interaction (left) during ethylene oxidation on IrO2 and of electrochemical promotion utilizing YSZ (right) [16, 85, 88].

Fig. 14

on ionic and mixed conducting supports. It therefore follows that promotion, electrochemical promotion and metal–support interactions on ion-conducting and mixedconducting supports are three different facets of the same phenomenon. All three are linked via the phenomenon of spillover–backspillover of the promoting species. Also, all three are due to the same underlying cause: the interaction of adsorbed reactants and intermediates with an effective double layer formed by promoting species at the metal/gas interface (Fig. 15). For time-scales shorter than that of a catalytic turnover (typically 10−2 –102 s) the three phenomena are indistinguishable. Looking at the Na-promoted Pt surface in Fig. 10c and imagining that CO oxidation is taking place on that surface, it is not possible to distinguish if this is a classically promoted surface where Na has been added

from the gas phase or an electrochemically promoted one where Na originated from β  -Al2 O3 interfaced with the Pt crystal or finally if it is the surface of a larger crystallite deposited on a porous β  -Al2 O3 carrier where Na has spontaneously migrated on the Pt surface (metal–support interaction). The oxidation of CO (Fig. 3) will be equally promoted in all three cases. Similar would be the situation on a Pt surface decorated with O2− , the only difference being the experimental difficulty of introducing O2− with classical promotion and its short lifetime on the catalyst surface, only  times longer than the catalytic turnover time. Consequently, the functional identity of classical promotion, electrochemical promotion and metal–support interactions should not lead to pessimistic conclusions regarding the practical usefulness of electrochemical

8.1.2 Electrochemical Modification of Catalytic Activity

∼1µm

O

2−

O

2−

O

2−

O

2−

O2−

YSZ or TiO2

O2− O2− O2− O

2−

O

2−

O

2−

O

2− 2−

O 2− O2− O O2−

Od− d+ + Od− d d− d+ O M d+ Od− d+ Od− + d Od− d+Od−

O2

UWR Φ E

∼1nm

Effective double layer O2− O

C2H4 + 3O2

2−

O2−

2CO2 + 2H2O

TiO2

µA

O

2−

O2− O2− O2−

Classical double layer

O2−

m e = EF mO2

UWR Φ mO2− Electrochemical promotion

1923

Od− d+ + Od− d d− d+ O + d− M d O + d Od− d+Od− d+ d− O

C2H4 + 3O2

2CO2 + 2H2O

O2 Φs

E

me = EF m O2

Φs

mO2−

Metal-support interaction

Schematic of a metal grain (∼µm) in a metal catalyst film deposited on YSZ or TiO2 under electrochemical promotion conditions (left) and of a metal nanoparticle (∼nm) deposited on a porous TiO2 support (right) showing the locations of the classical double layers formed at the metal/support interface and of the effective double layers formed at the metal/gas interface. The energy diagrams (bottom) indicate schematically the spatial constancy of the Fermi level EF (or electrochemical potential µ¯ e ) of electrons, of the chemical potential of oxygen and of the electrochemical potential of O2− . Note that under electrical bias application (left) µ¯ O2− remains spatially constant but µ¯ e and µO2 both bend in the solid electrolyte support (dashed lines). The Fermi level µ¯ e of the metal can be affected by varying UWR (left) or by varying via doping the Fermi level of the support (right) [16]. Reprinted with permission from Kluwer/Plenum Press. Fig. 15

promotion. Operational differences exist between the three phenomena and it is very difficult to imagine how one can use metal–support interactions with conventional oxidic supports to promote an electrophilic reaction or how one can use classical promotion to generate the strongest electronegative promoter, O2− , on a catalyst surface. Furthermore there is no reason to expect that a metal–support interaction-promoted catalyst is at its ‘‘best’’ electrochemically promoted state. Thus the experimental problem of inducing electrochemical promotion on very thin or fully dispersed catalysts remains an important one, as discussed in detail elsewhere [16]. Also important remains the issue of learning more, via surface spectroscopy, STM and ab initio quantum chemical calculations, about the exact state and geometry of O2− adsorption on metal surfaces. In the case, for example, of SMSI with TiO2 catalysts, it is well documented [1, 16] that the backspillover species is TiOx , where x is variable and about one for Pt. Although XPS investigations of Pt/YSZ [16] and Pt/TiO2 [16] NEMCA catalysts have not so far provided evidence for any such anion–cation pair electrochemically induced migration, this point is worth further investigation, particularly under reducing conditions. Regarding the adsorption geometry of O2− , STM has been used recently to follow the migration of O2− from YSZ on to Pt(111) under

atmospheric pressure conditions [67], using the same design and procedure as in Fig. 10, but now with YSZ as the solid electrolyte (Fig. 11). We found that, under opencircuit conditions, most of the Pt(111) surface is covered by the well-known (2 × 2)–O adlattice, although patches of bare Pt (111) and patches of a (12 × 12)–O overlayer are also visible. After anodic polarization at 1 V, the Pt(111) surface and (2 × 2)–O adlattice are covered almost entirely by the (12 × 12)–O adlayer which corresponds to the electrochemically migrating O2− species (Fig. 11). Each O2− adspecies appears to be ‘‘large’’, i.e., it perturbs the electronic cloud of at least 10 neighboring O atoms of the coexisting (2 × 2)–O adlattice, which remains clearly visible [67]. These observations are consistent with the ‘‘long-range’’ promoting effect of O2− and underline that, as Pritchard had correctly predicted [72], there is clearly much new surface chemistry to be explored. Having discussed the functional similarity of classical promotion, electrochemical promotion and metal–support interactions on O2− -conducting and mixed electronic–ionic conducting supports, it is useful also to address and systematize their operational differences. This is attempted in Fig. 16: the main operational difference is the promoter lifetime, τp , on the catalyst surface. References see page 1934

1924

8.1 Electrocatalysis

Metal-support Interactions Electrochemical promotion Classical Promotion 100

102

104

106

108

1010

tP / s Fig. 16 Operational range of classical promotion, electrochemical promotion and metal–support interactions in terms of the promoter lifetime, τp , on the catalyst surface [16].

One can also easily appreciate why the above mechanism of metal–support interactions was extremely difficult to detect with surface spectroscopic techniques: The O2− signal from the support can very effectively mask the O2− signal from the dispersed catalyst surface. Also, if one uses 18 O2 as the gaseous oxidant, then only a fraction f = 1/ of 16 O from the support will be found in the products. This fraction becomes significant only at elevated temperatures (>550 ◦ C), where  approaches unity and the phenomenon of electrochemical promotion with O2− supports disappears [15, 16], i.e. the double layer desorbs (Fig. 9) and the limits of pure electrocatalysis are reached [15, 16]. 8.1.2.4

For any practical classical promotion application in a fixed-bed catalytic reactor, τp must be longer than 1 year (∼3 × 107 s). But even for laboratory-scale classical promotion experiments, τp values in excess of 106 s are required (Fig. 16). On the other hand, electrochemical promotion is not subject to any such restrictions regarding τp (Fig. 16). Thus, when using O2− conductors or H+ conductors, τp is 102 –104 s, but when using Na+ conductors τp can be well in excess of 107 s at low temperature, but also in the range 104 –106 s for higher temperatures [16]. This is an important operational advantage of electrochemical promotion: it permits the use of a wide variety of sacrificial promoters (e.g. O2− , H+ ), which have too short lifetimes for classical promotion applications. 8.1.2.3.4 Why the ‘‘New’’ Supports? In view of the previous discussion, the answer to the third question posed in Section 8.1.2.2.3 becomes obvious: for electrophobic reactions, i.e. for reactions which are accelerated via positive potential application in NEMCA experiments or via supply of O2− on the catalyst surface, the ‘‘new’’ supports, i.e. YSZ, CeO2 , doped TiO2 , etc., are highly advantageous, since they offer continuous in situ promotion of the catalyst surface with backspillover O2− which is continuously replenished in the support by gaseous O2 . The catalyst support acts as a catalyst for transforming gaseous O2 to promoting O2− , which continuously migrates on the catalyst surface where it weakens the chemisorptive bond energy of normally chemisorbed oxygen, thus accelerating the catalytic reaction. Hence the frequently used ‘‘oxygen storage’’ mechanism [16, 86] to interpret the beneficial properties of CeO2 and of the other ‘‘new supports’’ is, in a broad sense, correct. It is only the two distinct types of oxygen present on supported metal surfaces, one promoting and the other highly active, which were additionally needed to complete the picture.

Double-Layer Approach to Catalysis

8.1.2.4.1 Why Double Layer? Rationalization and Prediction of Desired Types of Promoters and Supports One of the major advances following the discovery of electrochemical promotion, and the subsequent quest for understanding its molecular origin, was the observation that the work function, , of catalyst electrode surfaces changes with applied potential and, in fact, that over wide temperature and gas composition ranges (350–550 ◦ C for YSZ, 200–420◦ for β  -Al2 O3 ) the variation of  with catalyst potential UWR is given by the following simple equation:

 = eUWR

(11)

The ability to alter and control the work function of a catalyst surface via application of a potential led to strong interest among both leading surface scientists and electrochemists [44–46]. Equation (11) is now established as a basic relationship in solid-state electrochemistry and together with [16]: W − R = eUWR

(12)

allows for the definition of the absolute potential scale in solid-state electrochemistry [16, 66]. Since the absolute potential of an electrode is a property of the support and of the gas composition, but not of the metal, the same concept can be used to define the absolute potential of a catalyst support [16, 66]. This quantity, which equals /e, where  is the work function of the support, plays an important role in quantifying the promotional aspects of catalyst supports used to induce metal–support interactions, as already noted in Fig. 13. The experimental Eqs. (11) and (12) suggest by themselves the double-layer approach to electrochemical promotion and, in view of the already discussed mechanistic equivalence of electrochemical promotion, promotion and metal–support interactions [16], the double-layer approach to catalysis. The presence of an

8.1.2 Electrochemical Modification of Catalytic Activity

overall neutral double layer present at the metal/gas catalytic interface of catalysts is manifested simply by these equations, as follows [16, 66]. In general, by definition [15, 16], the following equation is valid for any metal catalyst or electrode: −µ¯ =  + e

(13)

where µ¯ is the electrochemical potential (which always equals the Fermi level EF of the metal) and  is the outer or Volta potential. Thus, for a working (W) catalyst electrode and a reference (R) electrode, one has UWR = W − R + e(W − R )

(14)

Comparison of the general theoretical Eq. (14) with the experimental Eqs. (11) and (12) gives W − R = 0

(15)

W = R = constant

(16)

which, in conjunction with Gauss’s theorem of electrostatics, gives for an overall neutral system W = R = 0 [66], i.e. the electrostatic Volta potential, , vanishes outside the double layer present at catalyst/gas interface, because this double layer is, as is every double layer, overall neutral. Consequently, one may conclude that promotion, electrochemical promotion and metal–support interactions are all catalysis in presence of a double layer. In the case of electrochemical promotion, this double layer is in situ tunable via variation of the applied potential. Hence one has the opportunity to study directly, at fixed temperature and gaseous composition, the effect of the catalyst work function, , on the kinetics of catalytic reactions. The Four Types of Rate–Work Function Dependence and the Promotional Rules As one would expect mathematically, four types of r vs.  dependence are observed on varying the catalyst potential UWR or work function  (Fig. 17, Table 1): 8.1.2.4.2

1. Rate increase with potential and work function, i.e. (∂r/∂θ)pA ,pD > 0, where pA and pD are the partial pressures of the electron acceptor (e.g. O2 , NO) and electron donor (e.g. C2 H4 , C6 H6 ) reactant, respectively. These reactions are enhanced/suppressed when the catalyst electrode is made positive/negative and thus have been termed electrophobic, from the Greek words electron + phobos (fear). Practically all oxidations under fuel-lean conditions are electrophobic reactions [16, 78–80]. 2. Rate decrease with work function, i.e. (∂r/∂θ)pA ,pD < 0. Practically all oxidations under fuel-rich conditions

1925

and the reduction of NO by most hydrocarbons are electrophilic reactions [16, 78–80]. 3. Volcano-type reactions, where the rate passes through a maximum with varying work function. Typical examples are the oxidation of CO and H2 on noble metals at low temperature [16, 78–80]. 4. Inverted volcano reactions, where the rate passes through a minimum with varying potential. Most catalytic reactions at elevated temperature exhibit inverted volcano behavior [16, 78–80]. In cases 1 and 2, the rate dependence on  is frequently found to satisfy the equation
 r α = exp (17) r0 kb T where α is positive for electrophobic and negative for electrophilic reactions. All classically or electrochemically promoted reactions can be grouped into these four categories (Table 1) [16, 78–80]. Fairly often, however, the same reaction changes its character as one varies the temperature or gaseous composition significantly. Can one predict in which of the four categories a given catalytic reaction on a given metal belongs? Although electrophobic and electrophilic reactions have been known and studied since the 1980s [13, 14] or even before, as the terms are synonymous with the terms electron acceptor and electron donor reaction introduced by Wolkenstein in the 1960s [83, 84], until recently, a positive answer to the above basic question appeared to be a very distant goal. Yet, as shown recently [16, 78–80], there exist simple and rigorous rules which enable one to predict the r vs.  behavior. In simple terms, a catalytic reaction is: 1. Electrophobic, if the electron acceptor is strongly bound on the catalyst surface and the electron donor is weakly bound (Rule G1, Table 1 (A)). 2. Electrophilic, if the opposite holds, i.e. if the electron acceptor is weakly bound and the electron donor is strongly bound (Rule G2, Table 1 (B)). 3. Volcano type, if both reactants are strongly bound on the catalyst surface (Rule G3, Table 1 (C)). 4. Inverted volcano type, if both reactants are weakly bound on the catalyst surface (Rule G4, Table 1 (D)). Conversely, if a reaction is electrophobic one can predict that the kinetics are positive order in D and zero or negative order in A. If a reaction is electrophilic, one can predict that its kinetics are negative or zero order in D and positive order in A. If a reaction is volcano type, then the rate vs. both pD and pA passes through a maximum, and References see page 1934

1926

8.1 Electrocatalysis

UWR /V

UWR / V −0.4

−0.8

80

0

−1.6 −1.2 −0.8 −0.4

0.4

pNO = 1.3 kPa pC2H4 = 3.7 kPa

r /r0

r /r0

5 10

pO2 = 0.25 kPa pCH4 = 10 kPa

5

pO2 = 0.50 kPa pCH4 = 10 kPa 1

−5

−10

0

1 0.6

5

Π

(a)

(b)

−25 −20 −15 −10 −5 Π

0

0.1

0.2

0

5

UWR /V

UWR / V −0.1 8

0.4

T = 450°C

10

T = 700°C

0

−0.8 −0.4

0.3

0

0.4

0.8

10

5 5

r /r0

r /r0

pO2 = 19 kPa pCH3OH = 0.9 kPa

T = 560°C pO2 = 0.2 kPa pCO = 11 kPa

1 −2 (c)

0

2 Π

T = 425°C

1 −15 −10 −5

4 (d)

0 Π

5

10 15

Fig. 17 Examples for the four types of global electrochemical promotion behavior: (a) electrophobic, (b) electrophilic, (c) volcano type and (d) inverted volcano type. (a) Effect of catalyst potential and work function change (vs. I = 0) for high (20 : 1 and 40 : 1) CH4 to O2 feed ratios, Pt/YSZ. (b) Effect of catalyst potential on the rate enhancement ratio for the rate of NO reduction by C2 H4 consumption on Pt/YSZ. (c) NEMCA-generated volcano plots during CO oxidation on Pt/YSZ. (d) Effect of dimensionless catalyst potential on the rate constant of H2 CO formation, Pt/YSZ ([16, 78–80] and original references therein).  = FUWR /RT(=/kb T).

if a reaction is inverted volcano type, then the kinetics are positive order in both A and D [16, 78–80]. These rules appear to have no exceptions [78–80]. They have been derived on the basis of more than 70 electrochemical and classical promotion studies [16, 78–80] and we have recently shown that there have been no exceptions to these rules during the last 10 years in publications in the Journal of Catalysis [80]. In mathematical terms, rules 1–3 can be expressed as [78–80]


∂r ∂




 pA ,pD

∂r ∂pD

 >0

(18)

,pA

i.e. the r vs.  dependence traces (has the same sign as) the r vs. pD dependence. Conversely, since in the presence of a double layer at the metal/gas interface  = −EF , the

above equation can be written as

 ∂r ∂r >0 ∂EF pA ,pD ∂pA EF ,pA

(19)

i.e. the r vs. catalyst Fermi level dependence traces (has the same sign as) the rate vs. pA dependence. The above rules 1–4 stem [88] from the following two fundamental rules: 
∂θD ≥0 (20) ∂ pA ,pD 
∂θA ≤0 (21) ∂ pA ,pD which express the fact that increasing the catalyst work function enhances the chemisorptive bond strength of electron donor adsorbates and weakens that of

8.1.2 Electrochemical Modification of Catalytic Activity

electron acceptor adsorbates. Both ‘‘fundamental’’ rules are in good agreement with the experimentally observed variation of adsorption enthalpies, Hj , with work function [16, 78–80]: |Hj | = αH,j 

(22)

where αH,j > 0 for electron donor adsorbates and αH,j < 0 for electron acceptor adsorbates. Equation (18) is also in excellent agreement with rigorous quantum mechanical calculations [70, 71]. The above four rules enable one to derive the following three ‘‘practical’’ rules [78–80] for promoter selection: 1. If a catalyst surface is predominantly covered by an electron acceptor reactant, e.g. O, then an electron acceptor promoter, e.g. O2− , is to be recommended. 2. If a catalyst surface is covered predominantly by an electron donor reactant, e.g. C6 H6 , C2 H4 , then an electron donor promoter (e.g. Na+ , K+ ) is to be recommended. 3. If both reactants are weakly adsorbed on the catalyst surface, then both electron acceptor and electron donor additives can enhance the rate. Clearly, all of the above rules are valid as long as site blocking by the promoter does not become a dominant factor [16]. Regarding metal–support interactions involving O2− conducting oxides, the following rule can be derived Metal–support interactions with oxidic ion conducting or mixed ionic–electronic conducting supports can enhance the rate of a catalytic reaction only when the reaction is electrophobic.

An example is shown in Fig. 13 for the case of C2 H4 oxidation on Rh supported on various supports of increasing absolute potential and work function. 8.1.2.4.3 Double-Layer Isotherms and Kinetics The experimentally proven existence of an overall neutral effective double layer at the metal/gas interface has been utilized recently [16, 78–80] to derive, starting from simple and rigorous thermodynamic and electrostatic principles, adsorption isotherms which account explicitly for the electrostatic interactions between the adsorbates and the double layer [78–80]. One starts from the equilibrium adsorption condition:

˜ AV µ¯ j (g) = µ¯ j (ad) = µj (ad) + P˜j EN

(23)

where µ¯ j is the electrochemical potential of adsorbed species j , µj is its chemical potential, P˜j , taken as a vector, is its dipole moment in the adsorbed state, E˜ is the local field strength in the double layer, assumed uniform, and NAV is Avogadro’s number.

1927

The equilibrium condition leads to the following effective double-layer (EDL) isotherm:   θj kj pj = (24) exp(−λj ) (1 − θj ) with




=  kj = exp

1 cos ω 2d kb T

 (25)

[µ0j (g) − µ0j (ad)]



RT

(26)

where  is the deviation of the work function, , from its value at the potential of zero charge (pzc) of the double layer, l is the dipole length, d is the doublelayer thickness [78–80], ω is the angle formed between the adsorbate dipole and the field strength and λj is the partial charge-transfer parameter. This parameter is zero for a truly covalent chemisorptive bond, positive for an electron donor adsorbate and negative for an electron acceptor adsorbate. Using Eq. (25) with cosω = 1 and the definition of the isosteric enthalpy of adsorption:   ∂[µ¯ j (ad) ] Had = T 2 T pj ,θj one can derive that 0 + Had,j = Had,j

λj l  2d

(27)

0 is the heat of adsorption for  = 0. where Had,j Assuming l ≈ d, one obtains
 λj 0 Had,j = Had,j +  (28) 2

Thus for an electron acceptor adsorbate (λj < 0), Eqs. (27) and (28) predict a linear decrease in Had with increasing , whereas for electron donor adsorbates (λj > 0) they predict a linear decrease in Had with decreasing . Both predictions are in excellent agreement with experiment (Fig. 6) [16] and with rigorous quantum mechanical calculations [70, 71]. One can use the effective double-layer isotherm [Eq. (24)] to derive analytical mathematical expressions for catalytic promotional kinetics [16, 79]. For the case of surface reaction rate control, the corresponding expression is r=

kR kA kD pA pD exp[(λD + λA )] [1 + kD pD exp(λD ) + kA pA exp(λA )]2

References see page 1934

(29)

8.1 Electrocatalysis

where kR = kR0 exp(λR ) and λR is the partial chargetransfer parameter of the transition state. In the limit of very weak adsorption (kA pA , kD pD 1) one can neglect repulsive interactions [16, 79] and consider only the attractive interactions. In this case, Eq. (29) becomes kR kD kA pD pA exp [max (0, λD ) + max (0, λA )]   1 + kD pD exp [max (0, λD )] 2 +kA pA exp [max (0, λA )] (30) where max(x, y) denotes x and x > y, y when x < y and x or y when x = y. The success of Eqs. (29) and (30) to describe the above recently derived promotional rules can be appreciated from Fig. 18, which shows the transition from electrophobic to electrophilic to volcano type and to inverted volcano type behavior by simply varying the values of the adsorption equilibrium constants kD and kA . UWR /V, ∆Φ/eV

UWR / V, ∆Φ/eV

0.5 1.0

JD, JA

lA = −0.15

1

0.4

ϑD

0.2

−10

kD = 100, pD = 1

0

10

1 JD JA

1.0

0 −20

20

0.5

0

10

20

Π

1.0

−1.0 −0.5

1 0.8

0

0.5

JD, JA

0.1

0.4

r /r0

ϑA

1.0

kD = 0.01, kA = 0.01 pD = 1, pA = 1 lD = 0.15, lA = −0.15,

ϑD

0.6

−10

0.01

UWR / V, ∆Φ/eV 1.0

0.8

JD, JA

0.1 JA

(b)

0

10

0.4

UWR /V, ∆Φ/eV −1.0 −0.5

100

lD = 0.15

0.01

Π

(a)

0.6

lA = −0.15

0.2

JD

0.5 1.0

kA = 0.01, pA = 1

0.1

ϑA

0

0.8 10

lD = 0.15

0 −20

−1.0 −0.5

100

kD = 0.01, pD = 1 kA = 100, pA = 1

0.6

1.0

JD, JA

0.8

0

r /r 0

1.0

−1.0 −0.5

r /r 0

r=

The success of double-layer kinetics can also be appreciated from Fig. 19, which compares model predictions (Fig. 19a and b) with some interesting and complex experimental results (Fig. 19c and d) obtained during C2 H4 oxidation on Pt/TiO2 [16]. As shown in Fig. 19c and d, the rate dependence on UWR and  shifts from inverted volcano (Fig. 19c) to purely electrophobic (Fig. 19d) as pC2 H4 (=pD ) is decreased by a factor of 10 at fixed pO2 . As shown in Fig. 19a and b, the model predicts the shift in global behavior in a semiquantitative manner and in fact with very reasonable λD and λA values (λD > 0, λA < 0). Finally, the success of the model can be judged from Fig. 20a and b, which show the experimental and model-predicted rate dependence on pCO and work function during CO oxidation on Pt/β  -Al2 O3 [16]. Note the transition from a classical Langmuir–Hinshelwood to a positive order rate dependence on pCO with decreasing work function. Also note that on every point of the experimental or model-predicted rate dependence,

0.6

10

JD JA

0.4

r / r0

1928

kD = 100, kA=100

0.2

0.2

pD = 1 pA=1 lD = 0.15 lA = −0.15

0 −20 (c)

−10

0

Π

10

1 20

0 −20

0.01 (d)

−10

0

10

20

Π

Fig. 18 Effective double-layer model predicted electrochemical or classical promotion behavior: (a) electrophobic, (b) electrophilic, (c) volcano type and (d) inverted volcano type [16, 78–80].

8.1.2 Electrochemical Modification of Catalytic Activity

UWR /V

UWR /V −2

3

−1

0

1

2

2

kA = 0.014, pA = 1.45 lD = 0.059 lA = −0.032

r /r0

r /r0 1 0

20

40

Π

−40

−20

3

2

0

20

40

Π

(b)

UWR /V −2

1

kD = 0.024, pD = 0.6

1 0.9 0.8 0.7 0.6 0.5

lD = 0.059 lA = −0.032 −20

0

3 2

−40

−1

4

kA = 0.014, pA = 1.45

(a)

−2

5

kD = 0.024, pD = 5.6

0.9

1929

UWR /V

0

−2

2 5

pC2H4 = 5.6 kPa

4

pO2 = 1.45 kPa

3

−1

0

1

2

pC2H4 = 0.6 kPa pO2 = 1.45 kPa

500°C

500°C

2

r /r0

r /r0

2

1 0.9 0.8 0.7 0.6

1 0.9

−40

−20

0

20

Π

(c)

0.5

40

−40

−20

0

20

40

Π

(d)

Experimentally observed [39] (c, d) and model predicted (a, b) transition from inverted volcano to electrophobic behavior upon increasing the O2 to ethylene (i.e. A : D) ratio by a factor of 10, C2 H4 oxidation on Pt/TiO2 [16, 78–80].

20 15 10 5 0 −10

−0.8 (a)

0

Π −0.4

0

UWR / V

0.4

10

5

4

3

2

0

1

r /10−7 mol s−1

r / 10−7 mol s−1

Fig. 19

20 15 10 5 0

Pa

−10

/k

p CO

−0.8 0.8

(b)

0 10

Π −0.4

0

UWR / V

0.4

5

4

3

p

0

1

2

Pa

/k

CO

0.8

Experimental (a) and model simulated (b) [16, 78–80] dependence of the rate of CO oxidation on Pt deposited on β  -Al2 O3 as a function of pCO , catalyst potential UWR and dimensionless catalyst work function (=/kb T) at pO2 = 6 kPa. Parameters used in equation [29]: kA = 9.133, kD = 8.715, λA = −0.08, λD = 0.09, λR = 0, kR = 6.19 × 10−6 mol s−1 .

Fig. 20

1930

8.1 Electrocatalysis

the basic promotional rule, Eq. (18), is strictly obeyed. The optimal λD and λA values are again reasonable (λD > 0, λA < 0). The large optimal kA and kD values (∼9) are also reasonable as they indicate strong adsorption of both CO(=D) and oxygen(=A), which is the necessary and sufficient condition (rule 3) for the appearance of volcano-type behavior. In general, Figs. 18–20 show, beyond any reasonable doubt, that the effective double-layer model of promotion, expressed mathematically by Eqs. (29) and (30), provides a satisfactory description of promotional kinetics. Electropromotion of Thin Films and Electropromoted Monolithic Reactors The main obstacles for the commercial utilization of electrochemical promotion are [16]: 8.1.2.5

1. expensive thick catalyst films (typically 0.1–5 µm thick) with metal dispersion below 0.01% 2. lack of efficient and compact reactor designs allowing for the utilization of electrochemical promotion with a minimum of electrical connections to the external power supply. As shown recently [90, 91], both of these limitations can be overcome via the use of thin sputtered noble metal electrodes with metal dispersion exceeding 15% in novel monolithic electrochemically promoted reactors (MEPRs) of the type briefly surveyed here. These studies [90, 91] confirmed the feasibility of electrochemically promoting thin (typically 40 nm) sputtered metal films with metal dispersion higher than 10–20%. This is an important practical breakthrough, because in these films the metal dispersion and utilization are comparable to those of state-of-the art conventional supported catalysts. Some new results are presented below in conjunction with the novel flat-plate electropromoted reactors. The recently developed novel MEPRs [90, 91], consist of flat (Figs. 21 and 22) or ribbed [90, 91] solid electrolyte plates, covered on both sides by appropriate thin, porous, conductive catalyst layers. The plates are inserted in appropriately carved grooves on the inside surfaces of the walls of the ceramic reactor casing. These surfaces are also used to create two current collectors, one establishing electrical contact among all catalyst films deposited on the top side of the plates (catalyst 1) and the other establishing electrical contact with all catalyst films deposited on the bottom side of the plates (catalyst 2, which can also be an inert conductive material). In this way, all catalyst films can be electrochemically promoted (anodically for catalyst 1 or 2, cathodically for catalyst 2 or 1) via only two external connecting wires (Figs. 21 and 22). This is a significant practical simplification and the MEPR can be considered as a hybrid between a classical monolithic honeycomb

reactor (of which it has all the geometric dimensions) and of a flat- or ribbed-plate solid oxide fuel cell. In contrast to fuel cells, where the fuel and air gas streams are kept separated, in the case of the MEPR there is only one gas stream containing all reactants and products, as in every classical catalytic reactor. An additional advantage of the MEPR is that it can be assembled and dismantled at will and its flat or ribbed plates can be replaced whenever necessary. Also, it is possible to use one of the plates as a gas-sensor element and utilize the potential signal generated by this element, under open circuit or at a fixed applied current, to control the current or potential applied to the electropromoted catalytic plates [90, 91]. In the first experimental investigation of an MEPR [90, 91], the following two reactions were studied: 1. The oxidation of ethylene to CO2 and H2 O, which exhibits strong electrophobic NEMCA behavior on both Rh and Pt [16] (Table 1). 2. The reduction of NOx by ethylene in the presence of O2 , a reaction which is electropromoted predominantly by positive current on Rh (electrophobic behavior) and predominantly by negative current on Pt (electrophilic behavior) [16] (Table 1). This dictated the choice of Rh and Pt as catalyst 1 and 2, respectively, in the MEPR (Fig. 21) [90, 91]. In this way, one obtains a synergistic effect of the electropromotion of Rh and Pt deposited on Y2 O3 -stabilized ZrO2 (YSZ) when the Rh film is made positive and the Pt film negative. Of course, with this choice of metals one does not expect a synergistic electropromotion effect in the case of C2 H4 oxidation, since the reaction is electrophobic on both metals. The solid electrolyte plates were produced by Bosch [90, 91] and had a thickness of 0.25 mm and dimensions of 50 × 50 mm. They were made of yttria-stabilized zirconia (YSZ, 8 wt.% Y2 O3 with a resulting molar composition Zr0.913 Y0.087 O1.957 ). The starting material had a mean particle size of 0.5 µm. The density in the sintered state was between 5.7 and 5.9 g cm−3 . The Rh/YSZ/Pt and Rh/YSZ/Au samples were prepared by metal sputtering [90, 91]. The surface area of the Rh catalyst films, and also the metal dispersion, were estimated using the galvanostatic transient technique [13, 14], by measuring the time constant, τ , required for the rate increase, r, in galvanostatic electropromotion rate transients to reach 63% of its steady-state value. In this way, one can estimate the reactive oxygen uptake, NG , of the Rh film and, assuming a 1 : 1 surface Rh:O ratio, the active catalyst surface area, NG , expressed in moles, from [13, 14] τ=

2F NG I

(31)

8.1.2 Electrochemical Modification of Catalytic Activity

1931

65 mm

70

mm 50

mm

0.5 mm

1.5 mm

65 mm

3 mm

46 mm Sensor element Au

Rh 4 mm Catalyst element Pt 1 mm 50

Rh

mm

1 mm 50 mm Au sensor element current collector

Rh sensor element current collector

Au sensor element (YSZ plate) Rh Pt catalyst element (YSZ plate) Rh Ag paste film

Pt catalyst current collector 1

Rh catalyst current collector 2

Fig. 21 Schematic and dimensions of the monolithic electropromoted reactor (MEPR) tested with 21 flat Rh/YSZ/Pt plates and one flat Rh/YSZ/Au plate [90, 91]. The bottom figure shows the geometry of the electrical connections on the MEP reactor ceramic walls.

In the first MEPR study [90], results were obtained with an MEPR using 1 or 22 Pt/YSZ/Rh plates, where the thin (40 nm) Pt and Rh films were deposited on the opposing sides of the thin (0.25 mm) YSZ plates.

More recent results in the same reactor with 10 similar Rh/YSZ/Au plates for the case of C2 H4 oxidation (Fig. 23) References see page 1934

1932

8.1 Electrocatalysis

Sensor element

Catalyst element Current collector (a)

Metal casing

Ceramic casing

(b)

(c)

Photographs of the MEP reactor–sensor showing (a) the machinable ceramic reactor walls, one of the Ag current collectors on the wall, the plate location for two-plate operation (top plate was used as sensor element) and part of the metal casing, (b) the 22-plate unit and (c) the assembled reactor with metal casing in the furnace. Also shown are the two thermocouple housings and the four shielded electrical connections for sensing and electropromotion [90, 91].

Fig. 22

and NO reduction by C2 H4 in presence of O2 are shown in Fig. 24. The Rh thin films were sputtered as described above, and paste Au electrodes were used as counter electrodes. As shown in these figures, both the MEPR and the thin (40 nm) sputtered Rh catalyst films operate very well and exhibit non-Faradaic performance even at fairly high reactant conversions. The same MEP reactor has recently been operated successfully at total gas flowrates of 25 L min−1 . This corresponds to a space velocity of 12 000 h−1 , which is of the same order of magnitude as practical exhaust treatment units. It is important to emphasize that this has been accomplished with similar noble metal dispersion (10–30%) as in state-of-the-art exhaust treatment catalysts.

Summary and Perspectives The search for understanding the phenomenon of electrochemical promotion at the molecular level during the last 10 years, by utilizing a wide variety of surface science and electrochemical techniques, has not only accomplished its initial goal, but perhaps more importantly, has also been particularly fruitful in defining, tackling and solving to a satisfactory extent several additional important problems in heterogeneous catalysis. This was due to the unique ability offered by electrochemical promotion to allow for in situ examination of the effect of promoters and of catalyst work function on catalytic activity and selectivity. Thus there is now compelling evidence that 8.1.2.6

8.1.2 Electrochemical Modification of Catalytic Activity

4 305 °C 0.9% O2 0.3% C2H4 Fv = 1000 cm3 min−1

40

30

rCO2 / 10−6 mol/s

x C2H4 , x O2 / %

2

3

U/ V

+20 mA

50

2

1.5

1

r = 2.3 = 34

1 20 0

50

100

150

200

0

t / min Ten-plate MEPR operation with C2 H4 oxidation on Rh: transient response of the rate of CO2 formation and of the conversions of O2 and C2 H4 upon application of a constant anodic current in the MEP reactor loaded with 10 catalyst plates (Rh/YSZ/Au) and one sensor element. T = 305 ◦ C [90, 91].

Fig. 23

1933

1. Electrochemical promotion is functionally identical with classical promotion, i.e. it is catalysis in the presence of a controllable double layer at the metal/gas interface. The main advantage of electrochemical promotion is that it also allows for the use of shortlived sacrificial promoters, such as O2− , which are continuously supplied to the catalyst/gas interface via electrochemically controlled reverse spillover from the solid electrolyte support. 2. Metal–support interactions of ZrO2 -, CeO2 -, Y2 O3 and TiO2 -based supports are due to a self-driven electrochemical promotional mechanism, i.e. continuous migration of sacrificial promoter O2− from the support to the metal/gas interface and continuous replenishment of O2− in the support from gaseous O2 . Electrochemical promotion itself is an electrically controlled metal–support interaction. Metal–support References see page 1934

+20 mA

x C2H4 /%

40 35 30

rCO2 / 10−6 mol s−1

45

305°C 3 0.36% C2H4 1.1% O2 1100ppm NO F V = 1200 cm3 min−1 2

2.5

2 rCO = 2.3

1

2

25

CO2 =

1.5

U/V

3

46.4

20

40

30

40

0.5 30

0.4

x O2 /%

−rNO / 10−6 mol s−1

x NO /%

50

0

0.6

60

r NO = 1.8 NO =

2.4

20

0.3 0

25

50

75

100

125

150

t / min Ten-plate MEPR operation with NO reduction by C2 H4 in presence of O2 : transient effect of constant applied anodic current (+20 mA) on the catalytic rates of CO2 production (rCO2 ) and NO reduction (−rNO ), on the NO conversion (xNO ) and on the Rh–Au potential difference (U). T = 305 ◦ C [90, 91].

Fig. 24

1934

3.

4.

5.

6.

7.

8.1 Electrocatalysis

interactions with these supports can only promote electrophobic reactions. Depending on their r vs.  dependence, catalytic reactions are grouped into four categories: electrophobic, electrophilic, volcano and inverted volcano. Rigorous rules have been derived which enable one to predict in which category a given catalytic reaction belongs, on the basis of its unpromoted kinetics. The same rules also enable one to predict the kinetics with respect to the electron acceptor or donor reactants (positive, zero or negative order) when the r vs.  dependence is known. The absolute potential of ionic or mixed ionic– electronic conducting supports has been defined and measured. It equals 0 /e, where 0 is the work function of the support under selected standard conditions. For electrophobic reactions, electronegative promoters and high work function supports enhance the catalytic activity significantly. For electrophilic reactions, electropositive promoters and low work function supports enhance the catalytic activity significantly. Electrochemical promotion, classical promotion but also metal–support interactions [16] can be modeled, similarly to electrocatalysis [16], by using simple and rigorous double-layer isotherms which utilize the fact that promotion, electrochemical promotion and metal–support interactions are different facets of the same phenomenon, i.e. catalytic reaction in the presence of a double layer, which for the case of electrochemical promotion is in situ controllable [16, 78–80].

Hence, aside from the, very likely, forthcoming technological applications [16, 90, 91], electrochemical promotion is a unique and efficient tool for studying the heart of classical catalysis, namely promotion and metal–support interactions. Outlook The recent advances in electropromoted reactor design and operation [90, 91] (Figs. 21–24) have shown that electrochemical promotion may soon find practical applications in exhaust treatment units and in chemical destruction or synthesis processes. Several aspects related to durability, useful lifetime, electrolyte and stack cost minimization, scale-up and scale-down have not yet been addressed in any detail, but there is already strong industrial interest and involvement (e.g. [90, 91]) and the next few years could potentially lead to the commercialization of some electropromoted reactors. 8.1.2.7

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56. C. G. Vayenas, S. Bebelis, I. V. Yentekakis, H.-G. Lintz, Catal. Today 1992, 11, 1. 57. T. Arakawa, A. Saito, J. Shiokawa, Appl. Surf. Sci. 1983, 16, 365. 58. S. Ladas, S. Kennou, S. Bebelis, C. G. Vayenas, J. Phys. Chem. 1993, 97, 8845. 59. W. Zipprich, H.-D. Wiemh¨ofer, U. V¨ohrer, W. G¨opel, Ber. Bunsges. Phys. Chem. 1995, 99, 1406. 60. S. G. Neophytides, C. G. Vayenas, J. Phys. Chem. 1995, 99, 17063. 61. D. Tsiplakides, C. G. Vayenas, J. Catal. 1999, 185, 237. 62. S. Neophytides, D. Tsiplakides, C. G. Vayenas, J. Catal. 1998, 178, 414. 63. D. Tsiplakides, S. Neophytides, C. G. Vayenas, Solid State Ionics 2000, 136–137, 839. 64. J. Poppe, A. Schaak, J. Janek, R. Imbihl, Ber. Bunsges. Phys. Chem. 1998, 102, 1019. 65. M. Makri, C. G. Vayenas, S. Bebelis, K. H. Besocke, C. Cavalca, Surf. Sci. 1996, 369, 351. 66. D. Tsiplakides, C. G. Vayenas, J. Electrochem. Soc. 2001, 148, E189. 67. C. G. Vayenas, D. Archonta, D. Tsiplakides, J. Electroanal. Chem. 2003, 554–555, 301. 68. D. Kek, M. Mogensen, S. Pejovnik, J. Electrochem. Soc. 2001, 148, A878. 69. A. D. Frantzis, S. Bebelis, C. G. Vayenas, Solid State Ionics 2000, 136–137, 863. 70. G. Pacchioni, F. Illas, S. Neophytides, C. G. Vayenas, J. Phys. Chem. 1996, 100, 16653. 71. G. Pacchioni, J. R. Lomas, F. Illas, Mol. Catal. A 1997 119, 263. 72. J. Pritchard, Nature 1990, 343, 592. 73. J. Nicole, C. Comninellis, J. Appl. Electrochem. 1998, 28, 223. 74. Ch. Papadopoulou, J. Vakros, H. K. Matralis, Ch. Kordulis, A. Lycourghiotis, J. Colloid Interface Sci. 2003, 261, 146. 75. C. G. Vayenas, S. Brosda, C. Pliangos, J. Catal. 2003, 216, 487. 76. C. Koutsodontis, A. Katsaounis, J. C. Figueroa, C. Cavalca, C. Pereira, C. G. Vayenas, Topics in Catalysis 2006, 39, 97. 77. C. G. Vayenas, S. Brosda, C. Pliangos, J. Catal. 2001, 203, 329. 78. S. Brosda, C. G. Vayenas, J. Catal. 2002, 208, 38. 79. C. G. Vayenas, S. Brosda, Stud. Surf. Sci. Catal. 2001, 138, 197. 80. S. Brosda, C. G. Vayenas, J. Wei, Appl. Catal. B 2006, 68, 109. 81. A. Katsaounis, Z. Nikopoulou, X. E. Verykios, C. G. Vayenas, J. Catal. 2004, 222, 192. 82. A. Katsaounis, Z. Nikopoulou, X. E. Verykios, C. G. Vayenas, J. Catal. 2004, 226, 197. 83. G.-M. Schwab, Adv. Catal. 1978, 27, 1. 84. F. Solymosi, Catal. Rev. Sci. Eng 1967, 1, 233. 85. J. Nicole, D. Tsiplakides, C. Pliangos, X. E. Verykios, C. Comninellis, C. G. Vayenas, J. Catal. 2001, 204, 23. 86. X. E. Verykios, in Catalysis, Electrocatalysis at Nanoparticles Surfaces, A. Wieckowski, E. R. Savinova, C. G. Vayenas (Eds.), Marcel Dekker, New York, 2003, p. 745. 87. C. A. Pliangos, I. V. Yentekakis, V. G. Papadakis, C. G. Vayenas, X. E. Verykios, Appl. Catal. B 1997, 14, 161. 88. J. Nicole, PhD Thesis, EPFL, Lausanne, 1999. 89. C. G. Vayenas, G. Pitselis, Ind. Eng. Chem. Res. 2001, 40, 4209. 90. S. Balomenou, D. Tsiplakides, A. Katsaounis, S. ThiemannHandler, B. Cramer, G. Foti, Ch. Comninellis, C. G. Vayenas, Appl. Catal. B 2004, 52, 181. 91. D. Tsiplakides, S. Balomenou, A. Katsaounis, D. Archonta, C. Koutsodontis, C. G. Vayenas, Catal. Today 2005, 100, 13.

1936

8.1 Electrocatalysis

8.1.3

Industrial Electrocatalysis .. Stefan Kotrel and Sigmar Brauninger∗

Introduction: From Industrial Electrochemistry to Electrocatalysis Electrochemical techniques have been successfully employed in the chemical industry for more than a century. Applications range from the large-scale production of commodities, such as chlorine and adiponitrile, to the efficient treatment of industrial wastewater. Even though electrochemistry has earned its respect as a very flexible tool for accessing reaction routes that are not available using methods more familiar to the chemical industry, the usefulness of applying a certain electrochemical process has to be proven in every single case. There have been many valuable papers reviewing the technical and commercial requirements for the successful implementation of electrochemical methods in industry. Table 1, for example, lists certain criteria that have to be met if electrochemical methods are to be successfully introduced in commercial processes [1, 2]. Roughly half of the prerequisites noted in Table 1 are influenced by the right choice of the electrode material and the properties of the electrode surface. The mechanism by which electrons are transferred across the electrode surface to the target molecule (‘‘charge transfer’’) very often determines the course of the electrochemical reaction. Since the chemical composition of the electrode is able to steer an electrochemical reaction in a favorable way without undergoing chemical reactions itself, it is justified to associate this electrode/electrolyte boundary with catalysis. Electrocatalysis in general can be defined as the modification of the reaction rate at an electrode surface depending on the electrode material used. Strictly, the suppression of a competing reaction is also an electrocatalytic effect since the relative reaction rate and therefore the selectivity are affected. Electrode effects can be of primary or secondary type [3]: primary effects are electrocatalytic as they involve the interaction of reactants, products and/or intermediates with the electrode surface having a direct effect on the activation energy of the rate-determining step; secondary effects are related to the structure of the electrical double layer and therefore, once corrected, the reaction rate will result independent of the electrode material [4] (Frumkin effect [5]). In essence, the electrode-related prerequisites for commercial success, as outlined in Table 1, follow two 8.1.3.1

∗

Corresponding author.

main purposes and electrocatalytic goals. First, it is essential to achieve high selectivity with respect to the desired product. Second, efficient electrocatalysis has to reduce electric power consumption as much as possible. The energy input is proportional to the potential difference across the cell, which is directly influenced by the electrode processes involved. The sum of products is proportional to the current passing through the cell (Faraday’s law). For a given current, I , it is therefore necessary to minimize the voltage drop, V . V consists of several components: V = E + η + Vω + Vt

(1)

where E is the thermodynamic (equilibrium) potential difference for the given electrode reactions, η is the sum of the anodic and cathodic overpotentials, Vω is the ohmic drop (‘‘IR-drop’’) of the inter-electrode gap, the electrodes and the connections and Vt is the so-called ‘‘stability’’, i.e. the drift of V with time due to degradation of the electrode performance. For new electrodes, Vt is zero by definition. Strictly, only the overpotential is a measure for electrocatalysis. However, in industrial processes, the objective is to reduce and keep V to a minimum and the overpotential is only one contribution to V . Hence other parameters such as electrode stability as a result of surface modification (Vt ) can have a much more pronounced influence on the electrode performance. The aim of this chapter is to discuss the importance of industrial electrocatalysis on the basis of some commercially relevant and educative examples. It will focus on electrode-related phenomena, which are important for the successful operation of an industrial electrochemical process. The examples are chosen from the fields of both organic and inorganic chemistry. Catalysis in Organic Electrochemistry Over recent decades, organic electrochemistry has grown into a mature field in chemical science and industry. A broad variety of reactions can be performed in all areas of organic chemistry using electrochemical methods [6–8]. One central argument in favor of electrochemistry deals with electrons being ‘‘massfree reagents’’. In a classical redox process, the redox reagent has to be employed at least in stoichiometric amounts and, thus, results in large quantities of spent reagent [6]. By contrast, electrons are completely consumed in the course of the reaction. If the counterreaction at the opposite electrode is chosen carefully, electrochemical methods can completely avoid unwanted products that have to be taken care of [9]. Furthermore, electrochemical reactions are typically carried out at low temperatures and ambient pressure. In this sense, organic electrochemistry can be considered not only as 8.1.3.2

8.1.3 Industrial Electrocatalysis

1937

Important preconditions for the successful implementation of an electrochemical, industrial process [1, 2]

Tab. 1

Factora

1988

High selectivity and product yields High current efficiencies Service life of the electrode Energy (electricity) consumption per kilogram of desired product Concentration of the end product in the electrolyte Membrane lifetime (divided cell) No environmentally problematic byproducts Simple work-up of the electrolysis products Simple removal and recycling of the electrolyte Low emissions

>80% >50% >1000 h 10% >2000 h

a Bold

1995

>70% >4000 h

>8000 h

entries indicate preconditions that are related to electrode properties.

an economically attractive but also as an environmentally friendly technique. Basically, there are two areas, in which organic electrochemistry contributes significantly to commerce and environment: First, organic electrochemistry may be used directly for organic syntheses. Here, in addition to the economical advantages that electrochemistry may offer, it also helps to protect the environment in a preventive way (preventive environmental protection). Harmful side products can be avoided; reactions are carried out at moderate temperatures and pressures, reducing the use of energy. Second, organic electrochemistry can help to deal with environmentally unfriendly substances once they have been produced. In this way, organic electrochemistry serves as a technique that protects the environment in a curative mode (curative environmental protection). Organic electrochemistry has been successfully applied in both the preventive and curative modes. It has found its way into numerous commercial processes and can be considered a competitive and environmentally friendly technique for the chemical industry. 8.1.3.2.1 Synthesis of Adiponitrile The best-known, large-scale electrochemical process to date is the cathodic hydrodimerization of acrylonitrile (ACN) to adiponitrile (ADN), a precursor to hexamethylenediamine, which is the amino constituent of nylon [10]. Monsanto carried out the basic work and scale-up of this process. It is now applied by Solutia, Asahi and BASF. The worldwide production figures of adiponitrile lie well over 400 000 t a−1 [11].

H2 C = CHCN + 2e− + 2H+ −−−→ CN(CH2 )4 CN (2) In the early days of electrosynthesis of ADN, the process was carried out in divided cells using Pb anodes

and 1 wt.% Ag/Pb cathodes. During the mid-1970s, the process performance was improved significantly by the introduction of second-generation technology that uses undivided, bipolar cells based on a stack of closely spaced steel plates and an improved electrolyte composition [12–14]. The electrochemical coupling of ACN to form ADN had been known since the 1940s; the breakthrough came only when Baizer found that the use of quaternary ammonium salts could increase the yield of ADN dramatically and could turn this process into a commercially viable one [15]. In the absence of the quaternary ammonium cations, the hydrodimerization of ACN has to compete with the undesired reduction of acrylonitrile to propionitrile. Since the reaction is carried out in an aqueous solution, the surface concentration of water is large and propionitrile is the main product of the cathodic reduction of acrylonitrile. CH2 = CHCN + 2e− + 2H+ −−−→ CH3 CH2 CN

(3)

If strongly surface-active ammonium ions are present in the electrolyte, the condition on the electrode surface changes towards a far less protic environment. The ammonium ions electrosorb on the electrode surface, replacing adsorbed water in that process. Acrylonitrile is co-adsorbed into the adsorbate layer of these cations. The primarily formed radical anion is effectively shielded from any protic attack, thereby favoring the desired hydrodimerization reaction to ADN [16]. CH2 = CHCN + e− −−−→ (CH2 = CHCN)•−

(4)

2(CH2 = CHCN)•− + 2H+ −−−→ NC(CH2 )4 CN

(5)

Wendt et al. called the catalytic action of adsorbed species ‘‘electrocatalysis of the second kind’’ [17]. The electrode surface itself does not control the course of the References see page 1954

1938

8.1 Electrocatalysis

reaction directly. However, upon adding a non-reacting electro-adsorbent to the system, the surface properties of the anode surface are effectively modified.

R C F

The Simons Process Electrochemical fluorination (ECF) of organic compounds is the prevailing preparation technique for perfluorinated organic compounds, which are industrially very important. Surfactants, refrigerants, lubricants, fire extinguishers, oxygen carriers and polymers are just a few of the most important applications. ECF was invented by the American chemist Joseph Simons at Pennsylvania State University [18]. 3M Company licensed the technology in the 1940s and led it to successful commercialization not long after. Even though 3M remained the most important entrepreneur in this field, other companies, such as Asahi Glass, Mitsubishi Metal and Bayer, succeeded in bringing ECF-based products to the market. In the Simons process, organic molecules are dissolved in anhydrous hydrogen fluoride and are fluorinated at the anode, which usually is made from nickel. Essentially all hydrogen atoms are substituted by fluorine. The electrode processes can be described as follows:

8.1.3.2.2

Cm Hn + nHF −−−→ Cm Fn + 2nH+ + 2ne−

(6)

2nH+ + 2ne− −−−→ nH2

(7)

During electrochemical fluorination, the anode develops a black film of nickel fluoride. In the absence of any organic molecules, the electrode would be polarized and progressively corroded. When an organic substrate is added to the system, the nature of the electrode surface changes and a secondary organic film grows on top of the nickel fluoride film (Fig. 1). The reaction mechanism had been a matter of discussion for many years. Recently, strong evidence emerged that the Simons process is a mediated process, in which NiF2 /NiF3 /NiF4 species catalyze the fluorination of the organic substrate [20, 21]. Bartlett et al. demonstrated that NiF3 and NiF4 generated in situ from K2 NiF6 and HF AHF electrolyte

HF

R

F−

C

F–

H F F

Ni

F F

F

Ni

F

NiF3 + NiF4 e−

e−

Nickel metal

Schematic representation of the fluorination process. After Ref. [25].

Fig. 2

are strong fluorinating agents capable of transforming partially fluorinated compounds into fluorinated products [22, 23]. High-valence nickel fluorides are not soluble in HF [24]. Hence the fluorination in HF is a heterogeneous process taking place on the surface of the Ni anode. The organic molecule is adsorbed on the anode and interacts with two NiF3 or NiF4 sites. The collapse of the organic molecule results in the generation of one HF molecule and one newly formed C–F bond. Figure 2 gives a schematic representation of this process. The entire anode process can be formulated as follows: R

C

H + 2 NiF3

R

F + 2 NiF2 + HF

C

(8) 2 NiF2 + 2 HF

2 NiF3

+

2 H + 2 e−

The Simons process is a step-by-step fluorination process which leads to the formation of all possible partially fluorinated compounds as intermediates on the way towards the final perfluorinated product.

Organic film

Nickel fluoride film Nickel metal

Anode film model for the Simons process [19]. The organic molecules are dissolved in anhydrous hydrogen fluoride (AHF).

Fig. 1

8.1.3.2.3 Bleached Montan Wax – Regeneration of Chromic Acid The technical production of montan wax requires brown coals, which are rich in wax/resin components (‘‘coal bitumen’’). These components, known as raw montan wax, are extracted from the predried coal with organic solvents, often toluene [26]. After the resin has been extracted, the crude montan wax is refined by oxidation/bleaching to yield a light yellow raffinate consisting primarily of long-chain wax acids. Since they

8.1.3 Industrial Electrocatalysis

are non-toxic, bleached montan wax has been approved by the BgVV [Bundesinstitut f¨ur gesundheitlichen Verbraucherschutz und Veterin¨armedizin (Federal German Institute for Consumer’s Health Care and Veterinary Medicine)] and the FDA (US Food and Drug Administration) and other national health authorities of industrial countries for almost all technical applications, especially as coating agents in the pharmaceutical industry. For more than 90 years, Clariant (formerly Hoechst) has been applying the bleaching of montan wax by electrochemically regenerated chromic acid at their plant in Gersthofen [27]. By this process, the dark-colored crude montan wax, which consists of esters of longchain alcohols and carboxylic acids, are hydrolyzed and the alcohols are oxidized by the chromic acid to the corresponding carboxylic acids. At the plant in Gersthofen, the spent reagent is electrochemically regenerated ‘‘ex-cell’’ in a divided cell using a ceramic diaphragm. The original process was developed by Le Blanc [28], who already realized that the efficiency of the anodic oxidation of Cr(III) to Cr(VI) is highly dependent on the choice of the anode material. Pt electrodes, for example, primarily generate oxygen from Cr(III) solution in sulfuric acid, whereas Pb/PbO2 electrodes produce almost exclusively Cr(VI). The first assumption that a less polarized Pt electrode never reaches the potential necessary to oxidize Cr(III) turned out to be insufficient to explain the differences between Pb and Pt electrodes. Gross and Hickling [29] demonstrated that, at sufficiently low current densities, the anodic overpotential of Pt electrodes exceeds the polarization of Pb electrodes under otherwise identical conditions. However, Pb electrodes show dramatically better selectivities for Cr(IV) generation even under the same conditions. It turned out that only Pb electrodes ensured the required Cr(VI) selectivities necessary to carry out an economically viable process. Apparently, the key property of Pb electrodes is their high ability to transfer oxygen generated by the oxidation of H2 O to solvated Cr(III) ions. The competing anodic formation of O2 is efficiently inhibited. At the start of operation, a thin PbO2 layer will form on the surface of the lead anodes: − Pb + SO2− 4 −−−→ PbSO4 + 2e

(9)

− PbSO4 + 2H2 O −−−→ PbO2 + 4H+ + SO2− 4 + 2e

(10) After the PbO2 layer has formed, the electrode surface undergoes constant reduction and oxidation processes in which oxygen is being shuffled from the solvent, H2 O, to the Cr species that is being oxidized. The reaction mechanism is described by the following steps [30–33]:

1939

2− 2Cr(H2 O)3+ 6 + 3PbO2 −−−→ Cr2 O7 + 3PbO

+ 8H2 O + 8H+ +

3PbO + 3H2 O −−−→ 3PbO2 + 6H + 6e

(11) −

(12) Equation (11) represents the chemical step in which the PbO2 transfers oxygen to the Cr(III). K¨appel [31] postulates that Cr ions electrosorb on the anode surface forming Cr(V)–O complexes as intermediates, which in a succeeding step will oxidatively desorb, forming chromates. In a second step [Eq. (12)] the reduced Pb oxide is then electrochemically reoxidized to the original PbO2 . In pure Cr(III)-containing solution this oxidation proceeds with current efficiencies above 90% and only a small amount of current results in oxygen evolution (2H2 O → 4H+ + 4e− + O2 ). A high Cr(III) concentration and elevated temperature (50–80 ◦ C) even increase the selectivity for Cr(III) oxidation. Under technically realistic conditions, current efficiencies above 90% are hardly ever achieved. Typically, an appreciable concentration of organic matter can be detected in the anode compartment of a cell. Organic acids, which are side products of the montan wax bleaching process, are carried over from the bleaching vessel with the spent Cr(III) liquor to the electrochemical cell [34]. The unfavorable effects of organic matter in the anodic chamber can be twofold. First, the carried over organic species can be totally oxidized to CO2 as a result of the prolonged contact with the Pb(PbO2 ) anodes. In contrast to the three-electron oxidation of Cr(III), the total oxidation of organics can easily involve 20–30 electrons. Second, organic matter can change the surface properties of the anodes in such a way that the once shut reaction pathway towards oxygen evolution becomes accessible again. Herbst et al. [35] determined the impact of dicarbonic acids on the current efficiency of chromic acid regeneration. Dicarbonic acids are typically present in the spent Cr(III) liquor that is to be regenerated in the electrochemical cell after the bleaching of montan wax. A careful analysis of the gases evolving at the anode of the electrochemical compartment revealed that the O2 content was always larger than the amount of CO2 . Hence, in the case of montan wax bleaching, the surface selectivity loss of the anode by adsorption of organic compounds seems to impair the efficiency of the regeneration process more than the mere competing oxidation of organic impurities. 8.1.3.2.4 Waste Water Treatment Industrial wastewater contains organic pollutants, which have to be removed before the water can be discharged [36]. Biological References see page 1954

1940

8.1 Electrocatalysis

treatment is the most economic process and is usually used for the treatment of ‘‘readily degradable’’ organic components present in the water. The situation is entirely different when the water contains toxic or/and refractory (non-biocompatible) organic pollutants. In these instances, organic electrochemistry serves as a powerful tool and makes an important contribution to the curative protection of the environment [37–39]. There are a variety of electrochemical methods that are available for the destruction of organics in aqueous wastes. Brillas et al. [40] give an overview of the electrochemical methods that are particularly relevant for this type of application. Principally, one can distinguish between direct and indirect electrochemical procedures. Indirect methods deal with the use of redox mediators. In those cases only the reagent is generated at the electrodes. The actual destruction of the organic waste occurs independently of the electrochemical process. The direct electrochemical destruction of organics proceeds at the electrode surface and is highly dependent on the properties of the electrode/electrolyte boundary. Hence, from an electrocatalytic point of view, the direct method offers a deeper insight into the nature of electrocatalysis. Therefore, the direct electrolysis of organic waste will be primarily discussed in the following section. A Anodic Oxidation of Toxic compounds Basically, there are two strategies that can be pursued. First, non-biodegradable compounds can be electrochemically oxidized into biocompatible organics, which are

then subjected to the regular biological treatment of wastewater. Second, the toxic organics can be completely combusted to CO2 . Figure 3 depicts the two alternative routes for wastewater treatment. The complete digestion of organic toxins at the anode appears to be the best natural choice if an electro-oxidative treatment of wastewater is to be implemented. However, since most wastewater streams carry only low concentrations of organic contaminants, the complete conversion to CO2 may require unsuitably large electrode areas, thus making the process economically impracticable. In some instances it is more viable to allow partial electro-oxidation of the non-biodegradable toxins. The final destruction of the electro-oxidized contaminants can then be carried out by applying conventional biological techniques. Typical waste materials that are treated following route 1 in Fig. 3 are phenols, benzoquinones and aniline. Figure 4 gives a typical reaction sequence for the electrochemical breakdown of these organic toxins. In many cases maleic and oxalic acid, which can be readily biodegraded, are the final products [41]. Electrochemical combustion is the method of choice in the following cases: • treatment of wastewater with a composition prohibitive for any, even partial, biological treatment (high amount of salt, low pH) [42] • treatment of organic wastes in closed life support systems [43] • water disinfection for domestic water treatment [44].

Route 1: Selective conversion Electrochemical Nonbiodegradable Conversion

Biodegradable

Biological Treatment

CO2

Route 2: Total combustion Electrochemical

Nonbiodegradable

Fig. 3

CO2

Combustion

The two routes for the electro-oxidative treatment of non-biodegradable toxic compounds in waste water.

COOH OH

H2O

H++e−

OH

H2O

H++e−

O

H2O

H++e−

COOH Maleic acid Fumaric acid

OH O

COOH COOH Oxalic acid

Fig. 4

Reaction sequence of the electrochemical combustion of phenol [41].

8.1.3 Industrial Electrocatalysis

a Creating Reactive Oxygen Surface Species The electrocatalytic requirements for the electrode material for each route are completely different from one another. For route 1 in Fig. 3, the ideal electrode material must catalyze a specific reaction that converts non-biodegradable into biodegradable substances with high selectivity, otherwise side products threaten to deactivate the electrode surface. If complete electrochemical combustion of the organic contaminants is intended, the electrode material has to facilitate the total electrochemical digestion of the organics to CO2 and water. To gain a better understanding of the mechanism that is responsible for the electrochemical behavior of a certain electrode, it is worthwhile considering the anodic O2 evolution in the absence of any organic material. In a first step, water is progressively dehydrogenated, leaving behind basically two types of reactive oxygen species on the electrode surface. The following equations describe how the anodic dehydrogenation of water proceeds on the electrode surface:

MOx + H2 O −−−→ MOx (· OH) + H+ + e−

(13)

MOx (· OH) −−−→ MOx+1 + H+ + e−

(14)

The initial products of the oxidative dehydrogenation of water are OH radicals loosely adsorbed on the electrode surface. In a subsequent step, these OH radicals can react with the electrode surface, bringing the electrode surface to a higher oxidation state [45]. Hence one can distinguish between two reactive oxygen surface species, which are both capable of producing O2 in a subsequent step [46]. Which oxygen species actually prevails on an electrode surface depends on the rates of hydroxyl radical formation and the rate of transition of the oxygen into the oxide lattice. If OH radicals are formed on the electrode surface much faster than the oxygen can be built into the oxygen lattice structure of the electrode, OH radicals will be the dominating active oxygen species on the electrode surface. If oxygen can be transferred to the oxygen lattice easily – for example because of a high concentration of oxygen vacancies on the electrode surface – and the OH radical formation becomes the limiting factor, the electrode will be rich in chemisorbed oxide species. Quadrants I and II in Fig. 5 give a generalized scheme of the reaction cycle that governs the anodic evolution of O2 from water [37, 40, 47]. Depending of the nature of the electrode, physisorbed OH radicals or chemisorbed oxide species may be the dominant oxygen species during electrolysis. b Selective Conversion vs. Total Combustion In the case of the electro-oxidative destruction of pollutants, the organic toxins have to participate in the electrode reactions on the

1941

½ O2

H++e− H++e− MO(·OH)

2 H++2e− MOx +1

MOx H2O z H+ + z e−

H2O

R

R

z CO2

RO

Generalized reaction scheme of O2 evolution (I and II) and electro-degradation of organics in aqueous media (III and IV). Quadrants I and III depict the situation for electrode surfaces rich in OH radicals. Quadrants II and IV represent the situation in the case of partially oxidized electrode surfaces.

Fig. 5

surface of the anode. If OH radicals are the prevailing reaction partners, the complete destruction of the organic molecule will be likely, mostly because the OH radical triggers a chain reaction, in which scission reactions continue rapidly until the complete toxic molecule is converted to CO2 and H2 O [48] (quadrant III in Fig. 5). Additionally, oxygen that is also generated at the anode can contribute to the combustion process by forming organic hydroperoxides, ROOH, which are relatively unstable and easily decompose, leading to molecular breakdown of the molecule [40]. Highly selective oxidation will occur only if the OH surface concentration on the electrode is close to zero and the organic reacts exclusively with active MOx+1 surface centers. For example, oxides such as PtO, IrO2 and PbO2 , having a high surface concentration of oxygen vacancies, can stabilize MOx+1 on the electrode surface [36] (quadrant IV in Fig. 5). To achieve a high OH radical concentration on the anode surface, the onset potential for O2 evolution has to be as large as possible. Figure 6 [42, 49] compares the onset of O2 evolution for some common electrode materials. Figure 6 is a helpful guide if one wants to choose the appropriate electrode material. Materials such as Pt and IrO2 have a relatively low overpotential towards O2 evolution. Hence the OH radical concentration during operation remains relatively low, making these materials the right choice for the selective electro-oxidation of organics. On the other hand, boron-doped diamond (BDD) electrodes possess a large overpotential of O2 References see page 1954

1942

8.1 Electrocatalysis

Diachem Si/BDD Ebonex PbSn PbO2 Graphite IrO2 Pt 0

0.5

1

1.5 2 Overpotential / V

2.5

3

Potential of oxygen evolution for some common electrode materials: commercial boron-doped diamonds (BDD) [50], Si-coated BDD [51], titanium oxide (Ebonex) [52], PbSn [50], PbO2 [53], graphite [54], IrO2 [50, 55] and Pt [50].

Fig. 6

evolution. Hence the high OH concentration on a BDD electrode surfaces will result in total combustion of the pollutants. c Common Electrode Materials Over recent decades, PbO2 and SnO2 electrodes have been intensively studied as promising materials for electrochemical combustion [56–59]. PbO2 electrodes are relatively cheap and effective in oxidizing pollutants. Unfortunately, there is a possibility that toxic Pb2+ ions will be generated, causing severe secondary pollution. Due to their semiconducting nature, pure SnO2 electrodes exhibit a very high resistivity at room temperature and therefore cannot be used as an electrode material directly. Hence doped SnO2 materials are typically used as electrodes. In electrochemical applications, Sb is the most common dopant for SnO2 . K¨otz et al. were the first to report on the anodic oxidation of pollutants on Sb-doped SnO2 -coated titanium electrodes (Ti/SnO2 –Sb2 O5 ) [60]. Since then, a number of publications have confirmed the high activity of SnO2 −Sb2 O5 electrodes for anodic wastewater treatment, exceeding even the activity of PbO2 [61–64]. However, despite the high efficiency for pollutant oxidation, SnO2 −Sb2 O5 electrodes lack sufficient electrochemical stability just like PbO2 . A more appropriate material for the electrochemical combustion was found in doped TiO2 electrodes. As with SnO2 electrodes, TiO2 is a poor electric conductor at room temperature. Doping TiO2 with Nb or Ta sufficiently lowers the resistivity of this material [65]. TiO2 -based electrodes are usually made of a Ti substrate that is coated with Nb- or Ta-doped TiO2 films. These electrodes are stable at low current densities, but suffer from a shortened lifetime at larger current densities [66, 67].

An alternative titanium oxide material that can be used as anode material is a non-stoichiometric Ti4 O7 /Ti5 O9 phase made by heating TiO2 to 1000 ◦ C in the presence of H2 . The resulting electrode material is stable in aqueous media. The brand name of this material is Ebonex [68]. In recent years, boron-doped diamond (BDD) thin-film electrodes have been studied intensively with regard to industrial applications. In addition to their outstanding chemical inertness, BDD electrodes show a very high onset potential for O2 evolution. Hence this material is very efficient in terms of OH radical generation. The diamond films are prepared using a variety of chemical vapor deposition (CVD) techniques. Ti is considered to be the best substrate for BDD electrodes [69]. Ti is fairly cheap and has advantageous attributes that distinguish it from other electrode materials. Ti is sufficiently electrically conductive, mechanically robust and electrochemically inert. Other relevant substrate materials are Si, Nb, Ta, Zr and W [70]. Nowadays, stable BDD films can be achieved by using a hot filament CVD (HTCVD) process. This technique can be applied to substrates with various geometries up to 0.5 m2 [71]. With HTCVD, BDD coatings of substrates can be produced fairly inexpensively with a deposition rate up to 1.2 µm h−1 and boron concentrations in the film of 500–8000 ppm. The thickness of the BDD layer on the substrate is typically in the range 1–10 µm [70]. B Cathodic Reduction of Toxic Compounds Cathodic electro-reduction is well suited for the dechlorination of pollutants in waste waters. Chlorinated compounds are produced in large quantities in the industry. Common products for chlorinated organics are refrigerants, pesticides and transformer oils (among others). Since most of such compounds are not biodegradable, they have to be removed from the main waste water stream. This separation is typically carried out by applying standard adsorption (active carbon) or extraction (with organic solvents). Once separated, the toxic compounds have to be processed and converted into harmless products by electrochemical means. One important prerequisite for optimal electrochemical conversion efficiencies is a high hydrogen overpotential of the chosen electrode [72–74]. Even though a variety of materials, such as graphite, Pb, Hg, Ni, Ti, TiO2 , meet this requirement, electrode degradation processes often reduce the number of applicable materials. Graphite, for example, often develops cracks and fractures along its basal planes. These fractures are caused by electrophoretic processes, which cause ions and organic molecules to diffuse between the graphite layers. Three-dimensional carbonaceous materials comprising partially graphitized amorphous carbon and graphite felts seem to lessen this

8.1.3 Industrial Electrocatalysis

problem [75, 76]. Pb is another popular electrode material, but is prone to stability problems during electro-reductive dehalogenation [76, 77]. The decomposition of CHCl3 and CCl4 is a very successful example of the degradation of chloroalkanes. These compounds react only reluctantly with OH radicals to give CO2 , making the electro-reductive route for detoxification particularly attractive [43, 78]. The reduction can be carried out at room temperature at potentials of about −760 mV vs. the standard hydrogen electrode (SHE). The general reaction can be written as follows: RCl + 2H+ + 2e− −−−→ RH + HCl

(15)

The electrode material has a significant impact on the end product of the electrolysis. Sonoyama et al. [79], for example, reported that the electro-reduction of chloroform on Ag, Zn, Pd and Cu cathodes yields almost 100% CH4 , whereas Pb electrodes show a high selectivity for H2 CCl2 . The highest conversion rates for the dehalogenation of organics can be achieved using Cu electrodes [43]. Gemmler et al. [80] ascribe the electrocatalytic effect of Cu electrodes to the fact that Cu surface atoms can readily change their state of oxidation between Cu0 , Cu+ and Cu2+ . At negative potentials of the cathode, these oxidized Cu species are only stable due to the formation of strong chloro complexes. Chlorinated phenols can be electro-reduced using platinized or palladized carbon electrodes [43, 81]. Basically, the mechanism resembles the conventional hydrogenation of aromatics over precious metal catalysts with the hydrogen being generated electrolytically from water. Figure 7 shows the mechanism of this process. The chlorinated aromatic is adsorbed on the carbon surface near the metal/C interface, where the actual hydrogenation takes place. The electro-reduction of substituted phenols yields cyclohexanols, which can be degraded by biological degradation. Chlorofluorocarbons (CFCs) are known to be highly destructive to the stratospheric ozone. Their production has been reduced gradually over the last decade. Since hydrofluorocarbons (HFCs) and fluorocarbons (FCs) are harmless to stratospheric ozone, there is currently no limitation on their production. Therefore, the electro-reductive conversion of CFCs to HFCs and FCs has recently attracted some interest. Since CFCs are not soluble in water, wet-proofed porous electrodes, which are also used for fuel cells, are used. Compared with conventional electrodes, they not only accelerate electrode reactions catalytically, but also lower diffusion limitations and simplify product isolation. Kornienko et al. [82] prepared wet-proofed electrodes consisting of PTFE, acetylene black and Cd powder. This electrode allows the conversion of CFCs at

1943

Cl− H

OH Cl

H

H+

OH

H

H+

Cl

Pd

OH

2e−

Mechanism of the electro-reductive hydrogenation of 4-chlorophenols on palladized carbon surfaces [43, 81]. In a first step. protons are cathodically reduced to hydride ions (I). In a second step, hydride ions react with the 4-chlorophenols, which are adsorbed on the carbon carrier adjacent to the Pd particles (II).

Fig. 7

35 ◦ C in the presence of 3 M LiCl. Another approach relies on metal-supported gas diffusion electrodes, which show up to 100% efficiencies without any current losses due to H2 production [83]. Catalysis in Inorganic Electrochemistry

8.1.3.3

8.1.3.3.1 Chloralkali Electrolysis Chlorine is an important chemical – more than 50 million tonnes per year are produced worldwide by chloralkali electrolysis [84]. Chlorinated intermediates are indispensable in various chemical syntheses, and poly(vinyl chloride) (PVC) is one of the most important polymers. Three different techniques are used in the chemical industry to produce chlorine by electrolysis: diaphragm cells, the amalgam and, the most sophisticated, membrane cells. All of them have advantages and disadvantages, and details can be found in the literature [84–87]. Among the innovations that have entered chlorine production technology during the last 50 years, two will be dealt with in detail in this chapter, and electrocatalysis is a key for the success of both: dimensionally stable anodes (DSAs) and oxygen-depolarized cathodes (ODCs).

A Electrocatalytically Activated Dimensionally Stable Chlorine-Evolving Electrodes (DSAs) Chloralkali electrolysis was technically realized more than 100 years ago. With respect to thermodynamics, two reactions could take place at an anode immersed in aqueous sodium chloride solution, oxygen and chlorine evolution:

3H2 O −−−→ 1/2O2 + 2H3 O+ + 2e− −

−

2Cl −−−→ Cl2 + 2e References see page 1954

(16) (17)

1944

8.1 Electrocatalysis

However, due to lower exchange current densities, oxygen evolution is kinetically hindered and chloride oxidation is the dominant reaction. Since the beginning of the 20th century, graphite anodes have been used. Graphite is chemically not stable as it suffers from oxidation. As a result, the inter-electrode spacing changes with operating time. This problem has stimulated intensive materials research in the chloralkali industry, and in the late 1950s Henry Beer found a solution, which today is called the dimensionally stable anode (DSA) [84, 88]. It consists of a titanium substrate coated with a catalytically active layer of a metal oxide mixture (TiO2 −RuO2 ). The introduction of DSAs revolutionized chloralkali technology [89], not least by allowing the replacement of amalgam technology by membrane technology [90], which is more energy efficient. B Electrocatalysis of Anodic Chlorine Evolution at RuO2 Anodes Graphite anodes used in the 1950s [91] had a high overpotential for chlorine evolution, and corrosion of the anode widened the inter-electrode gap [92]. As a consequence, the voltage drop V increased with time, and spent anodes had to be replaced periodically. High operating and maintenance costs stimulated the chloralkali industry to look for materials with better catalytic properties, i.e. lower overpotential (η), and better stability (Vt ). In Fig. 8, current–voltage curves of amalgam cells measured with graphite and RuO2 anodes in brine solution of a concentration which is typical for chloralkali electrolysis (310 gNaCl dm−3 ) are compared [93, 94]. Obviously chlorine evolution is less hindered at DSA electrodes. However, oxide coatings for DSAs from a pure electrokinetic point of view are not the best catalysts, i.e. they do not result in the lowest charge-transfer resistance [93].

5.1

Cell voltage / V

4.9 1

4.7 4.5

2

4.3 4.1 3.9 5

7

9

11

13

15

17

19

Current density / kA m−2

Cell potential difference using (1) graphite anodes and (2) DSAs. NaCl concentration, 310 g dm−3 ; temperature, 60 ◦ C [93, 94].

Fig. 8

De Dora mercury cells: comparison of performances with graphite and DSAsa

Tab. 2

Parameter Anode potential/V Cathode potential/V Anode ohmic drop/V Electrolyte ohmic drop/V Gas bubble effects/V Current efficiency/% Energy consumption/kWh t−1

Graphite 1.47 −1.85 0.15 0.60b 0.90 96 3910

DSA 1.37 −1.85 0.15 0.40c 0.13 97 3040

density 10 kA m−2 . distance 3 mm. c Anode–cathode distance 2 mm. a Current

b Anode–cathode

Table 2 gives a breakdown of the components of the voltage drop (V ) at 10 kA m−2 [94]. The applied V is 4.97 V with graphite and 3.90 V with DSA: the difference is 1.07 V. However, only 0.1 V is related to electrocatalytic effects, 0.2 V being due to a lower ohmic drop in the electrode gap and 0.77 V to the disappearance of gas bubble effects using a DSA. It is evident that the electrocatalytic activity is by itself a marginal reason for the good performance. However, chlorine evolution has a lower overvoltage at the DSA surface compared with oxygen formation and the nature of the catalysts allows cell engineering solutions which are not possible with other materials. Moreover, the stability (Vt ) is greatly improved. At first, precious metal alloy-plated Ti electrodes were used [95–97]. The use of a metallic material such as titanium as a substrate permitted dimensional flexibility, and the nets, meshes or expanded metals can be optimized with respect to hydrodynamic properties so that the chlorine generated can escape rapidly [98–101]. As a consequence, inhibitory gas bubble effects are minimized and the inter-electrode gap can be reduced compared with graphite electrode designs and thus the voltage drop is lowered. Furthermore, the concept of electrode activation by using small amounts of comparably expensive but stable coatings was introduced. Today, electrode activation has become common practice in the field of supported gas diffusion electrodes [102], e.g. for fuel cells. Conversion rates depend on electrokinetic effects (i.e. low activation energy/high reaction rates) and geometric parameters (i.e. high surface area). In an industrial process, it does not matter in principle whether high conversion rates are accomplished by reducing the activation energy or by increasing the actual surface per apparent unit area. Activated oxide electrodes show significantly higher surface areas than the geometric ones [103]. Roughness factors of 10–100, sometimes even 1000, comparable to Pt blacks, are achieved. Oxide electrodes

8.1.3 Industrial Electrocatalysis

have much higher catalytic activity than graphite or Pt–Ir alloys, comparable to that of the metal alloys [104]. If, however, the roughness factor is accounted for, the metal alloy turns out to be more active than the oxide. C Details of the Electrocatalytic Process The details of the electrocatalytic reaction on transition metal oxide electrodes have been analyzed in detail [105–107]. According to Krishtalik [106], chlorine evolution is characterized by a low Tafel slope of about 30 mV for RuO2 –TiO2 . For pure RuO2 the anodic slope is 40 mV and the cathodic slope is ca. 120 mV. The low slope of the anodic polarization curve in combination with the low surface coverage by intermediates has been interpreted in such a way that transfer of the first electron should be reversible, whereas the transfer of the second electron or the subsequent chemical reaction represents the slow step. Having proven by experiments that the reaction is first order with respect to chloride, Krishtalik concluded that the slow step of the anodic reaction consumes a total of two electrons per Cl− ion, i.e. it leads formally to the univalent positive chlorine state ‘‘Cl+ ’’. As two Cl− ions are required to form a Cl2 molecule, the univalent positive chlorine should react with chloride to give Cl2 . Krishtalik also concluded from the negative reaction order of the cathodic Cl2 reduction with respect to chloride, that chloride is formed in an equilibrium step preceding the slow step, i.e. the reaction of ‘‘Cl+ ’’ with Cl− being reversible. A three-step reaction mechanism with the rate-determining transfer of the second electron and the formation of the nominally unipositive chlorine as an intermediate has been formulated:

+

Cl− −−−→ Clads + e−

(18)

Clads −−−→ ‘‘Cl+ ’’ + e−

(19)

−

−−  ‘‘Cl ’’ + Cl −− − − Cl2

(20)

However, the nature of the positive univalent chlorine ‘‘Cl+ ’’ has not been elucidated unambiguously. Thus, it may be the Cl+ cation or some hydrochlorous acid derivative. In the latter case the oxygen necessary for its formation can come either from water or from RuO2 , i.e. the oxidic surface layer. Finally, the second step [Eq. (19)] can be interpreted not as the oxidation of Clads but as a change in the valence state of the adsorption site, e.g. an Run+ state. D Preparation and Formulation of the Coatings for DSAs Metal-plated Ti electrodes were found to be active but unstable in the chloralkali environment. However, the generation of oxidic layers by means of decomposition of

1945

suitable precursors turned out to be the method of choice with respect to chemical stability [95, 108]. Furthermore, these coatings made of activated mixed metal oxides reveal metallic conductivity [109, 110], which is an important property. The metals the oxides of which form the catalyst coating are dissolved in an appropriate manner (chlorides, organometallics) in an organic solvent of low volatility generating a so-called ‘‘paint’’ [17]. Typical paint examples of metal compounds are RuCl3 , H2 PtCl6 , CoCl2 (the last two may be added to modify the selectivity), titanium tetrabutylate and tantalum trichloride (these might be added to improve the adhesion and strength of the coating by forming dispersed TiO2 or Ta2 O5 particles or mixed oxides containing TiO2 ). Long-term adhesion and resistance against erosion are important for the stability of DSA coatings. Thus, as a first step of coating preparation, the TiO2 layer a few micrometers thick, which usually covers titanium metal and would prohibit the formation of a low-resistive and tight contact between support metal and coating, must be etched away by oxalic or hydrofluoric acid so that only a very thin oxide layer remains on to which the coating is applied. An oxide additive to RuO2 that assures good adhesion and improved conductivity of the interlayer between the titanium metal and the coating is necessary. It combines physical strength, hardness and erosion stability and simultaneously prevents the anodic oxidation of the titanium support, which would generate insulating, nonconducting titanium dioxide. TiO2 added to RuO2 by proper formulation of the catalyst paint, which crystallizes in the rutile-type structure, forms semiconducting mixed RuO2 −TiO2 phases and stabilizes the electrocatalyst mechanically and chemically without significantly affecting the catalytic activity [98, 111]. The TiO2 content of the coating may vary from 30 to 70 mol%. Also, Ta2 O5 is added to RuO2 in order to stabilize the coating. IrO2 , a catalytically active additive, also seems to improve the wear rate [112]. The ‘‘paint’’ is applied to the sandblasted, etched titanium electrode by brushing, dipping or spraying. The solvent of low volatility (e.g. butyl alcohol, turpentine oil) is evaporated at 100–200 ◦ C, and the metal compounds now covering the Ti electrode are converted to the respective oxides by heating the electrodes in air at temperatures ranging from 400 to 600 ◦ C [113, 114]. The temperature must be kept below 600 ◦ C because above this temperature non-conducting TiO2 inter-layers will grow on the substrate [115]. Trasatti [113] has shown that a curing temperature of approximately 450 ◦ C is optimal because below this References see page 1954

1946

8.1 Electrocatalysis

temperature oxide formation is incomplete, rendering Clcontaining, less active and more soluble coatings, and above 450 ◦ C the mixed oxides lose more and more residual chemically bound water and become more ordered and catalytically less active. The procedure of brushing and tempering is repeated several times, each time with thorough cleaning of the surface from oxidic dusts until the desired coating load of 10–40 mg cm−2 is achieved.

utilized either as a raw material for other chemical products or for chlorine recycling. Chlorine recycling is carried out by electrolysis as depicted in Fig. 9a. The cell reactions of the classical HCl electrolysis are as follows: Cathode: 4H+ + 4e− −−−→ 2H2 E0 = 0 V −

E Lifetime of Dimensionally Stable Chlorine-Evolving Anodes Stability basically has nothing to do with electrocatalysis. However, a factor which influences activity and at the same time limits stability is the formation of surface-soluble products. A redox reaction on the surface could enhance the catalysis of a reaction, but if the surface became soluble or the layer decomposed, the coating would be consumed. Electrode surface corrosion would be detrimental to electrocatalysis, and the potential drop would increase rapidly. In the case of RuO2 , however, up to a certain potential the surface is totally insoluble, but as soon as ruthenates or perruthenates are formed on the surface sites, they are dissolved [116–118]. Thus, at potentials below a certain limit, DSAs are electrochemically stable. It is normally O2 evolution which takes place with formation of soluble species because the electrode potential rises beyond the limit. Corrosion can be reduced by lifting the limit of the redox reaction. This is achieved by mixing RuO2 with other oxides, namely IrO2 . Mixtures of RuO2 and IrO2 are less corrodible because the intimate mixing of these two oxides lifts the limiting potential of surface oxidation [119]. IrO2 corrodes less than RuO2 by itself but also the corrosion rate of IrO2 is decreased in the presence of RuO2 [120]. The lifetime of RuO2 -coated anodes has been substantially improved by sophisticated coating formulations. At typical current densities, i.e. 2–3 kA m−2 for diaphragm cells and 4–6 kA m−2 for membrane cells [85], the catalytic coatings last at least 5 years. Even lifetimes of up to 10 years have been reported [121] which is equivalent to a production of 300 000 kg of chlorine per square meter. In mercury cells with current densities of 8–13 kA m−2 [85], the lifetimes amount to approximately 3–4 years [17]. F The Cathode Reaction – Oxygen-Depolarized Cathode (ODC) The solution of the anode problems in chloralkali cells led to emphasis on the cathode problems, either by activating steel or nickel [122, 123] or by trying to reduce losses by replacing H2 evolution with O2 reduction [124, 125].

Anode: 4Cl −−−→ 2Cl2 + 4e E0 = +1.36 V

(22)

4HCl −−−→ 2Cl2 + 2H2 E = +1.36 V

(23)

where E0 denotes the individual electrode potentials and E is the potential difference of the overall reaction. Unfortunately, HCl electrolysis is an energy-consuming process, and electrocatalysis is the key for improvements in efficiency. During the 1990s, Bayer, Uhdenora – a joint venture of Uhde (Dortmund) and De Nora (Milan) – and De Nora (USA) developed a method for reducing the energy input in industrial-scale electrolysis by applying gas diffusion electrodes (GDEs) [126–130]. The development took advantage of the progress which had been made in the development of GDEs for fuel cells. The ODC is in contact with an ion-exchange membrane (e.g. Nafion) on one side and on the back side oxygen is fed continuously (Fig. 9b). Unlike classical HCl electrolysis, where protons are reduced to form hydrogen on the cathode at a potential of theoretically 0 V vs. the normal hydrogen electrode (NHE), oxygen is reduced in the case of an ODC starting at +1.23 V (NHE): Cathode: O2 + 4H+ + 4e− −−−→ 2H2 O E0 = +1.23 V Anode:

−

4Cl −−−→ 2Cl2 + 4e E0 = +1.36 V

4HCl + O2 −−−→ 2Cl2 + 2H2 O E = 0.13 V

(24)

−

(25) (26)

By using this approach, the cell voltage theoretically can be reduced by 1.23 V (Fig. 10), which equals the thermodynamic open-circuit voltage of water splitting. The potential difference between oxygen reduction and hydrogen evolution is directly correlated with energy saving. In real systems, however, overvoltages occur at oxygen reducing electrodes. As a consequence, the observed depolarization is smaller. Preparation of Oxygen-Depolarized Cathodes Oxygen-depolarized cathodes (ODCs) consist of conductive webs such as carbon cloths or metal meshes coated with nanoporous carbon material and a supported 8.1.3.3.3

Electrolysis of Hydrochloric Acid In many industrial syntheses using chlorine, hydrochloric acid is produced as a by-product. Hydrochloric acid can be 8.1.3.3.2

(21)

−

8.1.3 Industrial Electrocatalysis

1947

O2

e− −

e

2 Cl H+

(a) Anode Fig. 9

−

1/2 O2

H2

Cl2

HCl

2

H+

Diaphragm

e−

Cl2

HCl

−

e

e− e−

HCl

2

(b) Anode

2 H+

Cl− H+

Cathode

e− e−

Membrane

H2O Cathode

H2O

HCl electrolysis: (a) with cathodic hydrogen evolution and (b) using an ODC.

Potential / V CI2 ODC

+1.36 +1.23

0V

Classical electrolysis

Ca. 1 V

O2

Current density / A.m−2

H2 Fig. 10 Schematic polarization curves for an HCl cell using classical (H2 -evolving) and oxygen-consuming (ODC) cathodes [85, 130].

precious metal catalyst (e.g. ELAT). Initial experiments were carried out using platinum [131]. However, platinum catalysts were not stable under the corrosive conditions of chloralkali electrolysis. In contrast, electrocatalysts based on rhodium and ruthenium sulfide are sufficiently active for oxygen reduction and stable operation could be demonstrated. Noble metal sulfides for use in electrocatalysis can be prepared by sparging hydrogen sulfide in an aqueous solution of a corresponding noble metal precursor, usually a chloride [87, 132]. The synthesis of noble metal sulfide catalysts with hydrogen sulfide in aqueous solutions is conveniently carried out in the presence of a conductive carrier, in most of the cases consisting of carbon particles. In this way, the noble metal sulfide is selectively precipitated on the carbon particle surface, and the resulting product is a carbon-supported catalyst, which is particularly suitable for incorporation into gas-diffusion electrode structures characterized by high efficiency at reduced noble metal loadings. High surface area electronconducting carbon blacks, such as Vulcan XC-72 from Cabot (USA) are commonly used.

A different procedure for the preparation of carbonsupported noble metal sulfide catalysts consists of an incipient wetness impregnation of the carbon carrier with a noble metal precursor salt, for instance a noble metal chloride, followed by solvent evaporation and gas-phase reaction under dilute hydrogen sulfide at ambient or elevated temperature, whereby the sulfide is formed in a stable phase [133]. In an alternative synthetic route, noble metal sulfide catalysts are obtained by reacting an appropriate precursor, preferably a chloride, with a thionic species (thiosulfates, thionic acids, etc.) in an aqueous solution and subsequent deposition on high surface area carbon black. After deposition, filtration and drying, a thermal treatment is carried out. In the case of rhodium, prior to its use, the noble metal sulfide catalysts obtained are subjected to a suitable stabilizing thermal treatment between 300 and 700 ◦ C [133]. First Industrial HCl Electrolysis Using ODCs At Bayer’s Brunsb¨uttel site, the first industrial ODC electrolysis unit for chlorine recycling from hydrochloric acid went on-stream at the end of 2003 with a capacity of 20 000 t of chlorine per year [134, 135] (Fig. 11). With a cell voltage of 1.5 V at 4–5 kA m−2 , 30% of electrical energy (600–700 kWh t−1 Cl2 ) can be saved. The cell structure is completely metallic and there is no part made of graphite [136]. Between the two chambers a membrane is interposed. The membrane is pressed against the oxygen-depolarized cathode by the differential pressure existing in the two compartments in addition to pressing the ODC against the cathodic current distribution mesh. The combination of the above concepts, i.e. ODC and metal electrolyzer, constitutes the core of the new technology. Key performance figures are as follows [136]: 8.1.3.3.4

• operating current density: 4–5 kA m−2 • HCl concentration: 14–15% References see page 1954

1948

8.1 Electrocatalysis

may become completely flooded with the NaOH solution, and the cathode, where the hydrogen evolution takes place, will operate at much lower potential [139]. The presence of oxygen and significantly higher potential of the oxygen-depolarized cathode compared with the hydrogen-evolving cathode can result in corrosion of the cathode hardware and other components in the cathode compartment, which would not be observed in the standard membrane cell thanks to the cathodic protection and reducing properties of hydrogen. Another difference of practical importance between the hydrogen-evolving cells and oxygen-depolarized cells is the relative stability of the intermediate products of the hydrogen evolution reaction (HER) and the oxygen reduction reaction (ORR). The HER does not generate any critical intermediates [140]: Electrolysis of hydrochloric acid using oxygen-depolarized .. cathodes at Bayer’s Brunsbuttel site. Source: Bayer MaterialScience AG. Fig. 11

2H2 O + 2e− −−−→ H2 + 2OH−

The ORR may follow the desired four-electron path: O2 + 2H2 O + 4e− −−−→ 4OH−

• operating pressure: chlorine at 1.2–1.25 bar, oxygen at atmospheric pressure • temperature: ≤ 60 ◦ C • power consumption: 1000–1100 kWh t−1 (Cl2 ) • chlorine quality: 99.8–99.9% Oxygen-Depolarized Cathodes for Chloralkali Electrolysis The implementation of ODCs in chloralkali electrolysis is still at the experimental stage since the cell design for sodium chloride electrolysis is much more complicated [135]. For example, the ODC and the ionexchange membrane cannot come into direct contact with each other. The reason is that sodium ions migrating through the membranes do not carry enough water to hydrate the membrane and no water is produced on the cathode. Thus, the membrane resistance increases and the membrane can even be destroyed under these conditions. The electrode has to be designed to facilitate the formation of the three-phase boundaries (gas/liquid/solid) that involve oxygen, water/sodium hydroxide solution and the catalyst particles. Moreover, it has to manage effectively the transport of oxygen to and sodium hydroxide from the catalyst layer. Although gas diffusion electrodes (GDEs) have been extensively used in fuel cells and optimized for this application, the conditions encountered in the cathode compartment of a chloralkali electrolyzer are different. The high viscosity of the concentrated NaOH and its strongly corrosive properties may have a negative influence on both the formation of the three-phase boundaries in the electrode pores and transport of the reagents and products to and from the reaction site [124, 137, 138]. In the worst-case scenario, the electrode pores 8.1.3.3.5

(27)

(28)

Alternatively, it can also occur according to the twoelectron mechanism, which results in the generation of peroxide: O2 + H2 O + 2e− −−−→ OH− + HO− 2

(29)

Although the peroxide eventually decomposes and produces an equivalent amount of the hydroxide and hence does not lower the overall caustic current efficiency, it is a rather troublesome by-product, because it produces gaseous oxygen upon decomposition and may also precipitate in the highly concentrated NaOH: 2Na+ + OH− + HO− 2 −−−→ Na2 O2 ↓ +H2 O

(30)

Precipitation of sodium peroxide [Eq. (30)] can cause liquid and gas flow maintenance problems, block the electrode active surface area and even destroy the microporous structure of the gas diffusion electrode [138, 141]. Although only few technical details about pilot-scale plants that utilize oxygen-depolarized cells for chloralkali electrolysis are known, it seems that finite-gap designs are the most prevalent [140]. In cells of this type, the gas diffusion electrode acts as a separator of the oxygen and caustic chambers. The advantage of such a design is that the whole surface area of the ion-exchange membrane remains in contact with NaOH solution of the optimum concentration at all times, which guarantees the best overall membrane performance. However, the design also has serious drawbacks. The layer of NaOH between the electrode and the membrane contributes unfavorably to the overall cell resistance. In addition, as the hydrostatic pressure of the NaOH solution changes

8.1.3 Industrial Electrocatalysis

along the electrode height, maintaining the uniform distribution and identical properties of the three-phase boundaries and preventing the electrode flooding may pose problems. A pressure compensation method was developed by Bayer [142, 143]. In these cells, the ODC is divided into several separate ‘‘gas pockets’’ located upon each other. During operation these pockets are filled with oxygen compressed by the hydrostatic pressure of the NaOH solution. As a result, the gas pressure is maintained within the optimal range over the entire height of the gas diffusion electrode. 8.1.3.3.6 Industrial Water Electrolysis Renewable energies such as solar, wind, geothermal and hydropower are currently being evaluated as electricity sources that could allow the large-volume production of hydrogen for use in transportation and distributed power applications. Many projects on hydrogen generation, storage and energy conversion in fuel cells have been initiated during recent years with billions of dollars of funding by governments and industry. Today hydrogen is not an energy carrier, however. In fact, most of the hydrogen is used for ammonia or syngas production and other chemical processes. Moreover, most of the 500 billion mN 3 (standard cubic meter) produced worldwide [86] is generated by reforming of fossil feeds such as natural gas or naphtha or partial oxidation of heavy oil. Less than 1% is produced by water electrolysis since this process is only competitive when cheap electricity is available, e.g. near hydroelectric power plants. The scenarios of the visionaries predicting a ‘‘hydrogen economy’’ may be too optimistic, but without doubt the trend towards renewable energy sources and the necessity to reduce the losses in the conversion chain of hydrogen are apparent. Furthermore, on-site hydrogen production may become more important in the future. Hence it is essential to employ sophisticated electrolyzer technologies together with efficient, cost-effective and long-term stable electrocatalysts. The requirements on industrial water electrolysis facilities generally can be summarized as follows:

• low volumetric energy consumption (99.9%) Tab. 3

• high availability (continuous operation) • dynamic behavior under intermittent operation in the entire current density regime • long service intervals. A The Electrochemistry of Water Electrolysis The electrochemical reactions of water electrolysis depend on the nature of the electrolyte. Three different technologies have been developed (Table 3) and all of them have their fuel cell counterpart where the corresponding reaction is reversed. The cathodic hydrogen evolution reaction has already been investigated by Tafel [144]. Despite intensive research, the mechanism of the electrode reaction is still not totally understood since it depends on the operating conditions and the electrode materials. The cathode reaction can be separated into several steps [145]: Transport of water molecules to the phase boundary and adsorption:

1. Charge transfer at metallic catalysts (M): a. Reaction at sites not occupied by H atoms according to the Volmer reaction; formation of adsorbed H atoms: Acidic : −  H3 O + + M + e− − −− − − M− H + H 2 O

Acid (e.g. Nafion) Base (KOH) Ceramic (ZrO2 /Y2 O3 )

(31)

Alkaline: − −  H2 O + M + e− − −− − − M−H + OH

b. Reaction at sites occupied by H atoms according to the Heyrovsky reaction; formation of H2 molecules: Acidic: −−  H3 O+ + M−H + e− −− − − M + H2 + H2 O (32) Alkaline:

m−3 N )

− −  H2 O + M−H + e− − −− − − M + H2 + OH

References see page 1954

Electrochemical reactions of water electrolysis

Electrolyte

1949

Anode reaction

Cathode reaction

3H2 O → 1/2O2 + 2H3 O+ + 2e− 2OH− → 1/2O2 + H2 O + 2e− O2− → 1/2O2 + 2e−

2H3 O+ + 2e− → H2 + 2H2 O 2H2 O + 2e− → H2 + 2OH− H2 O + 2e− → H2 + O2−

1950

8.1 Electrocatalysis

2. Recombination of adsorbed H atoms formed in Eq. (31) (Tafel reaction): −−  M−H + M−H −− − − 2M + H2

(33)

3. Desorption of H2 from the surface. 4. Migration of H2 by a. diffusion and convection b. gas-bubble formation. The mechanism on metal oxides has not been elucidated so far. A chemisorption (Volmer–Tafel or Volmer–Heyrovsky) or a redox mechanism can be assumed: Chemisorption mechanism (–MO is a trivalent metal oxide): Volmer reaction:

−MO−H + −MO−H −−  −− − − 2−MO + H2

(40)

It seems possible that such a mechanism is involved in processes with mixed catalysts such as Raney nickel/molybdenum oxide or platinum oxides. For the oxygen evolution reaction on the anode, a redox mechanism can be assumed as metallic catalysts are typically covered by an oxide layer at the high electrochemical potential [146]. A simplified mechanism assuming a bivalent oxide can be presumed [145]: Electron transfer:

+ − −M(OH) −−  −− − − −M (OH) + e

(41) Chemical reaction:

Acidic: −  H3 O+ + −MO + e− − −− − − −MO−H + H2 O

(34)

Acidic: + −  −M+ (OH) + 2H2 O − −− − − −M(OH)2 + H3 O (42)

Alkaline: − −−  H2 O + −MO + e− −− − − −MO−H + OH

−−  Alkaline: −M+ (OH) + OH− −− − − −M(OH)2

Heyrovsky reaction:

Electron transfer and chemical reaction are subsequently repeated until four adjacent hydroxyl groups can react:

Acidic: −−  H3 O+ + −MO−H+e− −− − − −MO + H2 + H2 O (35) Alkaline: − −−  H2 O + −MO−H + e− −− − − −MO + H2 + OH

Tafel reaction: −MO−H + −MO−H − −  −− − − 2−MO + H2

(36)

Redox mechanism: Electron transfer:

− −MO + e− − −  −− − − −MO

−−  H3 O+ + −MO− −− − − −MO−H + H2 O (38) −

−−  Alkaline: H2 O + −MO −− − − −MO−H + OH

−−  4−M(OH)2 −− − − 4−M(OH) + O2 + 2H2 O

(43)

Different slopes of the Tafel plots indicate other direct electron transfer reactions, e.g. a two-electron transfer: Acidic: −M(OH)2 + −M(OH) + H2 O + − − −  −− − − [−M(OH)]2 O + H3 O + e

Alkaline: −M(OH)2 + −M(OH) + OH

(44) −

− − −  −− − − [−M(OH)]2 O + H2 O + e

From adjacent sites oxygen evolution takes place: (37)

Volmer-analogous reaction: Acidic:

Tafel reaction:

−

Heyrovsky reaction: Acidic: −  H3 O+ + −MO−H + e− − −− − − −MO + H2 + H2 O (39) Alkaline: − −−  H2 O + −MO−H + e− −− − − −MO + H2 + OH

−−  2[−M(OH)]2 O −− − − 4−M(OH) + O2

(45)

B Alkaline Water Electrolysis In order to achieve high specific conductivity, KOH solution of 25–39 wt.% is used in alkaline electrolysis, which is the dominant technology today. The alkaline cell environment permits the use of non-noble metal catalysts which are relatively inexpensive, abundant and resistant to poisoning. Oxygen evolution – the necessary other electrochemical reaction in the cell – is inherently faster in alkaline solutions. Two areas that significantly impact cell efficiency are cell membranes/diaphragms and electrocatalysts. Traditionally, the diaphragm material was asbestos based; however, nowadays either woven polymers [147], inert ceramics such as titanates [148] or a composite consisting of polymer and ceramic powder [149]

8.1.3 Industrial Electrocatalysis

are used. Single modules with production rates up to 500 mN 3 h−1 H2 are state of the art. The volumetric energy consumption of advanced plants is in the range 4.0–4.2 kWh m−3 N . Reduction of the power demand can be accomplished particularly by advances in electrocatalysis, but also new diaphragm materials have contributed to improvements, namely by reducing the cell resistance. Electrocatalyst materials used in industrial water electrolysis must meet stringent requirements for long-term stability. Vandenborre et al. [150] have summarized the technical requirements for a good electrocatalyst system. These include: • good electrocatalytic properties • resistance to mechanical and electrical wear • good electrical conductivity (in the case of oxide surfaces on anodes) • good stability under open-circuit conditions/power interruptions. Successful electrode activation is an essential requirement for all advanced alkaline electrolyzer technologies. Early efforts were made by some groups to achieve high energy conversion efficiency through operation at high temperature alone, without electrode treatment [151, 152]. These attempts were abandoned for reasons related to economics, materials stability and thermodynamics [153–156]. The central importance of electrocatalysis stimulated intensive R&D efforts in the mid-1970s; several reviews are available [150, 157–160]. Anode activation systems that were given particular attention include porous nickel sinters [161], nickel cobalt spinels [162–167], cobalt oxides and lithiated cobalt oxides [166, 168, 169], nickel hydroxides [170–172], various perovskite oxides [164, 173–175] and plasma-sprayed alloys of nickel and stainless steel [176, 177]. Among the more successful cathode treatments studied are iron and nickel molybdates [158, 178], nickel borides [179–181], the traditional nickel sulfides [182], various Raney iron or nickel materials [175, 183–188] and nickel–cobalt thio-spinels or mixed sulfides [150, 189]. There has been a remarkable convergence in understanding which electrode treatment systems are effective and economic for use in alkaline water electrolyzers [190]. Later, the emphasis in most programs shifted from the identification of new materials to the demonstration of long-term stability under industrial operating conditions. Increasing the surface area of the electrode (thus decreasing the real current density and overpotential) can be achieved by selective dissolution of an element from a nickel alloy [191]. The electrode surface can also be activated by the cathodic deposition of an active alloy [192]. Recent investigations [193–195] have focused

1951

on using Ni−Mo alloys and it has been reported that the activity for hydrogen evolution increases with increasing Mo content [196]. Nickel–sulfur alloys have also been investigated [197]. Cathodes based on hydrogen-storing alloys (i.e. metal hydrides) have also been examined and show enhanced corrosion resistance during power interruptions [194, 196]. The oxygen evolution electrode overpotential is generally higher than the cathode overpotential and there are only a few inexpensive and effective electrocatalysts [198]. DSA-type (dimensionally stable anode, see Section 8.1.3.3.1) electrodes [108] have been widely used in the chloralkali industry; however, the highly active RuO2 -based coating is not stable in alkaline solutions during electrolysis [199]. Oxides of nickel [193], mixed Ni–Fe oxides [200] and mixed Co−Ni oxides [201] have been widely studied. Perovskite-type materials such as La1−x Srx CO3 [202] and perovskites based on LaMnO3 [203] have shown reduced anode overpotential. The performance of alkaline water electrolyzers is normally around 1.7–1.8 V at 300 mA cm−2 , but higher performance (1.5–1.7 V at 300 mA cm−2 ) can be achieved in advanced systems at high pressures and temperatures. Typical operation temperatures are in the range 70–90 ◦ C for ambient pressure and 90–100 ◦ C for pressurized operating conditions [198, 204]. As stated above, alkaline water electrolysis is a wellestablished technology and is the most feasible hydrogen production technology from renewable energy at the present time. C PEM Water Electrolysis The first commercial proton exchange membrane (PEM) water electrolyzer was developed by General Electric [205]. Proton-conducting membranes are also employed in fuel cells (see Chapters 13.20.1 and 13.20.2) as they combine high ionic conductivity, low electronic conductivity, gas tightness and good mechanical properties. Advanced water electrolyzers using PEMs are less common than conventional alkaline electrolyzers and generally utilize expensive materials such as noble metal electrocatalysts and sulfonated polymer membranes with perfluorocarbon chains of the Nafion type. The benefits of PEM-based electrolyzers over alkaline systems are the higher current density (1–3 compared with 0.2 A cm−2 [206, 207]), no circulating liquid electrolyte, wide range of power loadings and very fast response time. These systems also have the advantage of potentially being utilized as regenerative fuel cells (i.e. the cell works in electrolysis mode when renewable power is available and the fuel cell mode when renewable power is not available). The References see page 1954

1952

8.1 Electrocatalysis

PEM water electrolyzer uses a similar technology to the PEM fuel cell with the electrocatalyst layers, consisting of electrocatalysts and ionomer, coated directly on to the membrane. In this acidic system, the anode overpotential and the cell resistance contribute to the majority of the performance loss [208]. A high current density PEM water electrolyzer was investigated as a means of overcoming the disadvantage of the initial cost of these systems [207]. The cell consisted of a Nafion 117 electrolyte with a Pt electrode on the cathode side and a Pt–Ir electrode on the anode side. The electrocatalysts were deposited on the membrane by a chemical plating method. The performance was examined up to current densities of 13 A cm−2 , with free water circulation with gas-lift. The cell voltage at 13 A cm−2 was 5 V, which corresponds to an energy efficiency (εG) of around 25%, whereas the cell voltage at 1 A cm−2 was approximately 1.9 V. It was shown that the IR-drop contributed to about 90% of the cell voltage increase over the current density range 1–5 A cm−2 . Bifunctional electrocatalysts were examined for a regenerative PEM fuel cell [209]. The oxygen electrocatalyst was a 50 wt.% Pt black + 50 wt.% IrO2 and the hydrogen electrocatalyst a Pt black layer. A transfer printing technique was used to apply the catalyst layer at a loading of 0.4 mg cm−2 . A cell voltage of 1.71 V at 400 mA cm−2 at a temperature of 80 ◦ C was accomplished. Another regenerative PEM fuel cell was examined using a mixed Pt black + IrO2 anode electrocatalyst and a Pt black cathode electrocatalyst at a loading of 8–10 mg cm−2 and a Nafion 115 membrane [210]. The MEA (membrane electrode assembly) was constructed by hot-pressing the gas diffusion electrodes (porous catalyst and PTFE composite) to the Nafion membrane. The cell performance was 1.74 V at 1 A cm−2 . Ruthenium dioxide was examined as the anode electrocatalyst in a solid polymer electrolyte water electrolyzer [211]. The RuO2 used in this work was degraded fairly quickly, resulting in an increase in the cell voltage from 1.73 to 1.92 V (1 A cm−2 , 82 ◦ C) after 48 h. The deterioration of the RuO2 was thought to involve both electrical resistance and electrokinetic mechanism changes. A new Ru-based oxide anode was tested and showed that even after 3000 h of operation at 1 A cm−2 the cell voltage was comparable to that of a fresh RuO2 anode (1.76 V). MEAs for water electrolysis were developed using a plating method in which the noble metal salt precursors were reduced using sodium borohydride [212]. This method was found to give a good level of control in terms of the metal loading and plating adhesion. When the cell performance was tested using different anode catalysts and a Pt cathode catalyst, the anodic overvoltage

was seen to increase in the order Ir < Rh < Rh−Pt < Pt−Ru < Pt < Pd. Ruthenium showed superior initial activity, although it was found to corrode significantly during oxygen evolution. The best cell performance was obtained using an Ir anode and Pt cathode and resulted in a cell voltage of approximately 1.7 V at 1 A cm−2 and 90 ◦ C. Coatings of Pt and Ir were used in SPE (solid polymer electrolyte) water electrolyzers using a chemical reduction process [206]. The cell voltage using Pt as the anode and cathode (1.13 mg cm−2 ) was measured to be around 2.15 V at 1 A cm−2 and 80 ◦ C. On adding Ir to the anode at a loading of 0.2 mg cm−2 , the cell voltage was decreased to around 1.75 V under the same operating conditions. This preparation method was tested for electrode areas up to 150 cm2 , and over 5000 h no irreversible degradation could be detected. Further work showed that an Ru−Pt anode initially gave even better cell performance than the Pt–Ir anode, although the cell voltage quickly increased with time due to the dissolution of the Ru [213]. Millet et al. [214] discussed the design and performance of a PEM water electrolyzer. The electrode materials were plated on to and slightly within the membrane by reducing a Pt precursor salt with sodium borohydride. This method [206, 213] has been used to prepare MEAs with surface areas up to 150 cm2 [206]. Using a Pt/Nafion 117/Pt cell the voltage was 2.2 V at 1 A cm−2 and 80 ◦ C (cell area 10 cm2 ). For larger cell areas of 50 and 100 cm2 the voltage was approximately 2.2 and 5.5 V (two cell stack), respectively, at 0.5 A cm−2 . Rasten [208] reported a study of various electrocatalysts for water electrolysis using proton exchange membranes. The work focused on the oxygen evolution electrode where IrO2 , Ir–Ta mixed oxides and Ir–Ru mixed oxides were used. RuO2 and Pt were also investigated as catalysts for the hydrogen evolution reaction. The oxide catalysts were prepared using the Adams fusion technique [215]. The best performance for the anodic catalysts at current densities of 1 A cm−2 or below was observed using an Ir–Ta (85 mol% Ir) mixed oxide. At higher current densities the mixed Ir−Ru (60–80 mol% Ir) oxide electrocatalysts performed better. In terms of the hydrogen evolution reaction, Pt black was found to be most promising, with RuO2 showing low electrochemical activity and electrical conductivity. A cell voltage of 1.59 V was obtained at 1 A cm−2 and 90 ◦ C using a noble metal loading of less than 2.4 mg cm−2 and a Nafion 115 membrane. A solid polymer electrolyzer was developed using IrO2 as the anode catalyst and Pt black as the cathode catalyst with loadings of 3 mg cm−2 [216]. The IrO2 was prepared by pyrolysis of Ir in oxygen at various temperatures, and the specific surface area of the IrO2 was found to decrease with temperature. The cell voltage ranged from

8.1.3 Industrial Electrocatalysis

approximately 1.61 to 1.9 V at 1 A cm−2 and 80 ◦ C for the variously prepared IrO2 catalysts, with the IrO2 prepared at 200 ◦ C showing the best performance. The durability was demonstrated by operating the cells for 5000 h. Since 1987, Mitsubishi Heavy Industries has been developing solid polymer water electrolyzer technology [217]. A chemical plating technique was used to plate iridium metal on to each side of a Nafion membrane. The catalyst loading was varied between 0.5 and 2 mg cm−2 for cell areas of 50 cm2 and showed that the cells could be operated at 3 A cm−2 . Cell voltages were measured to be 1.7 and 2.01 V for current densities of 1 and 3 A cm−2 , respectively, at 80 ◦ C and an Ir loading of 1 mg cm−2 . Similar results were found for cell areas of 200 cm2 . A 50 m2 solid polymer electrolyte electrolyzer was developed using a hot-press method to adhere the catalyst film (catalyst + PTFE) to the membrane [218]. The effect of the hot-pressing temperature showed that the thickness of the MEA decreased with temperature, resulting in a decrease in the cell voltage when using an anode catalyst of IrO2 and a cathode catalyst of Pt black at loadings of 4 and 3 mg cm−2 , respectively. The effect of the anode catalyst type was also investigated over the current density range 0.5–4 A cm−2 . The results showed that the performance increased for the anode catalyst series RuO2 (pyrolyzed at 400 ◦ C) > RuO2 (pyrolyzed at 600 ◦ C) > Ir−Ru mixture (90%Ir) > IrO2 (pyrolyzed at 200 ◦ C) > Ir−Pt mixture (96.7%Ir) > Ir black > Ir2 O3 > Rh2 O3 > Pt black. All these catalysts were loaded at 4 mg cm−2 and Pt black was used as the cathode catalyst in all cases. The durability of the anode catalysts was investigated by monitoring the cell voltage over time. This showed the durability series as IrO2 (pyrolyzed at 200 ◦ C) > Ir−Ru mixture (90% Ir) = Ir−Pt mixture (96.7%Ir) > RuO2 (pyrolyzed at 600 ◦ C) > RuO2 (pyrolyzed at 400 ◦ C). Studies were also performed on the IrO2 loading for the anode catalyst layer. The results showed that increasing the loading beyond around 2.5 mg cm−2 gave little performance improvement with loadings of 3 mg cm−2 being suggested as adequate. Similar results were found for the Pt black cathode catalyst with little changes observed above 0.5 mg cm−2 . Overall the best cell performance was accomplished with the cell consisting of IrO2 (pyrolyzed at 200 ◦ C, 3 mg cm−2 ) and Pt black (0.5 mg cm−2 ) as anode and cathode catalysts, respectively. The membrane was 51 µm thick with an equivalent weight of 1000 and a hot-pressing temperature of 140 ◦ C was used. Cell voltages were 1.533 and 1.665 V for current densities of 1 and 3 A cm−2 , respectively, at 80 ◦ C. The current connectors were Pt-coated Ti sinter for the anode and gold-plated stainless-steel sinter for the cathode. Based on the results mentioned above, a 2500 cm2 solid polymer electrolyte electrolyzer was developed [219]. The

1953

IrO2 anode catalyst was prepared by reacting Na2 IrCl6 with an aqueous NaOH solution. The resulting Ir(OH)4 solution was purified, dried and pyrolyzed at 200 ◦ C. The catalyst layer was prepared by hot-pressing the catalyst film (catalyst + PTFE) to the solid polymer electrolyte membrane with the catalyst loading being 3.6 and 3 mg cm−2 for the anode (IrO2 ) and cathode (Pt black), respectively. At 80 ◦ C the cell voltage was measured to be 1.54 and 1.74 V at 1 and 3 A cm−2 , respectively. A high-pressure PEM water electrolyzer was developed with Pt and Ir black serving as the anode catalyst and Pt black as the cathode catalyst [61]. The electrolyzer showed a cell voltage of approximately 1.7 V at 1 A cm−2 and 90 ◦ C at atmospheric pressure. At an elevated pressure of 2.5 MPa a slightly lower voltage was observed. The improvement in performance was attributed to a decrease in the anodic overvoltage. MEAs for a reversible solid polymer fuel cell were examined using Pt, Rh, Ir and Ir–Ru mixed oxides, as the oxygen evolution electrocatalyst [220]. When operated as an electrolysis cell, the anodic overvoltage increased in the order Ir–Ru mixed oxide < Ir < Rh < Pt. Due to the poor hydrogen oxidation activity of Ir−Ru, Ir metal was used as a compromise between the oxygen evolution and hydrogen oxidation reactions. A cell voltage of approximately 1.52 V was achieved at a low current density of 0.1 A cm−2 when using Ir at the anode and Pt at the cathode for water electrolysis. During fuel cell operation a cell voltage of 0.8 V was measured at this same current density. From the above, it is obvious that good cell performance data (i.e. 1.6 V at 1 A cm−2 ) are achievable. This indicates that low-cost operation will be possible and this implies that PEM-based electrolyzers will probably replace the traditional alkaline systems. PEM electrolyzers are particularly suited to small-scale and on-site hydrogen production units such as at future hydrogen refueling stations. The main disadvantage with these units is the materials cost. However, it has been shown that the noble metal loading can be reduced significantly, leaving only the expensive membrane as the economically limiting factor. D High-Temperature Water Electrolysis High-temperature electrolysis makes use of the oxygen-ion conductivity of a ceramic material (ZrO2 /Y2 O3 ) at elevated temperatures (ca. 900 ◦ C). The electrodes are made from Ni/ZrO2 (cathode) and LaMnO3 (anode), basically the same materials as used in solid oxide fuel cells (SOFCs). High-temperature electrolyzers can potentially utilize excess heat from industrial plants or power stations References see page 1954

1954

8.1 Electrocatalysis

for hydrogen production. However, to date, prototype solid-oxide electrolyzer units have not achieved useful operational lives, and substantial engineering problems exist with respect to thermal cycling and gas sealing. Summary and Outlook Although electrochemical syntheses have been employed in the industry for more than 100 years, only a few large-scale processes have demonstrated their economic competitiveness. In this chapter, the most important electrochemical processes, such as chloralkali electrolysis and the synthesis of adiponitrile, were discussed in detail. As this chapter has demonstrated, many technological key factors, which decide whether an electrochemical process can be led to commercial success, are associated with the right choice of electrode materials. For this reason, electrode development is focused on issues such as longterm stability, high product selectivity, high turnover rate, energy efficiency and low material costs – topics not unfamiliar to scientists working in the field of heterogeneous catalysis. It goes without saying that the operating efficiency of electrochemical processes depends on the costs of electrical energy. However, this implies that one has to make the best use of the invested electrical energy by facilitating an effective and durable electron crossover from the electrode to the reactant. In this regard, the role of ongoing electrocatalyst development cannot be overemphasized. Some of the electrochemical processes that were discussed in this chapter are currently niche applications, but with a high potential to change profoundly the way we go about our daily life in the future. Water electrolysis is a good example. Today, our industrial society depends on the availability of fossil fuels. However, political instabilities, speculation about and fluctuations of the price of oil and the awareness that there has to be a solution for the time beyond oil are drivers for the growth of renewable energies. Photovoltaic systems and wind power plants are prospering, and in the near future they will generate large amounts of fluctuating electrical power, which has to be distributed and stored in an appropriate way. Some experts favor an energy infrastructure based on hydrogen as a secondary energy carrier. The electrolytic generation of hydrogen from water would occupy a key position in this coming hydrogen economy. Thus, water electrolysis could become one of the most important fields for electrocatalysis in the near future. 8.1.3.4

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G. Faita, G. Fiori, J. Appl. Electrochem. 1972, 2, 31. A. I. Krasil’shchikov, Zh. Fiz. Khim. 1963, 37, 531. L. I. Krishtalik, Electrochim. Acta 1981, 26, 329. L. J. J. Janssen, in Modern Chloralkali Technology, C. Jackson (Ed.), Vol. 2, Ellis Horwood, Chichester, 1983, p. 271. H. B. Beer, British Patent 1 147 442, 1966. D. B. Rogers, R. D. Shannon, A. W. Sleight, J. L. Gillson, Inorg. Chem. 1969, 8, 841. J. M. Honig, in Electrodes of Conductive Metallic Oxides, Part A, S. Trasatti, G. Lodi (Eds.), Elsevier, Amsterdam, 1980, p. 1. W. N. Brooks, D. A. Denton, N. M. Sammes, in Proceedings of the Symposium on Performance of Electrodes for Industrial Electrochemical Processes, Vol. 89-10, F. Hine, J. M. Fenton, B. U. Tilak, J. B. Lisius (Eds.), Electrochemical Society, Pennington, NJ, 1989, p. 39. H. J. Hansen, C. W. Krenk, T. P. M. Koster, A. Macker, in Electrochemische Stoffgewinnung–Grundlagen und Verfahrenstechnik, G. Kreysa (Ed.), DECHEMA-Monographien, Vol. 125, DECHEMA, Frankfurt am Main, 1992, p. 413. S. Trasatti, Electrochim. Acta 1984, 29, 1503. S. Kotowski, J. Parr, Titanium Anodes Used in Electrochemical Processes, Chemical Plants and Processes, Firmenschrift Heraeus Elektrochemie, Werk Freigericht, 1992. C. Comninellis, E. Plattner, in Proceedings of the Symposium on Performance of Electrodes for Industrial Electrochemical Processes, Vol. 89-10, F. Hine, J. M. Fenton, B. U. Tilak, J. B. Lisius (Eds.), Electrochemical Society, Pennington, NJ, 1989, p. 229. R. U. Bondar, A. E. Kalinovskii, Elektrokhimiya 1978, 14, 730. R. K¨otz, S. Stucki, D. Scherson, D. M. Kolb, J. Electroanal. Chem. 1984, 172, 211. A. Mills, H. Davies, J. Chem. Soc., Faraday Trans. 1990, 86, 955. R. K¨otz, S. Stucki, Electrochim. Acta 1986, 31, 1311. V. V. Gorodetskii, V. A. Neburchilov, M. M. Pecherskii, Russ. J. Electrochem. 1994, 30, 916. T. Shimanune, J. Kawanche, S. Nakamatsu, Y. Nishiki, R. Hayasi, in Proceedings of the Symposium on Performance of Electrodes for Industrial Electrochemical Processes, Vol. 89-10, F. Hine, J. M. Fenton, B. U. Tilak, J. B. Lisius (Eds.), Electrochemical Society, Pennington, NJ, 1989, p. 77. N. Yoshida, T. Morimoto, Electrochim. Acta 1994, 39, 1733. S. Trasatti, in Advances in Electrochemical Science and Engineering, H. Gerischer, C. W. Tobias (Eds.), VCH, Weinheim, 1992, p. 1. T. Morimoto, K. Suzuki, T. Matsubara, N. Yoshida, Electrochim. Acta 2000, 45, 4257. E. Yeager, P. Bindra, Chem. Ing. Tech. 1980, 52, 384. G. Faita, US Patent 5 770 035, assigned to De Nora, 1996. D. Oldani, P. Fabian, F. Fulvio, A. Fischer, L. Carrettin, WO Patent 2002, 068 718, assigned to Uhde-Nora Technologies, 2002. F. Gestermann, P. Fabian, US Patent 6 596 136, assigned to Uhde-Nora Technologies and Bayer, 2001. R. J. Allen, J. R. Giallombardo, D. Czerwiec, E. S. De Castro, K. Shaikh, F. Gestermann, H. D. Pinter, G. Speer, WO Patent 2002, 018 675, assigned to DeNora Elettrodi and Bayer, 2002. F. Gestermann, presented at the 12th International Forum on Electrolysis in the Chemical Industry, Clearwater, FL, 1998. T. G. Coker, R. M. Dempsey, A. B. LaConti, US Patent 4 191 618, assigned to General Electric, 1978.

132. R. J. Allen, J. R. Giallombardo, D. Czerwiec, E. S. De Castro, K. Shaikh, US Patent 6 149 782, assigned to De Nora, 1999. 133. R. J. Allen, A. F. Gulla, WO Patent 2005, 075 071, assigned to De Nora Elettrodi, 2005. 134. Anonymous, Breathing New Life into Electrolysis, Bayer News, Feb. 3, 2004. 135. Anonymous, Chlorine–Produced at the Oxygen Bar, Bayer Articles from Research, 16th Edn., 2004. 136. F. Federico, G. Faita, E. S. De Castro, F. Gestermann, H. Pinter, presented at the 207th Meeting of the Electrochemical Society, Quebec, 2005. 137. K. Hayashi, A. Sakata, N. Furuya, H. Aikawa, K. Aiki, in Proceedings of the R. B. MacMullin Memorial Symposium, Vol. 99-21, Electrochemical Society, Pennington, NJ, 1999, p. 209. 138. O. Ichinose, M. Kawaguchi, N. Furuya, J. Appl. Electrochem. 2004, 34, 55. 139. A. Sakata, N. Furuya, H. Aikawa, K. Saiki, in Proceedings of the R. B. MacMullin Memorial Symposium, Vol. 99-21, Electrochemical Society, Pennington, NJ, 1999, p. 223. 140. J. Chlistunoff, Final Technical Report: Advanced Chloralkali Technology, DOE Award 03EE-2F/ED 190403, Los Alamos National Laboratory, Los Alamos, NM, 2005. 141. M. Sugiyama, K. Saiki, A. Sakata, H. Aikawa, N. Furuya, J. Appl. Electrochem. 2003, 33, 929. 142. A. Bulan, F. Gestermann, T. Turek, R. Weber, P. Weuta, Chlorelektrolysen mit Gasdiffusionselektroden, GVC/DECHEMA-Jahrestagungen, Karlsruhe, 2004. 143. D. Hoormann, J. J¨orissen, H. P¨utter, Chem. Ing. Tech. 2005, 77, 1363. 144. J. Tafel, Z. Phys. Chem. 1905, 50, 641. 145. G. Sandstede, in Elektrochemische Stoffgewinnung–Grundlagen und Verfahrenstechnik, G. Kreysa (Ed.), DECHEMA-Monographien, Vol. 125, DECHEMA, Frankfurt am Main, 1992, p. 329. 146. D. Ohms, V. Plzak, S. Trasatti, K. Wiesener, H. Wendt, in Electrochemical Hydrogen Technologies, H. Wendt (Ed.), Elsevier Amsterdam, 1990, p. 1. 147. K. Andreassen, in Hydrogen Power: Theoretical and Engineering Solutions. Proceedings of the HYPOTHESIS II Symposium, Norway, 1997, T. O. Saetre (Ed.), Kluwer, Dordrecht, 1998, p. 91. 148. H. Wendt, H. Hofmann, J. Appl. Electrochem. 1989, 19, 605. 149. P. Vermeiren, W. Adriansens, J. Moreels, R. Leysen, Int. J. Hydrogen Energy 1998, 23, 321. 150. H. Vandenborre, R. Leysen, P. Vermeiren, Active Electrodes to Be Used in Advanced Alkaline–Water Electrolysis, Report DE ´ ´ 83900365, Centre dEtude dEnergie Nucl´eaire, Mol, Belgium, 1982. 151. G. Imarisio, Int. J. Hydrogen Energy 1981, 6, 153. 152. M. H. Miles, G. Kissel, P. W. T. Lu, S. Srinivasan, J. Electrochem. Soc. 1976, 123, 332. 153. R. L. LeRoy, Int. J. Hydrogen Energy 1983, 8, 401. 154. R. L. LeRoy, J. Electrochem. Soc. 1983, 130, 2158. 155. P. Combrade, in Hydrogen as an Energy Carrier, Proceedings of the 3rd International Seminar, Lyon, France, G. Imarisio, A. S. Strub (Eds.), Reidel, Dordrecht, 1983, p. 183. 156. P. Combrade, in Hydrogen Energy Progress IV, Vol. 1, T. N. Veziroglu, W. D. Van Vorst, J. H. Kelley (Eds.), Pergamon Press, Oxford, 1982, p. 355. 157. L. Martin, J. Diette, M. Prigent, J. Demarsy, C. Sellier, Mise ´ ´ au Point de Nouvaux Electrocatalyseurs pour l’Electrolyse Avanc´ee, Report EUR 7068 FR, Commission of the European Communities, Brussels, 1981.

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209. S. Zhigang, Y. Baolian, H. Ming, J. Power Sources 1999, 84(2), 82. 210. T. Ioroi, N. Kitazawa, K. Yasuda, Y. Yamamoto, H. Takenaka, J. Electrochem. Soc. 2000, 147, 2018. 211. J. Sedlak, R. Lawrence, J. Enos, Int. J. Hydrogen Energy 1981, 6, 159. 212. H. Takenaka, E. Torikai, Y. Kawami, N. Wakabayashi, Int. J. Hydrogen Energy 1982, 7, 397. 213. P. Millet, R. Durand, M. Pineri, Int. J. Hydrogen Energy 1990, 15, 245. 214. P. Millet, F. Andolfatto, R. Durand, Int. J. Hydrogen Energy 1996, 21, 87. 215. R. Adams, R. L. Shriner, J. Am. Chem. Soc. 1923, 45, 2171. 216. M. Yamaguchi, K. Yagiuchi, K. Okisawa, in Hydrogen Energy Progress XI, Proceedings of the 11th World Hydrogen Energy Conference, Vol. 1, T. N. Veziroglu, C.-J. Winter, J. P. Baselt, G. Kreysa (Eds.), DECHEMA, Frankfurt am Main, 1996, p. 781. 217. M. Nagai, H. Tazima, A. Sakanishi, N. Hisatome, S. Ohkura, in Hydrogen Energy Progress XI, Proceedings of the 11th World Hydrogen Energy Conference, Vol. 1, T. N. Veziroglu, C.-J. Winter, J. P. Baselt, G. Kreysa (Eds.), DECHEMA, Frankfurt am Main, 1996, p. 825. 218. M. Yamaguchi Shinohara, K. Okisawa, T. Nakanori, in Proceedings of the 32nd Intersociety Energy Conversion Engineering Conference, New York, NY, USA, 1997, p. 1958. 219. M. Yamaguchi, T. Shinohara, H. Taniguchi, T. Nakanori, K. Okisawa, in Hydrogen Energy Progress XII, Proceedings of the World Hydrogen Energy Conference, 12th, Buenos Aires, 1998, p. 747. 220. K. Ledjeff, F. Mahlendorf, V. Peinecke, A. Heinzel, Electrochim. Acta 1995, 40, 315.

8.2

Photocatalysis: Development of Highly Functional Titanium Oxide Photocatalysts Masato Takeuchi, Masaaki Kitano, Masaya Matsuoka, and Masakazu Anpo∗

environmentally friendly chemical processes and systems is an urgent concern not only for scientists but for all mankind. With these issues in mind, scientists have increasingly turned to photochemistry and photocatalysis as promising ‘‘environmentally harmonious systems’’. Especially titanium dioxide photocatalysts possess the potential to reduce CO2 into useful CH4 and CH3 OH, decompose NOx into harmless compounds and to oxidize various kinds of organic compounds into non-toxic CO2 and H2 O. In other words, photocatalytic systems can even be thought of as ‘‘artificial photosynthesis’’ [1–8]. However, since TiO2 is a semiconductor with a wide bandgap larger than 3.2 eV, corresponding to 388 nm in wavelength, irradiation with a UV light source is necessary for it to act as a workable photocatalyst. Many different approaches have been intensively investigated to develop TiO2 semiconducting photocatalysts sensitive to visible light. In 1991, O’Regan and Gr¨atzel reported on a dyesensitized solar cell using a TiO2 thin-film electrode and photofunctional dye to absorb visible light [9]. This finding opened the way for the conversion of solar energy into usable electric energy in the field of solar cells. However, there are still several issues that need to be resolved, such as improvement of the stability of the organic dyes and a more efficient sealing technique to prevent leakage of the liquid electrolyte. The development of visible lightresponsive photocatalysts will, thus, play an important role in future efforts to purify air, water and soil using methods that convert abundant and safe solar energy into useful and efficient chemical energy, a great challenge for chemical research scientists. In this chapter, we deal with an innovative application of ion engineering techniques for the preparation of well-defined, visible light-responsive TiO2 photocatalysts. 8.2.2

8.2.1

Ion Engineering Techniques for the Preparation of Well-Defined TiO2 Photocatalysts

Introduction

Mankind’s rapid economic growth in the past centuries has been achieved by consuming the Earth’s limited energy resources, such as petroleum. However, we are now paying for the depletion of these resources with serious environmental problems such as the greenhouse effect caused by uncontrolled CO2 emissions, acid rain caused by air pollutants from automobiles, power plants and industrial factories as well as polluted waterways on a global scale. We will also be confronted with serious energy issues as fossil fuel reserves become exhausted and prices rise critically. Taking these multiple and intertwined issues into consideration, the need to develop ∗

Corresponding author.

In recent years, advanced ion engineering techniques have been applied to modify the crystalline structure and electronic properties of semiconducting materials such as silicon. The schematic interactions between the accelerated ions with different energies and solid surfaces are shown in Fig. 1: (i) when ion beams with low energies of the order of magnitude of a few hundred eV impact on solid surfaces, these ions are accumulated on the solid surfaces to form thin, snow-like films; (ii) when solid surfaces are bombarded with ion beams of intermediate energies from a few hundred eV to several tens of keV, these ions sputter the atoms of the solid surfaces as secondary ions; (iii) when ion beams with energies higher than several tens of keV hit exposed solid surfaces, these ions are implanted within subsurface layers without any

8.2.2 Ion Engineering Techniques for the Preparation of Well-Defined TiO2 Photocatalysts Tab. 1

Ions with low energy Formation of thin film + + + (a) Ions with middle energy Sputtering +

CVD (chemical vapor deposition) Heat-assisted CVD CVD at ambient pressure CVD at low pressure Photo-assisted CVD Plasma-assisted CVD RF plasma-assisted CVD Microwave plasma-assisted CVD ECR plasma-assisted CVD DC plasma-assisted CVD

+

+ + + (b) Ions with high energy Ion implantation +

Various physical and chemical vapor deposition methods

PVD (physical vapor deposition) Vacuum deposition Heat-resistance Electron beam (EB) Molecular beam epitaxy (MBE) deposition Laser ablation deposition Sputtering deposition RF sputtering RF magnetron sputtering (RF-MS) Ion beam sputtering ECR sputtering Ion plating Ionized cluster beam (ICB) deposition

+ +

1959

+ +

+

+ +

+ + (c)

Schematic diagrams of the interaction between accelerated ions with different energies and solid surfaces.

Fig. 1

significant damage to the material, i.e. ion implantation is successfully carried out. The implantation of metal ions with high energy may sometimes induce the formation of an amorphous phase; however, it has been observed that the crystalline structure of the solid can be properly recovered by post-calcination processes in air. TiO2 thin films are generally prepared by wet processes such as sol–gel, dip-coating and spray-coating with titanium alkoxide in organic solvents as a precursor. Transparent TiO2 thin-film photocatalysts can also be prepared advantageously by dry processes because of the ease in controlling the various preparation conditions [42–46, 55]. The preparation of TiO2 thin films in a high-vacuum chamber has the following advantages: (i) contamination of the films with impurities can be prevented; (ii) since no organic solvents are used in the preparation, dry methods can be considered ‘‘environmentally friendly’’; (iii) thin films with high crystallinity and strong adhesion to the substrates can be easily prepared without calcination at high temperatures; and (iv) the various physical and chemical properties of the catalyst can be easily controlled.

Some physical vapor deposition (PVD) methods as dry processes are summarized in Table 1. Among these, we have successfully applied the ionized cluster beam (ICB) and RF-magnetron sputtering (RF-MS) deposition methods to prepare transparent and visible lightresponsive TiO2 thin films [42–46, 55] and schematic diagrams of the ICB and RF-MS deposition methods are shown in Fig. 2a and b, respectively. In the ICB method, titanium vapor obtained by heating titanium metal as the source material at 2200 K was introduced into a high-vacuum chamber to produce titanium clusters. The titanium clusters reacted with sufficient amounts of O2 molecules (O2 pressure: 2.7 × 10−2 Pa) in the high-vacuum chamber to form stoichiometric TiO2 clusters. These TiO2 clusters, ionized by electron beam irradiation, were accelerated by an electric field with an acceleration voltage of 500 V and bombarded on to the substrates to form transparent TiO2 thin films [42–46, 52–55]. In the RF-MS deposition method, a TiO2 plate with a rutile structure was used as the ion material. With a reactive sputtering method, the TiO2 thin films were prepared using a metallic Ti target in the presence of O2 as the reactive gas. However, since a stoichiometric TiO2 plate was used as the sputtering target, only Ar gas was used for the sputtering gas without the coexisting O2 as the reactive gas. When the magnetic field, which is orthogonal to the electric field, was applied in the presence of the sputtering gas, a References see page 1968

1960

8.2 Photocatalysis

High vacuum chamber (ca. 10−5 Pa) Substrates Heater

TiO2 thin film

Substrates

+ O2 atmosphere (2.7 × 10−2 Pa) e− e−

(a)

Ti metal

+

+

Electric field (500 V) e− Electron beam e−

crucible

TiO2 thin film

Gas plasma (Ar+)

(b)

N

S

N

S

N

S

Source material (TiO2 plate)

Magnet

Schematic diagrams of (a) ionized cluster beam (ICB) and (b) RF-magnetron sputtering (RF-MS) deposition methods to prepare TiO2 thin films.

Fig. 2

ring-state gas plasma was induced on the target material. This gas plasma sputtered the target surface significantly to produce sputtered particles such as Ti4+ and O2− and these particles produced by the gas plasma were accumulated on the substrate surface to form highly transparent TiO2 thin films [45, 46]. 8.2.3

Visible Light-Responsive TiO2 Photocatalysts Modification of the Electronic Properties of TiO2 Semiconducting Fine Powder Photocatalysts by an Advanced Metal Ion Implantation Method TiO2 semiconductors have a relatively large bandgap of 3.2 eV, corresponding to wavelengths shorter than 388 nm. In other words, TiO2 itself can make use of only 3–4% of the solar radiation that reaches the Earth. From this viewpoint, TiO2 photocatalysts which can operate efficiently under both UV and visible light would be ideal for practical and widespread applications. Since the 1970s, various approaches to developing visible light-responsive photocatalysts by adding additional components such as metal oxides or metal ions to TiO2 semiconductors have been evaluated [10–14]. The aggregation of the added metal oxides or ions on the catalyst surfaces acted as recombination centers for the photo-formed electron–hole pairs on these TiO2 semiconductors, enabling them to absorb visible light. It could be shown that metal ion implantation methods are the most effective engineering techniques to modify the electronic properties of various semiconductors and make them responsive to visible light energy [15–26]. TiO2 semiconductor photocatalysts were developed by 8.2.3.1

implantation of various transition metal ions such as V, Cr and Fe into their deep bulk as an atomic-level process [15–26]. A smooth red shift of the absorption edge of the TiO2 was observed, its intensity depending on the amount and kinds of metal ions implanted. These metal ion-implanted TiO2 powdered photocatalysts were found to show efficient absorption in visible light regions, as shown in Fig. 3a. Such a characteristic red shift in the absorption spectra could also be observed with other transition metals such as Mn, Ni, Co and Cu; however, the implantation of Mg, Ti and Ar ions failed to initiate any shifts. These results clearly indicate that the red shifts are attributable to the chemical interactions between the TiO2 semiconductor and implanted metal ions and are not the result of changes in the physical properties arising from the presence of lattice defects. It was also found that such shifts in the absorption edge were observed only after calcination of the metal ion-implanted TiO2 samples in O2 at around 723–823 K. Calcination under an O2 atmosphere in combination with metal-ion implantation was therefore concluded to be instrumental in the red shift of the absorption spectrum toward the visible light region. Figure 3b shows the UV–visible absorption spectra of TiO2 catalysts chemically doped with small amounts of Cr ions by a conventional impregnation method in a comparative study with the Cr ion-implanted samples. These catalysts showed a new absorption band at ca. 420 nm as a shoulder although the position of the absorption edge of the TiO2 itself at 388 nm did not change. The intensity of these shoulders was found to increase as the amount of Cr ions chemically doped was increased. The results showed the shoulders to be attributable to the formation of an aggregated Cr oxide

Solar spectrum

(B) (C) (D)

(A) 250 (a)

K. M. Absorbance/a. u.

K. M. Absorbance/a. u.

8.2.3 Visible Light-Responsive TiO2 Photocatalysts

350

450

1961

(E′) (D′)

(C′) (B′) (A′)

550

650

Wavelength/nm

250

350

450

550

650

Wavelength/nm

(b)

Diffuse reflectance UV–visible absorption spectra of (A, A ) TiO2 ; (a) (B–D) Cr ion-implanted TiO2 ; (b) (B –E ) Cr ion-doped TiO2 photocatalysts and the solar spectrum. Amount of Cr ions implanted (µmol g−1 ): (A) 0; (B) 0.22; (C) 0.66; (D) 1.3. Amount of Cr ions doped (wt.%): (A ) 0; (B ) 0.01; (C ) 0.1; (D ) 0.5; (E ) 1 (0.1 wt.% equals ca. 4.9 mmol g−1 TiO2 ). K.M. denotes Kubelka–Munk. Fig. 3

N2O 1.5

Amount of products/µmol g-TiO2−1

species, CrOx , on the TiO2 surface. These results also indicated that a chemical doping method causes the electronic properties of the TiO2 catalyst to be modified in completely different ways, thus confirming that only metal ion-implanted TiO2 catalysts were able to show unique shifts in the absorption band toward the visible region, even with a smaller amount of implanted ions compared with a chemical doping method. Visible light irradiation of these TiO2 catalysts prepared by metal ion implantation induced significant photocatalytic activity for reactions such as the decomposition of NO, isomerization of cis-2-butene, degradation of organic compounds in water and the hydrogenation of methylacetylene with water [20–26]. Figure 4 shows the reaction vs. time profiles for the photocatalytic decomposition of NO into N2 and N2 O under visible light irradiation (λ > 450 nm) over Cr ion-implanted TiO2 and the unimplanted original TiO2 photocatalysts. Under the same conditions of visible light irradiation, the chemically doped TiO2 (data not shown) and unimplanted original TiO2 catalysts did not exhibit any photocatalytic reactivity whereas the Cr ion-implanted catalyst showed efficient reactivity under visible light of up to 400–600 nm, hence, they are referred to as ‘‘second-generation TiO2 photocatalysts’’ [15–26]. It should also be emphasized that their photocatalytic performance under UV light could be retained. However, TiO2 incorporating metal ions doped by a chemical method showed drastically decreased performance under UV light irradiation owing to the rapid recombination of the photo-formed electron–hole pairs on the aggregated metal oxides doped. These results clearly show that when metal ions are physically implanted, they cannot work as electron–hole recombination centers but only in modifying the electronic properties of the TiO2 semiconductors themselves.

off

on

off

on

N2 1.0 Cr ion-implanted TiO2 0.5

Original TiO2 0 −2

0

2

4

6

8

10

Time/h Reaction vs. time profiles of the photocatalytic decomposition of NO over Cr ion-implanted TiO2 and the original untreated TiO2 photocatalysts under visible light (λ > 450 nm). Photocatalyst samples were placed in a flat bottom quartz cell (volume: ca. 33 cm3 ). UV light irradiation was carried out with a 100-W high-pressure Hg lamp (Toshiba, SHL-100UVQ-2) through a cutoff filter (Toshiba Glass, UV-27) at 275 K. In order to avoid the heating effect from the Hg lamp, the photocatalysts in the quartz cell were cooled in ice–water during the photoreaction. The reaction products were analyzed by gas chromatography (Shimadzu, GC-14A) using an apparatus equipped with TCD and FID detectors.

Fig. 4

Fieldwork experiments were subsequently carried out to test the photocatalytic reactivity of the newly developed metal ion-implanted TiO2 catalysts under actual solar beam irradiation [27]. Under outdoor light at ordinary temperatures, the Cr and V ion-implanted References see page 1968

1962

8.2 Photocatalysis

Rate of NO elimination/mol min−1

8

×10−9

6

4

2

0

TiO2

Cr/TiO2

V/TiO2

Catalysts Solar beam intensity : 38.5 mW cm−2 Amount of catalyst : 3.6 g Flow rate : 18 L min−1

Photocatalytic reactivity of the V ion- and Cr ion-implanted TiO2 and the original unmodified TiO2 photocatalysts for the decomposition of NO under solar light irradiation.

Fig. 5

TiO2 catalysts showed 2–3 times higher photocatalytic reactivity for the decomposition of NO, as shown in Fig. 5. It could therefore be seen that such advanced metal ion-implantation methods could be successfully applied to modify the electronic properties of the TiO2 semiconductor catalysts, enabling the absorption of visible light with wavelengths longer than 550 nm to initiate photocatalytic reactions several times more efficiently than with the unimplanted original TiO2 under both UV and visible light. Preparation of Highly Transparent TiO2 Thin-Film Photocatalysts Photocatalytic systems using fine powder catalysts are very useful for the purification of polluted water and air; however, it is difficult to filter and reuse the catalysts after each reaction. In order to overcome this problem, various methods to immobilize the TiO2 photocatalysts on substrates such as a sol–gel [28–32], metal organic chemical vapor deposition (MOCVD) [33, 34] and a direct deposition technique with an aqueous solution of (NH4 )2 TiF6 [35], TiF6 [36] or TiOSO4 [37] have been widely investigated. TiO2 powder is generally used as a white pigment or titanium pigment for its high refractive index [38]. However, TiO2 thin films show high transparency in visible light regions so that the TiO2 coating on the substrate does not obscure the design or texture of the substrate surface. Moreover, the film surface exhibits photoinduced hydrophilic conversion properties in which the contact angle of the water droplets on the TiO2 thin film reaches zero under UV light irradiation 8.2.3.2

[5, 39–41]. These thin-film photocatalysts are therefore widely used in various applications to eliminate offensive odors or to decompose volatile organic compounds such as formaldehyde and acetaldehyde into harmless CO2 and H2 O under UV light irradiation. At present, the coating of TiO2 thin films on various substrates is mainly carried out by sol–gel, dip-coating, spin-coating and spray-coating in wet media [28–37]. Since titanium alkoxides in organic solvents are used as precursor materials in such wet processes, postcalcination at high temperatures after coating the supports with the sol solution is necessary to obtain TiO2 oxide thin films with high crystallinity and strong adhesion. Hence substrates which are not heat resistant are not suitable for these wet-coating processes. In such cases, inorganic binders which can solidify at relatively low temperatures are used to immobilize the TiO2 fine-particles; however, these thin-films do not perform as well or show strong adhesion to the substrate. Ion engineering techniques can therefore be considered fairly useful for the preparation of well-defined TiO2 thinfilm photocatalysts and we shall present the synthesis of highly transparent TiO2 thin films on glass substrates by applying an ionized cluster beam (ICB) deposition method [42]. XRD measurements showed that films with a thickness of more than 300 nm had patterns attributable to the presence of anatase and rutile structures of the TiO2 . The ratio of anatase to rutile was around 70%. Thin films with a thickness of less than 100 nm did not show any XRD peaks of a crystalline structure due to their smaller film thickness; however, XAFS showed these thin films to have an anatase structure. Since the standard TiO2 powder P-25 (Degussa) shows a mixed structure of anatase and rutile and exhibits high photocatalytic reactivity, it could be expected that films prepared by this ICB deposition method would also show good performance. The optical properties of the thin films were investigated by UV–visible absorption measurements, as shown in Fig. 6. Clear interference fringes, which are a characteristic feature of thin films, could be observed in the visible regions, indicating that uniform and highly transparent films were formed on the substrates [42–46, 55]. The absorption edges were found to shift towards shorter wavelength regions as the film thickness decreased. This can be attributed to the quantum size effect caused by the presence of nano-sized TiO2 particles as the composition of the transparent thin films. In fact, UV light irradiation (λ > 270 nm) of the thin films in the presence of NO was found to lead to the photocatalytic decomposition of NO into N2 , O2 and N2 O at 275 K. The yield of the formation of N2 increased linearly with the irradiation time. Moreover, the formation of these reaction products did not occur under dark conditions. It was therefore clearly seen that TiO2 thin films prepared by this deposition method exhibits

8.2.3 Visible Light-Responsive TiO2 Photocatalysts

high photocatalytic reactivity for the decomposition of NO under UV light, their reactivity being comparable to or even higher than that of the films prepared by the sol–gel method. Also, the reactivities were found to depend strongly on the film thicknesses and Fig. 7 shows the effects of the film thickness on the photocatalytic reactivity, the BET surface areas and the wavelength of the absorption edge. Thin films having a small thickness showed high photocatalytic reactivity which decreased as the film thickness increased and ultimately leveled off at around 50% of the maximum yield. The BET surface areas and wavelength of the absorption edge showed the same tendency and these results indicate that the photocatalytic properties of the TiO2 thin films are strongly dependent on the film thickness. Although these thin films prepared by ICB deposition were found to show high reactivity under UV light,

(b) (c)

30 % (d)

200

400

600

800

Wavelength/nm UV–visible transmittance spectra of TiO2 thin films prepared by ICB deposition. Film thickness: (a) 20; (b) 100; (c) 300; (d) 1000 nm.

30

25 0.08 20 0.04 15

0

0

400

800

10 1200

280

300

320

340

Wavelength of absorption edge/nm

Amount of produced N2/µmol

0.12

BET surface area (TiO2 side only)/m2

Fig. 6

they did not work as photocatalysts under visible light because of their high transparency in visible light regions. Therefore, an advanced ion implantation method was applied to modify their properties as in the case with the powdered TiO2 system [15–26]. The UV–visible transmittance spectra of the Cr ion-implanted TiO2 thin films are shown in Fig. 8 and the absorption edge is shown to shift toward visible light regions as the amount of Cr ions increases, indicating that these thin films can be successfully modified to absorb visible light by Cr ion implantation [43]. In contrast, Cr ions chemically doped by an impregnation method did not produce such a shift into longer wavelength regions but instead caused a new shoulder due to the formation of aggregated Cr oxide (CrOx ) species on the TiO2 surface. Hence it can be seen that the absorption spectra of thin films containing physically implanted Cr ions are completely different from those containing chemically doped Cr ion species. These results indicate that the method of doping the second components determines whether the modification of the electronic properties of the TiO2 semiconductor can be realized. Visible light irradiation (λ > 450 nm) of the Cr ionimplanted TiO2 thin films in the presence of NO led to the efficient decomposition of NO at 275 K whereas the decomposition reaction did not take place either on the unimplanted original TiO2 thin films or on the thin films chemically doped with Cr ions, as shown in Fig. 9. Moreover, the reaction proceeded with good linearity versus the irradiation time. These results show that the Cr ions implanted into the TiO2 lattice do not work as electron–hole recombination centers but only modify the electronic properties of the catalyst, permitting absorption and operation under visible light irradiation.

Transmittance/a.u.

Transmittance/a.u.

(a)

1963

200

(a)

(b) (c)

300

360

400

500

600 650

Wavelength/nm

Film thickness/nm

UV–visible transmittance spectra of (a) TiO2 and (b, c) Cr ion-implanted TiO2 thin films. Amount of Cr ions implanted: (a) 0; (b) 0.22; (c) 0.66 µmol g−1 .

Fig. 8

Relationship between the photocatalytic reactivities for NO decomposition (circles), BET surface areas (diamonds) and wavelengths of the absorption edges (squares) of the TiO2 thin films prepared by ICB deposition.

Fig. 7

References see page 1968

8.2 Photocatalysis

Amount of produced N2/µmol

1964

0.08

off

on

off

on

Cr-implanted TiO2 thin film

0.04

Cr-impregnated TiO2 thin film Original TiO2 thin film 0 −2

0

2

6

4

8

Time/h Reaction vs. time profiles of the photocatalytic decomposition of NO over the Cr ion-implanted TiO2 thin-film, original unmodified thin-film and Cr ion-impregnated thin-film photocatalysts under visible light (λ > 450 nm). Photoreactions were carried out in the same way as shown in Fig. 4.

Fig. 9

(c)

(b)

FT of k 3c(k)/a. u.

(a)

Normalized absorption/a. u.

In order to understand the mechanisms behind the absorption of visible light after ion implantation, it is important to clarify the local structures of the Cr ions physically incorporated into the TiO2 catalysts. K-edge XAFS measurements on the Cr ions implanted were carried out [44] and Fig. 10 shows the Cr K-edge XANES spectra (a, b) and Fourier transforms of the EXAFS oscillations (c, d). The shape of the XANES region for the Cr ion implanted into the TiO2 photocatalyst in spectrum (a) is found to be similar to that of the reference Cr2 O3 in spectrum (b). Moreover, for the Fourier transforms of the EXAFS oscillation, a strong peak at around 1.6–1.7 A˚ attributable to the neighboring O atoms (Cr−O) could be seen in the first coordination sphere in addition to a peak at around 2.8–3.0 A˚ attributable to the Ti atoms (Cr−O−Ti) in the second coordination sphere in spectrum (c).

(d) 5970

6010

6050

Energy/eV

0

2

4

6

Distance/Å

Fig. 10 Cr K-edge XANES (a, b) and Fourier transforms of EXAFS oscillation (c, d) of the Cr ion-impregnated TiO2 (a, c) and the Cr ion-implanted TiO2 photocatalysts (b, d).

These results indicate that the Cr ions implanted into thin films having an octahedral structure are incorporated within the TiO6 lattice positions in place of the Ti ions. A modification of the electronic state of the TiO2 catalysts by metal ion implantation may also be closely associated with the strong and long-distance interactions between the TiO2 semiconductor and metal ions implanted and not with the formation of impurity energy levels within the bandgap of the catalysts. These findings suggest that highly dispersed and isolated Cr ions within the TiO6 lattice incorporated specifically by metal ion implantation play an important role in the design of TiO2 photocatalysts reactive under visible light irradiation. ESR investigations have also revealed the existence of characteristic reticular V4+ ions only on the V ion-implanted photocatalysts after calcination at around 723–823 K under an O2 atmosphere. The V ion-implanted TiO2 photocatalysts, on which there characteristic reticular V4+ ions could be observed, also showed an efficient red shift in the absorption spectra towards visible light regions. V ions having such a local structure or shifts in the absorption band could not be observed for the TiO2 catalysts chemically doped with the V ions. Hence it could be seen that the ion-implanted metal ions are highly dispersed within the TiO2 lattice positions and work to modify the electronic nature of the catalysts without any changes to the chemical properties of the catalyst surfaces. Preparation of Visible Light-Responsive TiO2 Thin-Film Photocatalysts by RF-Magnetron Sputtering (RF-MS) Deposition in a Single Process An RF-MS deposition method was used to prepare welldefined TiO2 thin-film photocatalysts using a TiO2 plate as the sputtering target and Ar as the sputtering gas. In the course of preparing highly transparent films, an alternative and more practical preparation method to synthesize visible light-responsive TiO2 thin films was developed [45, 46]. Figure 11 shows the UV–visible transmittance spectra of the thin films prepared by RF-MS deposition at different substrate temperatures. At relatively low temperatures (773 K), the films exhibited efficient absorption in visible regions with maximum absorption when prepared at 873 K. The origin of the efficient absorption of visible light was not associated with the contamination of the films with impurities since the impurity content of the sputtering target material was less than 0.1%. In fact, the highly transparent thin films prepared at temperatures lower than 473 K showed efficient photocatalytic reactivity for 8.2.3.3

8.2.3 Visible Light-Responsive TiO2 Photocatalysts

UV-type TiO2 thin film (473 K) 30

Conversion of CH3CHO/%

various significant reactions, such as the decomposition of NO, the complete oxidation of acetaldehyde with O2 and the decomposition of H2 O into H2 and O2 , only under UV light irradiation. However, the visible light-responsive TiO2 thin films exhibited good reactivity not only under UV but also visible light irradiation. The reactivity of these TiO2 thin films under visible light (λ > 450 nm) was completely coincident with the relative intensity of the absorbance at 450 nm in the UV–visible absorption spectra, indicating that they can operate efficiently under visible light irradiation with wavelengths longer than 450 nm, as shown in Fig. 12. Practical applications of these UV and visible lightresponsive TiO2 photocatalysts such as the complete oxidation of harmful organic compounds into CO2 and

20

Transmittance/a. u.

(d) (c) 20%

400

600

800

Wavelength/nm UV–visible transmittance spectra of TiO2 thin films prepared by an RF-MS deposition method. Substrate temperature: (a) 373; (b) 473; (c) 673; (d) 873; (e) 973 K.

0.8

l > 450 nm 1.0

0.4

0.5

0 500

700

Relative intensity at 450 nm of UV-VIS absorption spectra

Fig. 11

Amount of produced N2 /µmol m−2

l > 450 nm

Comparison of the photocatalytic reactivities of TiO2 thin films prepared at 473 and 873 K for the oxidation of acetaldehyde with O2 under UV (λ > 270 nm) and visible light (λ > 450 nm), respectively.

(e)

0 300

l > 270 nm

Fig. 13

(b)

200

Visible-type TiO2 thin film (873 K)

10

0

(a)

1965

900

Preparation temperatures/K Fig. 12 Relationship between the photocatalytic reactivity for the decomposition of NO over TiO2 thin films prepared by RF-MS deposition under visible light (λ > 450 nm) and the relative intensities at 450 nm of the UV–visible absorption spectra.

H2 O have been investigated [20–26, 45]. Figure 13 shows the results for the decomposition of acetaldehyde with O2 into CO2 and H2 O under both UV and visible light. The highly transparent thin films (sample size: 10 × 10 × 1 mm) prepared at 473 K were found to decompose about 30% of the acetaldehyde into CO2 and H2 O in the presence of sufficient amounts of O2 under UV light irradiation for 4 h, whereas visible lightresponsive thin films prepared at 873 K were found to decompose about 10% of the acetaldehyde under visible light. These results indicated that TiO2 thin films prepared by RF-MS deposition have the potential for applications in the removal of various air pollutants not only outdoors but also indoors. The photocatalytic decomposition of H2 O into H2 and O2 using these newly developed visible-light responsive TiO2 thin-film photocatalysts under visible light irradiation at wavelengths longer than 420 nm has been successfully carried out [46]. Moreover, these thin films (sample size: 10 × 20 × 1 mm) were applied to an H-type cell like an electrochemical cell in order to evolve H2 and O2 separately, as shown in Fig. 14. In fact, it was found that the separate evolution of H2 and O2 could be successfully achieved under visible light using a conventional solar light gathering system, as shown in Fig. 15. Although the evolution rates for H2 and O2 are dependent on the solar light intensities and weather conditions, these thin-films are promising for the effective References see page 1968

1966

8.2 Photocatalysis

Photocatalyst TiO2 thin film

Pt H+

O2 e−

h+ OH−

H2 Ti substrate

Quartz window Solar light gathering system

Optical fiber Solar light Stainless steel

Membrane

Stainless steel

H-type reaction cell

Schematic diagram of the experimental setup for the separate evolution of H2 and O2 by the H2 O splitting reaction over visible light-responsive TiO2 thin-film photocatalysts under solar light.

Fig. 14

utilization of visible light to produce H2 energy. Also, the preparation of such thin films could be realized by a single process with RF-MS deposition, allowing for their widespread availability at low cost [47–50].

Relative intensity of solar light/a. u.

Amounts of evolved H2 and O2/µmol

25 Solar beam intensity H2

20

15 O2

10

5

0 0

2

4

(9:00)

6

8 (17:00)

Time/h Fig. 15 Reaction time profiles of the photocatalytic splitting of H2 O into H2 (circles) and O2 (squares) over a visible light-responsive TiO2 thin film equipped with an H-type reaction cell under solar light and the relative light intensities (diamonds). Open squares were plotted for the photocatalytic reactivity of the UV light-responsive TiO2 thin film as a reference. Solar light irradiation: 7 h from 9.30 a.m. to 4.30 p.m.

The mechanisms behind the efficient absorption of visible light were investigated by observing the crosssectional SEM images of both the UV and visible-type TiO2 thin films. The highly transparent TiO2 thin-films prepared at 473 K, which work only under UV light, were found to consist of TiO2 fine particles randomly sintered with each other, as shown in Fig. 16. Meanwhile, for the visible light-responsive TiO2 thin films prepared at 873 K, the columnar TiO2 crystals (diameter: ca. 100 nm) were found to form an orderly alignment. The depth profiles of the O:Ti atomic ratios were also investigated by AES measurements (data not shown). These results showed that the O:Ti atomic ratio of the UV-type thin films was constant from surface to bulk at about 2.0, indicating the stoichiometric composition of the TiO2 , whereas the O:Ti atomic ratios of the visible-type thin films were found to decrease gradually from the surface to deep bulk, reaching about 1.93. Such a declined structure of the films was confirmed to be stable even after calcination in air at 773 K, since the stoichiometric TiO2 layers near the surface were found to protect the bulk as a passive layer. Moreover, the declined structure was found to be closely associated with the modification of the electronic properties of the TiO2 and to allow the absorption of visible light. It has already been reported that small amounts of an oxygen vacancy in the TiO2 lattice give rise to distortion of the TiO2 octahedral unit and weaken the Ti–O bonds, resulting in a reduction of the splitting between the bonding and non-bonding levels [51]; however, further investigations on the role of the oxygen vacancy is under way. Our

8.2.5 Conclusions

1967

TiO2

Quartz S4700-1 10.0kV 12.0mm x30.0k YAGBSE 1.00µm

(a)

Fig. 16

S4700-3 10.0kV 12.0mm x30.0k YAGBSE 1.00µm

(b)

Cross-sectional SEM images of a TiO2 thin film prepared at 473 K (a) and 873 K (b) by RF-MS deposition.

analytical investigations showed that the characteristic declined composition of the TiO2 thin films may be the determining factor in modifying the electronic properties so that they are able to absorb and operate as efficient photocatalysts under both UV and visible light irradiation. 8.2.4

Photo-induced Superhydrophilic Properties on the Surface of the TiO2 Thin-Film Photocatalysts

Since the discovery of photo-induced superhydrophilicity on TiO2 surfaces in 1997 [39], TiO2 thin-film photocatalysts prepared on various substrates have attracted much attention as a useful photo-functional material for selfcleaning, antifogging, antibacterial and stain-proofing applications. In order to improve the efficiency of the photo-induced hydrophilicity, titanium oxide-based mixed oxides such as TiO2 −SiO2 and TiO2 −B2 O3 with different TiO2 contents on glass substrates have been prepared and it was found that they can exhibit much higher photofunctionality than the pure untreated thin films [52–55]. The film surfaces showed high wettability under UV light irradiation, whereas under dark conditions this photoinduced wettability quickly disappeared. However, with TiO2 −SiO2 and TiO2 −B2 O3 binary mixture oxides or multi-layered thin films, enhanced surface wettable properties were observed under UV light irradiation which could be sustained even under dark conditions for prolonged periods. The high wettability of TiO2 photocatalyst was found to be achieved by irradiating the TiO2 surfaces only in the presence of O2 molecules [59]. This result showed that the O2 molecules play an important role for the photo-driven wettable conversion of TiO2 surfaces. Kn¨ozinger and coworkers have investigated the dynamics of the photo-formed electron–hole pairs trapped on TiO2 surfaces and also the interaction of O2 molecules with TiO2 surfaces by EPR and FT-IR measurements [56, 57]. Furthermore, detailed FT-IR [mid-infrared (MIR)] and near-infrared (NIR) spectroscopic measurements were also carried out to clarify the mechanisms behind

this photo-induced hydrophilicity for TiO2 photocatalysts [58, 59]. The following facts could be clarified: when the TiO2 surfaces were irradiated with UV light, the H2 O molecules on the surfaces desorbed due to the irradiation effect of the UV light, while the hydrocarbons adsorbed on the surfaces were simultaneously decomposed into CO2 and H2 O by the reactivity of the TiO2 . Also, when the amount of the adsorbed H2 O molecules decreased by UV light irradiation, the distribution of the intermolecular hydrogen-bonded H2 O clusters was also found to decrease, indicating that the shape of the H2 O clusters on the surfaces changes from a bulky round form to a spread-out layer form on irradiation with UV light. Such a decrease in the distribution of the hydrogen-bonded H2 O clusters was considered to play an important role as the driving force behind the photo-induced hydrophilicity. 8.2.5

Conclusions

The preparation of the well-defined TiO2 photocatalysts which can operate under visible or solar light irradiation by applying advanced ion engineering techniques such as metal ion implantation, ionized cluster beam deposition and RF-MS deposition was carried out along with characterization studies to determine the mechanism behind the reactions at the molecular level. The efficient utilization of visible and/or solar light energy has become of special interest in recent years since it is a safe and abundant energy source that is environmentally clean. Photocatalytic systems to harness the Sun’s energy for the purification of polluted air and water are being actively sought and the main challenges are to develop photocatalysts with ever higher reactivity and which can work under both UV and visible light. With these aims in mind, we have focused our attention on developing such visible light-responsive photocatalysts by advanced metal ion implantation and RF-MS deposition methods. In References see page 1968

1968

8.2 Photocatalysis

order to understand the reaction mechanisms, the direct detection and characterization of the reaction intermediate species at the molecular level have been investigated. Such visible light-responsive titanium oxide catalysts developed by incorporating a secondary compound have shown potential in such significant reactions as the reduction of CO2 into useful CH4 and CH3 OH, the decomposition of NOx into harmless compounds, the oxidation of various organic or toxic compounds into non-toxic CO2 and H2 O and the production of H2 in the photocatalytic splitting of H2 O under solar light irradiation. References 1. M. Anpo, in Green Chemistry, P. Tund, P. Anastas (Eds.), Oxford University Press, Oxford, 2000. 2. Y. Kubokawa, K Honda, Y. Saito (Eds.), Hikarishokubai (Photocatalysis), Asakura-shoten, Tokyo, 1988. 3. N. Serpone, E. Pelizzetti (Eds.), Photocatalysis Fundamentals and Applications, Wiley, New York, 1989. 4. D. F. Ollis, H. Al-Ekabi (Eds.), Photocatalytic Purification and Treatment of Water and Air, Elsevier, Amsterdam, 1993. 5. A. Fujishima, K. Hashimoto and T. Watanabe (Eds.), TiO2 Photocatalysis Fundamentals and Applications, BKC, 1999. 6. M. Anpo, H. Yamashita, in Surface Photochemistry, M. Anpo (Ed.), Wiley, Chichester, 1996. 7. M. Anpo, H. Yamashita, in Heterogeneous Catalysis, M. Schiavello (Ed.), Wiley, Chichester, 1997. 8. M. Anpo, in Handbook of Heterogeneous Catalysis, G. Ertl, H. Kn¨ozinger, J. Weitkamp (Eds.), Vol. 2, Wiley-VCH, Weinheim, 1997, p. 664, and references therein. 9. B. O’Regan, M. Gratzel, Nature 1991, 353, 737. 10. A. K. Ghosh, H. P. Maruska, J. Electrochem. Soc. 1977, 124, 1516. 11. E. Borgarello, J. Kiwi, M. Gr¨atzel, E. Pelizetti, M. Visca, J. Am. Chem. Soc. 1982, 104, 2996. 12. M. R. Hoffman, S. T. Martin, W. Choi, D. W. Bahnemann, Chem. Rev. 1995, 95, 69. 13. J. M. Jerrmann, J. Disdier, P. Pichat, Chem. Phys. Lett. 1984, 108, 618. 14. H. P. Maruska, A. K. Ghosh, Solar Energy Mater. 1979, 1, 237. 15. M. Anpo, Catal. Surv. Jpn. 1997, 1, 169. 16. M. Anpo, Y. Ichihashi, M. Takeuchi, H. Yamashita, Res. Chem. Intermed. 1998, 24, 143. 17. M. Anpo, Y. Ichihashi, M. Takeuchi, H. Yamashita, Science and Technology in Catalysis 1998, H. Hattori, K. Otsuka (Eds.), Kodan-sha, Tokyo, 1999, p. 305. 18. M. Anpo, M. Che, Adv. Catal. 1999, 44, 119, and references therein. 19. M. Anpo, H. Yamashita, S. Kanai, K. Sato, T. Fujimoto, US Patent 6 077 492, 2000. 20. M. Anpo, Pure Appl. Chem. 2000, 72, 1265. 21. M. Anpo, Stud. Surf. Sci. Catal. 2000, 130, 157. 22. M. Anpo, M. Takeuchi, Int. J. Photoenergy 2001, 3, 1. 23. M. Anpo, M. Takeuchi, K Ikeue, S. Dohshi, Curr. Opin. Solid State Mater. Sci. 2002, 6, 381. 24. M. Anpo, M. Takeuchi, J. Catal. 2003, 216, 505. 25. H. Yamashita, M. Anpo, Catal. Surv. Asia 2004, 8, 1. 26. M. Anpo, Bull. Chem. Soc. Jpn. 2004, 77, 1427, and references therein.

27. K. Takami, N. Segawa, H. Uehara, M. Anpo, Shokuba 2002, 41, 295. 28. I. Rosenberg, J. R. Brock, A. Heller, J. Phys. Chem. 1992, 96, 3423. 29. N. Negishi, T. Iyoda, K. Hashimoto, A. Fujishima, Chem. Lett. 1995, 841. 30. I. Sopyan, M. Watanabe, S. Murasawa, K. Hashimoto, A. Fujishima, Chem. Lett. 1996, 69. 31. Y. Ohko, K. Hashimoto, A. Fujishima, J. Phys. Chem. A 1997, 101, 8057. 32. Negishi, K. Takeuchi, T. Ibusuki, J. Mater. Sci. 1998, 33, 1. 33. H. Y. Lee, Y. H. Park, K. H. Ko, Langmuir 2000, 16(18), 7289. 34. D. Byun, Y. Jin, B. Kim, J. K. Lee, D. Park, J. Hazard. Mater. B 2000, 73, 199. 35. S. Deki, Y. Aoi, O. Hiroi, A. Kajinami, Chem. Lett. 1996, 433. 36. K. Shimizu, H. Imai, H. Hirashima, K. Tsukuma, Thin Solid Films 1999, 351, 220. 37. S. Yamabi, H. Imai, Chem. Lett. 2001, 220. 38. M. Kiyono, Sanka-Titan (Titanium Dioxide) – Bussei to Oyogijutu, Giho-do, Tokyo, 1991. 39. R. Wang, K. Hashimoto, A. Fujishima, M. Chikuni, E. Kojima, A. Kitamura, M. Shimohigoshi, T. Watanabe, Nature 1997, 388, 431. 40. N. Sakai, R. Wong, A. Fujishima, T. Watanabe, K. Hashimoto, Langmuir 1998, 14, 5918. 41. R. Wang, N. Sakai, A. Fujishima, T. Watanabe, K. Hashimoto, J. Phys. Chem. B 1999, 103, 2188. 42. M. Takeuchi, H. Yamashita, M. Matsuoka, T. Hirao, N. Itoh, N. Iwamoto, M. Anpo, Catal. Lett. 2000, 66, 185. 43. M. Takeuchi, H. Yamashita, M. Matsuoka, T. Hirao, N. Itoh, N. Iwamoto, M. Anpo, Catal. Lett. 2000, 67, 135. 44. M. Takeuchi, M. Matsuoka, H. Yamashita, M. Anpo, J. Synchrotron Radiat. 2001, 8, 643. 45. M. Takeuchi, M. Anpo, T. Hirao, N. Itoh, N. Iwamoto, Surf. Sci. Jpn. 2001, 22(9), 561. 46. M. Kitano, M. Takeuchi, M. Matsuoka, J. M. Thomas, M. Anpo, Chem. Lett. 2005, 34, 616. 47. H. Yamashita, M. Takeuchi, M. Anpo, Encycl. Nanosci. Nanotechnol. 2004, 10, 639. 48. M. Anpo, S. Dohshi, M. Kitano, Y. Hu, M. Takeuchi, M. Matsuoka, Annu. Rev. Mater. Res. 2005, 35, 1. 49. K. Iino, M. Kitano, M. Takeuchi, M. Matsuoka, M. Anpo, Curr. Appl. Phys. 2005, in press. 50. M. Matsuoka, Masaaki Kitano, M. Takeuchi, M. Anpo, J. M. Thomas, Mater. Sci. Forum 2005, 487, 81. 51. S. A. Bilmes, P. Mandelbaum, F. Alvarez, N. M. Victoria, J. Phys. Chem. B 2000, 104, 9851. 52. S. Dohshi, M. Takeuchi, M. Anpo, J. Nanosci. Nanotechnol. 2001, 1, 337. 53. M. Anpo, M. Takeuchi, K. Ikeue, S. Dohshi, Curr. Opin. Solid State Mater. Sci. 2002, 6, 381. 54. S. Dohshi, M. Takeuchi, M. Anpo, Catal. Today 2003, 85, 199. 55. M. Takeuchi, S. Dohshi, T. Eura, M. Anpo, J. Phys. Chem. B 2003, 107, 14278. 56. T. Berger, M. Sterrer, O. Diwald, E. Kn¨ozinger, D. Panayotov, T. L. Thompson, J. T. Yates Jr., J. Phys. Chem. B 2005, 109, 6061. 57. T. Berger, M. Sterrer, S. Stankic, Johannes Bernardi, O. Diwald, E. Kn¨ozinger, Mater. Sci. Eng. C 2005, in press. 58. M. Takeuchi, G. Martra, S. Coluccia, M. Anpo, J. Phys. Chem. B 2005, 109, 7287. 59. M. Takeuchi, K. Sakamoto, G. Martra, S. Coluccia, M. Anpo, J. Phys. Chem. B 2005, 109, 15422.

8.3.1 Introduction

8.3

Chemical Sensors Based on Catalytic Reactions Nicolae Barsan, Klaus D. Schierbaum∗ , and Udo Weimar

8.3.1

Introduction Definitions and Classifications A useful definition of a chemical sensor is ‘‘a small device which is capable of continuously converting a chemical state or, more precisely, a chemical or biochemical information in real time to an output signal’’ [1]. The output signal, commonly denoted simply the sensor response, is typically electronic, in some cases optical, in nature and correlates to the concentration, partial pressure or activity of atoms, molecules or ions; hence it is of analytical use. The definition is illustrated in Fig. 1 [2]. The selective binding (and in some cases, conversion) of a target molecule in a complex gas- or liquid-phase matrix is normally associated with a ‘‘recognition’’ element while the generation of the signal is achieved by a transducer. When the analyte interacts with the recognition element, a change in one or more physicochemical parameters occurs and the transducer converts these parameters into the output signal, which can be amplified, processed and displayed. When biomolecules and/or structures with biological significance or bioactivity are monitored, a more appropriate and common denotation of the device is ‘‘biosensor’’. Other different chemical sensors have been invented and developed during the last 100 years which are based on a variety of very distinct physicochemical quantities (e.g. heat conductance, heat of combustion, paramagnetism, refractive index) [3]. In contrast to conventional, accurate chemical solutions to chemical sensing, which involve complex and expensive laboratory methods, low-cost solutions are achieved by relatively simple solid-state or more advanced microelectronic technologies. Today, chemical sensors are used in a wide variety of disciplines, ranging from electrochemical analysis, through biomedical measurement, to pollution monitoring and industrial control. On the other hand, arrays of chemical sensors also offer new applications. In 1982, Persaud and Dodd introduced the concept of an electronic nose by proposing a system which comprises an array of essentially non-selective sensors and an appropriate pattern recognition system (e.g. an artificial neural network) [4]. Chemical sensors may be classified according to their detection principle (i.e. monitoring conductivities, 8.3.1.1

∗

Corresponding author.

1969

potentials, capacities, heats, masses or optical constants) with characteristic dependences of sensor responses x  upon changed analytical information pi [5]. If chemical sensors are to be used as analytical tools, timeindependent and stable calibration curves x  = f(pi ) are required. This implies the response x  to be a state function in a thermodynamic sense of reversibility. Over large ranges of pi this is usually not fulfilled. For limited ranges and restricted combinations of pi , however, and for a certain range of the temperature T , variations of x  are described by

 ∂x ∂x  dp1 + dp2 + . . . dx = ∂p1 p1=j,T ∂p2 p2=j,T
 ∂x dT (1) + ∂T pj and hence  x = 0

(2)

must hold. Here, γi = (∂x  /∂pi )pi=j,T denotes the partial sensitivity of the sensor with respect to changes of pi . Sensor performance is characterized by selectivity, sensitivity, stability, reliability and reusability. Selectivity, a measure of how well a chemical sensor discriminates between the analyte and compounds of similar or different chemical structure, is principally determined by the recognition component within the sensor device. Sensitivity is determined by the recognition element and the transducer. Depending on the signal-to-noise ratio, additional amplification steps can enhance sensitivity and lower the detection limit of the analyte. Further details of sensor parameters can be found elsewhere [6]. Maintaining long-term stability, withstanding harsh chemical environments and operating at high temperatures and/or pressures are severe challenges for sensors. Typical Examples Research and development of chemical and biochemical sensors are characterized by a huge gap between the large amount of new ideas and prototypes utilizing new principles of chemical sensing on the one hand and the very limited number of practically important sensors which are currently manufactured in large quantities on the other. Examples of commonly used chemical sensors are amperometric gas sensors, lambda probes, Taguchi SnO2 sensors, pellistors, pH ion-sensitive and Clark electrodes (to monitor oxygen in blood) and Pd-gate field effect transistors (GASFETs). 8.3.1.2

References see page 1985

1970

8.3 Chemical Sensors Based on Catalytic Reactions

Chemical sensor Heat electrons ions gases masses light

X′

pi

Particles

Filter

Recognition site

Transducer

Electronics

Electrical signal

Key–lock arrangement of a typical chemical sensor to detect atoms, molecules or ions (‘‘particles’’) in the gas or liquid phase. The sensor consists of the transducer, which transforms chemical input signals upon interaction of particles with recognition sites of the chemically sensitive layer into electrical output signals. Additional elements of ‘‘real’’ sensors are also indicated.

Fig. 1

Anode reaction: CO + H2O

CO2 + 2H+ + 2e−

Sample gas CO2

CO

2H+

I

2e−

Diffusion barrier Working electrode (anode) Electrolyte (H2SO4)

CO 2e−

CO2

Potentiostat

2e−

2e− 2H+

I

Counter electrode (cathode) ½ O2

½ O2

Diffusion barrier

Reference electrode

Reference gas Cathode reaction ½ O2 + 2H+ + 2e−

H 2O

Examples of liquid electrolyte sensors with two-electrode (left) and three-electrode arrangements (right) to monitor CO utilizing the amperometric detection principle. The sensor signal is the current I.

Fig. 2

Figures 2–8 illustrate schematically the various detection principles of ‘‘basic’’ sensors. In most of these examples, (electro)catalytic processes are involved during sensing – at least as intermediate or precursor steps. The detection of CO in air with two- and three-electrode liquid electrolyte sensors utilizing an amperometric operation mode is illustrated in Fig. 2. In this mode, the current I is monitored at a constant voltage V . The principle is based on the electrocatalytic reaction between CO and O2 molecules to form CO2 . A positive (anodic) bias voltage is applied between the working electrode (usually made from a noble metal such as platinum or gold) and the counter electrode. CO enters through a thin gas-permeable membrane (usually Teflon, which acts as

a diffusion barrier and limits the amount of CO reaching the anode) and is oxidized to CO2 : + − −  CO + H2 O − −− − − CO2 + 2H + 2e

(3)

This process consumes one H2 O molecule per CO molecule and simultaneously releases two protons, which diffuse through the electrolyte to the counter electrode. Corresponding anode reactions for other gases such as H2 S, NO and H2 are + − −−  H2 S + 4H2 O −− − − H2 SO4 + 8H + 8e

(4)

+ − −  NO + 2H2 O − −− − − HNO3 + 3H + 3e

(5)

+

−

−−  H2 −− − − 2H + 2e

(6)

8.3.1 Introduction

Ceramic-based protection layer (porous) Exhaust: Contact Anode reaction: v O2 + 4e− 2O2−

CO

Solid electrolyte ZrO2 (Y2O3)

(CO-O)d

O−

O−

e−

Heater

Air: Contacts Cathode reaction: O2 + 4e− 2O2−

CO2

− O− OH− O− O2

SnO2 (Pd, Al2O3, SiOx,…)

1971

Al2O3 substrate Au contacts

Ponous pt

Electrodes

Exhaust

Air

G

Electronic conductance sensor to monitor reducing gases such as CO. The sensor signal is the conductance σ .

Fig. 4

Solid-state electrolyte sensor to determine oxygen partial pressures (lambda probe) utilizing the potentiometric detection principle. The sensor signal is the voltage V.

Fig. 3

The reduction of O2 from the sampled gas stream (a few thousand ppm is normally sufficient) takes places at the cathode −−  1/2O2 + 2H+ + 2e− −− − − H2 O

(7)

so that the overall reaction for the CO detection is given by −−  CO + 1/2O2 −− − − CO2

(8)

resulting in a flow of electrons from the working electrode to the counter electrode through the external circuit. The rate-determining step in the overall reaction is controlled by the diffusion of CO molecules from the diffusion barrier to the surface of the anode where the fast oxidation

of CO molecules takes place. Since the concentration gradient between diffusion barrier and anode itself is proportional to the partial pressure of CO, the reaction rate and the electric current are proportional to the CO concentration; hence the sensor response x  is a linear function of pCO . The auxiliary or counter electrode therefore should not be polarized in order to maintain constant potential during the measurement. This problem is solved with a three-electrode setup in which a potentiostat keeps the potential between the working and reference electrodes constant. The three-electrode configuration allows for precise operation even with microelectrodes. In principle, this polarographic principle can be utilized for the detection of all electrochemically active species which can be oxidized or reduced at a certain potential. Since this potential is characteristic References see page 1985

Gases (H2, O2, H2O, CO,…) O H

O H

O HH

H H

O C

Pd SiO2

Vg

Id

Surface dipoles Interface dipoles

Vd

Pd-gate Source-contact

Drain-contact Thin SiO2 -coating

n

n p–Si Thin SiO2 -gate

Chemical field-effect sensor to monitor H2 . The sensor signal is the change in the drain-source voltage V or the change in the drain current.

Fig. 5

1972

8.3 Chemical Sensors Based on Catalytic Reactions

for a specific electrode and chemical component, its precise adjustment in a three-electrode arrangement (Fig. 2, right) provides greater specificity to target gases than the cheaper two-electrode arrangement (Fig. 2, left), resistance to over-exposure and stability in the presence of changing temperatures, pressures and humidity. Sensors involving a reduction reaction of the target gas – such as the reduction of nitrogen dioxide, chlorine and ozone – at the cathode produce water as a byproduct: −−  NO2 + 2H+ + 2e− −− − − NO + H2 O +

(9)

−

−  Cl2 + 2H + 2e − −− − − 2HCl

(10)

−−  O3 + 2H+ + 2e− −− − − O2 + H2 O

(11)

At the anode, water is simultaneously oxidized. Such sensors do not require the presence of oxygen to

CO

O2 O + O CO

Catalytically active metal atoms (Pt, Pd,…)

Qreact

Inactive oxide (ThO2, Al2O3,…)

Filament (Pt, Ir)

Calorimetric sensor to detect reducing gases. The sensor signal is the heat of reaction produced per unit time.

Fig. 6

function properly. A fourth auxiliary electrode can assist in overcoming cross-interference from other gases, for example when CO is monitored in the presence of H2 or if CO and H2 S are to be measured using one dualtype sensor. Liquid-electrolyte cells, which combine small size, low power, high sensitivity and accuracy (a typical value for CO is ±1 ppm at concentrations under 50 ppm), in addition to a relatively low price, are being used in many portable toxic and explosive gas applications for a broad range of analytes, including organic vapors such as alcohols, aldehydes or ketones [7]. Typical sensitivities are in the low-ppm range, but some sensors are capable of ambient ozone monitoring at ppb levels. However, temperature effects needs to be overcome for accurate detection at such low gas concentrations in order to avoid drifts of the baseline signal (which approximately doubles for every 10 ◦ C rise in temperature). An example for the use of electrochemistry in biosensing is given below. The potentiometric detection principle is demonstrated in Fig. 3 for the ZrO2 -based lambda probe (λ sond; the name originates from the shape of the voltage versus fuel-toair ratio), inserted prior to three-way catalytic converters in automobiles. Robert Bosch GmbH in Germany introduced the first automotive lambda probe in 1976. This solid-state electrolyte sensor measures the oxygen partial pressure pO2 and controls the air-to-fuel (A/F) ratio in the internal combustion engine. The stabilized zirconia ceramic shows high bulk ionic conductivity of O2− if suitably doped with Ca2+ or Y3+ at temperatures above approximately 300 ◦ C. Porous Pt electrodes make possible the catalytically activated dissociation of molecular Glucose oxidase MW 186 000

High moleclar substance (proteins, T-cells,...) MW > 100 000

V = const.

Counter electrode

Potentiostat R Fe R Fe

Glucose MW 186

e−

Gluconolactone MW 184

Interferent substances (ascorbate, paracetamol...) MW 100–1000

Fe

R Mediator: ferrocene derivatives MW 240–500

R Fe

Amino acids MW 800–500 R

R

Fe

Fe

MW = molecular weight

Gluconolactone + GODred.

Glucose + GODox. GODred. + Mediatorox.

GODred. + Mediatorred. or

GODred. + O2

Analyte

Semipermeable membrane

Reaction volume

GODox. + H2O2

Working electrode

Biochemical sensor to detect glucose in an analyte (e.g. blood) utilizing the potentiometric detection principle. The sensor signal is the voltage V.

Fig. 7

8.3.1 Introduction

Gold electrode

Sensitive coating ∆m

Quartz

Analyte (CO, CO2, NO2, Trl, Per, Octane)

Air

Support

Test chamber with constant temperature

Contacts

Quartz-microbalance sensor to monitor organic solvent molecules in air (Tri = trichloroethylene, Per = perchloroethylene). The sensor signal is the change in the oscillation frequency f . Fig. 8

O2 from the gas phase and electron transfer between the cathode and anode: 2− −−  O2,gas + 4e− −− − − 2O 2−

2O

−

− −  −− − − O2,gas + 4e

(cathode reaction)

(12)

(anode reaction)

(13)

At a constant temperature and a constant oxygen partial pressure pO2 in the reference gas (usually air), a potential V is generated in the concentration gradient between anode and cathode. The potential V is determined by the difference in the chemical potentials of O2 between exhaust and the constant reference phase (Nernst equation). Liquid and solid electrolyte sensors may be operated as potentiometric (I = 0), amperometric (V = constant) or conductometric devices. For car exhaust control, the ideal point is 0.45 V DC; this is where the quantities of air and fuel are in the optimum ratio, called the stoichiometric point, and the exhaust output will mainly consist of fully oxidized CO2 . Monitoring dissolved oxygen in molten steel is another application [8–10]. As an example of an electronic conductance sensor, Fig. 4 shows a metal oxide-based chemical sensor. These devices are capable of monitoring reducing gases at concentrations relevant to atmospheric monitoring applications. The first commercially available CO detector was developed by the Japanese company Figaro Engineering in the 1960s and are known as the Taguchi sensor, after its inventor Naoyoshi Taguchi [11]. The sensor consists of a semiconducting tin oxide pellet, heated to approximately 250 ◦ C. The response signal x  is the electronic conductance σ (in contrast to the ionic conductance of conductometric electrolyte sensors). It is usually measured between two ohmic contacts. Its value is determined by competitive electronic charge-transfer reactions between negative oxygen species at the surface or at grain boundaries (formed by chemisorption and/or dissociation of O2 from the gas phase at the n-type semiconductor) and the reducing molecules: The oxide acts as a catalyst

1973

for the oxidation of the molecules. The detection of poorly active reducing molecules requires catalytically activated reaction steps at metallic surface dopants. As a result of the electronic charge-transfer reactions, a change occurs in the density of free (conduction band) electrons and hence in the conductance G of the oxide. The definition of electronic conductance sensors also includes Schottky diode sensors with their voltage-dependent conductances and capacitance (dielectric) sensors with their capacitances C as sensor signals. The latter are monitored with alternating current (AC). The general sensor responses of these types of sensors may be characterized by a change in the frequency-dependent complex conductance. The detection principle of a gas-sensitive field-effect transistor (GASFET), first reported by Lundstr¨om’s group in the mid-1970s [12], is illustrated in Fig. 5. This sensor consists of a Pd gate on top of an oxide layer, typically SiO2 , and a p-type silicon base with n-doped channels on either side of the gate. For a GASFET with continuous Pd gate, used for the detection of hydrogen at operating temperatures around 150 ◦ C, a hydrogen atom-induced change in the electrical double layer at the Pd/SiO2 interface occurs which can be monitored by a changed drain-source current Ids through the n-conducting channel. The value of Ids is usually adjusted at a certain level by an electrical field perpendicular to the channel which is generated by a positive voltage Vg at the Pd gate. Pd is used because of its unique properties with respect to the catalytically activated dissociation of H2 and diffusion of H atoms to the Pd/SiO2 interface. This process leads to the formation of surface und interface dipoles which then modify the conductance threshold voltage according to √ Vt = Vmax /[1 + 1/(c pH2 )]

(14)

where Vmax is the maximum voltage shift (above 0.5 V), c is a constant and pH2 is the partial pressure of H2 in air [13]. Other devices use field-effect transistors with a suspended gate so that a gateless FET acts as the transducer which converts changes of the electrical potential in the area above the channel to an electrical signal, i.e. a change of the drain-source current Ids [14]. Ion-sensitive field-effect transistors (ISFETs), invented by the Dutch scientist Bergveld in 1970 [15], monitor the concentration of ions or adsorbed oriented dipoles [16, 17]. In these devices, the metallic gate of GASFETs is replaced by the liquid and a reference electrode is added. A pellistor primarily used as a methanometer to detect CH4 in coal mine surroundings at levels above 1% is illustrated schematically in Fig. 6. Such a device is called a calorimetric sensor [18]. A pellistor consists of a coil of fine platinum wire embedded in a bead of References see page 1985

1974

8.3 Chemical Sensors Based on Catalytic Reactions

alumina. Usually two of these pellistors are operated in a Wheatstone bridge arrangement, one being impregnated with a catalyst (Pt, Pd, etc.) and the other being used as an inactive reference. By passing current through the pellistors, a temperature is maintained (about 500 ◦ C) at which oxidation of the methane readily occurs at the surface of the active pellistor. The reaction heat H , resulting from the catalytic oxidation, produces a change in temperature T monitored by the changed resistance of a platinum filament. The latter leads to a deviation in the difference voltage in the balanced bridge. In a second operation mode, the sensor signal may be chosen as the changed power required to keep the pellistor temperature at a constant value. Oxidation of reducing gaseous species leads to a decrease in the electrical power consumption P . Diffusion barriers in front of the pellistor may lead to a linearization of the sensor signal (P ∼ pi ). The invention of biosensors dates back to the American scientist Leland C. Clark in the 1960s, who showed that dissolved oxygen can be monitored by its electrochemical reduction with a membrane-coated Pt electrode (‘‘Clark electrode’’) [19]. As an example of a biochemical sensor based on enzymatic catalysis, the detection of glucose with a glucose oxidase-coated electrode is shown in Fig. 7. Glucose is a target for medical applications; amperometric sensors were first developed in the 1960s. In the example displayed in Fig. 7, glucose diffuses through the semipermeable membrane, which prevents a variety of interfering compounds from passing to the reaction volume. Glucose is oxidized here by the enzyme glucose oxidase (GODox ) to give gluconolactone and GODred . The subsequent electron transfer to the mediator compounds (such as ferrocene) leads to reduced species Mediatorred , which diffuse to the electrode surface and are oxidized there at a potential which is low in order to prevent the oxidation of interfering substances in the reaction volume. The current I is monitored as the sensor signal. Alternatively, the oxidation of GODred by O2 may occur to produce H2 O2 , which can also be monitored in amperometric glucose sensors. Typical bulk mass-sensitive sensors are quartz microbalance sensors (Fig. 8), which are coated with polymeric or supramolecular compounds. These devices make possible the sensitive weighing of mass changes upon ad- or absorption of particles from the gas or liquid phase, a method which was introduced by Sauerbrey in 1959 [20]. The increase in mass leads to a shift of the resonance frequency f0 . This shift may be monitored as a difference signal x  compared with oscillating uncoated quartz. Alternatives are surface acoustic wave (SAW) devices. Often, in particular, if polymers are used as chemically sensitive layers, the mass increase is determined thermodynamically and apparently does not involve a catalytic reaction.

In the context of the present chapter on chemical sensors based on catalytic reactions, we will concentrate here on solid-state electronic conductance, calorimetric and electrochemical chemical sensors; for further reading, see Refs. [21–26]. 8.3.2

Calorimetric Sensors

The first calorimetric sensors were developed for mine applications to determine dangerous levels of methane [27]. For this purpose, the combustion took place in a pure butane flame housed in a modified flame safety lamp and the heat of combustion was measured with a thermocouple. Changes in the methane concentration lead to changes in the thermocouple voltage that is recorded as a function of time. Later, these early ‘‘methanometers’’ were replaced by sensors that comprised two filaments which are arranged in a Wheatstone bridge circuit. In one arm of the bridge, a housed electrically heated filament coil is used to provide a sufficiently high temperature to oxidize the methane in the air sample that passed over the coil. When this happens, the temperature of the filament raises which in turn increases its electrical resistance, so that the Wheatstone bridge is imbalanced since the filament in the second arm is not heated and hence is inactive. By using the concept of a heated filament, embedded in a bead of alumina with a suitable catalyst at the surface (such as Pt, Pd, Rh, Ir), the so-called pellistor (from pellet + resistor) was invented (Fig. 6), and nowadays is the most commonly used sensor for flammable gas detection [28]. Pellistors present an industry standard for this purpose. Because of the catalytic conversion of molecules taking place at their surfaces, they are sometimes called ‘‘catalytic gas sensors’’. Pellistors provide high sensitivity to most flammable gases and vapors but require at least 10% of oxygen to work properly. The operating temperature of an active pellistor is about 500 ◦ C. For most gases, the response time is within a period of a few seconds and the output signal is linear if the gas concentration is in the lower explosive limit (LEL) range. A typical output reading of a pellistor is displayed in Fig. 9. However, catalyst poisoning and inhibition may occur for high carbon content gases, and also after exposure to sulfur- and chlorine-containing compounds, heavy metals and silicones [29], so that pellistors cannot be used in locations where exposure to such compounds is likely to occur. Pellistors can in principle be used in two different operating modes, namely isothermally and non-isothermally, the latter being applied practically for flame gas detection. ˙ r , released during the catalytic The combustion heat Q

8.3.3 Electronic Conductance Sensors

UEL

LEL

c ′ [V]

10

5

0 5

10

15

20

40

100

pCH4 [% ] Typical sensor output vs. gas concentration of a pellistor. Adapted from [6].

Fig. 9

oxidation of a gas, leads to a temperature change T of the active pellistor. For small changes T , it holds that T ∝ Q˙ r


˙ −1 dQab dT

(15)

˙ ab /dT denotes the change in the power loss in which dQ ˙ r is given by the in the working point [30]. The value of Q reaction rate dr/dt (moles of converted gas per unit time and area), the effective area Aeff of the catalyst’s surface and the reaction enthalpy H per mole: ˙ r = Aeff (dr/dt)H Q

(16)

The total reaction rate dr/dt shows a strong temperature dependence. At low temperatures in the kinetically controlled region I (Fig. 10), the value of dr/dt is controlled by the reaction rate at the pellistor surface which in turn is characteristic for a specific gas. In region III, in contrast, all reducible species diffusing to

Reaction rate

Transition region

1975

the pellistor surface are oxidized immediately. The total reaction rate is completely determined by the diffusion of molecules to the pellistor surface and is nearly constant for small variations in temperature. As a consequence of the diffusion law, a linear dependence between P and the partial pressure pi of the gas is usually observed in this region, which makes it possible to determine pi quantitatively. In the transition region II, the temperaturedependence of the reaction rate is characteristic for a specific gas and also for a specific catalyst with high reaction rates and therefore large pellistor responses. To obtain the LEL value of the detected gas, correction values must be taken into account, depending on the gas used for calibration, e.g. pentane or methane. In the isothermal mode, in which the temperature of the active pellistor is kept constant electronically by controlling the operation voltage of the Wheatstone bridge, one measures the change in the electrical power ˙ r . This mode consumption P , which is proportional to Q is rarely used in practical applications [29]. Micromachined ‘‘hot-plate’’ planar sensor structures, where a supported thin etched SiO2 or Si3 N4 membrane carries a platinum track on one side and a catalyst layer (e.g. Pd) on the other, have been fabricated and explored as ‘‘micro’’-pellistors [31–33]. Such a device is shown schematically in Fig. 11, allowing for low power consumption (e.g. 800 K for 15–25 mW [33]). In addition, these microcalorimetric devices are suitable for integration in sensor arrays with different catalytically active coatings. One of the main advantages of the calorimetric detection principle is the possible application of selective catalysts, which have been developed for a certain compound to be catalyzed in a specific reaction, for the detection of that compound. Other calorimetric sensors are pyroelectric sensors, Seebeck effect sensors and thermal conductivity sensors, the last type being used to detect gases at high concentrations [e.g. above the upper explosive limit (UEL) where pellistors cannot be used] [18]. 8.3.3

Diffusion controlled region

Kinetically controlled region

Electronic Conductance Sensors Sensing Process and Signal Transduction The development of electronic conductance sensors traces back to 1953, when Brattain and Bardeen reported that surrounded gases could modulate the conductances of Ge samples [34]. Later, the groups of Heiland [35], Bielanski [36] and Seiyama [37] found sensing effects, mainly with reducible gases such as H2 , CO and carbon–hydrogen compounds, on metal oxides such 8.3.3.1

I

II

III

Temperature / °C Fig. 10 Reaction rate of the catalytic oxidation of reducing gases such H2 and CO versus pellistor temperature. Typical temperature regions are marked for different operating modes of the pellistors.

References see page 1985

1976

8.3 Chemical Sensors Based on Catalytic Reactions

Catalytic reaction

Chemisorption

Porous catalyst

CO

Thin-film heater

NO2

CO2

O2

Grain boundary

Si3N4 membrane

Surface

Metal cluster

Si substrate Fig. 11 Schematic cross-section of an Si-based micromachined pellistor. The temperature-sensitive Si3 N4 membrane has a thickness of 150 nm and a size of 1 × 1 mm.

Electrode e−

as ZnO and SnO2 . These materials are always nonstoichiometric compounds Mn Om−δ (M = Zn, Sn or other metals, m, n = integers) with a small deviation δ from the ideal stoichiometry n : m. From the preparation point-ofview, the value of δ may be adjusted thermodynamically at high temperatures through equilibration with gasphase oxygen or by adding small amounts of other metals with different valence states [38]. The n-type, wide-bandgap semiconductor SnO2 can be considered as ‘‘prototypical’’ metal oxide for gas sensing, exhibiting a gap of about 3.5 eV, which makes stoichiometric samples electrically insulating. The n-type conductance is caused by intrinsic defects, predominantly oxygen vacancies, as well as extrinsic defects, i.e. foreign atoms in the crystal lattice [39]. The simple donor/acceptor space charge layer models that were developed in the 1950s to describe the charge transfer upon chemisorption and catalysis at semiconductor surfaces [40] explain the conductance changes of SnO2 and, correspondingly, of electronic conductance sensors at the qualitative level [41]. Since then, gas-sensing effects on metal oxides have been intensively studied and the acquired knowledge presented in an impressive number of articles in various journals and monographs; a few examples of the latter can be cited [42–46]. An electronic conductance sensor normally comprises a sensitive layer deposited on a substrate, the latter being provided with electrodes for the measurement of the electrical characteristics (Fig. 12). The sensitive layer consists of the metal oxide (such as SnO2 ) and added catalysts (such as Pt, Pd). The device is generally heated by its own heater, which is separated from the sensing layer and the electrodes by an electrical insulator. The elementary reaction steps of gas sensing (such as chemisorption and catalytic conversion) will be transduced into an electrical signal, i.e. a change in conductance, measured by appropriate electrode structures. Typically, one observes a change in the conductance G of the sensor on exposure to one gaseous

Bulk

e−

Substrate Heater

a.c., d.c. Fig. 12 Schematic representation of elementary steps during the detection of molecules (free particles in the gas phase) with electronic conductance sensors. Surface and bulk reactions lead to changes in the overall direct current (DC) or alternating current (AC) conductance, measured between the electrodes, which may contain contributions from the surface, bulk and grain boundaries.

species, which follows the relation G ∝ p i ni

(17)

where pi is the partial pressure of the reducing gas species i in air and ni is a characteristic exponent, lower than 1. The value of ni depends on the metal oxide and the added surface catalysts, the temperature, the gas and also the gas composition. A typical result of gas exposure experiments with an SnO2 thick-film sensor is given in Fig. 13 [47]. Here the relative resistance change is plotted as a function of temperature for different partial pressures of CH4 , CO and C2 H5 OH; in each case exposure is to the pure gas. This figure demonstrates (1) the ‘‘volcano’’ type of sensor responses with broad maxima centered around a gasspecific temperature and (2) the pronounced differences in the sensitivity for the individual gas molecules i (with respect to an interpolation to the same partial pressures). The temperature dependence allows optimum operating conditions (e.g. adjusted by the heating power) to be used to tune for the highest sensitivity of a selected gas. The major problem with all electronic conductance

8.3.3 Electronic Conductance Sensors

1977

G / 10−4 Ω−1

SnO2 thick film

dG = 0

10 40 ppm C2H5OH

G /Gair

30

100 200 300

8000

pCO / ppm

13600

5 300 ppm CO

18700

pH2O / ppm

Fig. 14 Conductance G of a commercial Taguchi sensor as a function of pCO and pH2 O at a constant temperature of 760 K [49].

2000 ppm CH4 1

200

300

400

500

600

700

T / °C Fig. 13 Relative resistance change (ratio of the conductance G in the presence of pi and G in air) for a thick-film SnO2 sensor for different gases i. Adapted from [47].

sensors, the lack of selectivity (called cross-sensitivity), is also obvious from Fig. 13; adding noble metals during the fabrication process of the sensors is a commonly used approach to modify the characteristics of the SnO2 . Formally, the cross-sensitivity can be described in many cases by introducing exponents ni in Eq. (17) that depend on the partial pressure pj of an interfering gas j [48]. In addition to the cross-sensitivity with respect to reducible gases, a problem with electronic conductance sensors, and a challenge to minimize its influence, is the sensitivity towards humidity. Earlier results have demonstrated that the conductance of SnO2 sensors increases when the water partial pressure pH2 O increases, even though the conductance still maintains its ‘‘statefunction’’ character, as displayed in Fig. 14, in the presence of H2 O [49]. The CO sensitivity, on the other hand, decreases on reduction of the humidity in air. The sensitivity of SnO2 towards H2 O has been associated with the formation of hydroxyl groups upon adsorption of H2 O [50–52], whereas molecular water is no longer present at the surface above 200 ◦ C, according to temperature-programmed desorption and infrared spectroscopic data. The experimentally proven increase in conductance in the presence of water vapor has been explained by

electron donation from OH groups, formed by capture of an H atom by a lattice oxygen OO [53]: + − −−  H2 Ogas + SnSn + OO −− − − (SnSn − OH ) − + (OH)+ O +e

(18)

− where (Sn+ Sn − OH ) represents an OH group attached to an Sn surface site. Only the (OH)+ O , bound to an O site, becomes ionized due to a lower electron affinity and consequently injects an electron into the conduction band. The second mechanism, also proposed by Heiland and Kohl [53], involves a reaction between the hydrogen atom and the lattice oxygen and the binding of the resulting hydroxyl group to the Sn atom. The resulting oxygen vacancy will produce, by ionization, the additional electrons: + − −  H2 Ogas + 2SnSn + OO − −− − − 2(SnSn − OH )

+ VO 2+ + 2e−

(19)

A further refinement of the interaction mechanisms has been made by Morrison [54] and Henrich and Cox [55]. Henrich and Cox suggested that the pre-adsorbed oxygen could be displaced by water adsorption; recently, experimental proof for this mechanism was provided by IR measurements [56]. In any of these mechanisms, the particular state of the surface has a major role, because it is considered that steps and surface defects will increase the dissociative adsorption. The surface dopants could also influence these phenomena. Details can be found elsewhere [57]. References see page 1985

1978

8.3 Chemical Sensors Based on Catalytic Reactions

Gas

Product

Volume not accessible to gases

z

Surface band bending qVS

z0 Current flow

Conducting channel

zg z (a)

Energy

x qVS

q∆V S

z zg

(b)

Ec,bulk

Energy

Fig. 15 Schematic representation of a compact sensing layer with geometry and energy band representations; z0 is the thickness of the depleted surface layer, zg is the layer thickness and qVs is the band bending. (a) A partly depleted compact layer (‘‘thicker’’); (b) a completely depleted layer (‘‘thinner’’).

Large grains

Product Gas

z

Energy

x

Current flow

qVs

Eb xg > l D

xg

Small grains

x

2x0

Current flow

Energy

As already shown in Fig. 12, electronic conductance sensors rely on materials and structures of high complexity of the sensing layer with respect to bulk, surfaces, grain boundaries and interfaces with the contacts and the substrate. It is obvious that, depending of the fabrication technique of the metal oxide, a simple distinction can be made between compact layers, produced by thin-film techniques, and porous layers [58], produced by thick-film techniques (dealing with preprocessed powders), RGTO (rheotaxial growth and thermal oxidation) [59] and were fabricated recently by FSP (flame spray pyrolysis) [60]. Our current understanding of sensing processes and of signal transduction will be outlined, applying a semiconductor physics approach combined with chemical kinetics. Because of the strong influence of the morphology of the sensing layer, a distinction is commonly made between compact layers, for which the interaction with gases takes place only at the geometric surface (Fig. 15), and porous layers. The volume of porous layers is also accessible to the gases and the active surface is much higher than the geometric surface (Fig. 16). For compact layers, one may further distinguish between completely or partly depleted layers, depending on the ratio between layer thickness and the Debye length λD of the electrons. For partly depleted layers, when surface reactions do not influence the conduction in the entire layer (zg > z0 ), the conduction process takes place in the bulk region (of thickness zg − z0 , much more conductive that the surface-depleted layer). Formally two resistances are connected in parallel, one influenced by surface reactions and the other not; the conduction is parallel to the surface and this explains the limited sensitivity. Such

q∆VS < kBT flat band condition

qVs

Eb xg < l D

x Extended surface influence

Fig. 16 Schematic representation of a porous sensing layer with geometry and energy band. λD , Debye length; xg , grain size. Adapted from [61].

a case is generally treated as a conductive layer with a reaction-dependent thickness. For the case of completely depleted layers in the absence of reducing gases, it is possible that exposure to reducing gases acts as a ‘‘switch’’ to the partly depleted layer case (due to the injection of additional free charge carriers). It is also possible that exposure to oxidizing gases acts as a switch between partly depleted and completely depleted layer cases. For porous layers, the situation may be complicated further by the presence of necks between grains (Fig. 17), establishing conduction paths through the sensitive layer [61]. They are very much affected by changes in the surface potential in close vicinity to the necks except for small grains and narrow necks. Again, the possible switching role of reducing gases may also occur for porous layers. For small grains and narrow necks, when the mean free path of free charge carriers becomes comparable to the dimensions of the grains, a surface influence on mobility should be taken into consideration. This happens because the number of collisions experienced by the free charge carriers in the bulk of the grain becomes comparable to the number of surface collisions; the latter may be influenced by adsorbed species acting as additional scattering centers (see discussion in Ref. [62]).

8.3.3 Electronic Conductance Sensors

1979

z

x

z

z

qVS Current flow

zn

z0 x

zn - 2z0 Energy

Conduction channel

(a)

zn neck diameter z0 depletion layer z

z qVS

zn

x Energy

q∆VS

(b)

Schematic representation of a porous sensing layer with geometry and surface energy band case with necks between grains; zn is the neck diameter and z0 is the thickness of the depletion layer. (a) The case of only partly depleted necks; (b) large grains where the neck contact is completely depleted. Adapted from [61].

Fig. 17

Figure 18 illustrates the way in which the metal–semiconductor junction, built at electrode-sensitive layer interfaces, influences the overall conduction process. For compact layers they appear as a contact resistance (RC ) in series with the resistance of the SnO2 layer. For partly depleted layers, RC could be dominant and the reactions taking place at the three-phase boundary, electrode/SnO2 /atmosphere, control the sensing properties. In porous layers the influence of RC may be minimized because it will be connected in series with a large number of resistances, typically thousands, which may have comparable values (Rgi in Fig. 18). Transmission line measurements (TLM) performed with thick SnO2 layers exposed to CO and NO2 did not result in values of RC clearly distinguishable from the noise [63], whereas in the case of dense thin films the existence of RC was proved [64]. Again, the relative importance played by different terms may be influenced by the presence of reducing gases because one can expect different effects for grain/grain interfaces compared with electrode/grain interfaces. Modeling the Sensor Response Only for some instances, such as the ‘‘simple’’ case of CO detection in air, has the theoretical modeling of the sensor response, described by relations of the 8.3.3.2

type G ∝ pi ni , been achieved. The modeling is based on the determination of charge carrier concentrations in the depletion layer at the surface for large and small SnO2 grains and crystallites and the conduction process in the sensing layer itself, which also depends on the morphology of the layer [58, 65]. Figure 19 shows the calculated dependence of the conductance G on the CO partial pressure for O2− surface species reacting with CO. The black squares describe a situation corresponding to either (i) compact thin films or (ii) completely depleted small grains where, in both cases, the difference in band bending is lower that the thermal energy (flat band case). The cross-hatched area indicates the range of exponents in the power law between 0.45 and 0.5, which is valid for thin, compact layers (completely depleted but not in flat band condition). The largest variation (simple hatched area) of the exponent between 0.36 and 0.6 corresponds to porous layers with large grains (interconnected by close necks or in point contacts). Role of Contacts The importance of contacts to the overall sensor resistance is due to its electrical contribution (interface between semiconducting sensitive layer and electrode) and its possible chemical influence (e.g. the catalytic activity of the 8.3.3.3

References see page 1985

1980

8.3 Chemical Sensors Based on Catalytic Reactions

Compact layer

zg z0

Energy

EF Metal

Layer

Metal

Rc Rl1 Rc x Rl1 large because of depletion of surface

Porous layer

Energy

EF Energy

EF

qVC

qVS

Eb

qVC

Eb Rc

qVS

Eb

qVC

Rgi Rc

Rl

Rc

Rl2

Rc

x

Rl2 low due to lack of surface influence R1 || R2; R2 R ≈ R2 R2 = Rc + Rl2

x

R = Rc + Σ Rgi

Resistance can be dominated by electrode contact properties

Resistance is less influenced by electrode contact properties

Schematic representation of compact and porous sensing layers with geometry and energetic bands, which shows the possible influence of electrode-sensing layers contacts. RC is the resistance of the electrode–SnO2 contact, Rl1 is the resistance of the depleted region of the compact layer, Rl2 is the resistance of the bulk region of the compact layer, R1 is the equivalent series resistance of Rl1 and RC , R2 is the equivalent series resistance of Rl2 and RC , Rgi is the average intergrain resistance in the case of porous layer, Eb is the minimum of the conduction band in the bulk, qV s is the band bending associated with surface phenomena on the layer and qV C also contains the band bending induced at the electrode–SnO2 contact. Fig. 18

Signal / arb. units

70

pCO0.6

60

pCO0.5

50

pCO0.45 pCO0.36

40

pCO0.33

30 20 10 0 0

200

400

600

800

1000

Concentration / ppm Fig. 19 Summarized calculated power law dependence of CO interaction with doubly ionized oxygen (O2− ).

contact material in the region close to the contacts). It has been shown by semiconductor approaches, applied to the interface between the semiconducting sensitive layer and the metallic electrode, that gas exposure in most cases does not affect the sensing properties, as the value of the contact resistance is established during the preparation. This

value, in contrast to its change, might be important in the overall resistance of the sensor and could even decrease the sensor response by being a ‘‘dead’’ series element, especially for the case of compact films where zg > z0 . More important are chemical sensitization and catalytic effects of the electrodes due to their chemical nature (Pt or Au). Several effects could be taken into consideration, as follows. Surface species, which can be more easily adsorbed on the electrode metal, may diffuse rapidly to the three-phase boundary where they can react with the partner adsorbed on the metal oxide sensitive layer. Hence the increased diffusion will lead to a higher catalytic conversion rate, which will be monitored by the electrical readout. Another effect is the increased production (‘‘catalysis’’) of reaction partners by the metal electrode material (for Pt but not for Au). This can happen by, e.g., cleavage of hydrocarbons into more active radicals. Hence the reaction partners can diffuse to the three-phase boundary (electrode/sensing layer/gas phase) and consequently this region becomes ‘‘more active’’ in gas detection. In contrast to the above-mentioned effects that will enhance the sensor response, it is possible that an increased catalytic reaction on the electrode material, with direct desorption from there, will lead to gas consumption

8.3.3 Electronic Conductance Sensors

that is not monitored by the electrical readout. In consequence, this gas consumption may lead to an overall lowering of the analyte concentration (depending on the given setup) and may thus even lead to a lowering of the sensor signal.

1981

toxic gas monitoring. The second-placed material is WO3 while the other industrially used materials are practically not investigated any longer. Belows, short overviews on the most relevant gas-sensing metal oxides are given. Tin Oxide (SnO2 ) Even if tin oxide is the best known material for gas sensors and an impressive number of publications are available (see, e.g., Refs. [72–74] and the monographs cited above), there is still a clear need for a deeper understanding of both gas sensing and signal transduction. This is mostly due to the fact that most of the basic knowledge was obtained by using model systems loosely linked to real-world sensors [75] (single crystals, RT and UHV for modeling polycrystalline layers operated at elevated temperatures and in humid air). This is especially true due to the lack of solutions for some of the most annoying problems encountered when one uses SnO2 -based gas sensors, namely the lack of selectivity and the influence of water vapor. Figure 20 presents results obtained with pure, Pt and Pd surface-activated tin oxide layers screen printed over alumina substrates provided with electrodes and heaters (the operating temperature for the sensors was 300 ◦ C; the represented sensor signal is the relative change in resistance on gas 8.3.3.4.1

Materials Nowadays, many companies offer semiconducting gas sensors, such as Figaro, FIS, CityTech, AppliedSensors and MICS [66–70]. Also, since they found an additional mass market application in the automotive area (cabin air quality control), one can say that they successfully spread out from their traditional field and became relevant for mass market applications [71]. Most of the companies use SnO2 -based sensing materials; alternative materials are WO3 , Cr2−x Tix O3 (with x = 0.05–0.4) and Ga2 O3 . Table 1 presents an overview of the metal oxidelinked contributions at the 10th International Meeting on Chemical Sensors, Tsukuba (Japan), 11–14 July 2004, together with the gases they were used to detect. One can easily observe that even after more than 30 years of commercial use, SnO2 is still the most investigated gas-sensing material and CO, NO2 and volatile organic compounds (VOCs) the main target gases; this indicates an additional application field in which metal oxide-based gas sensors are considered as good candidates, namely 8.3.3.4

References see page 1985

Tab. 1 Statistical data derived from the contributions presented at the 10th International Meeting on Chemical Sensors, Tsukuba (Japan), 11–14 July 2004 (VOC = volatile organic compound, CWA = chemical warfare agent)

ZrO2 TiO2 SnO2 WO3 MoO3 ZnO In2 O3 CuO Fe2 O3 SrTiO3 LaMeFeO3 BaTiO3 e-nose MOX LaOCI BiMeVOx V2 O5 NdCoO3 CeO2 Ga2 O3 SmFeO3 NiO Co3 O4 Total

O2

Ethanol

H2 S

CO

H2

VOCs

Methane

NOx

H2 O

NH3

CO2

cwa

O3

SO2

6 1 3 1 1 0 0 0 0 3 4 0 0 0 0 1 0 0 2 0 0 0 0 22

0 2 7 1 0 0 0 0 0 0 0 0 0 1 0 0 0 1 1 0 0 0 0 13

0 0 2 0 0 0 0 0 0 0 0 0 1 0 0 0 0 0 0 0 0 0 0 3

0 5 21 4 2 2 0 0 3 2 0 2 0 0 0 0 0 1 0 0 0 1 1 44

0 2 8 2 1 1 1 0 1 0 0 1 0 0 0 0 0 0 0 1 0 1 1 20

0 2 14 5 4 1 1 1 0 1 1 0 1 3 0 0 0 1 0 0 1 0 0 36

0 0 4 0 0 0 0 0 0 0 0 0 1 0 0 0 0 0 0 0 0 0 0 5

3 0 4 6 2 0 1 0 2 1 1 0 0 1 0 0 0 0 0 0 1 0 0 22

0 0 5 0 0 0 0 1 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 6

0 0 2 4 1 0 0 0 0 0 0 0 0 0 0 0 0 1 0 0 0 0 0 8

0 0 0 0 0 0 1 0 0 0 0 0 0 0 1 0 0 0 0 0 0 0 0 2

0 1 3 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 4

0 0 0 0 0 0 0 0 0 0 0 0 0 1 0 0 0 0 0 0 1 0 0 2

2 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 2

Amines Total 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 1 0 0 0 0 0 0 1

11 13 73 23 11 4 4 2 6 7 6 3 3 6 1 1 1 4 3 1 3 2 2 190

1982

8.3 Chemical Sensors Based on Catalytic Reactions

0% r.h.

Pt

1

50% r.h.

for CO for NO2

Pd 102

Pd

101

+ + + + undoped ++

100 0.01

+ + 0.1

+++ ++ 1

+

+++

Sensor signal S

Sensor signal S

10

3

50% r.h. 0.1

270°C 10

100

1000

CO concentration / ppm

0.01 0

Sensor signal as a function of the CO concentration for Pd-doped sensors (0.2 wt.%, open symbols), Pt-doped sensors (0.2 wt.%, full symbols) and undoped sensors (cross in center) for relative humidity 0% (squares) and 50% (circles).

exposure, Rair /RCO ). Both advantages and disadvantages of SnO2 are obvious: they are extremely sensitive, offering reasonable signals even at 100 ppb CO; one can tune their response by using additives; the presence of humidity has a marked effect. To the advantages one can add the fact that it is possible to obtain very stable sensors with the present state of the art knowledge [76]. 8.3.3.4.2 Tungsten Oxide (WO3 ) As suggested by the statistics and its use in commercial gas sensors, tungsten oxide (WO3 ) is a promising material. Gas-sensing effects recorded as changes in the resistivity of WO3 have long been known [77], but research on this material really took-off only in the last 10 years. The reason is related to the need to find additional gas-sensing materials that will provide the additional performance that it is not possible obtain with SnO2 alone. The material proved to be extremely suitable for the detection of oxidizing gases (NO2 and O3 ) with little or no cross-sensitivity to reducing gases, e.g. CO. This is illustrated in Fig. 21 for pure WO3 layers screen printed over alumina substrates provided with electrodes and heaters identical with those used for the SnO2 sensors presented in Fig. 20 (the operating temperature for the sensors was 340 ◦ C; the represented sensor signal is the relative change in the resistance upon gas exposure, Rair /RCO ). The level of basic understanding for gas sensing is far more limited than in the case of SnO2 and there are no comprehensive reviews for WO3 as a gas-sensing material; it is known, however, that the material is non-stoichiometric – the semiconducting effect is linked, as also for SnO2 , to the donor levels in the bandgap associated with the oxygen vacancies. It is also known that the crystal chemistry of WO3 is extremely complex, with two different crystalline phases in the range of

2

4

6

CO/NO2 concentration/ppm

Fig. 20

Gas-sensing performance of WO3 thick-film layers on exposure to CO and NO2 . Apart from oxidizing gases, sensing of VOCs and NH3 are the most targeted applications for WO3 -based gas sensors.

Fig. 21

temperatures where the gas sensors are normally operated (transition temperature from monoclinic to orthorhombic around 330 ◦ C) [78]. 8.3.3.4.3 Chromium Titanium Oxide (Cr2−x Tix O3 ) In the quest to find gas-sensing materials without the high crosssensitivity to water vapor, Moseley and Williams identified as a candidate the compound chromium titanium oxide (Cr2−x Tix O3 , CTO) [79]. It was found that the effect of humidity on both baseline and sensitivity is reduced compared with SnO2 and that the material shows high sensitivity to NH3 and H2 (CH4 , NO2 , CO and H2 O can also be detected). This material was used for the development of commercial gas sensors by Capteur, which is now part of City Technology [68]. In contrast to SnO2 and WO3 , which are both n-type semiconductors, CrO shows p-type conductivity due to chromium vacancies [80]. 8.3.3.4.4 Gallium Oxide (Ga2 O3 ) The influence of water vapor on the gas-sensing properties of metal oxides is known to decrease when the operation temperature increases. The problem with such an approach to minimize the cross-sensitivity to humidity is that, on the one hand, the sensitivity for most gases also decreases and, on the other, most metal oxides tend to become unstable at high temperatures (also the electrical conduction process is complicated by the contribution of ions). This is the reason why this strategy did not work for most applications with the notable exception of natural gas sensing. In order to eliminate the downside of the approach, materials able to operate at high temperatures without losing the

8.3.4 Solid-State Electrochemical Sensors

sensitivity towards target gases were investigated. Hightemperature Ga2 O3 gas sensors were first reported by Fleischer [81]. Ga2 O3 is an n-type semiconductor (due to the presence of oxygen vacancies) that has electronic conduction even at 600–700 ◦ C and shows good sensitivity to many gases, e.g. CO and CH4 . Commercial sensors based on this material are manufacture by Steinel Solutions [82]. 8.3.4

Type I

where pBref and pB denote the partial pressure of the gaseous species under detection at the left- and righthand side, respectively, kB is Boltzmann’s constant, T the absolute temperature, z the charge number of the mobile ions B and e the elementary charge. According to Eq. (20), the sensor output voltage is proportional to the logarithm of partial pressure pB if the reference gas partial pressure is kept constant. It is obvious that this principle is restricted mainly to oxygen if the reference phase consists of ambient air and, strictly, an oxygen reference electrode is used. This electrode is formed at the three-phase boundary O2 /Pt/ZrO2 . However, the detection of gases B (e.g. Cl2 ) with a solid electrolyte AB with A+ as the mobile ion is possible if AB is formed in a surface reaction between the gas B and A+ so that the gas equilibrates with a component of the electrolyte that is different from the predominantly mobile species. This results in a type II sensor. An example is

B−

Gas B

(AB)

V

Type II

Solid-State Electrochemical Sensors

Solid-state electrochemical sensors are based on ionic conduction, a phenomenon which is well known for a variety of inorganic and polymeric compounds. Ceramictype solid-state electrochemical sensors, often called solid electrolyte sensors, have been widely used for decades as potentiometric zirconia sensors for oxygen under hightemperature operation conditions [83, 84]. This detection principle, which has been already illustrated in Fig. 3, yields a cell voltage which is determined thermodynamically by the concentrations of the same gas (for the lambda probe O2 ) which forms the mobile ions (here O2− ) in the solid electrolyte. The situation is shown in Fig. 22 in a general scheme where the gas B yields the mobile ions B− of the B-sublattice of the ion conductor. The notation type I refers to a commonly used classification of potentiometric sensors into three categories, introduced by Weppner [85]. Since the voltage, generated at the electronic junctions of the metal/solid electrolyte interfaces of this cell, follows Nernst’s law, a so-called Nernstian-type sensor signal results:   kB T pB V = (20) ln ze pBref

Reference B

1983

Reference A

A+

Gas B

(AB)

V AP

Type III

Reference A

A+ (AB)

Or

Reference A

Gas C AC

V AP A+ (AB)

Gas C AC

V Fig. 22 Three types of solid-state ionic devices in a schematic survey. AP denotes an auxiliary phase with elements AC.

the detection of Cl2 with the Ag+ ion conductor AgCl. As for type I sensors, only electrons are exchanged across the interface between the electrode and the electrolyte and thermodynamic equilibrium is established between the gas and the immobile species of the electrolyte. Again, the chemical potential of electrons is measured at the interface between the electrolyte and the electrode and hence it is not relevant which type of ions are mobile within the ionic conductor. The relationship between the gas partial pressures and the open-circuit voltage can be derived from Nernst’s law and the Gibbs–Duhem equation. For details, the reader is referred to Ref. [86]. Finally, for the type III sensor the gas is being equilibrated with the solid electrolyte through an auxiliary phase AP on top of the electrolyte. The auxiliary phase should preferably be a mixed ionic–electronic conductor and must contain both the mobile species of the solid electrolyte and the species to be detected [86]. This approach extends the applicability of potentiometric References see page 1985

1984

8.3 Chemical Sensors Based on Catalytic Reactions

sensors and makes it possible to detect gases even if related ions are not mobile or present in the solid electrolyte. In the two bottom representations in Fig. 22, the auxiliary phases consist of a compound AC, allowing molecules C to be detected, since AC is formed by a reaction of C with ions A+ of the bulk component AB. An example is the detection of C = CO2 with Nasicon (Na1+x Zr2 P3−x Six O12 , where x may vary from 0 to 3) as Na+ ion conductor [87] and Na2 CO3 –BaCO3 mixtures as intermediate phase AC. Again, the relationship between the cell voltage (sensor response) and the partial pressure pC of the gas C can be derived from Nernst’s law and the Gibbs–Duhem equation. This concept apparently overcomes the limitations of type I or II sensors but, due to the complex structure, a more advanced sensor technology is required by which stable ionic junctions can be fabricated. These junctions exchange both electrons and ions across the interface between solid ionic conductors and the contacts. In potentiometric sensors of any of these types, the equilibrium potential drop at the electrode–electrolyte interfaces is related to the activity or partial pressure of species under detection. The sensor output voltage is proportional to the logarithm of the partial pressure of the gas. Another concept of gas sensing through solid-state electrochemistry is based on a non-Nernstian or mixedpotential response, deviating from the equilibrium potential. Such an effect occurs if the chemical equilibrium is not fully established at the electrode, which then results in an ‘‘anomalous’’ potential, given by electrode kinetics. For example, a modified type I cell (such as the lambda probe that monitors oxygen O2 by electrochemically coupling it to the mobile ions O2− of the solid electrolyte) exhibits a Nernstian potential, determined by the partial pressure of O2 , if there is no interference with other gases that may change the partial pressure pO2 at the electrode. In other words, the cell probes the actual value of pO2 in the gas phase, equal to the equilibrium value. If the value of pO2 is significantly higher than the partial pressures of the interfering gases (e.g. pNO if a single reaction NO + 1/2O2 → NO2 to give gaseous molecules of NO2 takes place) the difference between the actual and the equilibrium partial pressures of O2 is so small that the deviation from the Nernstian potential is below its noise level. Under these conditions, the potential of the cell is controlled by the equilibrium O2 + 4e−  2O2− . However, if the interfering equilibrium reaction is not fully established at the electrode, e.g. due to a low temperature or a catalytically inactive electrode coating, then competitive electrochemical reactions, involving NO and NO2 , may exhibit an influence on the potential. In this case the potential is controlled by the reaction rates of the cathodic processes (O2 + 4e− → 2O2− and NO2 + 2e− → NO + O2− ) and the anodic processes

(2O2− → O2 + 4e− and NO + O2− → NO2 + 2e− ). At the mixed potential the anodic and cathodic reactions have the same rate. This occurs when the anodic (ia ) and cathodic (ic ) currents are equal, making the overall net current zero, ia + ic = 0. The concept may be used to measure NOx in the exhaust gas of an internal combustion engine. The output potential of a mixed-potential sensor becomes larger with respect to the equilibrium potential of O2 + 4e− → 2O2− in the presence of NO2 and smaller for NO; the opposite effects cause problems, however, if both NOx species are present. For mixed-potential sensors, the selectivity is primarily to be controlled by tuning the electrode composition, e.g. by covering one of the catalytically very active Pt electrodes of a YSZ-based sensor with a porous oxide such as CdO and SnO2 [88], LaMnO3 [89], La0.8 Sr0.2 CrO3 [90], Ta-doped TiO2 [91] or binary alloys of Pt with Au [92], to produce weakly active electrodes. The concept offers novel opportunities for CO, H2 , hydrocarbon and NOx sensing, especially in connection with the control of pollutant emission and improvement of efficiency of internal combustion engines [84]. However, the operational range, under which mixed-potential sensors can be employed for gas detection, is limited and, as in the case of NOx reverse sensitivities may occur. Amperometric solid electrolyte sensors are usually operated under diffusion-controlled limit current conditions. Figure 23 illustrates schematically the design of such a sensor, based on zirconia. The voltage, applied across the two electrodes (approximately 300 mV), results in a reduction of O2 at the cathode to give O2− ions, which are electrochemically transferred through the cell to the anode and oxidized to O2 ; this process is called ‘‘electrochemical pumping’’. According to Faraday’s law, the flux through the small opening jO2 = I /4F is given by the current I ; F denotes Faraday’s constant. It decreases the partial pressure of oxygen within the inner cavity, hence the oxygen activity at the cathode is very small with respect to the oxygen activity in the exhaust gas for a sufficiently large pumping voltage. The difference results in a driving force and hence to oxygen diffusion from the surrounding atmosphere (e.g. the exhaust gas) through the orifice into the cell cavity. The O2− flux is limited, governed by a linear diffusion law so that the sensor signal is proportional to the oxygen partial pressure (Fig. 24), in contrast to the potentiometric lambda probe. Amperometric devices can be used to determine high partial pressures of oxygen. A recent approach towards novel sensors in combustion control and monitoring applications is based on the use of arrays of amperometric solid-state sensors. An example is shown in Fig. 25. Here, a conventional oxygen electrode (Pt for reduction of O2 ) is combined with an NO

References

8.3.5

Gas

Conclusions

Diffusion barrier e−

Pt − +

O2−

Y−ZrO2

VC = const

IC Fig. 23

Sketch of an amperometric solid-state electrolyte sensor.

Control through electrode kinetics

Diffusion control

3 % O2

Ic

2 % O2 1 % O2

0 Working potential

Vc

Current IC as a function of voltage VC , applied across the amperometric solid-state electrolyte sensor for different oxygen partial pressures.

Fig. 24

O2− NO, O2, CO, HC

O2

O2 O2−

O2− O2 O2−

IC1 VC1 Reduction of oxygen IC1 ~po2

O2− NO

O2−

N2 HC CO2' H2O O2−

The chemical sensors presented in this overview utilize characteristic physicochemical properties of particle– solid interactions to sense molecules in the gas-phase, including temperature effects and electronic and ionic conduction effects. Catalysts are typically used with metal oxide sensors, such as electronic conductance and calorimetric devices, to enhance the sensitivity and reduce the required operating temperature. Special situations occur in solid-state electrochemical sensors in which catalytic reactions affect potential-determining processes at three-phase boundaries. References

4 % O2

0

1985

ZrO2

O2−

IC2 VC2 Reduction of NO, IC2 ~p NO Oxidation of HC, IC2 ~p HC

Fig. 25 Principle of an array of amperometric solid-state electrolyte sensors with different electrodes.

electrode (Pt alloys for NO reduction). The replacement of the Pt of the oxygen electrode by Au converts its high catalytic activity for hydrocarbon oxidation into an inactive contact but maintains the oxygen reduction capability so that hydrocarbons can be electrochemically oxidized (and hence measured separately) at the second Pt electrode, operating at a suitable potential [93].

1. J. R. Stetter, W. R. Penrose, S. Yao, J. Electrochem. Soc. 2003, 150, S11. 2. W. G¨opel and K. D. Schierbaum, in Sensors – A Comprehensive Survey, Vol. 2: Chemical and Biochemical Sensors (Part I), W. G¨opel, J. Hesse, J. N. Zemel (Eds.), VCH, Weinheim, 1991, pp. 2. 3. W. G¨opel, T. A Jones, T. Seiyama, J. N. Zemel, in Sensors – A Comprehensive Survey, Vol. 2: Chemical and Biochemical Sensors (Part I), W. G¨opel, J. Hesse, J. N. Zemel (Eds.), VCH, Weinheim, 1991. 4. K. Persaud, G. Dodd, Nature 1982, 299, 352. 5. W. G¨opel, K. D. Schierbaum, in Sensors – A Comprehensive Survey, Vol. 2: Chemical and Biochemical Sensors (Part I), W. G¨opel, J. Hesse, J. N. Zemel (Eds.), VCH, Weinheim, 1991. 6. J. Chou, Harzardous Gas Monitors, McGraw-Hill, New York, 2000, pp. 1. 7. Z. Cao, W. J. Buttner, J. R. Stetter, Electroanalysis 1992, 4, 253. 8. D. J. Fray, Mater. Sci. Technol. 2000, 16, 237. 9. E. T. Turkdogan, R. J. Fruehan, Can. Metall. Q. 1972, 11, 371. 10. M. Iwase, Y. Waseda, High Temp. Mater. Proc. 1986, 7, 123. 11. N. Taguchi, British Patent 1 280 809, 1972. US Patent 3 631 436, 1971. 12. I. Lundstr¨om, S. Shivaraman, C. Svensson, L. Lundkvist, Appl. Phys. 1975, 46, 3876. 13. S. Middelhoek, S. A. Audet, Silicon Sensors, Academic Press. London, 1989. 14. M. Burgmair, M. Zimmer, I. Eisele, Sens. Actuators B 2003, 93, 271. 15. P. Bergveld, Thesis, University of Twente, 1973. 16. N. F. de Rooij, Thesis, University of Twente, 1978. 17. L. J. Bousse, Thesis, University of Twente, 1982. 18. T. A. Jones, P. Walsh, in Sensors – A Comprehensive Survey, Vol. 2: Chemical and Biochemical Sensors (Part I), W. G¨opel, J. Hesse, J. N. Zemel (Eds.), VCH, Weinheim, 1991. 19. L. C. Clark, US Patent 2 913 386, 1959. 20. G. Sauerbrey, Z. Phys. 1959, 155, 206. 21. O. Wolfbeis (Series Ed.), Springer Series on Chemical Sensors and Biosensors, Methods and Applications, Springer-Verlag, Berlin. 22. D. Diamond, Principles of Chemical and Biological Sensors, Wiley, New York, 1998. 23. C. Di Natale, A. D’Amico (Eds.), Sensors and Microsystems, World Scientific, Singapore, 1996.

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24. J. W. Gardner, Microsensors: Principles and Applications, Wiley, Chichester, 1994. 25. W. G¨opel, T. A. Jones, M. Kleitz, I. Lundstrom, T. Seiyama (Eds.), Chemical and Biochemical Sensors (Part I), in Sensors, A Comprehensive Survey, Vol. 2, W. G¨opel, J. Hesse, J. N. Zemel (Eds.), Wiley-VCH, Weinheim, 1991. 26. W. G¨opel, T. A. Jones, M. Kleitz, I. Lundstrom, T. Seiyama (Eds.), Chemical and Biochemical Sensors (Part II), in Sensors, A Comprehensive Survey, Vol. 3, W. G¨opel, J. Hesse, J. N. Zemel (Eds.), Wiley-VCH, Weinheim, 1991. 27. R. Baker, Min. Eng. (London) 1969, 643. 28. R. Baker, British Patent 892 530, 1962. 29. S. J. Gentry, T. H. Jones, Sens. Actuators 1969, 10, 141. 30. J. Riegels, Fortschr. Ber. VDI-Reihe 1989, 8(194). 31. J. W. Gardner, S. M. Lee, P. N. Bartlett, S. Guerin, D. Briand, N. F. de Rooij, presented at Transducers ’01. Eurosensors XV, 11th International Conference on Solid-State Sensors and Actuators, Munich, 10–14 June 2001. 32. P. N. Bartlett, S. Guerin, Anal. Chem. 2003, 75, 126. ´ am, A. Morrisey, Cs. D¨ucs˜o, 33. P. F¨urjes, Zs. Vizv´ary, M. Ad´ I. Bars´ony, Sens. Actuators A 2002, 3276, 1. 34. W. Brattain, J. Bardeen, Surface Properties of Germanium, Bell Telephone System Technical Publications, Monograph 2086, Bell Telephone System, 1953, Murrey Hill, New Jersey, p. 1. 35. G. Heiland, Z. Phys. 1954, 138, 459. 36. A. Bielanski, J. Deren, J. Haber, Nature 1957, 179, 668. 37. T. Seiyama, A. Kato, K. Fujiishi, M. Nagatani, Anal. Chem. 1962, 34, 1502. 38. O. T Sørensen, in Nonstoichiometric Oxides, O. T Sørensen (Ed.), Academic Press, New York, 1981. 39. Z. M. Jarzebski, J. P. Marton, J. Electrochem. Soc. 1976, 132, 229C. 40. K. Hauffe, in Semiconductor Surface Physics, R. H. Kingston (Ed.), University of Pennsylvania Press, Philadelphia, PA, 1957, p. 259. 41. W. G¨opel, Prog. Surf. Sci. 1985, 20, 9. 42. P. Moseley, B. Tofield, Solid State Gas Sensors, Adam Hilger, Bristol, 1987, 245 pp. 43. M. J. Madou, S. R. Morrison, Chemical Sensing with Solid State Devices, Academic Press, Boston, 1989; 556 pp. 44. P. T. Moseley, J. O. W. Norris, D. E. Williams, Techniques and Mechanisms in Gas Sensing, Adam Hilger, Bristol, 1991, 408 pp. 45. W. G¨opel, J. Hesse, J. N. Zemel (Eds.), Sensors: A Comprehensive Survey, Vols. 1–3, VCH, Weinheim, 1991. 46. G. Sberveglieri (Ed.), Gas Sensors, Kluwer, Dordrecht, 1992, 420 pp. 47. B. Licznerski, Bull. Pol. Acad. Sci. Tech. Sci. 2004, 52, 37. 48. K. D. Schierbaum, U. Weimar, W. G¨opel, Sens. Actuators B 1990, 2, 71. 49. K. D. Schierbaum, U. Weimar, R. Kowalkowski, W. G¨opel, Sens. Actuators B 1991, 3, 205. 50. K. Morishige, S. Kittaka, T. Morimoto, Bull. Chem. Soc. Jpn. 1980, 53, 2128. 51. A. Guest, PhD Thesis, University of Nottingham, 1985. 52. E. W. Thornton, P. G. Harrison, J. Chem. Soc., Faraday Trans. 1975, 71, 461. 53. G. Heiland, D. Kohl, in Chemical Sensor Technology, Vol. 1, T. Seiyama (Ed.), Kodansha, Tokyo, Ch. 2, 1988, pp. 15–38. 54. S. R. Morrison, The Chemical Physics of Surfaces, 2nd Ed., Plenum Press, New York, 1990. 55. V. E. Henrich, P. A. Cox, The Surface Science of Metal Oxides, Cambridge University Press, Cambridge, 1994.

56. D. Koziej, N. Bˆarsan, U. Weimar, J. Szuber, K. Shimanoe, N. Yamazoe, Chem. Phys. Lett. 2005, 410, 321. 57. M. Egashira, M. Nakashima, S. Kawasumi, J. Phys. Chem. 1981, 85, 4125. 58. N. Bˆarsan, U. Weimar, J. Electroceram. 2001, 7, 143. 59. G. Sberveglieri, Sens. Actuators B 1992, 6, 239. 60. L. M¨adler, A. Roessler, S. E. Pratsinis, T. Sahm, A. Gurlo, N. Barsan, U. Weimar, Sens. Actuators B, 2006, 114, 283. 61. M. Schweizer-Berberich, Thesis, Universit¨at T¨ubingen, 1998. 62. N. Bˆarsan, Sens. Actuators B 1994, 17, 241. 63. M. Bauer, N. Bˆarsan, K. Ingrisch, A. Zeppenfeld, I. Denk, B. Schuman, U. Weimar, W. G¨opel, in Proceedings of the 11th European Microelectronic Conference, Venice, May 14–16, 1997. 64. U. Hoefer, K. Steiner, E. Wagner, Sens. Actuators B 1995, 26–27, 59. 65. A. Rothschild, Y. Komem, J. Appl. Phys. 2004, 95, 6374. 66. http://www.figarosens.com. 67. http://www.fisinc.co.jp. 68. http://www.citytech.com. 69. http://www.appliedsensors.com. 70. http://www.microchem.com. 71. J. Marek, H.-P. Trah, Y. Suzuki, I. Yokomori (Eds.), Sensors for Automotive Technology, VCH, Weinheim, 2003, 562 pp. 72. K. Ihokura, J. Watson, Stannic Oxide Gas Sensors, Principles and Applications, CRC Press, Boca Raton, FL, 1994, 187 pp. 73. N. Bˆarsan, M. Schweizer-Berberich, W. G¨opel, Fresenius’ J. Anal. Chem. 1999, 365, 287. 74. D. Kohl, J. Phys. D: Appl. Phys. 2001, 34, R125. 75. N. Bˆarsan, U. Weimar, J. Phys.: Condens. Matter 2003, 15, R813. 76. N. Bˆarsan, J. R. Stetter, M. Findlay, W. G¨opel, Anal. Chem. 1999, 71, 2512. 77. P. J. Shaver, Appl. Phys. Lett. 1967, 11, 255. 78. Landolt–B¨ornstein: Numerical Data and Functional Relationship in Science and Technology, Vol. 17, Semiconductors, Subvol. g, Physics of Non-Tetrahedrally Bonded Binary Compounds II, Springer-Verlag, Berlin, 1982, pp. 166–200, 446, 594. 79. P. T. Moseley, D. E. Williams, Sens. Actuators B 1990, 113, 1. 80. D. Williams, Sens. Actuators B 1999, 57, 1. 81. M. Fleischer, Sens. Actuators B 1998, 6–7, 257. 82. http://www.steinel.ch. 83. M. Kleitz, E. Siebert, P. Fabry, in Sensors – A Comprehensive Survey, Vol. 2: Chemical and Biochemical Sensors (Part I), W. G¨opel, J. Hesse, J. N. Zemel (Eds.), VCH, Weinheim, 1991. 84. N. Docquier, S. Candel, Prog. Energy Combust. Control 2002, 28, 107. 85. W. Weppner, in Proceedings of the 2nd International Meeting on Chemical Sensors, J.-L. Ancouturier, J.-S. Cauhap´e, M. Destriau, P. Hagenmuller, C. Lucat, F. Menil, J. Porter, J. Salardenne, (Eds.), Bordeaux, 1986. 86. E. D. Tsagarakis, Thesis, Universit¨at Kiel, 2004. 87. J. B. Goodenough, H. Y.-P. Hong, J. A. Kafalas, Mater. Res. Bull. 1976, 11, 203. 88. N. Miura, T. Raisen, N. Yamazoe, Sens. Actuators B 1998, 47, 84. 89. R. Sorita, T. Kawano, Sens. Actuators B 1997, 40, 29. 90. R. Mukundan, E. L. Brosha, F. H. Garzon, J. Electrochem. Soc. 2003, 150, H279. 91. E. Traversa, M. L. DiVona, S. Licoccia, M. Sacerdoti, M. C. Carotta, M. Gallana, G. Martinelli, J. Sol–Gel Sci. Technol. 2000, 19, 193. 92. H. Okamato, H. Obayashi, T. Hudo, Solid State Ionics 1980, 1, 319. 93. S. I. Somov, G. Reinhardt, U. Guth, W. G¨opel, Sens. Actuators B 2000, 65, 68.

8.4.2 Non-Conventional Solvents

8.4

Heterogeneous Catalysis in Non-Conventional Solvents .. Roger Glaser∗

8.4.1

Introduction and Scope

1987

part, the most intensely studied types of conversions in these reaction media are introduced. Selected examples are presented to illustrate the potential and limitations of heterogeneous catalysis in non-conventional solvents. Complete coverage of the literature in the field is not attempted. The cases discussed here are restricted to conversions over solid catalysts. The utilization of non-conventional solvents for multiphase catalysis with organometallic complex catalysts [5, 6], although a special case of heterogeneous catalysis in a rigorous sense, will not be discussed in detail here. Phase-transfer catalysis [7] and catalysis using biocatalysts in non-conventional solvents also offer attractive opportunities [8–11]. These will, however, not be explicitly treated here either. Nonconventional solvents, in particular supercritical fluids, have also been utilized for the preparation of catalytically active solids [12–14]. Such applications will also not be covered in this chapter.

Solvents play an important role in numerous catalytic processes, particularly in the production of bulk, fine and specialty chemicals [1–3]. The most abundantly used solvents are liquid organic compounds. Organic solvents contribute ca. 85% of the non-aqueous mass in pharmaceutical production processes with a recovery rate currently in the range 50–80% [1]. In most instances, the solvent serves to contact reactants with each other and with the catalyst of a given conversion. However, solvent molecules may also interact directly with the reaction intermediates or the catalytically active site and, thus, have a strong influence on reaction rate and selectivity. The selection of a reaction medium for a catalytic process is guided not only by its bulk or molecular solvent properties, but also by safety considerations and the suitability for product and/or catalyst separation after the reaction step. The separation, recycling and disposal of a process solvent are often cost- and energy-intensive. Especially in the case of toxic, hazardous or ecologically problematic solvents, it is necessary to avoid product contamination by or emissions of the solvent. For these reasons, the demand for alternative solvent systems is continuously becoming more urgent from both economic and environmental points of view. Although, certainly, ‘‘the best solvent is no solvent’’ [4], the utilization of a solvent cannot always be avoided. It may, as will be shown below, through the tunability of the solvent properties, even offer the advantage of selectivity or rate enhancement with respect to a solvent-free conversion. This chapter focuses on non-conventional (also referred to as ‘‘neoteric’’ [3]) as opposed to conventional liquid or gaseous reaction media for heterogeneously catalyzed conversions. The non-conventional solvents treated here include supercritical fluids (SCFs), ionic liquids (ILs) and gas-expanded liquids (GXLs). Other recently studied options for reaction media such as polyethylene glycols (PEGs), perfluorohydrocarbons (‘‘fluorous liquids’’) and thermoregulated solvent systems have only rarely been exploited in heterogeneously catalyzed reactions and are, therefore, only briefly mentioned. In the first section, a short introduction of the general properties of non-conventional reaction media with respect to heterogeneous catalysis is given. In the second

Supercritical Fluids (SCFs) A substance is present as a supercritical fluid at conditions above its critical point, i.e. above the critical temperature Tc and the critical pressure pc , but below the pressure for condensation into the solid state. At the critical point, the gas and liquid phases in equilibrium at subcritical conditions become indistinguishable. In the supercritical state, a separation into two phases by an isothermal pressure increase or an isobaric temperature decrease cannot occur. The interest in supercritical fluids as solvents and reaction media is largely based on the fact that their behavior and properties can be continuously changed from liquid- to gas-like without the occurrence of phase transitions (see below). For mixtures, the term ‘‘supercritical’’ is not as clearly defined as for pure components. Although a critical point also exists for mixtures, there may be more than one phase present at conditions above a mixture critical point. The critical temperatures and pressures of several pure substances are shown in Fig. 1. For critical data for a larger number of substances, see Ref. [15]. Many substances have a critical temperature in a range that is of practical interest for organic chemical syntheses (0–150 ◦ C) and a critical pressure below 10.0 MPa. For these substances, supercritical conditions are readily accessible for experimental laboratory studies or technical applications. For water and ammonia with more polar and hydrogen-bonding molecules, however, significantly higher temperatures and pressures are needed to reach the supercritical state. By far the most frequently studied

∗

References see page 2004

Corresponding author.

8.4.2

Non-Conventional Solvents 8.4.2.1

8.4 Heterogeneous Catalysis in Non-Conventional Solvents Properties of supercritical fluids compared with those of liquids and gases

Tab. 1

22 20 12

H2O NH3

Gasa

Property 10 CO2 N2O

8

Xe C H CHF3 2 6

6 N2

4 2 0

C3H8 C4H10 C5H12

SF6

H2

CH4 −100

0

100

200

C6H14 C7H16

a At

400

Critical temperature Tc / °C

Critical temperature Tc and pressure pc for selected pure substances.

Fig. 1

and used supercritical solvents are carbon dioxide (CO2 ), water and lower hydrocarbons such as ethane and propane. Its non-toxic and non-flammable character, the chemical inertness under most synthesis conditions and the low cost, abundant availability and low environmental impact often make CO2 the first choice as a solvent or reaction medium [16]. The interest in supercritical fluids as solvents and reaction media was strongly promoted by the application of supercritical CO2 for natural products extraction such as decaffeination of coffee since the 1960s [17]. Other substances are occasionally applied, mainly for research purposes, as highly polar (CHF3 , N2 O) and highly polarizable (Xe, SF6 ) supercritical media. 8.4.2.1.1 Properties of SCFs As mentioned above, the properties of supercritical fluids are intermediate between those of typical liquids and gases (Table 1). In the vicinity

0.6–1.6 0.2–3.0 0.0002–0.002

of the critical point, i.e. at 1.05–1.20Tc and 0.9–2.0pc , the so-called ‘‘near-critical region’’, the properties of supercritical fluids may vary strongly with only small changes in pressure or temperature. This is exemplified for the density ρ, the dynamic viscosity η and the selfdiffusion coefficient D11 of CO2 in Fig. 2. The resulting tunability of the physicochemical properties in the nearcritical region is the basis for numerous applications of supercritical solvents. Although the properties of supercritical fluids can be continuously varied, some of them are closer to those of liquids (such as density) and some closer to those of gases (such as viscosity). Immediately at the critical point, some physical properties show an anomalous behavior. For instance, the enthalpy of vaporization vanishes, while partial molar volumes and the isothermal compressibility approach infinity. As a result of the high compressibility, strong density fluctuations and spatial inhomogeneities [18] occur. These lead to the well-known phenomenon of critical opalescence. In supercritical solutions near the critical point, the density of the solvent molecules around a solute molecule may be much higher than the average density in the bulk. This ‘‘local density enhancement’’ (also referred to as ‘‘clustering’’ [19] or 0.10

0.08

h

0.6 2.4

0.4

1.8 0.2

D11 r 0

10

20

30

0.8 40

D11 r 104 / g cm−1 s−1

r

0.8

r /g cm−3

0.2–0.5 0.01–0.03 0.07

pGas = 0.1 MPa.

1.0

0.0

Liquid

ρ/g cm−3 0.0006–0.002 η/mPa s 0.01–0.3 D/10−6 m2 s−1 10–40

He −200

Supercritical fluid

0.06

0.04

h / mPa s

Critical pressure pc / MPa

1988

0.02

0.00

p /MPa

Dependence of density ρ, dynamic viscosity η and self-diffusion coefficient D11 of CO2 on pressure at 40 ◦ C, a temperature in the vicinity of the critical point (Tc = 31.1 ◦ C, pc = 7.38 MPa).

Fig. 2

8.4.2 Non-Conventional Solvents

‘‘molecular charisma’’ [20]) may also occur between different dissolved molecules and may lead to higher reaction rates than those expected from the average bulk concentration [21]. It should be expected that a local density enhancement can also occur at solid surfaces. However, corresponding effects on the rate or the selectivity of heterogeneously catalyzed reactions near the critical point have, so far, not been unambiguously proven. Despite these anomalies at the critical point, it should be noted that the attractiveness of supercritical fluids generally does not evolve from their fundamentally different properties compared with those of conventional liquids and gases. It rather lies in the unique combination of typical liquid- and gas-like properties and the strong dependence of these properties on pressure and temperature in the vicinity of the critical point. Due to the liquid-like densities, the solubility of higher molecular weight organics can be much higher in supercritical fluids than in gases. For instance, 10 mol% of naphthalene can be dissolved in supercritical CO2 at 60 ◦ C and 25.0 MPa [22], whereas only 0.2 mol% is present in a gas at 0.1 MPa. Since the solubility in supercritical fluids increases exponentially with density, it can be sensitively tuned by pressure at a given near-critical temperature. This often allows a significant facilitation of the downstream separation or even fractionation of dissolved reactants or products after a chemical reaction simply by pressure variation. Nevertheless, the solvent power of supercritical fluids is mostly not as high as in conventional liquid organic solvents. Especially polar or hydrogen-bonding compounds such as polyalcohols, polyacrylates or even salts are only sparingly soluble in supercritical CO2 , the polarity of which can be compared with that of liquid n-hexane [23], toluene or tetrachloromethane [24]. To improve the solubility of functionalized and/or polar solutes, a co-solvent is often added to a lowpolarity supercritical fluid as a ‘‘modifier’’ or ‘‘entrainer’’. Typically, volatile compounds such as methanol, ethanol, acetone or alkylamines are used that can undergo specific dipolar interactions and/or hydrogen bonding with the solute molecules. Co-solvent concentrations are usually in the range of 1–5 mol% [25]. The nature and the amount of a co-solvent provide an additional tool for tuning the solvent properties of supercritical fluids. As a more polar solvent than CO2 , supercritical water is an attractive medium for chemical reactions [26]. Liquid water changes its properties considerably upon approaching the critical point. Under autogenous pressure, the dielectric constant of liquid water at 250 ◦ C is similar to that of acetone at ambient conditions. This allows the dissolution of both organic and ionic compounds in significant concentrations in liquid, near-critical water (NCW,

1989

250–350 ◦ C, 4.0–9.0 MPa). In addition, the dissociation constant of liquid water, KW , reaches a maximum at 250 ◦ C, where it is three orders of magnitude higher than at ambient conditions. The increased concentration of H+ and OH− ions in near-critical water can be utilized to carry out acid- or base-catalyzed conversions in the absence of added catalysts [26, 27]. Above the critical point, however, the ionization constant of water decreases strongly. Supercritical water is, therefore, not a suitable medium for chemical reactions involving ionic species. In addition to the ability to solubilize higher molecular weight solutes in significant concentrations, supercritical fluids are mostly miscible with permanent gases such as oxygen, hydrogen or carbon monoxide over a wide range of compositions. Thus, chemical conversions may be carried out in a single, supercritical phase whereas, under conventional conditions, a multiphase reaction system would be involved. The possibility of operating in a single homogeneous phase rather than in a heterogeneous multiphase reaction system is particularly advantageous for reactions that are frequently limited by mass transfer of a gas into a liquid reaction phase, e.g. hydrogenations, oxidations, hydroformylations or Fischer–Tropsch synthesis (see Section 8.4.3.1). Moreover, mass transfer restrictions on the reaction rate may be alleviated in supercritical media due to their lower viscosity and, thus, higher diffusivity than in conventional liquids (see Table 1). Also, the heat transfer properties of supercritical fluids are favorable for chemical reactions. For instance, the heat capacity of CO2 goes through a maximum in the near-critical region. This allows one to significantly reduce the adiabatic temperature rise in strongly exothermic reactions such as the combustion of hydrogen [28]. To capitalize fully on the favorable solvent properties of supercritical fluids, it is imperative to know the phase behavior at higher pressures and at temperatures relevant for an envisaged application. Apparently, the phase behavior becomes increasingly complex with the number of components in a mixture. The effects of temperature and pressure on the nature and number of the phases present may, therefore, be much more pronounced than in a single-component system. Additionally, the phase behavior, including the conditions of the critical point and the single-phase region, may be drastically shifted as the composition of the mixture changes during a chemical reaction [29–31]. Thorough and reliable studies of the phase behavior in view of the nature and fractions of all components comprising a reaction mixture are, therefore, indispensable before carrying out reactions in supercritical fluids. A treatment of the phase behavior of binary systems based on the classification by Scott and van References see page 2004

8.4 Heterogeneous Catalysis in Non-Conventional Solvents

Konyenburg [32] can be found in a recent review [33]. This review also covers the experimental observation of highpressure phase equilibria and their relation to catalytic reactions. Opportunities for SCFs in Heterogeneous Catalysis From their unique properties, the advantages of utilizing supercritical fluids for heterogeneously catalyzed conversions are obvious. The most prominent ones can be summarized as follows: 8.4.2.1.2

• Rate enhancements may be achieved due to the favorable phase behavior, i.e. higher reactant concentrations, due to improved heat and mass transfer within the supercritical reaction phase or by eliminating interphase mass transfer resistances. In addition, the reaction rate may be influenced by the kinetic pressure effect: according to transition-state theory, the mole fraction-based rate constant kx is largely determined by the activation volume v = , i.e. the difference in the partial molar volumes of the transition state and the reactants [Eq. (1)]: 
v = ∂ ln kx =− (1) ∂p RT T 
vR ∂ ln Kx =− (2) ∂p RT T Since the partial molar volumes diverge at the critical point, the activation volumes and the rate changes for homogeneous reactions at near-critical conditions may be several orders of magnitude higher than in conventional liquids [34]. Likewise, the position of the chemical equilibrium as expressed by Kx may be strongly pressure dependent near the critical point [Eq. (2) with vR as the reaction volume, i.e. the difference in the partial molar volumes of the products and the reactants]. Similarly pronounced pressure effects on the rate of heterogeneously catalyzed conversions have, however, not been reported, possibly due to the lack of pronounced solvation at solid surfaces.

100

Selectivity/%

1990

80

kB

B

kC

C

A Assumptions: kB /kC = 1

60

∆ν ≠B = − 40 cm3 mol−1 ∆ν ≠C = + 30 cm3 mol−1

40 20 0

0

20

40

100

Selectivity for two products (B, C) formed in irreversible parallel reactions from the same reactant (A) as a function of pressure (after Ref. [35]).

• The tunability of the solvent properties of supercritical fluids by pressure, temperature or co-solvents offers the opportunity to control the catalytic selectivity. This can emerge from the tuning of heat or mass transfer, e.g. through the transport of a reaction intermediate away from the catalyst surface before it can undergo a consecutive reaction. Selectivity tuning may also be based on the kinetic pressure effect, e.g. when two parallel reactions involve opposing activation volumes (Fig. 3) [35]. • The catalyst lifetime may be significantly enhanced in the case of deactivation by poisoning, e.g. through deposition of higher molecular weight compounds such as coke on the catalyst surface. These deposits may be efficiently removed during the reaction (‘‘in situ extraction’’) due to the liquid-like solubility and the gas-like transport properties of supercritical fluids. Improved activity maintenance may be particularly achieved for meso- and macroporous catalysts (Fig. 4) [36]. While the removal of high-boiling deposits in gases is desorption limited (volatility-driven) and in liquids diffusion limited (solubility-driven), the balance between desorption and diffusion is more favorable in supercritical fluids. Thus, an optimum density exists in the supercritical

Supercritical phase

Liquid phase

Volatility-driven

Desorption/diffusionbalance

Solubility-driven

Diffusion

80

Fig. 3

Gas phase

Desorption

60

p /MPa

Catalyst pore Higher molecular weight deposit

Balance of desorption and diffusion for removal of higher molecular weight deposits from a catalyst pore in gaseous, liquid and supercritical reaction phases (after Ref. [37]).

Fig. 4

8.4.2 Non-Conventional Solvents

region at which the solubility and the rate of transport of the deposits out of the catalyst pores are maximized [37]. • Due to the rate enhancement and the higher densities of supercritical vs. gaseous reaction phases, reactors may be designed with smaller volumes to achieve the same performance. This obvious advantage for continuously operated reactors offers significant potential for process intensification by applying supercritical fluids. Despite the obvious advantages of supercritical fluids as media for catalytic reactions, there are some serious drawbacks that have often prevented their industrial application. Most of all, it is the high capital costs associated with the proper equipment for safely operating at the higher pressures required to reach and maintain supercritical process conditions. Ionic Liquids (ILs) Ionic liquids have been intensively studied as solvents for chemical reactions since the early 1980s [38–40]. All these solvents are composed of ions only and are present in the liquid state at temperatures below 100 ◦ C, in many cases also at room temperature. In other words, ionic liquids are salts with a low melting point. Typical constituents of ILs are shown in Fig. 5 and Table 2. The large number of different cation–anion combinations gives rise to a broad variety of ILs with most diverse properties. Moreover, the properties of these solvents can be sensitively tuned by variation of the chemical structure of the cation or the anion. For instance, this can be achieved by changing the alkyl chain length in the ammonium, phosphonium, imidazolium or pyridinium cations (Fig. 5). Due to this possibility of adjusting the physical and chemical properties to the requirements of a chemical reaction, ILs are sometimes referred to as ‘‘designer solvents’’ [41]. 8.4.2.2

8.4.2.2.1 Properties of ILs A comprehensive overview of the physicochemical properties of ILs is given in +

+ R

N

N

R′

R′ R

N

[BF4]− R

[AlCl4]−

R +

+ R′ N R

R

P R

[PF6]−

R

[(CF3SO2)2N]− [(C2F5)3PF3]− [CF3SO3]− [n-C8H17OSO3]

−

Typical cations and anions as constituents of ionic liquids (after Ref. [42]).

1991

Abbreviations for typical cations and anions as constituents of ionic liquids

Tab. 2

Abbreviation

Full name

Cations [bdmim] [bmim] [bmpyr] [n-Bu4 N] [Cn mim] [C4 mpip] [Cn py] [emim] [mim] [N8881 ] [P66614 ]

1-Butyl-2,3-dimethylimidazolium 1-Butyl-3-methylimidazolium 1-Butyl-3-methylpyrrolidinium Tetrabutylammonium 1-n-Alkyl-3-methylimidazolium N-Butyl-N-methylpiperidinium 1-n-Alkylpyridinium 1-Ethyl-3-methylimidazolium 3-Methylimidazolium Methyltrioctylammonium Trihexyltetradecylphosphonium

Anions [BF4 ] Br Cl [EtOSO3 ] [NTf2 ] [OTf] [PF6 ]

Tetrafluoroborate Bromide Chloride Ethylsulfate Bis(trifluoromethanesulfonyl)imide Trifluoromethanesulfonate (triflate) Hexafluorophosphate

Ref. [42]. Most importantly, ILs have no measurable vapor pressure. Besides the obvious advantage over volatile organic solvents regarding negligible emissions to the atmosphere, this allows efficient product removal from ionic liquids by distillation. ILs are typically stable up to temperatures of ca. 200 ◦ C, and some may be stable even above 400 ◦ C. This often results in a much broader temperature range of the liquid state for ILs compared with water or conventional organic solvents. Product separationw from ionic liquid solutions can, in many instances, also be accomplished by liquid–liquid extraction. This relies on the high densities of ionic liquids in the range 1.12–2.40 g cm−3 , which allows phase separation from most common liquid organic solvents. Several ILs are also immiscible with water over a broad range of mixture compositions. They may, therefore, provide a polar, yet non-aqueous phase for separations with or from organic liquids. Nevertheless, most ILs are not extremely polar. Their solvent polarity is fairly comparable to that of alcohols with short to medium alkyl chain lengths, e.g. methanol and butanol. As may be expected, however, the polarity of ILs may vary considerably depending on the chemical nature of the cation and the anion. Consequently, interesting combinations of solubilities for both organic and inorganic compounds may be found in ILs that are often uncommon in conventional organic solvents.

Fig. 5

References see page 2004

1992

8.4 Heterogeneous Catalysis in Non-Conventional Solvents

However, the influence of polarity or polarizability on the solvent properties of ILs is much less pronounced than that of the hydrogen bonding abilities of the ions comprising the IL. Moreover, in the case of non- or weakly coordinating anions, the polar nature of ILs is combined with a low nucleophilicity of the solvent. ILs are, therefore, interesting media for highly electrophilic catalysts or reagents. Opportunities for ILs in Heterogeneous Catalysis As a drawback for heterogeneously catalyzed conversions, ILs posses an inherently high viscosity, typically between 10 and 500 mPa s. These values are much larger than those for water (0.89 mPa s) or ethylene glycol (16.1 mPa s) and are of the same order of magnitude as that of glycerol (934 mPa s), all as liquids at room temperature. Since, in addition, gases such as hydrogen, oxygen and carbon monoxide are scarcely soluble in ILs, heterogeneously catalyzed conversions involving these gases may readily be limited by mass-transfer processes. ILs are, therefore, not the first choice as media for heterogeneously catalyzed oxidations, hydrogenations or hydroformylations in multiphase gas–liquid–solid systems. On the other hand, carbon dioxide is highly soluble in many ILs. Mole fractions of carbon dioxide in ILs may reach 70% or higher at room temperature and elevated pressures. Likewise, lower alkanes such as ethane and propane are well soluble in ILs. As a consequence, carbon dioxide and lower alkanes in their supercritical state may be applied as solvents to transport reactants into and/or products out of ionic liquid reaction phases. This approach 8.4.2.2.2

is particularly promising for multiphase catalysis with transition metal complexes ‘‘immobilized’’ in the IL [43]. Another benefit for the application of ILs in heterogeneous catalysis is their low surface tension and the correspondingly low capillary forces. Wetting of solid surfaces is, therefore, facilitated. Moreover, ILs possess a reasonably high electrical conductivity in their temperature range of existence in the liquid state. Hence, applications in heterogeneous electrocatalysis appear to be an attractive goal. Finally, the increasing commercial availability of a broad range of different ILs, also on a larger scale, will foster the consideration of these solvents for industrial catalytic conversions. A limitation may, however, be their higher costs in comparison with most conventional organic liquid solvents. A further limitation was attributed to the lack of knowledge of the toxicity of ILs. First systematic studies indicate, however, that ILs are not more toxic than common organic liquids such as methanol, acetone and acetonitrile [44]. Gas-Expanded Liquids (GXLs) Gas-expanded liquids (GXLs) are closely related to conventional organic liquids on the one hand and to supercritical fluids on the other. GXLs are, by definition, multicomponent mixtures. At subcritical pressures, a gas such as carbon dioxide may dissolve in an organic liquid in considerable quantities. As shown for CO2 in Fig. 6a, the fraction of the dissolved gas increases with pressure depending on the temperature and the nature of the organic liquid. The increasing fraction of the dissolved CO2 (for some solvents over 80 wt.% below 10.0 MPa) leads to a volume increase (‘‘expansion’’ or ‘‘swelling’’) 8.4.2.3

12

T = 40 °C 1200

10

900

V/V0 / %

pCO2 / MPa

8

6

4

(a)

600

300

2

0 0.0

n-Decane T = 30 − 70 °C Cyclohexane Benzene Ethanol 2-Propanol Methanol 1,4-Dioxane Acetonitrile Acetone

0.2

0.4

0.6

XCO2

0.8

0 0.0

1.0 (b)

0.2

0.4

0.6

0.8

1.0

XCO2

Influence of pressure on the weight fraction of CO2 , xCO2 , in organic solvents (a) and volume increase of CO2 -expanded liquids, V/V0 , as a function of CO2 weight fraction (b).

Fig. 6

8.4.2 Non-Conventional Solvents

and a density and viscosity decrease of the liquid phase. This volume expansion with respect to the liquid in absence of the gas, V /V0 , depends, for most organic solvents and temperatures, on the fraction of the dissolved gas only (Fig. 6b). The volume may readily increase by more than 1000% of its initial value. Similarly to SCFs, the properties of GXLs can, thus, be tuned to a large extent by pressure and temperature (see below). In most cases, carbon dioxide is used to create GXLs due to its high solubility in many liquid organic solvents. The high fraction of CO2 in both the liquid and gas phases may shift the flammability or explosion limits away from those of the pure organic solvent. Operation in CO2 -expanded liquids may, therefore, be considered to be intrinsically safe. Moreover, the interaction of CO2 with alcohols or water in GXLs can lead to the formation of carbonic acids that may act as catalysts for chemical conversions [45, 46]. Properties of GXLs An apparent advantage of GXLs over supercritical fluids is that they typically exist at lower, subcritical pressures (p < 10 MPa). The occasionally high pressures needed for obtaining a singlephase supercritical mixture for a catalytic conversion, often above 20.0 MPa, can therefore be avoided. Moreover, the higher density of the GXLs makes them generally more powerful solvents for higher molecular weight compounds than supercritical fluids. GXLs have, consequently, been utilized as solvents for the extraction of high-boiling solutes from matrices such as sediment [47, 48]. By variation of the pressure and/or the composition of the gas phase, the solubility of solutes in the GXL phase may be changed to afford dissolution or precipitation of individual compounds. The latter effect, also referred to as ‘‘gas antisolvent precipitation’’ [49], can be used for the separation of products or soluble catalysts from reaction mixtures. Also, the solubility of permanent gases in GXLs is largely increased with respect to the conventional liquid organic solvents. For instance, the solubility of oxygen in CO2 -expanded acetonitrile at 30 ◦ C is higher by two orders of magnitude than in the absence of CO2 and is comparable to that of liquid CO2 [50, 51]. Hence GXLs are particularly attractive as solvents for catalytic hydrogenations, oxidations or hydroformylations where mass transfer of gases into an organic liquid phase often limits conversion and/or selectivity (cf. Section 8.4.3.3). Due to the higher content of organic solvent in GXLs compared with co-solvent-modifiedw supercritical fluids, a much wider range of polarities, dielectric constants and hydrogen-bonding behavior is accessible in this type of solvent system. In addition to the choice of organic solvent, the solvent polarity of GXLs may depend strongly on the pressure and the mole fraction of the 8.4.2.3.1

1993

gas dissolved in the liquid phase. For CO2 -expanded solvents, the solvatochromic parameter ET (30), a measure of solvent polarity [3], was measured for GXLs based on several organic liquids [52]. Generally, the solvent polarity decreases with temperature and with pressure, i.e. with increasing CO2 fraction in the liquid phase. This is particularly pronounced for polar solvents such as acetone: Thus, CO2 -saturated, i.e. expanded, acetone at 35 ◦ C and a pressure of ca. 6.0 MPa has the same polarity as liquid cyclohexane at ambient pressure [52]. For high CO2 mole fractions towards unity, the polarity assumes the same value similar to that of pure liquid CO2 . Opportunities for GXLs in Heterogeneous Catalysis As a result of the viscosity decrease by gas expansion, the diffusion within GXLs may be significantly faster than in conventional liquids. For instance, the selfdiffusion coefficient of acetone in a liquid mixture with CO2 may be about a factor of three higher than in the single-component liquid [53]. Evidently, the diffusivity enhancement is directly related to the CO2 content in the GXL phase [54]. In combination with the higher solubility for permanent gases, this makes GXLs attractive media for heterogeneously catalyzed conversions that are conventionally carried out in multiphase reaction systems. If the presence of an organic solvent can be tolerated or is even required, GXLs may be a particularly interesting alternative to supercritical fluids as non-conventional solvents. 8.4.2.3.2

Other Non-Conventional Solvents Other non-conventional solvents that have recently been studied include low-melting polyethylene glycols (PEGs) [55], perfluorohydrocarbons (‘‘fluorous liquids’’) [56] and thermoregulated solvents [1]. Since they are often immiscible with other organic or aqueous phases, these solvents are attractive for multiphase systems as, for instance, in multiphase catalysis with molecular transition metal complexes [57]. In some cases, the reaction mixtures may be biphasic before and after reaction, whereas they are single-phase under reaction conditions. The transition from a twoto a single-phase system may be achieved by a temperature increase for thermoregulated solvents or by pressurization with CO2 for fluorous liquids [27]. Evidently, interphase transport limitations are avoided and product and/or catalyst separation are facilitated in these systems. However, most of these other non-conventional solvents have, so far, hardly found any application in heterogeneous catalysis. 8.4.2.4

References see page 2004

1994

8.4 Heterogeneous Catalysis in Non-Conventional Solvents

8.4.3

Heterogeneously Catalyzed Conversions in Non-Conventional Solvents Conversions in Supercritical Fluids The number of heterogeneously catalyzed reactions that have been studied in supercritical fluids is very large and still growing. They cover a most diverse variety of chemical conversions over a broad spectrum of solid catalysts. An overview of the most intensively studied types of reactions and pertinent examples are given in Table 3. Several reviews have summarized the state-of-the art in the field of heterogeneous catalysis in supercritical fluids [33, 36, 113–117]. For an exhaustive treatment of heterogeneously catalyzed conversions in supercritical fluids, the papers by Baiker and coworkers [33, 113] can be particularly recommended. One of them focuses especially on in situ spectroscopy and on monitoring of high-pressure phase behavior [33]. The similarly recommendable reviews by Subramaniam et al. are devoted to the activity enhancement of porous catalysts in supercritical fluids [36] and to the rational design of reactors for catalytic conversions in dense-phase CO2 [115]. Savage gives an overview of heterogeneous catalysis in water [117]. In addition to the strong research interest in the area, heterogeneous catalysis with supercritical fluids is applied in several industrial processes [15]. The classical Haber–Bosch synthesis of ammonia and the early methanol synthesis over zinc–chromium oxide catalysts are just two well-known examples. In recent years, the range of industrial applications of supercritical fluids in heterogeneous catalysis has continuously expanded. Some of the most prominent cases will be mentioned in the following sections. 8.4.3.1

8.4.3.1.1 Hydrogenations and Dehydrogenations Solidcatalyzed hydrogenations belong to the most extensively studied conversions in supercritical fluids. Conventional processes often involve three-phase reactors where gaseous hydrogen is contacted with a slurry of the catalyst in a liquid containing the dissolved or pure reactant. Since hydrogenations are mostly fast, the reaction rate is frequently limited by the low hydrogen solubility in the liquid and by the hydrogen transfer to the active sites in the catalyst pores. The advantages of an environmentally benign supercritical, single-phase reaction system with high hydrogen solubility, no interphase transport restrictions and efficient transfer of the exothermic heat of reaction away from the catalyst are evident. Moreover, the enhanced rates and the dense reaction phase can overcome the large reactor volumes often encountered in conventional multiphase hydrogenation processes.

A large variety of unsaturated organic compounds are accessible to heterogeneously catalyzed hydrogenation in supercritical fluids. This ranges from unsaturated hydrocarbons, fats and oils (see below) through aromatic and aliphatic aldehydes and ketones, epoxides and nitriles to substituted nitroaromatics or even polymers (cf. Table 3). A short compilation of rate data for catalytic hydrogenations in supercritical fluids can be found in Ref. [118]. The influence of the high-pressure phase behavior on the semihydrogenation of phenylacetylene to styrene over an amorphous Pd81 Si19 catalyst was systematically studied by Tschan et al. [59]. In the near-critical region, the conversion at 55 ◦ C increased strongly with pressure. It reached a maximum close to the transition of the reaction mixture to supercritical conditions around 13.0 MPa. This was attributed to the improved mass transfer of hydrogen to the catalyst under single-phase supercritical conditions. A high selectivity for styrene was reached which decreased with pressure from 100% at subcritical to ca. 80% at supercritical conditions due to overhydrogenation. Similar results were also found for the enantioselective hydrogenation of ethyl pyruvate to ethyl (R)-lactate over a cinchonidine-modified Pt/Al2 O3 catalyst in supercritical ethane [67, 68]. In the supercritical single-phase region at high pressure, it is possible to tune the reaction conditions, e.g. pressure, temperature and composition, independently. This advantage was used by Hitzler et al. [61] to optimize the selectivity for a desired product within a series of successive hydrogenations. The Pd-catalyzed selective conversion of anisole to methylcyclohexanol or methylcyclohexane in supercritical CO2 and that of nitrobenzene to aniline in supercritical propane are two examples. For the ethyl pyruvate hydrogenation in supercritical ethane, an increase of the enantiomeric excess (ee) was found on increasing the catalyst:substrate ratio. Interestingly, the same change causes a decrease of the ee in liquid solvents where mass transfer of hydrogen becomes rate limiting at higher catalyst:substrate ratios [67]. When the hydrogenation of ethyl pyruvate was carried out in supercritical CO2 instead of ethane, deactivation of the Pt/Al2 O3 catalyst was observed [67]. This was explained by platinum poisoning by CO formed by the reverse water-gas shift (RWGS) reaction of the supercritical CO2 with hydrogen. However, no deactivation of supported Pd catalysts was reported by Hitzler et al. [61] for the hydrogenation of several different unsaturated organic substrates in supercritical CO2 . A detailed study of Arunajatesan et al. [58] showed that the deactivation of a Pd/C catalyst in the hydrogenation of cyclohexene could be avoided when the concentration of peroxides in the feed was reduced from an initial value of 180 ppm to 0.96 [125]. In a detailed study, Elbashir and Roberts [77] related the extent of this deviation to the physical properties of the reaction mixture at near-critical conditions: the enhanced solubility of higher molecular weight products results in more vacant sites at the catalyst surface for readsorption and further incorporation of the middle-distillate alkenes into the chain-growth reaction. 8.4.3.1.3 Hydroformylation Hydroformylation, i.e. the conversion of alkenes with carbon monoxide and hydrogen to aldehydes (also referred to as ‘‘oxo synthesis’’), is conventionally catalyzed by soluble transition metal complexes, mainly of rhodium or cobalt, often in biphasic reaction systems. New concepts for hydroformylation are based on the combination of an ionic liquid for immobilization of a transition metal complex with a separate reaction phase containing the reactants and the products in a supercritical fluid (see Section 8.4.2.2). As for hydrogenations, the opportunity to supply all reactants (and products) in a single phase makes supercritical fluids interesting media for hydroformylation reactions. Early reports by Abraham’s group [79] showed that the hydroformylation of propene over a supported rhodium catalyst in supercritical CO2 was unselective due to the strong adsorption of the aldehyde product. Catalysts prepared by impregnation of rhodium complexes on

1997

modified silica were subject to leaching of the active metal into the supercritical reaction phase [80]. As shown by Meehan et al. [85] for the hydroformylation of 1-octene, a stable catalyst can be obtained by covalent bonding of a rhodium complex to a silica support. Following this strategy, other rhodium and platinum complexes were anchored to different supports to give stable and selective catalysts for the hydroformylation of 1-hexene [81]. Their activity in supercritical CO2 was lower than that that for homogeneous catalysis in liquid solution, but higher than that for heterogeneous gas-phase catalysis [82]. An increase in the n-aldehyde yield with pressure could by be explained by the larger activation volume of hydroformylation compared with that of isomerization [83]. Interestingly, the highest activity was found for a catalyst supported on an ordered mesoporous MCM-41-type material [84]. For this catalyst, the yield ratio of linear and branched aldehydes was 8.8 (100 ◦ C, 24.0 MPa). With zeolite MCM-20 as the support, the selectivity for the linear aldehyde was even higher (yield ratio 15.8), but the activity was lower due to diffusion limitations in the zeolitic micropores. 8.4.3.1.4 Oxidations The advantages of supercritical fluids as reaction media for solid-catalyzed oxidations are similar to those for hydrogenations. Above all, the elimination of oxygen mass transfer resistances through operation in a single phase for all reactants and products, the efficient removal of the exothermic reaction heat away from the catalyst and the potential to desorb partially oxidized products prior to an unwanted consecutive or even total oxidation may lead to enhanced rates and/or selectivities. Furthermore, supercritical CO2 or water, due to their non-flammability and chemical inertness, may shift the explosion limits for reactants and, thus, contribute to an inherently improved process safety. Despite the obvious advantages, heterogeneously catalyzed oxidations in supercritical fluids became a focus of research only in recent years. The most intensively studied reactions are partial oxidations of hydrocarbons and alcohols in supercritical CO2 (cf. Table 3). Catalytic total oxidation in supercritical CO2 has received less attention. However, total oxidation in supercritical water was investigated in more detail (see below). In the majority of the studies, oxygen was used as the oxidizing agent. As for hydrogenations, the high-pressure phase behavior also plays an important role for selective oxidations. The effect of pressure on the rate of benzyl alcohol oxidation with molecular oxygen over a Pd/Al2 O3 catalyst in dense CO2 is shown in Fig. 8 [94]. The rate increases by a factor of ca. 2 between 14.0 and 15.0 MPa, where a References see page 2004

8.4 Heterogeneous Catalysis in Non-Conventional Solvents

2000

100

1600

95

1200

90

800

85

400

8

10

12

14

16

Selectivity /%

Reaction rate/mol mol−1 h−1

1998

80 18

p / MPa Reaction rate and benzaldehyde selectivity in the continuous conversion of benzyl alcohol over 0.5 wt.% Pd/Al2 O3 as a function of pressure (T = 80 ◦ C, W/Falcohol = 20 g h mol−1 ; nO2 /nalcohol = 0.5) (after Ref. [94]).

Fig. 8

transition from a two- into a single-phase system occurs. At the same time, the selectivity for benzaldehyde decreases only slightly (Fig. 8). The beneficial effect of supercritical reaction conditions on the oxidation rate is particularly evident from the turnover frequency at 15.0 MPa (1585 h−1 ), which is almost 80 times higher than that for the conversion in liquid toluene (20 h−1 ). In situ studies by extended X-ray absorption fine structure (EXAFS) showed that the Pd catalyst deactivates at high oxygen concentrations due to overoxidation of the noble metal surface [93]. This deactivation can occur even at temperatures as low as 40 ◦ C [92]. The high-pressure phase behavior is utilized in an innovative concept for the epoxidation of propene with hydrogen peroxide over the bifunctional microporous catalyst Pd,Pt/TS-1 [90, 126]. In a single, supercritical CO2 phase, the hydrogen peroxide can first be generated from the elements on the noble metal (see also Section 8.4.3.1.1) and, then, be converted with propene to propene oxide and water on titanium-substituted silicalite-1 (TS-1). Currently, propene conversions do not reach values above 10%, whereas the selectivity to propene oxide is higher than 90% [90]. Also, the more challenging, heterogeneously catalyzed oxidation of alkanes by molecular oxygen was attempted in supercritical CO2 . With propane as the reactant, the yields of oxyfunctionalization products did not exceed 20% [87, 88]. Nevertheless, the conversion and oxygenate selectivity increased on changing from gaseous to supercritical conditions. This selectivity increase was explained by an accelerated desorption of the intermediate oxygenates from the catalyst surface under supercritical conditions. Low conversions of 135 ◦ C (Tc for isobutane), unwanted side-reactions, above all cracking and oligomerization, are favored over the desired alkylation. As shown by Clark and Subramaniam [100], the critical temperature of an isobutane–1-butene mixture (molar ratio 9:1) could be decreased to 40 ◦ C by addition of a fivefold molar excess of CO2 . Whereas a USY zeolite deactivates in the conversion of the undiluted feed at 140 ◦ C, it is stable under supercritical conditions at 50 ◦ C in the presence of CO2 over 2 days on-stream (Fig. 9). However, the 1-butene conversion was 20% and the alkylate fraction in the product was only 5–10%.

8.4.3 Heterogeneously Catalyzed Conversions in Non-Conventional Solvents

Time-on-stream / h 4

8

16

20

140°C, 6.1 MPa 140°C, 5.1 MPa 50°C, 15.5 MPa

0.4

yalkylate /y C5+

12

Supercritical

0.3

Near-critical

0.2

Supercritical with CO2

0.1 0.0 0

1

2

3

4

bulkier 2,7-diisopropylnaphthalene [103]. Since, however, the isomer distribution was almost identical with that for the liquid-phase reaction at the same temperature, a direct influence of the supercritical reaction medium on the shape-selective conversion inside the zeolite pores was rendered unlikely. Such an influence was, however, discussed to explain, e.g. shape-selectivity effects on the formation of coke precursors extracted from the acidic zeolite catalysts H-ZSM-5 and H-mordenite during the conversion of supercritical ethylbenzene [131] or the threefold higher rate for the hydroxyalkylation of phenol over a zeolite H-mordenite in supercritical CO2 than in liquid toluene [105].

5

m1-butene /mcat / g g−1 Alkylate fraction in the C5+ products from the alkylation of isobutane with 1-butene over zeolite H-USY at near-critical (140 ◦ C, 5.1 MPa) and supercritical conditions in the absence (140 ◦ C, 6.1 MPa) and presence of CO2 (50 ◦ C, 15.5 MPa) (after Ref. [100]).

Fig. 9

Moreover, the product was strongly diluted with CO2 . At higher butene conversion, an H-USY zeolite is subject to deactivation even under supercritical conditions [129]. A cyclic regeneration of the zeolite catalyst in supercritical isobutane allows to remove most of the coke and to obtain a stable butene conversion of >92% for more than 210 h on-stream [130]. Nevertheless, the alkylate composition does not comply with industrial standards. Several studies have been directed towards the influence of the phase behavior on the alkylation of aromatics with light alkenes. For instance, Shi et al. [102] systematically varied the phases present during benzene alkylation with ethene over the acidic zeolite H-Beta by changing the pressure and feed composition in the temperature range 240–260 ◦ C. They found a maximum for the reaction rate slightly above the critical pressure and a decreasing rate with increasing pressure. An additional, less pronounced rate maximum was observed for a two-phase, liquid–gas feed mixture. These results again show that applying single-phase, supercritical conditions alone does not necessarily result in high reaction rates, and that sensitive tuning is required for rate optimization. Shape-selectivity effects have been observed for a number of alkylations of aromatics over acidic zeolites or related microporous materials at supercritical reaction conditions [103–105, 131, 132]. The question of whether the physical state of the reaction medium surrounding the catalyst has a direct influence on the processes occurring inside zeolitic micropores is discussed controversially. For instance, the shape-selective conversion of naphthalene with 2-propanol over zeolite H-mordenite in supercritical CO2 led to a high yield ratio of the slender 2,6- vs. the

8.4.3.1.6 Isomerizations The double-bond isomerization of 1-hexene to cis- and trans-2- and -3-hexene has been used as an example to study the effects of supercritical reaction conditions in conversions over porous catalysts. In particular, the favorable balance between efficient desorption of hexene oligomers that can lead to catalyst coking at low gas-like densities and rapid pore diffusion that can limit the reaction rate at high liquidlike densities was demonstrated for this reaction (see Section 8.4.2.1) [36, 37, 107, 108]. Moreover, the yield ratio of cis- and trans-2-hexene over a γ -Al2 O3 shell-type catalyst was found to increase with pressure for temperatures above the critical point, whereas it decreased with pressure for subcritical temperatures of a liquid reaction phase (Fig. 10) [108]. This was attributed to a faster desorption of cis-2-hexene as the primary product from the catalyst surface under supercritical than under subcritical, liquid-phase conditions. Thus, high, liquid-like densities favor the consecutive reaction of the cis- to the thermodynamically favored trans-isomer. 2.00

Ycis-2-hexene / Ytrans-2-hexene

0.5

0

1999

250 °C 240 °C 230 °C 220 °C

1.75

Tc = 231 °C pc = 3.1 MPa

1.50

1.25

1.00

0

20

40

60

80

100

p / MPa

Yield ratio of cis- and trans-2-hexene in the isomerization of 1-hexene over a γ -Al2 O3 shell-type catalyst at different super- and subcritical temperatures as a function of pressure (after Ref. [108]).

Fig. 10

References see page 2004

2000

8.4 Heterogeneous Catalysis in Non-Conventional Solvents

1-Hexene isomerization is one of the rare cases for which the effective diffusivity of a reactant in the pores of a catalyst at near-critical conditions has been determined experimentally [106]. It was calculated from the pore effectiveness factor at isothermal, diffusionlimited conditions and its relation to the Thiele modulus (see Chapter 6.3). A maximum for the effective diffusivity was found close to the critical point of 1-hexene (Tc = 231 ◦ C, pc = 3.17 MPa). Although the effective diffusivity in the pores was much lower than in the bulk, it could be varied over two orders of magnitude by relatively small pressure changes in the near-critical region. Several studies have focused on the isomerization of nto isobutane under supercritical conditions (cf. Table 3). The highest reaction rate on a sulfated zirconia catalyst was found at 215 ◦ C and 4.0 MPa close to the critical point of n-butane (Tc = 152 ◦ C, pc = 3.80 MPa). Under these conditions, deactivation due to coking was much lower than in the gas phase and a stable conversion was observed [109, 110]. Although this conversion was slower than in the gas phase, higher production capacities for the isomerization products were achieved as a result of the higher feed density under supercritical conditions. 8.4.3.1.7 Miscellaneous Numerous other heterogeneously catalyzed conversions have been carried out in supercritical fluids as reaction media. These include disproportionation of alkylaromatics [108, 131, 133], hydrocarbon cracking [134, 135], etherifications [136], esterifications [137, 138], carbonylations [139, 140] or biomass gasification [141], to name just a few. Especially worthy of mention are reactions where the supercritical fluid acts both as a solvent and as a reactant. Examples for this type of conversion are the amination of alcohols with supercritical ammonia over unsupported Co–Fe [142, 143], the synthesis of cyclic carbonates from epoxides and supercritical CO2 over an immobilized Co complex [144], also as intermediates to dimethyl carbonate [145, 146] or the synthesis of N, N -dimethylformamide from supercritical CO2 , hydrogen and dimethylamine over a rutheniumcontaining silica aerogel [147].

Conversions in Ionic Liquids Ionic liquids have, so far, mainly been investigated as non-conventional solvents for homogeneously catalyzed conversions [38–40, 43]. In particular, multiphase (homogeneous) catalysis with ionic liquids represents a rapidly developing area (see Section 8.4.2.2). Another emerging field is the utilization of ionic liquids for the immobilization of molecular catalysts on porous solid supports [supported ionic liquid phases (SILPs); see Chapter 2.4.11] [148]. The application of ionic liquids 8.4.3.2

as reaction media for heterogeneous catalysis over solids, however, is still in its infancy. Nevertheless, several studies of such applications have been reported in the literature. Table 4 gives an overview of the types of solid-catalyzed reactions applying ionic liquids as reaction media. Selected cases will be discussed in more detail in the following to illustrate the role and potential advantages or limitations of ionic liquids in heterogeneous catalysis. 8.4.3.2.1 Hydrogenations and Oxidations The hydrogenation of cinnamaldehyde over 10 wt.% Pd/C was the topic of a systematic study using a variety of ionic liquid solvents (cf. Tables 4 and 5) [149]. First, the reactant is hydrogenated at the C−C double bond to hydrocinnamaldehyde before complete hydrogenation to 3-phenylpropanol takes place. Hydrogenation of the carbonyl group to cinnamyl alcohol in the first step was not observed in any solvent used in this study. The selectivity to hydrocinnamaldehyde is generally higher in the ionic liquids compared to the conventional organic solvents such as toluene or dioxane, even at complete conversion. This was attributed to the stronger interaction of the ionic liquid with the carbonyl bond lowering the reactivity of the C−O- vs. the C−C-double bond in cinnamaldehyde towards hydrogenation. The selectivity for hydrocinnamaldehyde is strongly dependent on the nature of the anion of the ionic liquid (Table 5). With respect to the conversion in the conventional organic solvents, the rate of the cinnamaldehyde hydrogenation was much lower in the ionic liquids (Table 5). Mass transfer limitations due to the higher viscosity of the ionic liquid and also due to the lower hydrogen solubility in the liquid reaction phase are most likely responsible for this observation. This is also evident from the lower conversions found for increasingly viscous ionic liquids (cf. Table 5). The diffusion coefficient in the ionic liquid [bmim][BF4 ] was estimated to be 20 times lower than in liquid toluene [149]. Similarly, lower reaction rates in ionic liquids compared with conversions in liquid organic solvents were also found for the hydrogenation of halonitrobenzenes to the corresponding anilines over Raney nickel or carbon-supported platinum or palladium [150] and for the hydrogenation of benzene over rhodium, iridium or platinum nanoparticles [163, 164]. Upon recycling of the Pd/C catalyst in the cinnamaldehyde hydrogenation, the activity dropped after the first run, but remained constant after that [149]. This was explained by blocking of the catalyst pores by the ionic liquid. Even after extensive washing with acetonitrile, the specific surface area of the catalyst after contact with the ionic liquids was reduced, depending on the nature of the ionic liquid, to 27–50% of that of the fresh Pd/C catalyst. Whether or not pore blocking can

8.4.3 Heterogeneously Catalyzed Conversions in Non-Conventional Solvents Tab. 4

2001

Heterogeneously catalyzed reactions carried out in ionic liquids as solvents TR / ◦ C

Catalyst

Ionic liquida

Cinnamaldehyde to hydrocinnamaldehyde or citral to citronellal

Pd/C

20–120

[149]

(Di)chloronitrobenzenes to (di)chloroanilines

Raney Ni, Pt/C, Pd/C

100

[150]

Nitriles to amines or dinitriles to aminonitriles

Ru/C

Combinations of the cations [bmim], [emim], [C6 py], [C8 py], [N8881 ] and the anions [BF4 ], [PF6 ], [OAc], [OTf], [NTf2 ] Combinations of the cations [bmim], [emim], [C6 mim], [C8 mim] and the anions [BF4 ], [PF6 ] ethylimidazolium chloride, [mim][HSO4 ]

100

[151]

Pd/Al2 O3

[bmim][NTf2 ]

50–85

[152]

H3 CReO3 /polymers

[bmim][PF6 ], [emim][NTf2 ]

25–60

[153]

Zeolite [Sn]Beta [Ti,Ge]MCM-41

[bmim][BF4 ] Combinations of the cations [emim], [bmim], [bdmim], [C8 mim], [Cn py] (n = 4, 6, 8), [bmpyr] and the anions [BF4 ], [PF6 ], [OTf], [NTf2 ], [EtOSO3 ].

25 40

[154] [155]

Friedel–Crafts alkylation of phenol with tert-butanol Friedel–Crafts acylation of anisole with benzoic anhydride

H-ZSM-5, H-Beta, H3 PW12 O40 /MCM-41 H-ZSM-5, H-mordenite, H-USY, H-Beta

60

[156]

80

[157]

Heck-reaction of iodobenzenes with acrylic esters, acrylonitrile or 2-methyl-prop-2-en–1-ol Prins cyclization of homoallyl alcohols with aldehydes Knoevenagel and nitroaldol condensation

Pd/C, Pd/SiO2 Pd/C Pd/chitosan

[bmim][PF6 ], [C6 mim][BF4 ], [C8 mim][BF4 ] Combinations of the cations [emim], [bmim], [edmim], [C4 mpip], [C8 py], [N8881 ], [P66614 ] and the anions [BF4 ], [PF6 ], [OTf], [NTf2 ]. [bmim][PF6 ] [bmpyr][NTf2 ], [n-Bu4 N]Br [n-Bu4 N]Br, [n-Bu4 N][OAc], [bmim][BF4 ], [C4 py][BF4 ] [bmim][PF6 ], [C6 mim]Cl, [n-Bu4 N][BF4 ] [bmim][PF6 ], [bmim][BF4 ]

100 100 130

[158] [159] [160]

25

[161]

25–60

[162]

Reaction

Ref.

Hydrogenations

Oxidations Cinnamyl alcohol to cinnamaldehyde with O2 Alkanes to alcohols or alcohols to ketones with H2 O2 Baeyer–Villiger oxidation Thioethers to sulfoxides and sulfones with H2 O2 C−C bond formation

a Abbreviations

H-ZSM-5, Amberlyst-15 Mg–Al hydrotalcites

as in Table 2.

occur obviously depends on the size of the molecular species in the ionic liquid relative to the pore size of the catalyst. Effects of the ionic liquid solvent on the selectivity of solid-catalyzed hydrogenations were also observed in the conversion of halonitrobenzenes [150] and of dinitriles [151], both over noble metal catalysts on carbon supports (cf. Table 4). In the former example, less dehalogenation was observed than in liquid methanol as the solvent. This was rationalized by (i) a strong interaction of the ionic liquid with the haloaniline resulting in rapid transport of the product away from the catalyst surface before consecutive dehalogenation or (ii) a lower activity of the catalyst towards dehalogenation in the ionic liquid. In the latter case, the addition

of an ionic liquid to toluene as the solvent afforded a higher selectivity for the partial hydrogenation of adipodinitrile to aminocapronitrile than in toluene as the single solvent. As for the hydrogenations, the aerobic oxidation of cinnamyl alcohol over 5 wt.% Pd/Al2 O3 was slower in an ionic liquid ([bmim][NTf2 ]) than in a conventional liquid solvent (toluene) [152]. This was, however, not attributed to mass transfer effects, but to more efficient solvation of the polar reactant in the ionic liquid than in the organic solvent. The resulting lower reactant concentration on the catalyst surface was assumed be responsible for the lower rate. References see page 2004

2002

8.4 Heterogeneous Catalysis in Non-Conventional Solvents

Tab. 5 Conversion X and selectivity S for the hydrogenation of cinnamaldehyde (CA) to hydrocinnamaldehyde (HCA) over 10 wt.% Pd/C in ionic liquids and in conventional organic solvents (η, solvent viscosity; TR , reaction temperature; pH2 , hydrogen pressure; tR , reaction time) (after Ref. [149])

Solventa

η/mPa s TR / ◦ C pH2 /MPa tR /h XCA /% SHCA /%

[bmim][BF4 ] [bmim][OTf] [bmim][OAc] [bmim][NTf2 ] [C6 py][BF4 ] [C8 py][BF4 ] [N8881 ][BF4 ] Toluene Cyclohexane Dioxane a Abbreviations

15.4 39.0 23.1 16.7 38.1 56.3 405.5 0.6 1.1 1.4

60 60 60 60 60 60 60 30 30 30

4.0 4.0 4.0 4.0 4.0 4.0 4.0 0.3 0.3 0.3

4 4 4 4 4 4 4 1 1 1

100 100 100 100 77 71 10 100 100 100

100 91 78 88 100 100 90 85 79 89

as in Table 2.

In contrast, for the heterogeneously catalyzed oxidations with hydrogen peroxide reported so far, the rates for the conversions in ionic liquids are higher than those in conventional organic solvents. One interesting example is the Baeyer–Villiger oxidation of ketones to esters over a tin-substituted zeolite Beta [154]. Whereas the conversion of substituted acetophenones proceeds well in liquid dioxane at 80 ◦ C, it does not occur at all at room temperature. However, yields of the ester product above 70% can be obtained over the same catalyst at room temperature if the ionic liquid [bmim][BF4 ] is used as the solvent. It is a promising result that the catalyst can be reused at least three times without a measurable loss of product yield. Higher reaction rates than in liquid dioxane were also reported for the selective oxidation of alkyl thioethers to the corresponding sulfoxides and sulfones with hydrogen peroxide over titanium- and germanium-substituted MCM-41-type materials as catalysts [155]. In this study, it could be proven that the leaching of active catalyst component into the reaction solution is significantly less for the conversion in the ionic liquid [emim][BF4 ] than in liquid ethanol (loss of titanium with respect to the fresh catalyst: [emim][BF4 ] 18 mol%, ethanol 40 mol%). Moreover, the authors systematically investigated the influence of the cation and the anion of the ionic liquid on conversion and selectivity (Table 6). With increasing carbon chain length in the 1-alkyl-3-methylimidazolium cation, the conversion decreased with almost unchanged sulfoxide selectivity. Two explanations were given for this result: (i) an increase in the viscosity of the ionic liquid and (ii) an increasingly pronounced reduction of the accessibility of the catalyst mesopores with increasing

Conversion and sulfoxide selectivity for the oxidation of 2-thiobenzylpyrimidine with hydrogen peroxide over [Ti, Ge]MCM-41 in ionic liquids at 40 ◦ C (after Ref. [155])

Tab. 6

Ionic liquida

X2−thiobenzylpyrimidine /%

[emim][BF4 ] [bmim][BF4 ] [C8 mim][BF4 ] [emim][OTf] [emim][EtOSO3 ] a Abbreviations

77 73 27 49 25

Ssulfoxide /% 95 96 100 98 100

as in Table 2.

size of the cation. However, the anion also has a considerable effect on the thioether conversion (cf. Table 6). Coordination of the anion to the catalytically active site and, thus, a direct influence on the transition state in the rate-limiting reaction step were held responsible for this finding. This coordination is particularly strong for the ethylsulfate anion. These results demonstrate that both the conversion and selectivity of a heterogeneously catalyzed reaction may be improved by a proper selection of an ionic liquid solvent with respect to the molecular nature of its constituents. 8.4.3.2.2 C−C Bond Formation The Heck reaction of aryl iodides with substituted alkenes was one of the first examples of a heterogeneously catalyzed conversions to be studied in ionic liquid solvents (cf. Table 4). In an early publication by Hagiwara et al. [165], a Pd/C catalyst was found to be reusable for six cycles without a considerable loss in the yield of the Heck coupling product. From the absence of Pd in the reaction mixture after the catalytic experiments, the authors concluded that the reaction occurred heterogeneously. However, later studies by Okubo et al. [158] and Forsyth et al. [159] unambiguously showed that the Heck coupling is homogeneously catalyzed by metallic palladium that was leached from the catalyst support. Thus, 9 mol% of the palladium on the fresh catalyst was lost to the reaction mixture after 24 h in the conversion of 1-tert-butyliodobenzene with 2-methylprop-2-en-1-ol at 100 ◦ C in the ionic liquid [bmpyr][NTf2 ] [159]. This leaching was not due to the action of the ionic liquid, but only occurred in the presence of the reactants and products. Nevertheless, the conversions in the ionic liquid solvents were faster than those in conventional organic solvents such as dimethylformamide [158] and N -methylpyrrolidinone or under solventless conditions [159]. Leaching of the noble metal was also observed in the Heck reaction over palladium nanoparticles supported on the natural polymer chitosan [160]. Moreover, in this case, an appreciable conversion was achieved in

8.4.3 Heterogeneously Catalyzed Conversions in Non-Conventional Solvents

ionic liquids with tetraalkylammonium or -phosphonium cations, whereas the conversion was negligible in ionic liquids, e.g. with the [bmim] cation. This was attributed to stabilization of the palladium nanoparticles by the tetraalkylammonium or -phosphonium cations. The metal leaching was explained by partial decomposition of the cations at the relatively high reaction temperature of 130 ◦ C (cf. Table 4). In case of [NBu4 ][OAc] as the ionic liquid, the anion acts both as a reducing agent for the Pd(II) complex initially present on the catalyst support and as the base for the Heck coupling. For the Friedel–Crafts alkylation of phenol with tertbutanol over the acidic zeolites H-ZSM-5 and H-Beta, higher conversions were achieved in ionic liquids than in liquid n-hexane as a conventional solvent [156]. With a selectivity for the para-substituted product in the range 50–99%, this conversion proves that shape selectivity effects over zeolite catalysts can also be exploited in ionic liquid solvents. Similarly enhanced conversions with respect to organic solvents as for the tert-butylation of phenol were also found for the Prins cyclization of allylic alcohols with aldehydes over zeolite H-ZSM-5 [161] and in the Friedel–Crafts acylation of anisole with benzoic anhydride over several acidic zeolites [157] in ionic liquids (cf. Table 4). Conversion and selectivity in the anisole benzoylation depended strongly on the nature of the cation in the ionic liquid. Several solvent cations were found to undergo an ion exchange for the protons in the acidic zeolite catalyst. The protons liberated into the ionic liquid reaction phase form Brønsted acids by interaction with the solvent anions. These are, then, the catalysts for the acylation in homogeneous solution. Consistently, for ionic liquids with cations too small to diffuse into the zeolite pores, only low or no conversion occurred. The dark orange color of the liquid reaction phase after the conversion due to the formation of polyacylated products further supports a homogeneously catalyzed conversion. Although no coke deposition on the zeolite was observed, the catalyst deactivated as a result of continuous loss of protons by ion exchange with the cation of the ionic liquid. Another consequence of the proton exchange into the ionic liquid is the formation of HF in the case of [PF6 ] as the solvent anion. The HF may lead to partial damage of the zeolite framework during the conversion. Conversions in Gas-Expanded Liquids Despite the obvious advantages of GXLs for catalytic conversions, reports on their application in heterogeneous catalysis are rare so far. To date, the examples of such conversions have been limited to hydrogenation, 8.4.3.3

2003

hydroformylation and oxidation reactions. Some of those will be briefly discussed in the following. The first report of the utilization of GXLs in heterogeneous catalysis was published in 1997 by Bertucco et al. [166]. They investigated the continuous selective hydrogenation of a mixture of unsaturated ketones with hydrogen in the presence of CO2 over a 1 wt.% Pd/Al2 O3 egg-shell catalyst at pressures of 12–17.5 MPa and temperatures of 423–473 K. At a constant reactant:CO2 ratio, the reaction rate increased with pressure. This was attributed to the higher mole fraction of hydrogen in the liquid reaction phase with increasing expansion by CO2 . However, the dilution of the reactants by increasing amounts of CO2 relative to the reactant in the feed led to lower reaction rates, when compared at constant pressure. In another example of a hydrogenation, α-pinene was converted in the presence of CO2 over a 10 wt.% Pd/C catalyst at 50 ◦ C [167]. The reaction occurred at a higher rate when the mixture was present as a CO2 -expanded liquid than when it was present as a single supercritical phase (above ca. 9.3 MPa). This was explained by the higher concentration of the reactant in the GXL phase (10 mol%) than in the supercritical mixture (mole fraction of 1 mol%) and, consequently, a higher reactant concentration at the catalyst surface. More recently, however, two groups showed that the rate of metal-catalyzed hydrogenations can be related to the hydrogen solubility in the GXL reaction phase. Phiong et al. [168] performed a detailed kinetic analysis of the hydrogenation of α-methylstyrene over a carbon-supported palladium catalyst. Xu et al. [169] studied the hydrogenation of the aromatic rings in polystyrene with 5 wt.% Pd/BaSO4 and 65 wt.% Ni/Al2 O3 –SiO2 as catalysts. Both the hydrogen solubility and the hydrogenation rate were higher in CO2 -expanded decahydronaphthalene than in the pure liquid solvent at a constant hydrogen partial pressure in the gas phase (Table 7). However, the rate increase was not directly proportional to the hydrogen solubility. In a subsequent publication, the authors investigated the deactivation behavior of the two metal catalysts in the GXL phase conversion [170]. In both cases, poisoning of the catalyst occurred. The Pd-based catalyst was poisoned by CO formed by the reverse water-gas shift reaction from CO2 hydrogenation as also reported for many other hydrogenation reactions in supercritical CO2 (see Section 8.4.3.1.1). Although much less pronounced, the Ni-based catalyst was subject to poisoning by water as a product of CO2 methanation in the GXL reaction phase. References see page 2004

2004

8.4 Heterogeneous Catalysis in Non-Conventional Solvents

Tab. 7 Volume expansion V/V0 , hydrogen solubility and rate constant k for the hydrogenation of polystyrene in decahydronaphthalene without and with expansion by CO2 over 5 wt.% Pd/BaSO4 and 65 wt.% Ni/Al2 O3 –SiO2 (weight ratio 3 : 100) at 150 ◦ C (pH2 = 5.1 MPa, (pH2 + pCO2 ) = 20.4 MPa) (after Ref. [169])

Gas H2 H2 + CO2

V/V0

H2 solubility/mol cm−3

k/cm3 g−1 s−1

1.0 1.3

2.4 × 10−4 4.3 × 10−4

6.2 × 10−3 9.7 × 10−3

Hemminger et al. [171] compared the hydroformylation of 1-hexene over a rhodium complex immobilized on a phosphine-modified silica support in liquid toluene, CO2 expanded toluene and supercritical CO2 . Interestingly, the conversion rate was highest in the CO2 -expanded liquid phase. However, substantially higher amounts of doublebond isomerization products were formed in the GXL phase. Thus, the aldehyde yield achieved per rhodium atom was similar in the expanded liquid and in supercritical CO2 . As a further result of the isomerization in the expanded liquid, the selectivity for the desired n-aldehyde relative to isoaldehydes was lower in the CO2 expanded toluene. Supercritical CO2 was, therefore, the most favorable medium for the heterogeneously catalyzed hydroformylation. GXLs may also aid in reducing leaching of active catalyst components into a liquid reaction mixture. Improved resistance of an iron porphyrin chloride complex supported on the ordered mesoporous material MCM41 towards leaching of the metal complex into liquid acetonitrile was reported when the solvent was expanded by CO2 [172]. Thus, in two subsequent experiments on cyclohexene epoxidation with iodosylbenzene using the fresh and the used catalyst, the iron contents of the reaction mixture were 1.19 and 0.81 ppm for the CO2 expanded solvent, whereas 2.67 and 2.52 ppm of iron were found after reaction in neat acetonitrile, respectively. This was attributed to the lower polarity of the expanded vs. the pure organic solvent. The lower polarity of the acetonitrile with increasing CO2 content was also used to explain the decreasing conversion of cyclohexene oxidation with oxygen over the same catalyst at increasing levels of expansion. Nevertheless, at an expansion degree of ca. 1.4, the cyclohexene conversion could be roughly doubled with respect to the reaction in pure acetonitrile. 8.4.4

Outlook

The application of non-conventional solvents in heterogeneously catalyzed conversions provides several promising opportunities. Although the effect of solvents in

heterogeneous catalysis is generally less diverse and mostly less pronounced than in homogeneous reaction systems, there are sufficient examples that render further efforts towards the design of innovative solvents for heterogeneous catalysis worthwhile. Among the major benefits of non-conventional solvents such as supercritical fluids, ionic or gas-expanded liquids in heterogeneous catalysis are the following: (i) Conventional solvents may be replaced, resulting in intrinsically safe, more environmentally benign and economically more attractive processes. (ii) Non-conventional solvents can be truly enabling tools. Thus, a significant enhancement of rates and selectivities may be achieved by sensitive tuning of the solvent properties. (iii) Innovative solvents may aid in enhancing catalyst lifetime, e.g. by extraction of catalyst poisons such as coke deposits under reaction conditions. (iv) The combination of solvents may offer additional advantages, e.g. for facilitated and productionintegrated separations or for process intensification. In that respect, the concomitant application of supercritical fluids and ionic liquids appears particularly attractive. Despite these advantages, several challenges lie ahead to permit more rational exploitation of innovative solvent systems in heterogeneous catalysis. Above all, the effects of the nature and the properties of the solvents on the catalytically active sites or on the adsorption and mass transfer, especially in porous catalysts, are not well understood. The application of increasingly powerful in situ spectroscopic methods and computer-based modeling or simulation on the molecular level may aid in achieving a deeper understanding of these processes in non-conventional solvents. Innovative solvent systems are one approach to design eventually a tailor-made reaction environment for a given catalytic conversion. This is especially useful for the manufacture of fine and specialty chemicals where both efficient solvation and high selectivities at low temperatures are indispensable prerequisites. It is, therefore, obvious that non-conventional solvents can make a significant contribution towards the goal of ‘‘green’’ and sustainable catalytic processes. References 1. R. A. Sheldon, Green Chem. 2005, 7, 267. 2. H.-U. Blaser, M. Studer, Green Chem. 2002, 5, 112. 3. C. Reichardt, Solvents and Solvent Effects in Organic Chemistry, 3rd Ed., Wiley-VCH, Weinheim, 2003, 629 pp. 4. R. A. Sheldon, Pure Appl. Chem. 2000, 72, 1233. 5. W. Leitner, Pure Appl. Chem. 2004, 76, 635. 6. W. Keim, Green Chem. 2003, 5, 105.

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2007

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8.5

Sonocatalysis Kenneth S. Suslick∗ and Sara E. Skrabalak

8.5.1

Introduction and the Origins of Sonochemistry

Research on the chemical effects of ultrasound has undergone a renaissance during the past few decades and has had a significant impact in a variety of areas [1, 2]. Applications of sonochemistry have been developed in virtually all areas of chemistry and related chemical and materials technologies [3–5]. We can conceptually divide the effects of ultrasonic irradiation on heterogeneous catalysis into those that alter the formation of heterogeneous catalysts, those that perturb the properties of previously formed catalysts and those that affect catalyst reactivity during catalysis. In practice, these three classes of effects are often deeply intertwined in reported experimental results. No direct coupling of the acoustic field with chemical species on a molecular level can account for sonochemistry. Ultrasound spans frequencies from roughly 20 kHz to 10 MHz, with associated acoustic wavelengths in liquids of roughly 100–0.15 mm: these are not on the scale of molecular dimensions. Instead, the chemical effects of ultrasound derive from several non-linear acoustic phenomena, of which cavitation is the most important. Acoustic cavitation is the formation, growth and implosive collapse of bubbles in a liquid irradiated with sound or ultrasound. When sound passes through a liquid, it consists of expansion (negative pressure) waves and compression (positive pressure) waves. These cause bubbles (which are filled with both solvent and solute vapor as well as dissolved gases) to grow and recompress. Under proper conditions, acoustic cavitation can lead to implosive compression in such cavities, producing intense local heating, high pressures and very short lifetimes. As discussed elsewhere, these hot spots have temperatures of roughly 5000 ◦ C, pressures of about 1000 atm (1 atm = 101.325 kPa) and heating and cooling rates above 109 K s−1 [6–10]. Cavitation is an extraordinary method of concentrating the diffuse energy of sound into a chemically usable form. References see page 2015 ∗ Corresponding author.

2008

8.5 Sonocatalysis

When a liquid–solid interface is subjected to ultrasound, cavitation still occurs, but with major changes in the nature of the bubble collapse. If the surface is significantly larger than the cavitating bubble (∼100 µm at 20 kHz), spherical implosion of the cavity no longer occurs, but instead there is a markedly asymmetric collapse which generates a jet of liquid directed at the surface, as seen directly in high speed micro-cinematographic sequences shown in Fig. 1. The tip jet velocities have been measured by Lauterborn to be greater than 100 m s−1 [11]. The origin of this jet formation is essentially a shapedcharge effect: the rate of collapse is proportional to the local radius of curvature. As collapse of a bubble near a surface begins, it does so with a slight elliptical asymmetry, which is self-reinforcing and generates the observed jet. The impingement of this jet can create localized erosion (and even melting), surface pitting and ultrasonic cleaning. A second contribution to erosion created by cavitation involves the impact of shock waves generated by cavitational collapse. The magnitude of such shock waves is thought to be as high as 104 atmospheres, which will easily produce plastic deformation of malleable metals [12]. The relative importance of these two effects depends heavily on the specific system under consideration. Enhanced chemical reactivity of solid surfaces is associated with these processes. Cavitational erosion generates unpassivated, highly reactive surfaces; it causes short-lived high temperatures and pressures at the surface; it produces surface defects and deformations; it forms fines and increases the surface area of friable solid supports and it ejects material in unknown form into solution. Finally, local turbulent flow associated with acoustic streaming improves mass transport between the

liquid phase and the surface, thus increasing observed reaction rates. In general, all of these effects are likely to be occurring simultaneously. In contrast, the effects of ultrasound on slurries of fine particles do not come from microjet formation during cavitation. Distortion of bubble collapse requires a solid surface several times larger than the resonance bubble size. Thus, for ultrasonic frequencies of ∼20 kHz, damage associated with jet formation cannot occur for solid particles smaller than ∼200 µm. In these cases, however, shockwaves created by homogeneous cavitation can create high velocity interparticle collisions, with impact speeds of several hundred meters per second and local effective transient impact temperatures of roughly 3000 K [13, 14]. The turbulent flow and shockwaves produced by intense ultrasound can drive metal particles together at sufficiently high speeds to induce effective melting at the point of collision, as shown in Fig. 2. The high-velocity interparticle collisions produced in slurries of malleable materials cause smoothing of individual particles and agglomeration of particles into extended aggregates [5, 15]. Surface composition depth profiles of sonicated powders show that ultrasonic irradiation effectively removes surface oxide coatings. The removal of such passivating coatings dramatically improves reaction rates for a wide variety of reactions. With larger flakes of brittle materials, interparticle collisions cause shock

1 µm

Scanning electron micrograph of 5 µm diameter Zn powder after ultrasonic irradiation of a slurry in decane. Neck formation from localized melting is caused by high-velocity interparticle collisions. Similar micrographs and elemental composition maps (by Auger electron spectroscopy) of other metal powders and mixed metal collisions have also been made. Reproduced with permission [25].

Fig. 2

Cavitation near a liquid–solid interface. High-speed micro-cinematographic sequence of laser-induced cavitation near a solid surface, showing the formation of a microjet impact; 75 000 frames s−1 . The sequence is from left to right, top to bottom; the solid boundary is at the bottom of each frame. Photograph courtesy of W. Lauterborn; reproduced with permission [11].

Fig. 1

8.5.2 Effects of Ultrasound on Heterogeneous Catalysts

fragmentation instead, which can increase surface areas dramatically and contribute to increased activity [15–17]. The term sonocatalysis should be restricted in its use to refer only to the creation of a catalytically competent intermediate by ultrasonic irradiation. One should not refer to a simple sonochemical rate enhancement of a reaction by this term, just as one would use the term photochemistry and not photocatalysis, to describe a stoichiometric reaction caused by light. In this chapter, the symbol shown in Eq. (1) will be used to indicate ultrasonic irradiation or ‘‘sonication’’ of a solution leading to a sonochemical reaction. )))

−−−→

(1)

8.5.2

Effects of Ultrasound on Heterogeneous Catalysts

Ultrasonic irradiation can alter the reactivity observed during the heterogeneous catalysis of a variety of reactions [5, 18]. In addition to the more recent work described in this chapter, there is extensive (but little recognized) past literature in this area, particularly from Eastern Europe [19]. The effects of ultrasound on catalyst formation can be far reaching; changes in crystallization, dispersion and surface properties are all possible. Alteration of properties of pre-formed catalysts can also have substantial effects. Oxide or other passivating coatings can be removed and increased dispersion can occur, sometimes from the fracture of friable supports. Irradiating operating catalysts often improves mass transport. Metal Powders The use of ultrasound for syntheses involving liquid–solid heterogeneous reactions has been a matter of intense investigation [20–22]. In general, ultrasonic treatment of these metals promotes reaction pathways favoring single 8.5.2.1

2009

electron transfers [23], probably through the removal of thin oxide coatings which are often dominated by acid–base activity. Ultrasonic activation of commercial transition metal powders has also received substantial attention [24–27]. 8.5.2.1.1 Modification of Bulk Metals The effect of ultrasonic irradiation on bulk metals has been studied extensively [28–30]. In particular, ultrasonic pretreatment on hydrogenation catalysts has been studied and impressive rate accelerations have been reported. The hydrogenation of alkenes by ordinary Ni powder is enormously enhanced (>105 -fold) by ultrasonic irradiation [24]. The surface area of the catalyst, however, did not change significantly even after lengthy irradiation. Rather, both surface smoothing (Fig. 3) and particle agglomeration were observed, due to interparticle collisions caused by cavitation-induced shock waves. Auger electron spectroscopy revealed a decrease in the thickness of the oxide coat after ultrasonic irradiation; the removal of this passivating layer is likely responsible for the increased catalytic activity. Ultrasonic treatment of Raney Ni enhances hydrogenation and hydrogen–deuterium exchange rates [31]. Hydrogen isotopes have been selectively introduced into aromatic compounds by the reaction of haloaromatic compounds with basic deuterated (or tritiated) aqueous solutions over Raney catalysts under ultrasound. Carbohydrates and glycosphingolipids have also been deuterated [32, 33]. Treatment of ultrasonically prepared Raney nickel with tartaric acid results in a highly efficient enantioselective catalyst for the hydrogenation of 1,3-diketones to 1,3-diols [34] with a similar catalyst preparatory method having recently been patented [35]. Ultrasound is also being used to regenerate Raney Ni catalysts in situ [36, 37]. References see page 2015

80 µm

The effect of ultrasonic irradiation on the surface morphology and particle size of Ni powder. Initial particle diameters before ultrasound were ∼160 µm; after ultrasound, ∼80 µm; the micrographs are on the same scale. High-velocity interparticle collisions caused by ultrasonic irradiation of slurries are responsible for the smoothing and removal of passivating oxide coating. Reproduced with permission [15].

Fig. 3

2010

8.5 Sonocatalysis

As an extension of the Raney Ni work, Boudjouk prepared an efficient, recyclable Ni hydrosilation catalyst from the reduction of NiI2 with Li under ultrasound [38, 39]. The reaction of acrylonitrile had yields >95% at 0 ◦ C, whereas commercial Ni powder was not active even after extensive sonication. More recent work has looked at the hydrogenation of α, β-unsaturated ketones with high yields and chemoselectivity [40]. Ultrasound also influences the properties of platinum and palladium blacks prepared by the reduction of H2 PtCl6 or PdCl2 in formaldehyde [41]. For Pt black, slight increases in hex-1-ene hydrogenation and ethanol oxidation were observed and explained by an increase in surface area. H2 O2 decomposition rates, however, were much greater than the corresponding surface area enhancement, suggesting that the amorphous phase generated from ultrasonic irradiation supports H2 O2 decomposition whereas the hydrogenation and oxidation reactions require more ordered structures not abundantly formed under such conditions. The allylation of ketones and aldehydes by allylic alcohols [Eq. (2)] has been improved using ultrasonic irradiation of a palladium–tin dichloride catalyst in less polar solvents [42]. Inverted regioselectivity was observed compared with homogeneous carbonyl allylation in polar solvents. )))

H3 CCH=CHCH2 OH + H5 C6 CHO −−−−−→ Pd/SnCl2

H3 CCH=CHCH2 CH(OH)C6 H5

(2)

Ultrasound has also been used to prepare Fischer– Tropsch catalysts. Liquid-phase hydrogenation of carbon monoxide was accomplished with ultrafine particles (106 K s−1 ) of molten metals is necessary to prevent crystallization. Acoustic cavitation can induce extraordinary local heating in otherwise cold liquids and can provide enormous cooling rates (>109 K s−1 ), thus providing a new synthetic route to amorphous metal powders. From work on the sonolysis of volatile Co, Mo and W precursors [50], it appears that this is a general phenomenon and extension to the synthesis of amorphous intermetallic alloys has proven successful. Sonochemically synthesized amorphous powders may have important catalytic applications, especially given their very high surface areas and nanometer cluster size. For example, sonochemically prepared nanophase iron is an active catalyst for the Fischer–Tropsch hydrogenation of CO and for hydrogenolysis and dehydrogenation of alkanes, in large part due to its high surface area (>120 m g−1 ). Rates of conversion of CO and H2 to low molecular weight alkanes were approximately 20 times higher per gram of Fe than for fine particle (5 µm diameter) commercial iron powder at 250 ◦ C and more than 100 times more active at 200 ◦ C. Selectivities are similar. The reactions of cyclohexane are interesting because of their inherent catalyst surfacestructure sensitivity. In this manner, the nature of the catalytic process can be useful as a chemical probe of 8.5.2.1.2

200 nm

Scanning electron micrograph of amorphous nanostructured iron powder produced from the ultrasonic irradiation of Fe(CO)5 . Reproduced with permission [47].

Fig. 4

8.5.2 Effects of Ultrasound on Heterogeneous Catalysts

the effect of ultrasound on the catalytically active surface. Catalytic studies were performed on nanophase Fe/Co alloys produced sonochemically. The ratio of cyclohexane dehydrogenation to hydrogenolysis depended on alloy composition. The 1 : 1 alloys gave nearly exclusively benzene, in stark contrast to either pure metal [51]. Building on this work, Gedankan and coworkers have looked at the aerobic oxidation of cycloalkanes with Fe, Co and Fe/Ni alloys produced sonochemically [52]; with cyclohexane as a substrate, conversions of 40%, with 80% selectivity for cyclohexanone and cyclohexanol, were achieved under mild conditions. Several groups have reported the sonochemical synthesis of metal (Cu, Pt, Pd, Au) or bimetallic (Au/Pt, Au/Pd, Co/Ni) nanoparticles; increased rates of hydrogenation for 4-pentenoic acid with Au/Pd catalysts [53] and the Ullman reaction for Cu nanoparticles [54] have been reported. Otherwise, applications of sonochemically generated discrete metallic nanoparticles to catalysis are scant. Metal Oxides as Catalysts There are many reports on the effects of ultrasound on metal oxide catalyst preparation [55]. Mixed Cr–Mo and Cr–Fe oxide catalysts have been prepared with ultrasonic treatment and examined for the oxidation of methanol to formaldehyde [56, 57]. CuO catalysts were prepared with ultrasound and tested them for de-NOx reduction [58]. Mokryi and Starchevskii examined the vapor-phase oxidation of a number of organic compounds after ultrasonic activation of Fe–Te–Mo and Cs–Pb–Mo oxide catalysts [59]. Isobutylene, methanol and ethanol were examined; modest increases in specific surface areas and catalytic activity were obtained, but selectivity towards the desired products decreased. Gedankan and coworkers have published numerous papers on the sonochemical synthesis and catalytic properties of metal oxides. Mesoporous cobalt, nickel and iron oxides have been prepared by incorporating structure-directing agents such as cetyltrimethylammonium bromide (CTAB) in an inorganic precursor solution. Ultrasonic irradiation was supplied under air [60, 61]. After solvent extraction, high surface area materials (α-Fe2 O3 274, Co3 O4 72.83, NiO 39.84 m2 g−1 ) were produced. Slightly better conversions of cyclohexane to cyclohexanone and cyclohexanol were observed for the calcined oxides compared with other nanostructured forms of the corresponding metal oxides, with the best conversions, ∼40%, being obtained for the sonochemically produced Fe2 O3 . The same metal oxides have also been deposited into the pores of mesoporous oxide carriers [62, 63] and incorporated into composites [64] of traditional support materials; the conversions and selectivities were comparable 8.5.2.2

2011

to those for unsupported, sonochemically generated bulk materials. Sahle-Demessie and coworkers have reported selective hydrocarbon oxidation using sonochemically generated vanadium phosphorus oxide (VPO) [65]. It is well known that alumina itself, acting as a solid acid or base, can be an active catalyst for a variety of organic reactions. Early work in this area was conducted by Ando’s group [66]. Their initial discovery was the improvement made by ultrasonic irradiation of the liquid–solid two-phase synthesis of aromatic acyl cyanides from acid chlorides and solid KCN in acetonitrile [67]. The extension of this reaction to benzyl bromides led to an unusual observation of reaction pathway switching [68]. With mechanical agitation (i.e. stirring), the reaction of benzyl bromide and KCN in aromatic solvents, catalyzed by alumina, yields diarylmethane products from Friedel–Crafts attack on the solvent [Eq. (3)], whereas with ultrasonic irradiation, one obtains benzyl cyanide [Eq. (4)]. Apparently, the ultrasonic irradiation of alumina deactivates the Lewis acid sites normally present that are responsible for the Friedel–Crafts reactivity. It is thought that this poisoning is accomplished by the added solid basic salts (e.g. KCN) with ultrasound, perhaps through solid–solid contacts or through increased access of dissolved bases to the alumina surface. stirred

C6 H5 CH2 Br + C6 H5 CH3 + KCN −−−→ Al2 O3

C6 H5 CH2 C6 H4 CH3

(3) )))

C6 H5 CH2 Br + C6 H5 CH3 + KCN −−−→ C6 H5 CH2 CN Al2 O3

(4) Catalysis by alumina in the presence of ultrasound is a generalizable class of reactions, e.g. the sonocatalysis of aldol condensations and related reactions, Michael additions [69, 70] and Knoevenagel condensations [71]. Substantial improvements in yields were observed, with greatly diminished reaction times, for a variety of substrates. In the same vein, a useful synthesis of α-aminonitriles, which are important intermediates in amino acid synthesis, has been reported using alumina with ultrasound [72]. Yu et al. studied the effect of ultrasound on the synthesis and resulting photocatalytic activity of titania. They found that ultrasonic treatment of titania sols accelerated the hydrolysis and crystallization of titania [73]. The dried gels had an increased brookite to anatase ratio compared with titania sols prepared without ultrasound treatment; they also showed a modest increase in photocatalytic activity compared with that of the standard Degussa P25 titania. References see page 2015

2012

8.5 Sonocatalysis

80 nm

Transmission electron micrograph of mesoporous titania with a wormhole-like internal structure prepared with high intensity ultrasound. Photograph courtesy of J. Yu; reproduced with permission [74].

Fig. 5

Similarly, they found that dropwise addition of titanium isopropoxide–glacial acetic acid–ethanol solution to water under high intensity ultrasonic irradiation generated thermally stable, mesoporous TiO2 with an unusual wormhole-like structure (Fig. 5) [74]; they attributed this structure to the controlled condensation and agglomeration of TiO2 sol particles unique to ultrasound irradiation. Calcination resulted in a completely anatase phase TiO2 network; its photocatalytic activity was found to be slightly greater than that of Degussa P25 titania, probably due to increased surface area and enhanced diffusion of reactants and products. Variations on these procedures are published with similar results [75, 76]. A few reports on the ‘‘sonophotocatalytic’’ properties of titania have also been reported in which cooperative effects between irradiation with sound and light increase the degree of water splitting observed over commercial titania [77, 78]. Metal Carbides and Sulfides Molybdenum and tungsten carbides have been examined as catalysts because their activity is often similar to that of the platinum group metals. For catalytic applications, high surface area materials are needed, but the preparation of molybdenum and tungsten carbides with high surface areas has proven difficult. Suslick and coworkers have synthesized high surface area forms of these carbides [79, 80] from ultrasonic irradiation of 8.5.2.3

Mo(CO)6 or W(CO)6 in hexadecane under an argon atmosphere. Surface areas of 188 m2 g−1 for Mo2 C and 120 m2 g−1 for W2 C were obtained. Carburization of the resulting materials induced crystallization and the surface areas decreased slightly to 130 and 60 m2 g−1 , respectively. The suppression of hydrocracking during dehydrogenation remains a significant challenge for non-platinum catalysts. Sonochemically generated Mo2 C was tested for the dehydrogenation of cyclohexane [79] and found to be highly selective for benzene. In fact, like platinum, no hydrogenolysis products were observed. The overall activity of the sonochemical Mo2 C was comparable to that of Pt. Sonochemically generated Mo2 C and W2 C were also tested for the hydrodehalogenation (HDH) of monohalobenzenes [80]. Previous studies have emphasized the possibility of using noble metals (primarily Pd, Pt and Rh) for HDH but the use of such metals has been unsatisfactory, as hydrogenation tends to be favored over HDH. Both sonochemically generated Mo2 C and W2 C demonstrated high selectivity for the HDH of monohalobenzenes. For example, no chlorocyclohexane and cyclohexane were observed for the HDH of chlorobenzene. A previous study with conventional Mo2 C reported the hydrogenation of benzene to cyclohexane. The sonochemically generated carbides were also active for the HDH of CFCs, PCBs and their brominated analogs. Sonochemically generated Mo2 C was also tested for the catalytic hydrodenitrogenation (HDN) of indole [81], a common organonitrogen compound in crude oil. From 300 to 400 ◦ C, slightly higher conversions of indole to aryl amines and hydrocarbons were obtained compared with standards; this result was attributed to more randomly distributed crystallites, a direct consequence of the sonochemical synthesis. Above 400 ◦ C, substantial sintering of the sonochemically generated Mo2 C resulted in a drop in activity below that of the standards. Given the increasingly strict regulation of the sulfur content of fuels, improved hydrodesulfurization (HDS) catalysts are needed. Molybdenum sulfide is traditionally used; however, as HDS activity is due only to the edges of this layered material, the preparatory method greatly affects its catalytic performance. High-intensity ultrasonic irradiation of Mo(CO)6 and sulfur in isodurene under argon yields nanostructured MoS2 with a very high edge content [82]. TEM images (Fig. 6) show highly disordered MoS2 with much greater edge and defect content. This MoS2 was tested for the HDS of thiophene at atmospheric pressure and found to be highly active: three times more active than conventional MoS2 at 375 ◦ C and also more active than ReS2 and RuS2 . Recently, Mo2 N [83], hollow MoS2 nanospheres [84] and MoS2 microspheres [85] have also been prepared by ultrasound; the last two both display a higher degree of disorder with higher edge surface

8.5.2 Effects of Ultrasound on Heterogeneous Catalysts

1 µm

1 µm

2013

10 nm

TEM of sonochemical MoS2

SEM of conventional MoS2 (a)

SEM of sonochemical MoS2 (b)

(c)

Scanning electron micrograph of conventional molybdenum sulfide (a) compared with the SEM (b) and TEM (c) of nanostructured MoS2 produced from the ultrasonic irradiation of Mo(CO)6 and elemental sulfur as a slurry in isodurene. Reproduced with permission [82].

Fig. 6

exposure than conventional MoS2 and enhanced HDS activities. Catalyst Supports: Sonogels and Zeolites Support materials are used to disperse catalytically active phases. There has been a flurry of research studying the effect that ultrasound has on the synthesis and activation of catalyst support materials. Inorganic supports such as SiO2 and TiO2 are typically produced via sol–gel processing. In 1984, Tarasevich described an approach to sol–gel processing that eliminated the need for additional solvent and reduced the preparation time by exposing a gel precursor solution to intense ultrasonic irradiation [86]. Since then, the groups of Esquivias [87] and Zarzycki [88] have studied extensively the effect that ultrasound has on the kinetics and morphology of the resulting ‘‘sonogels’’. The term ‘‘sonogel’’ refers to either a solventcontaining gel or a gas-containing gel (a xerogel) made in the presence of ultrasound. SiO2 , TiO2 , SiO2 −TiO2 , SiO2 −Al2 O3 −MgO, SiO2 −P2 O5 and ZrO2 sonogels have been synthesized [87] and also ormosils (ORganically MOdified SILicates) [89]. The sonogels, upon solvent removal, appear to have finer porosity and greater reticulation of the network than gels prepared without ultrasound. Interestingly, chemical characterization of TiO2 −SiO2 sonogels showed improved dispersion of Ti in the SiO2 network [90]. Likewise, a patent was issued for a sonochemical process for preparing Zr containing aluminoxanes [91]. Upon sonication of a toluene solution of (CH3 )3 Al and Cp2 ZrCl2 with water (Cp = cyclpentadienyl), an aluminoxane gel was formed that was an active catalyst for oligomerization of 1-octene. Rhodium metal has been dispersed on TiO2 −SiO2 aerogels [92]. Ultrasound was utilized in two different preparations. In the first case, a ‘‘sonogel’’ was obtained by hydrolysis of Ti and Si alkoxides in the presence of 8.5.2.4

ultrasonic irradiation, which was then impregnated with a rhodium nitrate solution. In the second, a mixture of the alkoxides and a rhodium nitrate solution were exposed to ultrasound, thus leading to a ternary Rh−TiO2 −SiO2 sonogel. The behavior of these catalysts was compared with that of an Rh/TiO2 −SiO2 system obtained by conventional impregnation methods, starting with a commercial silica support. The first example gave a ca. 10-fold increase in catalytic activity for the hydrogenation of benzene whereas the second sample was not active for benzene hydrogenation due to poor Rh dispersion. Other metals have also been dispersed on to sonogel carriers, with similar results being obtained [90, 92, 93]. Carbon sonogels have been produced by Tamon’s group [94, 95]. Irradiation with high-intensity ultrasound promotes the sol–gel polycondensation of resorcinol and formaldehyde, typical precursors for carbon supports. As with the oxides, gelation occurs more rapidly when exposed to ultrasonic irradiation. The resulting carbon gels have increased mesoporosity compared with those prepared without ultrasound and moderate surface areas (500–800 m2 g−1 ). Such materials could be ideal supports for Pt-based fuel cell catalysts. Ordered, mesoporous SiO2 structures such as MCM-41 have been prepared via ultrasound with a drastic reduction in fabrication time being reported, down from several days to 3–6 h [96]. The resulting product has thicker pore walls, improving its thermal stability. Although some investigations of the effects of ultrasound on aluminosilicates and their syntheses have been published, this area still remains relatively unexplored. The best characterized study is that of Lindley [97], who examined the sonochemical effects on syntheses of zeolite NaA. Several-fold reductions in nucleation time and rates of formation during hydrothermal synthesis were monitored References see page 2015

2014

8.5 Sonocatalysis

by X-ray diffraction. Scanning electron micrographs showed significant changes in morphology also, with ultrasound producing a more agglomerated product made up of finer, micron-sized crystallites. Delaminated zeolites [98] have also been prepared by exposure to ultrasound, resulting in disordered, individual sheets of crystalline zeolitic materials; the delaminated zeolites were compared with conventional zeolites for a variety of acid-catalyzed organic reactions and found to have superior catalytic activity. Supported Catalysts The use of ultrasound in the preparation of supported metal catalysts has been examined primarily for hydrogenation reactions. For example, ultrasonic irradiation during the deposition of Pt on silica produces an 80% increase in Pt dispersion [99]. UV–visible, pH and transmission electron microscopy measurements have been conducted on supported catalysts prepared in the presence of ultrasound [100]; these studies indicated that the enhanced dispersion is probably due to a faster rate of metal reduction due to radical production from the sonolysis of water or other solvent molecules. Under appropriate conditions, it is believed that ultrasound can assist the insertion of metal particles into support pores due to microjet and/or shockwave formation accelerating metal agglomerates into the support material [101]. Several Japanese patents make use of ultrasound to improve the dispersion and reliability of supported noble metal for fuel cells [102, 103]. The general process described involves the reduction of H2 PtCl6 in a carbon carrier, often colloidal, in the presence of ultrasonic irradiation. Ultrasound can also alter the reactivity of already formed supported catalysts. Han and coworkers examined the acceleration of hydrosilation reactions of alkenes and alkynes catalyzed by Pt/C in the presence of ultrasound [104, 105]. Various substrates, including 1-hexene, styrene and phenylacetylene, work effectively even at −30 ◦ C with various silanes. The separation of products from catalyst by filtration, however, is not possible as ultrasonic treatment generates a fine colloidal suspension of support material, thus defeating one of the primary advantages of heterogeneous catalysis. These researchers extended the use of this system to the hydrogenation of alkenes using formic acid as a hydrogen transfer agent and Pd/C catalyst [106]. In this case, filtration was still effective for removal of the catalyst, but rate enhancements were no greater than with heating. Replacing formic acid with hydrazine yielded similar results [107]. T¨or¨ok and coworkers have developed highly chemoselective and enantioselective hydrogenation catalysts that incorporate a sonochemical pretreatment to supported Pt and Pd catalysts [108, 109]. For example, a two-fold 8.5.2.5

increase in the rate of hydrogenation of cinnamaldehyde was observed for Pt/SiO2 catalysts treated with ultrasound [108]; the selectivity for cinnamyl alcohol increased by 45%. A similar rate enhancement was observed for the hydrogenation of prochiral carbonyl compounds to their corresponding (R)-hydroxyl derivatives [110]. The addition of cinchonidine, a chiral modifier, to the pretreatment solution, increased enantioselectivity for a Pt/Al2 O3 catalyst with various substrates including ethyl pyruvate (97% ee), methyl pyruvate (95% ee), ethyl 4-phenyl-2-oxobutyrate (95% ee) and ethyl benzoylformate (92% ee). These are the best ees ever achieved for this heterogeneous system; these results are probably due to more effective surface modification achieved with ultrasonic pretreatment. The influence of ultrasound on supported catalyst preparation has been extended beyond noble metal deposition by Suslick’s group. A nanostructured, bifunctional catalyst, Mo2 C/ZSM-5, was prepared by irradiation of Mo(CO)6 and HZSM-5 in a slurry with hexadecane [111]. As the event responsible for the formation of metal clusters is in the gas phase of the collapsing bubbles and therefore separate from the oxide support, eggshell catalysts are formed; ∼2-nm Mo2 C particles decorate the surface of the ZSM-5 support. Studies of the dehydroaromatization of methane to benzene were performed. Several supported HDS catalysts have been prepared using ultrasound. Suslick and coworkers [112] prepared Coand Ni-promoted MoS2 supported on alumina through high–intensity ultrasonic irradiation of isodurene slurries containing Mo(CO)6 , Co2 (CO)8 , elemental sulfur and Al2 O3 or Ni−Al2 O3 under an Ar flow. The sonochemically prepared catalysts are extremely active catalysts for the HDS of thiophene and dibenzothiophene with activities several times those of comparable catalysts under identical conditions. Moon’s group has also prepared MoS2 /Al2 O3 [113]. Through the use of ultrasound, higher loadings of Mo can be achieved, resulting in a more active catalyst. A CoMoS/Al2 O3 catalyst has also been prepared by combining sonochemical and CVD techniques [114]. The effect of ultrasound on the gas–solid heterogeneous catalytic decomposition of cumene to benzene and propylene was examined with a silica–alumina cracking catalyst where the entire reaction bed was subjected to ultrasound [115]. Rate improvements of up to 160% were observed. Because cavitation cannot occur in such a system, these results must come simply from improved mass transport between the gas and surface. Loss of activity during extensive use is a common industrial problem with any catalyst, but especially with supported metal catalysts. The deactivation process varies depending on the catalyst and conditions of use and includes coking, oxidation of metal surfaces and neutralization of surface acid sites. There is an extensive patent literature over the past 20 years describing the use

References

of ultrasound to regenerate spent catalysts. Although the mechanism of action has not been examined, it is likely that improved mass transport and increased fine-pore penetration are significant contributors. The selectivities of these systems were greatly enhanced. An early disclosure of the use of ultrasound to reactivate a deactivated hydrocarbon conversion catalyst goes back to Exxon Research and Engineering in 1978 [116]. Highly deactivated hydrocracking catalysts could be reclaimed by oxidizing the catalyst at elevated temperatures followed by ultrasonic irradiation of the catalyst in a non-reactive liquid. A variety of similar applications of ultrasound to clean or reactivate various catalysts have also been reported. The most common carrier/cleaning liquid phase has been either aqueous [117] or standard feedstock flow. Commercial noble metal catalysts supported on alumina used either for NOx removal or hydrogenation of hydrocarbons have been regenerated efficiently with ultrasound [118, 119] with nearly complete restoration of specific surface area, porosity and activity. In the same manner, substantial regeneration has been disclosed for deactivated TiO2 −V2 O5 catalysts for oxidation of o-xylene to phthalic anhydride [120] and for flue gas denitration [121]. Ultrasonic reactivation is also useful for a partially deactivated BF3 –graphite intercalate catalyst used in an alkylation process [122]. Polymerization Catalysts Ultrasound is commonly used to accelerate rates of polymerization or to modify the structure of existing polymers [5, 123]. There are, however, only a few examples of the use of ultrasound to modify heterogeneous catalysts for polymerization. The first was in a 1961 patent [124], which found a substantial decrease in catalyst particle size and a consequent increase in activity due to diminished aggregation of Ziegler–Natta catalysts [125]. Further investigations of Ziegler–Natta polymerization under high-intensity ultrasound of styrene using a TiCl4 −Et3 Al catalyst have been published [126, 127]. The polymers are produced in better yield and with more control over the molecular weight distribution than in the conventional, unsonicated process. For example, polystyrene produced under the same conditions as stirring yields polydispersities above 10, whereas with ultrasound they are 2.5, with comparable mean molecular weights of ∼50 000. In part, this may be due to preferential cleavage of the longer chains by the ultrasound [128]. 8.5.2.6

8.5.3

Conclusion

In principal, ultrasound is well suited to industrial applications. Since the reaction liquid itself carries the

2015

sound, there is no barrier to its use with large volumes. In fact, ultrasound is already heavily used industrially for the physical processing of liquids, such as emulsification, solvent degassing, solid dispersion and sol formation. It is also extremely important in solids processing, including cutting, welding, cleaning and precipitation. Ultrasonic spray pyrolysis (USP), an aerosol synthesis technique, is also commonly used in industry for the production of fine powders. Recently, USP has been used to generate catalytic materials [129], although cavitation plays no role in such syntheses other than contributing to the formation of the liquid aerosol. Ultrasound has already become a common laboratory tool for nearly any case where a liquid and a solid must react. The production of heterogeneous catalysts involves high value-added materials, where processing costs are not always economically limiting. In this context, ultrasound is a viable method for the preparation and treatment of heterogeneous catalysts. The ability of ultrasound to create highly reactive surfaces, to improve mixing even in viscous media and to increase mass transport makes it a particularly promising technique to explore for catalyst preparation, activation and regeneration. References 1. L. A. Crum, T. J. Mason, J. Reisse, K. S. Suslick (Eds.), Sonochemistry and Sonoluminescence, NATO ASI Series C, Vol. 524, Kluwer, Dordrecht, 1999, p. 424. 2. K. S. Suslick (Ed.), Ultrasound: Its Chemical, Physical and Biological Effects, VCH, New York, 1988, p. 336. 3. T. J. Mason (Ed.), Advances in Sonochemistry, Vols. 1–6, Elsevier, New York, 1990–2001. 4. T. J. Mason, J. P. Lorimer, Applied Sonochemistry, WileyVCH: Weinheim, 2002, p. 314. 5. K. S. Suslick, G. Price, Annu. Rev. Mater. Sci. 1999, 29, 295–326. 6. K. S. Suslick, D. A. Hammerton, R. E. Cline Jr., J. Am. Chem. Soc. 1986, 108, 5641–5642. 7. E. B. Flint, K. S. Suslick, Science 1991, 253, 1397–1399. 8. K. S. Suslick, K. A. Kemper, in Bubble Dynamics and Interface Phenomena, J. R. Blake, N. Thomas (Eds.), Kluwer, Dordrecht, 1994, pp. 311–320. 9. W. B. McNamara III, Y. Didenko, K. S. Suslick, Nature 1999, 401, 772–775. 10. Y. Didenko, W. B. McNamara III, K. S. Suslick, J. Phys. Chem. A 1999, 103, 10783–10788. 11. W. Lauterborn, H. Bolle, J. Fluid Mech. 1975, 72, 391–399. 12. C. M. Preece, I. L. Hansson, Adv. Mech. Phys. Surf. 1981, 1, 199–254. 13. K. S. Suslick, S. J. Doktycz, Science 1990, 247, 1067–1069. 14. T. Prozorov, R. Prozorov, K. S. Suslick, J. Am. Chem. Soc. 2004, 126, 13890–13891. 15. K. S. Suslick, S. J. Doktycz, in Advances in Sonochemistry, T. J. Mason (Ed.), Vol. 1, JAI Press, New York, 1990, pp. 197–230. 16. K. S. Suslick, M. L. H. Green, M. E. Thompson, K. Chatakondu, J. Chem. Soc., Chem. Commun. 1987, 900–901.

2016

8.5 Sonocatalysis

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9

Laboratory Testing of Solid Catalysts 9.1

Laboratory Catalytic Reactors: Aspects of Catalyst Testing1

In the development of catalysts for new processes or the improvement of existing catalytic systems, various stages can be distinguished, as exemplified by Fig. 1. This development process covers the whole range from the new idea for a process or catalyst via catalyst preparation, catalyst screening, establishing reaction networks, kinetic studies and life tests to scale-up to pilot-plant level before a new or modified process is introduced. During the development, there will be continuous feedback from one activity to another to optimize the catalyst and/or the process. The number of catalyst formulations will decrease during development, but the equipment size and man-hours required increase and thus the costs involved. This demands for an efficient and proper approach for laboratory-scale experimentation. The objectives of the different development stages vary:

automation and miniaturization of the test equipment. This field is often referred to as ‘‘combinatorial chemistry’’ or ‘‘high-throughput experimentation’’. It started in the pharmaceutical industry, but in the past decade also has proven to be a useful approach in catalyst development studies. • Establishing the reaction network by the wealth of techniques at the disposal of catalyst researchers gives insight into how the catalyst works and provides the basis for the kinetic modeling studies. • The time-consuming kinetic studies are indispensable for process design, operation and control. A description is needed of the reaction rate as a function of the process variables, i.e. temperature, pressure and composition of the reaction mixture. • Life studies are intended to test the catalysts during a longer time-on-stream, often on a bench or pilot scale, with real feed and recycle streams. The latter are included to investigate the effect of trace impurities or accumulated components, not observed in laboratoryscale experiments. It is often desired to test the catalyst in the shape used for practical application. Here the need exists for experimental results that can be directly linked to commercial applications.

• Screening must provide the first, rough data on the activity and selectivity of the various catalyst formulations as a function of their composition and preparation and pretreatment history. As there are many variables, a large number of catalysts should be screened at a high throughput rate. The catalysts can then be compared based on their activity per unit mass, active phase or volume, depending on the specific goals. Often a first insight into the deactivation behavior, i.e. The catalyst stability, is obtained simultaneously. In this stage, potentially interesting catalysts are identified. Because of the large number of measurements that must be carried out, there is a continuing trend towards

Although they may be part of a catalyst testing [1–3] program, investigations focused on revealing the reaction mechanism, e.g. in situ FTIR in transmission or reflection mode, NMR, XRD, EXAFS, XPS, EM, ESR, UV–visible, and the reaction cells used are not treated in this chapter. For the correct interpretation of the results, however, this chapter may also provide a worthwhile guide. Various types of reactors can be applied and it is of primary importance to select the proper reactor to obtain the required information, which generally means that one should not mimic the reactor type in which the process will be carried out (no ‘‘Dinky

1 A list of symbols used in the text is provided at the end of the chapter. ∗ Corresponding author.

References see page 2043

Freek Kapteijn and Jacob A. Moulijn∗

9.1.1

Introduction

Handbook of Heterogeneous Catalysis, 2nd Ed. .. .. Edited by G. Ertl, H. Knozinger, F. Schuth, and J. Weitkamp Copyright  2008 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim ISBN: 978-3-527-31241-2

2020

9.1 Laboratory Catalytic Reactors: Aspects of Catalyst Testing

Mechanism Combinatorial stage

Screening

n tio iza tim Op

Prepration

Kinetics

Quantification stage

Reaction network Kinetics Increasing: Time Money Reality Fig. 1

Stability

Catalytic reaction engineering

Stability tests Scale-up

Stages in a catalyst development program.

Reactor engineering

Transport phenomena Catalyst

Toy’’ approach). The criteria for the selection of a reactor for catalyst testing are different from those for the selection of an industrial reactor. Scaling down as far as possible is desirable for reasons of lower equipment costs, lower materials consumption, less waste formation, lower utility requirements, reduced demands on laboratory infrastructure and increased intrinsic safety (reduced hazards of toxic emissions, explosions and fires). A smaller scale, however, requires a more accurate experimentation and the use of representative (catalyst) samples. In the laboratory, catalysts are often tested using a packed bed contained in a reactor through which a reaction mixture flows (Fig. 2). On the reactor level the fluid seems to flow as a front through the bed (‘‘plug flow’’), but due to the flow around the catalyst particles mixing (‘‘dispersion’’) may occur on the particle size level. This may affect the reactor performance. Along the length of the bed, reactants will be gradually converted. Reactants have to diffuse from the bulk of Mixing dispersion

Plug flow

Diffusion Reaction Transport phenomena

Transport phenomena in a catalytic packed bed reactor on different levels.

Fig. 2

Fig. 3

Catalytic reaction engineering.

the fluid through a stagnant layer around the particles to the exterior surface and subsequently through the pore system to an active site where they can react. Products proceed in the opposite way. Heat produced or consumed will have to be removed or supplied at a sufficient rate, otherwise temperature gradients may develop locally to an undesired level. The description of these phenomena is the subject of catalytic reaction engineering [4–10], which integrates catalysis, reactor analysis and transport phenomena to aid in procuring and interpreting data on catalyst activity, stability, kinetics and reaction mechanisms (Fig. 3). With the goal of obtaining intrinsic catalyst properties (reaction kinetics and selectivities) from experimental data without being disguised by the above-mentioned phenomena, the following conditions should be fulfilled: • effective contact between reactants and catalyst • absence of mass and heat transport limitations inside and outside the catalyst particles • good description of reactor characteristics, with welldefined residence time distributions under isothermal conditions (ideal systems). Generally, this implies the use of ideal reactor types such as the plug-flow reactor (PFR) and the continuously stirred tank reactor (CSTR). By-passing of part of the catalyst by channeling in a packed bed or uneven flow distributions must be avoided. In three-phase systems (gas–liquid–solid), the even distribution of both fluid phases over the catalyst is crucial, while the gas to liquid mass transfer should not be limiting.

9.1.2 Reactor Systems

In this chapter, criteria will be derived that can be used to check whether one operates under conditions that result in deviations of not more than 5% in reaction rate from the ideal situation. This can be expressed as follows: rateobserved = 1 ± 0.05 rateideal

(1)

The various aspects that are to be considered to achieve a proper and efficient catalyst testing approach are presented for heterogeneous systems in which the catalyst is the solid phase and the reactants are in the gaseous and/or liquid phase. The presence of a solid phase introduces complicating phenomena on which this chapter focuses. The solid catalyst can be present in a variety of forms. Fixed-bed configurations are often the most convenient. They can be based on random packings or on structured packings, such as washcoated monoliths. Systems in which the catalyst is not in a fixed position are also widely applied. The most commonly used reactors of this type in catalyst testing are slurry reactors.

2021

Classification Catalytic reactors can globally be classified according to their mode of operation: under steady-state or transient conditions, as indicated in Fig. 4, or according to the contacting/mixing mode, as indicated in Fig. 5 for a fluid–solid system. In both schemes good contact between catalyst and reaction mixture is assumed. Figure 4 represents a classical classification based on mode of operation. Steady-state reactors, especially packed-bed reactors, are most widely used in catalyst testing, predominantly because of the ease of operation and the low costs. Transient operation is less common and has some disadvantages for the mere goal of catalyst testing. In batch reactors, possible deactivation during the extent of the experiment cannot be established, except by verification afterwards by repeating the experiment. Pulse reactors, such as the TAP (temporal analysis of product) and the Multitrack, operate under catalyst conditions that are completely different from those in steady-state operation. An exception is transient operation by using isotopically labeled species under steady-state 9.1.2.1

9.1.2

Laboratory catalytic reactors

Reactor Systems Transient

Steady state

Various laboratory reactors have been described in the literature [3, 11–14]. They depend mainly on the type of reaction system that is investigated: gas–solid (GS), liquid–solid (LS), gas–liquid–solid (GLS), liquid (L) and gas–liquid (GL) systems. The first three are intended for solid or immobilized catalysts, whereas the last two refer to homogeneously catalyzed reactions, which will not be discussed here. The presence of two reaction phases (gas and liquid) should be avoided as much as possible for ease of data interpretation and experimentation. Premixing and saturation of the liquid phase with gas can be an alternative in this case. One must be sure that in the analysis, the catalyst does not further affect the composition of samples taken from the reactor contents or product stream.

Fig. 5

Batch

Continuous flow

Mixed flow

Plug flow Integral

Semi-batch

Discontinuous Step

Pulse

Differential

Single pass

Recycle

External Riser reactor

Fluidization Internal

Berty reactor

Thermobalance Packed bed

Fluid bed Slurry

TAP Multitrack

Classification of laboratory reactors according to mode of operation.

Fig. 4

References see page 2043

1

2

3

4

5

6

7

8

Classification of laboratory reactors according to contacting mode.

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9.1 Laboratory Catalytic Reactors: Aspects of Catalyst Testing

conditions, such as the positron emission profiling technique (see Chapter 6.5). Transient operation is mainly applied for obtaining mechanistic information and for establishing reaction networks [15], but the steadystate isotope transient kinetic analysis (SSITKA) [16–21] approach yields both steady-state information and that of individual reaction sequences. The classification in Fig. 5 serves for the description of the reactors used. Here two ideal contacting types are used, the plug-flow mode and the ideally mixed mode, both for the fluid and the solid phase. By application of the design equations of these ideal reactor types, the experimental results are interpreted in a straightforward manner. For two phases, two contacting types and two operation modes (batch and flow), eight combinations arise. Some laboratory reactors are easily recognized; the packed-bed plug-flow reactor (PFR) (type 1), the internalor external-recirculation reactor (CSTR) with fixed catalyst bed (type 2), the (circulating) fluid-bed reactor (FBR) (type 3), the fluid bed with recycle (type 4), the slurry reactor (type 4), the riser reactor (type 5), the riser reactor with recycle (type 6), the (circulating) fluid bed reactor with continuous catalyst feed (type 7) and the slurry reactor or fluid bed with recycle of fluid and continuous catalyst feed (type 8). Types of Laboratory Reactors There is usually no single best laboratory reactor that could be used for all types of reactions and catalysts. In this section, we discuss the various types of reactors that can be chosen for catalyst screening and/or obtaining the kinetic parameters for a specific reaction system. Table 1 gives 9.1.2.2

Tab. 1

a generalized comparison of frequently used laboratory reactors. More advanced laboratory reactor systems are discussed in Section 9.1.8. 9.1.2.2.1 Gas–Solid and liquid–Solid Reactions The most commonly used (‘‘the workhorse’’) and simplest type of laboratory reactor for gas–solid reaction systems is the packed-bed tubular reactor (Fig. 6). It consists simply of a reaction tube in which the catalyst is held between plugs of quartz wool on a sintered frit or wire mesh gauze. Depending on the reaction conditions, the tube consists of glass, quartz, steel or ceramic (SiC, alumina). For low temperature and pressure glass can be used, high pressure requires steel, high temperature requires quartz or ceramics. The latter are especially suited for high-pressure–high-temperature applications [22] where most steel alloys lose their strength. Its optimal internal diameter amounts to 4–6 mm, predominantly to ensure good heat transport to or from the catalyst [23]. The reactor may have a straight or U-shape or a concentric geometry, with the catalyst in the inner tube [22]. The porous catalyst may be diluted with an inert material with good heat conductivity (SiC) to improve heat transfer and to satisfy plug flow criteria; see Section 9.1.2.5. An inert material may be placed before and after the bed for preheating the fluid phase and avoiding fluidization in up-flow configurations. The packed-bed reactor can also be used for liquid–solid reactions. For low conversions, the packed-bed reactor is operated in the differential mode, and for high conversions in the integral mode. By recirculation of the reactor exit flow one can approach a well-mixed reactor system, the CSTR.

Summary of relative reactor ratings (L = low, M = medium, H = high)

Aspect

Ease of use Ease of construction Cost Ease of sampling and analysis Approach to ideal type Fluid−catalyst contact Isothermicity Temperature measurement Kinetics Deactivation noticed GLS use

Reactor type PFR Differential fixed bed

PFR Integral fixed bed

PFR Solids transport (riser)

CSTR External recirculation

CSTR Internal fluid recirculation

CSTR Spinning catalyst basket

Batch Internal fluid recirculation

H H L M

H H L H

H L H L

M−H M L−M H

M M M−H H

L−M L−M M−H H

M M M M

H H H H

H H M−H H

M M H H

H H M−H H

M−H M−H H M−H

L−M L−M M L

H H H H

H H L−M

H H M−H

M L L

M M M−H

M−H M M−H

L−M M M

M L H

9.1.2 Reactor Systems

2023

Union

O-ring seal

Ceramics coupling

Nut

Swagelok coupling

Inert Catalyst

Quartz wool

Thermocouple

Quartz tube U-shaped Fig. 6

Concentric

Packed bed reactor designs. Examples of geometries and connections.

Internally mixed

Externally mixed

(a) Berty-type

(b)

Carberry-type

Basket designs for Carberry-type reactors: (a) flat blade; (b) pitched blade.

Fig. 8

Schematics applications.

Fig. 7

of

CSTRs

for

heterogeneous

catalysis

An interesting concept is the catalytic wall reactor, in which – as the name suggests – the catalyst is applied to the reactor wall [24, 25]. This type of reactor is particularly suitable for kinetic studies of fast, highly exothermic (or endothermic) reactions, because rapid heat removal (or addition) is possible, ensuring isothermal conditions. A disadvantage of the catalytic wall reactor may be the existence of radial mass transport limitations, which can partly be solved by filling the reactor tubes with inert particles of good heat conductivity [25]. Various designs exist to approach the ideal CSTR or gradientless reactor (Fig. 7) for gas–solid and liquid–solid reactions [4, 11, 26, 27]. A number of configurations are commercially available. Two approaches are followed: either the catalyst is placed in baskets and rotated, the Carberry type [4], or the fluid phase is circulated internally at a high rate, the Berty type [11], or externally by means of a gas pump (membrane, piston, plunger or centrifugal types) or liquid pump [11]. A major problem in gas–solid reactors is achieving good contact between the two phases.

Various basket designs have been proposed to optimize gas–catalyst contact. Figure 8 shows the original Carberry basket in the form of a four-bladed stirrer and a more recent pitched-blade stirrer, claimed to have good mixing properties [7, 26, 28]. The temperature of the catalyst bed cannot be measured in the Carberry-type reactor and highly endo- or exothermic reactions are to be avoided. Figure 9 shows an example of the Berty-type reactor, the RotoBerty reactor. The catalyst is packed in the inner tube to avoid local maldistribution and catalyst by-passing. Monolithic catalysts are ideal for use in these reactors because of their intrinsic low pressure drop. Various stirrer designs exist to create good mixing and even twostage impellers are applied. These stirrers, with speeds up to 10 000 rpm, are usually actuated by feedthrough of the stirrer axis or by magnetic coupling. Since most magnets have their Curie point, i.e. lose their magnetism, at relatively low temperatures, this coupling occurs outside the oven section. Feedthroughs may result in leakages, References see page 2043

2024

9.1 Laboratory Catalytic Reactors: Aspects of Catalyst Testing

Riser oven Products

Catalyst N2

Oil heating

Oil feed

Preheating catalyst Fig. 10 Schematic of a laboratory-scale microriser reactor set-up for FCC investigations.

Two-stage rotor

Shaft Feed

Fig. 9

Product

Layout of the RotoBerty reactor.

and magnetic coupling may give rise to non-ideal mixing due to the presence of areas without convection. As a result of the low conversion over the catalyst bed, isothermal conditions are approached more easily than in packed-bed reactors. In LS systems, mixing is less problematic and the catalyst is not necessarily held in a packed bed, but can be suspended in the fluid with measures to prevent washing out of the reactor and accumulation at the exit: the slurry reactor (see Section 9.1.2.2.2). Compared with the simple packed-bed reactor, the reactor volume of CSTRs is much larger, often >100 mL, while the amount of catalyst per unit volume is at least two to three times smaller. This implies that on changing the conditions, one should take into account the physical residence time of the reactor before considering the system to have reached a new steady state. This may be appreciable, depending on the pressure and feed flowrates used. Also, deactivation of the catalyst is not so directly noticed for the same reason. It is self-evident that all CSTR systems may operate as batch reactors, in which the composition changes as a function of time (transient operation), when the inlet and outlet lines are closed. An example is the determination of the cis–trans equilibrium for 2-butene by means of the metathesis reaction in an externally recirculated (membrane pump) system [29].

A disadvantage of batch operation is that catalyst deactivation cannot be followed, unless the experiment can be repeated with the same catalyst sample. A special case in kinetic studies is the modeling of catalyst deactivation. For moderately decaying systems, sequential experimental design is recommended (see Chapter 6.1). Here, new experiments are carefully planned on the basis of previous results, using a packedbed reactor. In the case of rapidly deactivating catalysts, these have to be fed continuously to the reactor together with the feed. A relatively new development is the laboratory riser reactor for fluid catalytic cracking (FCC) [30, 31]. It consists of a folded tubular reactor to which the catalyst and evaporated reactants are co-fed and both travel through the system in plug flow (Fig. 10). Various reactor elements can be connected to vary the length. This system results in more representative data than the so-called micro-activity test (MAT) of ASTM [3], in which a certain amount of liquid feed is pushed through a packed bed of catalyst. The reason is that the time-scale of the reactions and especially that of coke deposition – in the order of milliseconds – is much better mimicked in the riser reactor than in the fixed-bed reactor. Another approach for FCC kinetic studies was described by Kraemer et al. [32]. They use an internally recirculated reactor containing the catalyst in a basket, in which it is fluidized by the recirculating gas. A small pulse of reactant is allowed to react for a short time, after which the reactor is purged. This procedure is repeated a number of times to map the catalyst activity as a function of time. The aim is to give a mapping along the length of a riser reactor. 9.1.2.2.2 Gas–Liquid–Solid Reactions Under conditions of good wetting (see Section 9.1.2.5 and Ref. [33]), a packed-bed reactor can be used in trickle flow, i.e. with downflow of gas and liquid, for gas–liquid–solid systems.

9.1.2 Reactor Systems

Laboratory trickle-flow reactors are not very good tools for quantifying reaction kinetics, because fluid dynamics and reaction kinetics are so closely interlinked that their effects on conversion are practically inseparable [34]. Stirred-tank reactors are the predominant reactors used for studying heterogeneously catalyzed gas–liquid reactions. Stirred-tank slurry reactors (usually autoclaves), in which small catalyst particles (1–200 µm) are suspended in the liquid either by mechanical agitation or by the gas flow, can be operated both batchwise and (semi-)continuously. Temperature control is relatively easy due to the large amount of liquid present. Also, the liquid–solid contact area is relatively large and the intraparticle diffusion distances are small. Disadvantages are the potential non-uniform distribution of the catalyst particles in the reactor and the poor gas–liquid mass transfer. Also, filtering the catalyst from the liquid mixture may be problematic. A commonly used stirred-tank reactor system for studying gas–liquid and gas–liquid–solid reactions is the autoclave equipped with gas-inducing impeller. The gas is sucked in through the impeller, creating gas bubbles in the liquid [35]. Most laboratory units are dead-end autoclaves with only feed of the reactant gas by keeping the pressure constant (semi-batch operation). This implies that it is relatively difficult to determine gas–liquid mass transfer [36–39] (see also Section 9.1.3.4). An elegant solution for internally mixed reactors that avoids the use of feedthroughs or voluminous magnetic couplings is the swinging capillary stirrer [40] (Fig. 11). A capillary is welded to the lid of the reactor and sealed at the bottom and equipped with a stirrer blade. From the top a bent rod (piano string) is inserted in this capillary and turned around. This yields a swinging movement and effective mixing and can be used in LS and GLS reactions under high pressures in autoclave reactors. The compact design reduces the reaction volumes to 10 mL and enhances the safety. An alternative to slurry reactors for solid catalyst particles with sizes suitable to be retained by a wire mesh (>1 mm) are the basket reactors discussed in Section 9.1.2.2.1, but adapted for gas–liquid–solid reactions. A popular gas–liquid–solid gradientless laboratory reactor is the Robinson–Mahoney reactor (Fig. 12), in which the gas–liquid dispersion is forced through a stationary catalyst bed. Impellers draw fluid into the center of an annular catalyst basket. The reactor is suitable for high-pressure and high-temperature reactions. Two alternative types of internal recirculation reactors have been proven successful for kinetic investigations: the turbine reactor, in which the gas–liquid mixture is transported back to the top by turbine blades located in the

Fig. 11

2025

Schematic of a microautoclave with swinging capillary

stirrer.

outer annulus of the reactor (Fig. 13) [41], and the screw impeller stirred reactor, having the screw in the central shaft (Fig. 14) [42]. Design Equations To relate the reaction rate or conversion, pressure drop and temperature variation in a catalytic reactor with the operating variables of a reactor, flow-rate, catalyst amount etc., so-called mass, heat and momentum balances are being used in catalytic reaction engineering [4, 8]. In this chapter, it is assumed that the catalyst bed is isothermal 9.1.2.3

References see page 2043

2026

9.1 Laboratory Catalytic Reactors: Aspects of Catalyst Testing

Fig. 12 Schematic of a Robinson–Mahoney reactor (Autoclave Engineers).

and the pressure drop over the bed is negligible. This leaves only mass balances for each reactant or product to be considered. For a component i this mole balance can be written for part of or the whole catalyst bed as inputi − outputi + productioni = accumulationi

(2)

where the production term refers to both production (positive) and consumption (negative) contributions. Under steady-state conditions, the accumulation term vanishes, simplifying the resulting expression. This approach yields the following expressions for a component i and a single reaction: PFR :

CSTR :

dxi
 = −νi r W d Fi0 W Fi0

=

xiout − xiin −νi r

(3)

(4)

As the conditions in the CSTR are the same everywhere, the mass balance is considered over the whole reactor, resulting in an algebraic expression relating the reaction rate directly to the measured conversion. For the PFR, only a small slice of the bed can be considered to have constant conditions, leading to a differential equation which describes the conversion as a function of the amount of catalyst. For low conversions, the rate can be

Fig. 13 Schematic of turbine reactor for multiphase kinetic measurements.

considered to be constant over the bed and the expression can be simplified to that of a CSTR (differential reactor). At high conversion, the rate varies over the bed length and the reactor operates under integral conditions. Variation of the space time W/Fi0 yields the corresponding conversion data. By numerical differentiation, rates can be estimated that can be compared for different catalysts. Differential conditions at high conversions, especially of interest for practical applications, can be achieved by using an upwind reactor to attain the desired conversion level, followed by a differential reactor. This requires highly accurate analyses to determine the minor composition changes over the catalyst bed. In general, catalyst activities are compared by comparison of the conversion levels at equal space times, preferably at low conversion levels (‘‘initial reaction rates’’). At high conversions, the reaction rate is much less sensitive and wrong conclusions may be drawn, especially with respect to selectivities (see later).

9.1.2 Reactor Systems

2027

equation (=W/Fi0 ). This exponential behavior can easily be verified experimentally. In recirculation reactors, frequently a recirculation ratio Rc (ratio of the flow returned to the inlet and the flow leaving the reactor) larger than 20 is recommended [11, 26]. This criterion is not complete, as is apparent from literature [43–45]. One can demand that the rate over the catalyst bed may not change more than 5%, thus for an nth-order isothermal reaction n
r(cin , T in ) cin rate at inlet of bed = = rate at outlet of bed r(cout , T out ) cout  n x = 1+ (1 − x)(1 + Rc ) = 1 ± 0.05

(6)

which yields the following criterion: x 0.05 < (1 − x)(1 + Rc ) n

Fig. 14 Schematic of screw impeller stirred reactor for multiphase kinetic measurements.

The CSTR and PFR are the two extremes of continuous reactor types, one with extremely good mixing and the other without. A number of criteria or methods exist to judge whether one can use one of these reactor models. CSTR The most commonly used CSTRs in heterogeneous catalysis are recirculation reactors (gas–solid, liquid– solid, gas–liquid–solid systems) and well-stirred vessels (liquid–solid, gas–liquid–solid systems). Injection of a δ-pulse of a tracer should produce an exponentially decaying output concentration: 9.1.2.4

c(t) = c0 exp(−t/τ )

This criterion implies that the necessary recirculation ratio is not a fixed value, but depends on the reaction under consideration and the conversion level. At low conversions, the recirculation rate does not need to be high according to this criterion [of course, good mixing, Eq. (5), is still required to avoid dead zones in the reactor] and in fact a differential PFR model can be used. At high conversions, the recirculation rate must increase. It can easily be seen that a recirculation ratio of 20 limits the conversion for a first- and second-order reaction to 50% and 25%, respectively. The problem with verification of these criteria is the calculation of the recirculation ratio. No general relation is valid and it must be estimated for each reactor configuration by experimentation; see, e.g., Refs. [26, 44]. PFR PFR systems can be described by an infinite number of CSTRs in series. A packed bed can be considered as a finite number of CSTRs, each corresponding to a certain height of a slice of the catalyst bed, also referred to as the equivalent height of mixing. Reducing the bed length will increase the deviation from plug flow towards more axial mixing. This is also described by the axial dispersion model, in which the dimensionless P´eclet number, Pe, is used as parameter. The relation between Pe and the number N of mixers in series for large N is 9.1.2.5

(5)

where τ represents the average physical residence time in the reactor system (=Ntot,R /Ftot,out ), which should not be confused with the space-time used in the mass balance

(7)

2N = Pe =

Lb Lb u = Pep Dax dp

References see page 2043

(8)

9.1 Laboratory Catalytic Reactors: Aspects of Catalyst Testing

For a sufficiently close approach to ideal plug flow, N or Pe should exceed a certain value that depends on the degree of conversion and on the reaction order. Gierman [46] refined the criterion of Mears [47] and arrived at Eq. (9), which can also be used for monoliths. In the latter case, the axial dispersion coefficient should be replaced by the smaller molecular diffusivity, thus relaxing the criterion for monolithic elements as compared to packed beds.   1 (9) Pe > 8n ln 1−x Evaluation of literature data on the correlation between the particle P´eclet number (also referred to as the Bodenstein number) and the particle Reynolds number for single-phase and trickle flow yielded Fig. 15. At low Rep ( ln (10) dp Pep 1−x Taking Pep = 0.5 for the low Reynolds region of interest for laboratory-scale operation, it is possible to calculate the minimum reactor length for a given reaction order, conversion and particle size. Alternatively, a maximum allowable particle diameter can be calculated for a reactor of a given length. Figure 16 gives the maximum allowable particle diameters for a first-order reaction and various reactor lengths. For trickle-flow reactors the reactor lengths should be 10 times larger at the same particle diameter, due to the lower Pep value for this case. In PFRs, not only the reactor length but also the diameter plays a role in its plug-flow performance. Close

100

Pep

Single phase flow 10−1

10

dt = 50 mm

dp,max / mm

2028

dt = 5 mm Lb = 10 mm 0.1

dt = 1 mm Lb = 1 mm 0.01 0.01

10−1

100

101

102

103

Rep Fig. 15 Relation between the P´eclet particle and Reynolds particle numbers for single-phase and trickle flow. Adapted from Ref. [46].

1

Fig. 16 Maximum allowable particle diameter as a function of (1 − x) for plug-flow behavior (first-order reaction, single phase) at different bed lengths and for a flat velocity profile for different reactor diameters.

to the wall the packing density is lower than that in the interior because of the presence of a flat wall surface. The higher voidage implies a lower resistance to flow and higher local velocities are to be expected. On the other hand, the radial mixing effects in the packed bed may relax this influence. To avoid these phenomena, as a rule of thumb the following condition should be satisfied [48]: dt > 10 dp

(11)

The maximum allowable particle diameters for typical reactor diameters are included in Fig. 16. Especially at low conversion levels, the last criterion is more severe at a given reactor length, while with increasing conversion the axial dispersion criterion should be satisfied. In ideal trickle-flow reactors, all particles in the catalyst bed take part in the overall reactor performance, since each is surrounded, ‘‘wetted’’, by the liquid phase that flows around it. Situations in which the liquid preferentially flows through a certain part of the bed, while the gas phase predominantly flows through another part, should be avoided [34]. In such cases, part of the bed is not contacted by the liquid reactant at all and does not contribute to the overall conversion. To avoid this maldistribution, Gierman [46] proposed the following criterion for the wetting number Wtr for concurrent downflow operation: Wtr =

10−2

0.1

(1– x)

Trickle flow 10−2 −3 10

Lb = 100 mm

1

vl u l > 5 × 10−6 dp 2 g

(12)

The main variables that dictate the uniformity of catalyst irrigation are the liquid velocity, the particle diameter and the kinematic viscosity of the liquid. Figure 17 shows the maximum allowable particle diameter for several reactor

9.1.2 Reactor Systems

of the reaction rate constant from the ideal situation:

10

≡ Lb = 1000 m

1

=

Lb = 100 m Lb = 10 m

0.1 10−7

10−6

10−5

nl /m2 s−1 Fig. 17 Maximum allowable particle diameter as a function of the kinematic viscosity for complete wetting in trickle-flow reactors.

lengths as a function of the kinematic viscosity at constant liquid hourly space velocity. From the foregoing, it will be clear that the particle size is a major factor to be considered in the application of small reactors. Obviously, catalysts in the commonly used form of extrudates of e.g. 1–3 mm in diameter can only be tested in a large reactor at the pilot-plant scale. Although these particles can be crushed to the size suitable for testing in laboratory reactors, this is not always desired. Sometimes performance data are needed under conditions where internal diffusion gradients exist, to be able to predict industrial operation. The catalyst may also be of the eggshell type or have another defined concentration profile of the active phase that is not to be destroyed. A way to satisfy the criteria of small particle size and yet to use the full-size catalyst is to dilute the large catalyst particles with fine inert particles. In this way, the hydrodynamics and kinetics are decoupled [34]. The fine particles dictate the hydrodynamic performance, i.e. The plug-flow and wetting behavior, while the catalyst particles determine the kinetic behavior. Needless to say that bed dilution is also applicable in the case of small catalyst particles to obtain a sufficient bed length. This dilution technique has been successfully applied for years in hydrotreating research using trickle-flow reactors [34]. The optimal and most reproducible way of packing a reactor with large and fine particles is according to the dry method. First, the reactor is packed with the large particles, followed by packing this prepacked bed with the fine particles while vibrating the reactor, using a sufficiently low flow-rate to avoid segregation [49]. One should be careful, however, with the degree of dilution. A too dilute system, especially in the case of equal sized particles, may lead to uneven distribution and bypassing of the catalyst. Berger et al. [50] evaluated this problem and derived an expression to estimate the deviation ()

xundiluted bed − xdiluted bed b nxdp ≈ xundiluted bed 1 − b 2Lb bnxdp < 0.05 2L0

(13)

where L0 is the undiluted bed length and b the volume fraction of diluent. This criterion is valid at low conversion (x < 0.5) and reaction orders n ranging from 0 to 1. It imposes constraints on the maximum particle size as a function of bed dilution and vice versa, as represented graphically in Fig. 18 for various bed lengths. In general, samples should not be diluted more than 5–10-fold. In kinetic investigations, the measured conversions are used to calculate the rate coefficient, e.g. to determine the apparent activation energy of the catalytic reaction. For an irreversible first-order reaction, the reaction coefficient can be calculated directly from the conversion using   1 1 (14) ln k= 1−x (W/Fi0 ) pi,0 Therefore, in most cases it is more useful to calculate the deviation of this k instead of the deviation of xi . If this is done using Eq. (13), the following expression for the deviation of the rate coefficient (k ) is obtained, defined similarly to  in Eq. (13): kundiluted bed − kdiluted bed kundiluted bed   1 − x(1 − ) ln 1−x   = 1 ln 1−x

k =

(15)

10

dp,max / mm

dp,max / mm

2029

Lb = 1000 mm

1

Lb = 100 mm Lb = 10 mm

0.1

0.01 1

0.1

Fraction of catalyst (1 – b) Fig. 18 Maximum allowable particle diameter as a function of the catalyst fraction (1 − b) in a diluted bed; nx = 0.2. References see page 2043

2030

9.1 Laboratory Catalytic Reactors: Aspects of Catalyst Testing

This imposes limits on the particle size at large fractions of inerts. It is recommended, however, that the preparation method should be improved. Visual inspection of the sample batches often already indicates its (in)homogeneity.

0.99 Axial dispersion criterion 0.20

0.98

b

0.95

0.10

0.90

0.05

0.80 0.50

∆k = 0.005

0.00 0

0.2

0.4

0.6

Summary of Laboratory Systems Usually small amounts of catalyst are applied in laboratory reactors, ranging from 10 to 1000 mg, with gas flow-rates between 10 and 1000 mL min−1 (STP). This depends largely on the reactor type used. PFRs cover the whole spectrum with diameters ranging from a few to 20 mm, while CSTRs generally need more catalyst and higher flow-rates due to their dimensions. Using packed beds in the latter necessitates the use of larger particles to overcome the pressure drop problem to obtain high recirculation rates [42]. This is not always acceptable. In this respect, monolithic catalysts are easier to use in CSTRs due to their intrinsic low pressure drop. Some practical advantages/disadvantages of commonly used reactors are: 9.1.2.6

0.02 0.01

0.8

1

x Calculated effect of conversion and dilution on k , the relative deviation in the first-order rate coefficient calculated from the conversion, using Eq. (15) for an Lb /dp ratio of 100. Dashed lines represent the criterion for neglecting the effect of axial dispersion at Wcat /F 0 = 5.58 × 105 g s mol−1 and Wcat + Wdil = 400 mg.

Fig. 19

Figure 19 shows the results of this calculation for several selected values of k as a function of the dilution and the conversion for a constant Lb /dp = 100. The figure should be read as follows: an experiment at x = 0.5 and b = 0.95 (and Lb /dp = 100) will result in a deviation of the calculated rate coefficient due to the dilution (k ) of approximately 7%. The shaded area in Fig. 19 corresponds to conditions where the deviation exceeds the criterion of 5% deviation and it should be avoided in particular if the aim is to measure intrinsic reaction kinetics. Generally, one should try to avoid the combination of a high degree of bed dilution and high conversion levels. It should be noted that Eq. (13) and hence also Eq. (15) have not been validated at conversion levels above 0.75 and dilutions below 0.8, but they can be used as a safe rule of thumb. The plug-flow criterion for 5% deviation of the conversion due to axial dispersion is also included in Fig. 19. The small dark shaded area corresponds to conditions where this deviation exceeds the criterion of 5%. This criterion is easily satisfied due to the high Lb /dp ratio used as a consequence of the application of a high degree of dilution. This is not generally valid; dilution is used to satisfy this criterion, amongst other uses, but on the other hand it creates a deviation in the observed conversion. Hence there should be a dilution degree at which both criteria will yield similar constraints. This dilution degree depends strongly on the system in question. One should also be cautious with too much dilution for reasons of sample inhomogeneity. Catalyst batches may contain particles of different activities, due to incomplete wetting during preparation or, more frequently, during impregnation. If too few particles are being used, the sample may not be statistically representative.

• PFR: deactivation noted directly, small amounts of catalyst needed, simple, yields primarily conversion data not rates • CSTR: larger amounts of catalyst and flows needed, deactivation not immediately determined, direct rate from conversions • FBR: no ideal reactor behavior, continuous handling of solids possible • batch: catalyst deactivation hard to detect, yields quickly conversion and selectivity data over wide conversion range. More detailed discussions of laboratory reactors can be found elsewhere [3, 12, 13, 51]. 9.1.3

Mass and Heat Transfer

Due to the consumption of reactants and the production or consumption of heat, concentration and temperature profiles can develop in the stagnant zone around the catalyst particle and in the particle itself (Fig. 20). In this section, criteria are derived to ensure that the effect of these gradients on the observed reaction rate is negligible [4, 23, 52]. In gas–liquid–solid slurry reactors the mass transfer between the gas and liquid phase also has to be considered [9, 53]. Here, the focus will be on solid-catalyzed single-phase reactions. Extraparticle Gradients Extraparticle mass and heat transfer are most conveniently analyzed by means of the so-called film model, in which 9.1.3.1

9.1.3 Mass and Heat Transfer

orders and exothermic reactions. For an isothermal, nthorder irreversible reaction, this results in the following criterion:

Gas film

Ts T c

Cs

Ca < Tb

T

Cb

0.05 |n|

(20)

Analogously to mass transfer, the following equation is used for external heat transfer:

c

Bulk gas Exothermic

2031

− a  h(Tb − Ts ) = rv,obs (−Hr )

(21)

Endothermic

Combination with Eq. (17) yields Temperature and concentration gradients in and around a catalyst particle for exo- and endothermic reaction. Fig. 20

the flux from the bulk fluid to particle is defined in terms of mass and heat transfer coefficients kf and h, respectively. A mass balance over the film layer yields Ap kf (cb − cs ) = Vp rv

(16)

This expression shows that the mass transfer rate is proportional to the concentration difference over the film. The observed reaction rate can then be expressed as follows: rv,obs = r(cs ) = a  kf (cb − cs )

(17)

where a  = Ap /Vp is the specific particle area (m−1 ). No transport limitations exist if cs ≈ cb . But how can we determine the concentration at the surface cs ? In order to relate cs to observable quantities, a dimensionless number Ca, the Carberry number [8], is introduced as follows: Ca =

rv,obs a  kf (cb − cs ) cb − cs = =  a kf cb a  kf cb cb

(18)

where a  kf cb is the maximum mass transfer rate, which is obtained when the surface concentration equals zero (this corresponds to the best catalyst possible). Ca relates the concentration difference over the film to procurable quantities and is therefore a so-called observable [4]. A criterion for the absence of extraparticle gradients in the rate data can be derived from the definition of an effectiveness factor for a particle. This should not deviate more than 5% from unity as criterion: observed reaction rate rate at bulk fluid conditions rv,obs (cs , Ts ) = = (1 − Ca)n = 1 ± 0.05 rv,chem (cb , Tb )

ηe =

(19)

The − sign applies to positive reaction orders and endothermic reactions and the + sign to negative reaction

Ts − Tb kf cb (−Hr ) cb − cs = = βe Ca Tb hTb cb

(22)

where βe is called the external Prater number. This represents the maximum relative temperature difference over the film or the ratio between the maximum heat production and heat transfer rates. For rather general kinetics, the effectiveness factor definition can be approximated as rv (cs , Ts ) kv (Ts ) f (cs ) kv (Ts ) = · ≈ rv (cb , Tb ) kv (Tb ) f (cb ) kv (Tb )    Ea Tb = exp − − 1 = 1 ± 0.05 RTb Ts

ηe =

(23)

with the assumption that heat effects dominate the transport limitations, so for small Ca values. Series expansion of the exponential leads to a simple result:      Ea rv,obs  (−Hr )kf cb  |γb βe Ca| =  < 0.05 RTb hTb kf a  cb  (24) A result like this is rather logical since it contains the three groups that determine the overall process. The Carberry number determines the concentration drop over the film, the Prater number determines the maximum temperature rise or drop (exothermic or endothermic reaction) and the dimensionless activation energy expresses the sensitivity of the reaction towards a temperature change. Both criteria for extraparticle gradients contain observables and can be calculated based on experimental observations of reaction rates. For heat and mass transfer coefficients in packed beds, various correlations exist in terms of dimensionless numbers. Table 2 summarizes the most appropriate ones for laboratory reactors [5, 7, 54, 55]. Values of kf and h for gases in laboratory systems range between 0.1 and 10 m s−1 and between 100 and 1000 J K−1 s−1 m−2 , respectively. In References see page 2043

2032

9.1 Laboratory Catalytic Reactors: Aspects of Catalyst Testing Correlations to calculate mass and heat transfer coefficients in packed beds [5, 7, 54, 55] and monoliths [56–58]

Tab. 2

Packed beds: Mass transfer

Heat transfer

Sh =

kf dp D1f

Nu =

Sc =

µf ρf D1f

Pr =

Range of validity

hdp λf

Rep =

ρf udp µf

µf Cˆ pf λf

Gases Sh =

1 0.357 Rep 0.641 Sc 3 εb

Sh ≈ 0.07Rep

Nu =

1 0.428 Rep 0.641 Pr 3 εb

Nu ≈ 0.07Rep

3 < Rep < 2000 0.1 < Rep < 10

Liquids Sh =

1 0.250 Rep 0.69 Sc 3 εb

Nu =

1 0.300 Rep 0.69 Pr 3 εb

55 < Rep < 1500

Sh =

1.09 13 1 Rep Sc 3 εb

Nu =

1.31 13 1 Rep Pr 3 εb

0.0016 < Rep < 55

Monoliths:   dH 0.45 Sh = Nu∞ 1 + B1 ReSc L   dH 0.45 Nu = Nu∞ 1 + B1 RePr L where ρ udH Re = f µf √ εM δ w 4Ach dH = = √ 1 − εM Och and Nu∞ = 2.976(square) 3.657 (circular) 3.660 (hexagonal) B1 = 0.095 (catalytic monoliths)

the case of monoliths, other correlations should be used because of the different geometry [56–58]. Intraparticle Gradients Since the reactants have to diffuse to the active sites in the interior of the catalyst particles and meanwhile they can react, a concentration profile can develop over the particle diameter. Mathematically, this is described by a differential equation derived from a mass balance over a small shell of the catalyst particle. This reads, for a slab geometry and first-order irreversible reaction: 9.1.3.2

Deff

d2 c − kv c = 0 dz2

(25)

where Deff is the effective diffusivity and kv the rate constant based on unit volume of catalyst. The effective diffusivity depends on the morphology of the porous material. For large pores (>∼10 nm) the diffusivity is determined by the molecular diffusivity (∼10−4 − 10−5 m2 s−1 for gases), and in the case of small pores by the Knudsen diffusivity (∼10−6 − 10−7 m2 s−1 ). The molecular diffusivity of species i in the fluid can be calculated by correlations [5, 6] or by extrapolation from a known value at other conditions of temperature and pressure, using the dependence Df ,i ∝

T 1.75 ptot

(26)

9.1.3 Mass and Heat Transfer

The Knudsen diffusivity of species i in a cylindrical pore of radius r0 equals  T (27) DK,i = 97r0 Mi In the transition region, the Bosanquet equation, which accounts for both contributions, can be used: 1

1 1 = + D D D K,i f ,i

(28)

In zeolitic materials, in which the molecules are of pore size dimensions, activated configurational diffusion takes place, for which even lower diffusivity values, below 10−8 m2 s−1 , hold [59]. In this case, the diffusion description resembles that of surface diffusion – migration over the surface in the adsorbed phase – which explains the sometimes unexpected high diffusional flux [60]. A more rigorous approach for the combined effect of the various diffusion modes is based on the Maxwell–Stefan approach [60], for porous materials often referred to as the ‘‘dusty gas model’’. In addition, due to the presence of solid material, the volume in which diffusion can take place is reduced by a factor εp , the particle porosity. The tortuous path increases the diffusion length for a molecule relative to the spatial coordinate by a factor τp . The effective diffusivity can then be expressed as Deff =

εp ≈ 0.05 − 0.1D D τp

(29)

The form of the steady-state mass balance [Eq. (25)] will depend on the particle shape (sphere, cylinder, trilobe, slab, etc.). For any geometry, the characteristic length should be defined as L=

Vp 1 =  Ap a

(30)

to obtain essentially the same result as for the slab geometry of thickness 2L given below. The solution of the concentration profile is z

cosh φ c L = (31) cs cosh(φ) in which φ is known as the Thiele modulus and is the square root of the ratio of the reaction rate and the diffusional rate in the particle:  kv (32) φ=L Deff When the Thiele modulus is small, no internal concentration profile exists. In the case of large values,

2033

due to the existence of a concentration profile, the catalyst is not effectively used and an effectiveness factor is defined as observed reaction rate rate without internal gradients  rv (c, T ) dV rv,obs = = rv,chem (cs , Ts ) rv,chem (cs , Ts ) Vp

ηi =

(33)

Only for simple reaction kinetics can an analytical expression for the effectiveness factor be given, such as for a first-order reaction in a slab, Eq. (34). In comparing catalyst activities and in kinetic studies, one needs data that are not disguised by concentration gradients. For a 5% tolerance level, the criterion for the effectiveness factor for the absence of internal diffusion limitations reads ηi =

tanh(φ) = 1 ± 0.05 φ

(34)

Since during a kinetic study one measures the observed reaction rate and not the intrinsic rate, one cannot determine whether this criterion is satisfied. Therefore, the so-called Wheeler–Weisz modulus is introduced, which yields a procurable quantity [4, 8]. From series expansion and since ηi is close to 1, the following criterion follows for an nth-order reaction [23]:   rv,obs L2 n + 1 < 0.15 (35) = ηi φ 2 = Deff cs 2 Note that cs is not an observable quantity. However, its value may be calculated from Eq. (18). Alternatively, using cb instead is often justified. Here, the generalized form of the Thiele modulus is used, in order to be independent of the reaction kinetics and which is defined such that for the limit φ → ∞, ηi → 1/φ [8, 61] Lrv (cs ) φ=

cs

2 D r (c) dc  eff v

(36)

ceq

For an nth-order irreversible reaction, Eq. (36) yields  kv n−1 φ=L · n+1 (37) 2 cs Deff Figure 21 shows the effectiveness factor as a function of the Wheeler–Weisz modulus for a number of reaction orders, indicating that criterion (35) holds for the References see page 2043

2034

9.1 Laboratory Catalytic Reactors: Aspects of Catalyst Testing

profiles over a particle:

1st order 1

T − Ts cs − c = βi Ts cs

0th order 3rd order

with

hi

2nd order

βi =

0.1 0.1

1

10

hif2 Fig. 21 Isothermal internal effectiveness factor as a function of the Wheeler–Weisz modulus for different reaction orders.

generalized Thiele modulus. Due to the definition of L, it is fairly independent of the catalyst geometry. The derivation of the internal temperature gradient can only be performed numerically, even for simple kinetics. Again, the 5% criterion is used for the internal effectiveness factor. Assuming that the rate can be simplified into a temperature- and a concentrationdependent part, this yields rv (c, T ) kv (T ) f (c) kv (T ) = ≈ rv (cs , Ts ) kv (Ts ) f (cs ) kv (Ts )    Ea Ts = exp − − 1 = 1 ± 0.05 RTs T

ηi =

(38)

where the concentration effects are assumed to be negligible in order to derive a criterion for the temperature effects only. In the non-isothermal case, the reaction temperature in the center of the particle will be higher than Ts at the external surface for exothermic reactions and lower for endothermic reactions. Since we focus on small deviations from the isothermal behavior, the rates in the particle center will be higher or lower, respectively, than at the surface. An expression for the temperature rise in the particle is obtained from the mass [Eq. (25)] and heat balance for a particle: − λp,eff

d2 T = rv (−Hr ) dz2

(40)

(39)

Effective thermal conductivity values of porous materials, λp,eff , in gaseous atmospheres range from 0.1 to 0.5 J s−1 K−1 m−1 [6] and are only slightly larger than those of the gas phase. Straightforward combination of Eqs. (25) and (39) and integration lead to a simple general result that relates the temperature and concentration

(−Hr )Deff cs Tmax = λp,eff Ts Ts

(41)

the internal Prater number, which represents the maximum relative temperature difference across the particle or the ratio of the heat production rate and the heat conduction rate in the particle. Series development and combination of Eqs. (38) and (39) yield the criterion for absence of temperature effects on the experimental data [47, 23]:       E  (−H )D c a r   eff s 2 γs βi ηi φ  =   RTs λp,eff Ts 
rv,obs L2  ×  < 0.1 Deff cs 

(42)

This criterion also contains ‘‘observables’’ and so an estimation can be made of the presence of temperature gradients. This result is not unexpected and is similar to that for external transport. The Wheeler–Weisz parameter represents the concentration profile, the Prater number the maximum heat production relative to heat removal by conduction and the dimensionless activation energy the sensitivity of the rate towards a temperature change. It can be shown that |βi |ηi φ 2 represents the relative temperature gradient over the particle, just as |βe |Ca represents the relative temperature gradient over the film. When using the criterion, cs and Ts are not always known. Then, using the values for the bulk phase generally yields a good estimate. They can also be calculated from the external transport balances, Eqs. (18) and (22). It will be clear that in the case of exothermic reactions, the effectiveness factor can become larger than one. This is indicated by the numerical solution for the effectiveness factor in Fig. 22 for various values of βi and γs . Situations may arise in which concentration effects counterbalance the temperature effect, resulting in an effectiveness factor of around one. On the other hand, the temperature effect in the case of endothermic reactions only adds to the lowering of the effectiveness factor by a concentration profile. Catalyst Bed Gradients In the catalyst bed, temperature gradients may develop, analogously to those in a particle. In the axial direction 9.1.3.3

9.1.3 Mass and Heat Transfer

where the reaction rate rV ,obs based on the whole bed volume can be expressed as

10 bi

rV ,obs = rv,obs (1 − εb )(1 − b)

0.6 hi

0.4 1 0.2 −0.2 0.1 0.1

1 f

2035

10

Fig. 22 Internal effectiveness factor as a function of the Thiele modulus for non-isothermal reactions at different values for the Prater number and γs = 10 (numerical solutions for a first-order reaction).

the conversion increases, causing a temperature gradient. These temperature gradients should be kept as small as possible, because otherwise the rate data may not have the value attributed to them. The heat produced or consumed by the reaction can be removed or supplied, respectively, by the fluid phase flowing through the bed or by axial and radial conduction through the bed itself. Especially in the case of gas-phase reactions, the heat capacity of the fluid may be insufficient to keep the temperature change within acceptable limits. Axial conduction contributions are negligible in the case of packed beds due to the large length-to-diameter ratio imposed by the various criteria. Radial conduction to the wall of the reactor has the largest effect in minimizing temperature gradients. Heat conduction through packed beds, however, is poor. The porous particles themselves have poor conductivity and heat transfer between the particles occurs only through their contact points and via the fluid phase. Radiation is only important at very high temperatures, in reactions such as the reforming of methane by steam or carbon dioxide or the methane coupling reaction. Additionally, due to the lower porosity in the wall region the heat transfer from the bed to the reactor wall might be even worse. Mears analyzed this problem for a plug-flow reactor [52] and obtained an approximate solution for the maximum temperature rise in the catalyst bed. Based on this result, a criterion for a (T − grad)ext > (c − grad)int , (T − grad)int > (c − grad)ext

(56)

Effect of Transport Limitations on Observed Behavior

Bim = 10–104 Bih = 10−4 –10−1

gas–solid

The observed reaction rate can be expressed as follows: rv,obs = ηrv,chem (cb , Tb ) = ηe ηi rv,chem (cb , Tb )

gas–liquid

This indicates that for liquid–solid reactions, the external gradient is negligible, whereas in gas–solid reactions it will depend on the specific parameter values. Unlike in industrial operation (high flow-rates, relatively large particles), where as a rule of thumb the internal temperature gradient is assumed to be the largest [66], in laboratory operation the external gradient usually is the largest. Comparing the criteria for intraparticle and for bed behavior, Eqs. (42) and (43), neglecting the wall contribution and bed dilution, one sees immediately that their ratio results in   rv,obs (−Hr )dt 2 (1 − εb ) γw 32λb,eff Tw  ≈

γb βe Ca 2   2  λp,eff dt s (1 − εb ) dp

λb,eff

8

>> 1

(55)

The criterion for a flat velocity profile requires that the first term be >100. The second term is >1 since the effective conductivity of the bed is smaller than that of a particle and the last term is only slightly < 1. Therefore, the temperature gradient in the catalyst bed will be the first to develop to an unacceptable extent and should be verified first. In the case of too large temperature gradients, bed dilution with a wellconducting chemically inert material will improve the situation considerably. Whether the criteria for intra- or extraparticle heattransport limitations are more severe than the corresponding ones for mass transport depends on the absolute values of the products γs βi and γb βe . In some cases the former product exceeds unity but more often it does

(57)

If external mass transport limitations strongly dominate, the rate becomes equal to the mass transfer rate, Eq. (17), with cs = 0. Hence a first-order dependence is observed. Furthermore, the mass transfer coefficient is fairly independent of temperature. On the other hand, the observed rate constant does depend on the flow-rate and particle size (through correlations in Table 2). If internal diffusion limitations dominate ηi → 1/φ, so for an nth-order reaction:    n+1 Ea rv,obs ∝ L−1 kv Deff cs n+1 ∝ L−1 c 2 exp − 2RTs (58) Then the reaction is dependent on the particle size, has an apparent order of (n + 1)/2 and an apparent activation energy that is half the true activation energy. At low temperatures reaction rates are generally kinetically controlled. With increasing temperature, one first enters the diffusion-limited region and at still higher temperatures the film (extraparticle) diffusion-limited region. In so-called Arrhenius plots this is nicely seen as a changing slope of the rate versus 1/T plot, which can be used as an indication for the presence of limitations. A number of examples of this behavior can be found in the literature, e.g. for the Fischer–Tropsch synthesis of middle distillates [67] and the catalytic gasification of carbon [68]. These examples also prove the applicability of the theory. Changing activation energies, however, are not always indicative of the presence of limitations. The approach of thermodynamic equilibrium in the case of exothermic reactions can cause this phenomenon, References see page 2043

2038

9.1 Laboratory Catalytic Reactors: Aspects of Catalyst Testing Tab. 4

Apparent catalyst rate behavior depending on rate-controlling regime (isothermal case)

Controlling process Kinetics Internal diffusion External mass transfer a

Apparent order

Apparent activation energy

Dependency L

Dependency u

n

Ea (true)

–

–

(n + 1)/2

1 Ea (true) 2

1/L

–

1

∼0

Lm−2[a]

um[a]

m represents the power of the Reynolds number in the Sherwood correlation in Table 2.

for example in hydrogenation reactions [68]. A change in rate-determining steps or deactivation of the catalyst might also be causes. The same holds for reaction orders. Table 4 surveys the various observations that can be made when mass transfer affects the isothermal kinetic behavior of catalyst particles. Temperature gradients in endothermic reactions amplify the effects of concentration gradients. In exothermic reactions the thermal effects can compensate the concentration effects or dominate completely. In the latter cases the apparent activation energy may become higher than the true one due to a kind of ‘‘light-off’’ [5]. 9.1.6

Diagnostic Experimental Tests

Apart from the criteria based on the observable quantities derived in Section 9.1.4, some other experimental tests exist to verify the presence or absence of transport limitations. Extraparticle Concentration Gradients Figure 23a shows a test to determine whether external mass transport limitations exist. The test consists of varying the flow-rate and amount of catalyst simultaneously, while keeping the space-time W/Fi0 constant. As the mass transfer coefficient kf depends on the fluid velocity in the catalyst bed, external limitations show by a change in conversion. If no external limitations exist, the resulting conversions should be the same. However, if temperature effects also interfere, these might (over)compensate for concentration effects, so this method should be used with caution. 9.1.6.1

Intraparticle Concentration Gradients More attention should be paid to the diffusion limitations, as was shown before. Under strongly limited conditions, the observed rate becomes particle size dependent, ∝1/L. Variation of the particle size, e.g. by crushing and sieving the catalyst and performing the tests 9.1.6.2

under identical conditions, should give a proper answer to whether diffusion interferes or not. Figure 23b exemplifies this. At small particle size, the reaction is chemically controlled and independent of the particle size. Only for larger particles does a decrease in the observed rate occur. One should be aware of the fact that extraparticle limitations also induce a particle size dependence. The flow-rate through the bed should not have any effect (Table 4). Also in this case temperature effects should be avoided. Koros and Nowak [69] designed a more complex test based on the fact that the intrinsic reaction rate is proportional to the concentration of active sites in the kinetic regime, to its square root in the diffusion-limited regime and independent of this concentration in the external transport-limited regime. The idea is to crush a catalyst and dilute it with inert support to various concentrations and pelletize to identical particle sizes. The observed rate as a function of the site concentration should indicate the regime in which one operates. This test is not always conclusive: some reactions are structure sensitive or diffusion limitations occur on a very small scale, as in zeolite crystals. Madon and Boudart discussed this test in detail [70]. Temperature Gradients The best way to investigate the possible disguise by temperature gradients is to dilute the catalyst bed with an inert material with good heat conduction properties such as SiC. Upon dilution, lower conversions should result for exothermic reactions and higher for endothermic reactions, if gradients are present. The presentation of crude rate data in ‘‘Arrhenius plots’’ to inspect the temperature behavior can give indications of the presence of limiting transport processes by the changing slope (∝ activation energy). However, one should be aware of the fact that other phenomena can also induce this. Examples are a changing rate-determining step, catalyst deactivation or fouling, approach of thermodynamic equilibrium and ignition phenomena. 9.1.6.3

9.1.7 Proper Catalyst Testing and Kinetic Studies

F 0A

F 0A1

F 0A5

F 0A

F 0A

F 0A

Increasing flow rate

W1

Wi /F A0 i constant

X1

Increasing dp1 particle size

W5

X5

dp5

X1

X

2039

X5

X

X5 X1

X1 X5 F 0A1

(a)

Fig. 23

F 0A5 Flow rate

dp1 Reference catalyst (b)

dp5 Particle size

Diagnostic tests for (a) extraparticle limitations and (b) intraparticle limitations.

9.1.7

1

During the experimental investigations, one should carry out a proper testing of the catalysts, in order to obtain the relevant information with regard to intrinsic activity, selectivity, deactivation and kinetic behavior. The following guidelines should be applied: • adhere to criteria: ideal reactor behavior: plug flow or CSTR; isothermal bed; absence of limitations: observables, diagnostic tests • compare at low conversions to get insight in real activity differences (PFR) • compare selectivities at same conversion levels • make sure steady-state catalysis is really achieved. In diagnostic tests, crushing of the particles will not always be conclusive. Eggshell catalysts or other types, zeolites and washcoated monoliths are examples. In washcoated monoliths, the layer thickness is generally already small (1. This problem especially occurs at low reactant concentrations (e.g. exhaust catalysis, water purification) and/or low flowrates. The number of molecules passed over the catalyst per unit time may be much smaller than the number of available active sites in the catalyst. It may take a long time before each site has converted sufficient molecules to consider it as catalysis. Sometimes, the occurrence of a temperature profile along the reactor axis cannot be avoided, as in the study of exhaust gas monolithic catalysts. A possible approach can be to measure the temperature profile along the catalyst length and to use this for the data interpretation. It requires, however, a complete analysis of the mass and heat balances for a correct interpretation [71]. 9.1.8

Current Trends in Catalyst Testing

Figure 25 illustrates the historical trends towards scaling down and automation in catalyst development in terms of manpower needed since the 1960s [72]. The conventional approach towards solid catalyst development involved the evaluation of a single or a few catalyst compositions over a range of reaction conditions in a single dedicated reactor system, first in pilot plants, later in non-automated, then semi-automated bench-type reactor systems. Further benefit came from scaling down to micro- and nano-flow with full automation. In the past decade, further improvements have resulted from so-called parallelization or highthroughput experimentation (HTE) and from the development of so-called microchannel reactors or microreactors. Parallelization/High-Throughput Experimentation Although since the early days of catalyst development test rigs have been used in parallel for screening series of new catalysts [73, 74], the first reference in the open literature referring to the use of a single dedicated setup using parallel reactors to study solid catalysts appeared in 1980 [75]. In 1986, Creer et al. [76] published a detailed account of the design, implementation and verification of a system for testing of solid catalysts with six microreactors 9.1.8.1

2 Pilot plant (non-automated)

1

Manhour per reactor hour

2040

Bench scale (non-automated)

0.5

Bench scale (semi-automated)

0.1 Micro-and nano-flow (automated)

0.05

Parallellization

0.01

1960

1970

1980

1990

2000

Year Fig. 25 Manpower needed for different reactor units as a function of the approximate introduction year for general use in oil processing research.

in parallel. Several parallel catalyst testing systems have now been commercialized. In the early days of HTE, experimentation was mainly aimed at a qualitative comparison of activities of a large number of catalyst compositions to obtain ‘‘hits’’, socalled primary screening. However, the development of new and better experimental methods and techniques to control gas supply, fast analysis and computerization has led to more quantitative results, similar to or better than those of conventional laboratory reactors [77, 78]. A typical setup of the so-called six-flow reactor for catalyst comparison is shown in Fig. 26 [78–80]. Basically, a feed, a reactor and an analysis section are required. The six parallel reactors are placed in one oven. Nowadays, mass flow controllers for both liquid and gas result in stable molar flow-rates, ideal for kinetic studies. Pressure controllers maintain a constant feed pressure for the flow controllers, while backpressure controllers maintain the pressure in the reactors. Various methods of product analysis are available, which depend considerably on the system under investigation. An example of the successful implementation of the sixflow reactor was the kinetic study of the selective catalytic reduction of NO over alumina-supported manganese oxide [80]. One reactor served as the reference filled with inert particles, the others being filled with different amounts of catalyst. Product analysis was performed with a mass spectrometer. This yielded data at five different space times. The equipment was PC-controlled and predesigned kinetic runs (pressure and temperature variation) could be carried out automatically within a short period. The feed composition to all reactors is the same and only the cost of additional flow controllers to divide the feed mixture and a selection

9.1.8 Current Trends in Catalyst Testing

Feed control

Reactor

2041

Analysis

P MFC MFC MFC

Vent

MFC

SV

MFC MFC MFC MFC

BPC

MFC

BPC

Fig. 26

Six-flow reactor setup for simultaneously testing of different catalyst samples.

Tab. 5

Examples of parallel catalyst testing systems

Reactor system

Dimensions

Catalyst

Test reaction

Tubular plug flow (six-flow) Ceramic monolith

6 reactors 2.2 mm diameter, 150 mm length, 16 columns, 16 rows 2.6 mm diameter, 75 mm length, 16 columns, 8 rows 16 teeth 6-mm circular holes 15 autoclaves, 45 mL

Particles Powder, particles

Catalytic reduction of NO Methane oxidation

Catalytic coating

Total oxidations

[82]

Granular Powder

Gas-phase reactions Hydrogenation of citral (G−L reaction)

[83] [84]

Ceramic monolith with impermeable wall Rotating sample wheel Multibatch mini-autoclaves

valve are required. The tremendously increased data generation speed outweighs completely the relatively small additional investment. Since then, this basic approach of parallelization has been extended to multiples of reactors in various forms; see Table 5 for examples. Automated equipment, especially when using a parallel reactor system, yields a wealth of data that should be properly processed and interpreted. The time involved in the latter should not be underestimated and one should not lose oneself in experimentation only. Miniaturization The successful use of parallel multi-tube fixed-bed reactor systems for catalyst optimization has increased the demand for fast primary screening devices and methodologies, with further miniaturization being one of the key issues [77]. Recent advances in automation, robotics, microfabrication and instrumentation have allowed the development of new types of reactors, namely microchannel reactors or microreactors [85–87]. Such 9.1.8.2

References [78–80] [81]

reactors consist of a very large number of parallel channels, having diameters between 10 and several hundred micrometers. These small dimensions result in very high surface-to-volume ratios and thus very large heat and mass transfer rates, so that they are suitable not only for the initial screening of catalysts, but also for catalyst testing under kinetically controlled conditions. In Table 6, a few examples of the use of microreactors in catalyst testing are given. More about microreactors and the microfabrication techniques can be found in, e.g., Refs. [77] and [84–88] and in Chapter 10.8. The disadvantage of microchannels filled with catalyst powder is the large pressure drop in the small-diameter channels. To overcome this problem, cross flow can be used [93, 94] or the catalytic material can be deposited on the channel walls [87, 95]. The latter is most commonly used. The catalytic-wall microreactor is particularly suitable for studying fast, highly exothermic (or endothermic) References see page 2043

2042 Tab. 6

9.1 Laboratory Catalytic Reactors: Aspects of Catalyst Testing Examples of microreactors

Reactor system

Channel dimensions

Catalyst

Test reaction

References

Tubular plug flow

1 mm wide, 0.3 mm deep, 6 mm long, six channels

Granular

CO oxidation

[89]

Tubular plug flow (SWITCH 16)

16 channels

Powder

CO oxidation

[90–92]

Cross-flow packed bed

25.5 mm wide, 500 µm deep, 400 µm long

Particles

CO oxidation

[93, 94]

Stacked plates

300 µm wide, 250 µm deep, 20 mm long, 34 channels/plate, 10 plates

Catalytic coating

Oxidative dehydrogenation of propane

[95]

Cross-flow stacked plates

145 µm wide, 145 µm deep, 6.5 mm long

Catalytic coating

Ammonia oxidation, partial oxidation of methane

[96–98]

Single-bead

7 × 15 channels (or more)

Particles, usually spherical

Partial hydrogenation, partial oxidation

reactions, because very rapid heat removal (or addition) is possible, ensuring isothermal conditions. Microreactors have also been developed for heterogeneously catalyzed gas–liquid reactions. Examples are given by, e.g., Jensen [87]. The operation of many coated channels in parallel assumes that each channel receives equal flows by using special geometries or contains equal amounts of catalyst coating. This is not always the case, however, and one should consider the effect of deviations on the kinetic interpretation of data or on the design of microreactor systems. Delsman et al. [99] present relations to approximate these deviations and showed that especially distributions in the channel diameter (flow-rate) cause the largest deviations. List of Symbols

a a Ap Ach b

c p C p C dH dp

area of gas bubbles per unit volume of liquid specific surface area of catalyst particle external surface area of particle cross-sectional area of monolith channel volume of inert particles as fraction of total solids volume concentration molar heat capacity mass heat capacity hydrodynamic channel diameter monolith particle diameter

m−1

dt Dax DK,i Dif D Deff Ea f (c) Fi0 Ftot g h h0 kv

m−1

kf

m2

kL

m2

L

– mol m−3 J mol−1 K−1 J kg−1 K−1 m

Lb L0 Mi n N Ni Och p

m

diameter of reactor tube axial dispersion coefficient Knudsen diffusivity of component i molecular diffusivity of i in fluid phase f average diffusivity effective diffusivity in particle activation energy concentration function in rate expression molar flow of i at reactor inlet total flow through catalyst bed gravitational acceleration heat transfer coefficient static heat transfer coefficient rate constant per unit particle volume mass transfer coefficient (fluid-to-solid) mass transfer coefficient (gas-to-liquid) characteristic catalyst dimension bed length undiluted bed length molar mass of component i reaction order number of tanks in series number of moles periphery of monolith channel cross section pressure

[77]

m m2 s−1 m2 s−1 m2 s−1 m2 s−1 m2 s−1 J mol−1 mol s−1 mol s−1 m s−2 J m−2 s−1 K−1 J m−2 s−1 K−1 m s−1 m s−1 m m m kg kmol−1 – – mol m Pa

References

r0 rv rv,chem rv,obs rV ,obs rW R Rc s t T u Vp W x y0 z

average pore radius reaction rate per unit particle volume intrinsic reaction rate per unit particle volume observed reaction rate per unit particle volume observed reaction rate per unit bed or liquid volume reaction rate per unit catalyst mass universal gas constant recirculation ratio geometry parameter time temperature superficial velocity particle volume catalyst mass conversion mole fraction of reactant in reactor feed gas coordinate

m mol s−1 m−3 p mol s−1 m−3 p mol s−1 m−3 p mol s−1 m−3 b mol s−1 kg−1 J mol−1 K−1 – – s K m s−1 m3 kg – – m

Dimensionless numbers Bim Biot number of mass Bih Biot number of heat Biot number heat transfer at Biw wall Ca Carberry number Nu Nusselt number Pep P´eclet particle number Pr Prandtl number Re Reynolds number Rep Reynolds particle number Sc Schmidt number Sh Sherwood number Wetting number in trickle Wtr flow Greek letters β Prater number γ dimensionless activation energy wall thickness in monolith δw Hr heat of reaction ε porosity εM porosity of monolith η overall effectiveness factor external effectiveness factor ηe ηi internal effectiveness factor λf thermal conductivity of fluid phase

λeff λ0eff µ νl νi ρ τ τp φ φG ϕ

effective thermal conductivity stationary effective thermal conductivity dynamic viscosity kinematic viscosity of liquid stoichiometric coefficient of i in reaction density average residence time CSTR particle tortuosity factor Thiele modulus generalized Thiele modulus parameter in effective thermal conductivity equation

2043

J m−1 s−1 K−1 J m−1 s−1 K−1 kg m s−1 m2 s−1 – kg m−3 s – – – –

Subscripts app apparent b in bulk phase; also bed chem chemically controlled e external eff effective f fluid phase i intraparticle i of species i M monolith obs observed p particle s at external particle surface t tube w at the reactor wall Superscript eq equilibrium References

– – m J mol−1 – – – – – J m−1 s−1 K−1

1. Appl. Catal. 1988, 43, 211. 2. Catal. Today 1991, 11, 1. 3. J. R. Anderson, K. C. Pratt, Introduction to Characterization and Testing of Catalysts, Academic Press, Sydney, 1985. 457 pp. 4. J. J. Carberry, in Catalysis: Science and Technology, J. R. Anderson, M. Boudart (Eds.), Vol. 8, Springer-Verlag, Berlin, 1987, p. 131. 5. L. K. Doraiswamy, M. M. Sharma, Heterogeneous Reactions: Analysis, Examples and Reactor Design, Vol. 1: Gas–Solid and Solid–Solid Reactions, Wiley-Interscience, New York, 1984, 538 pp. 6. C. N. Satterfield, Mass Transfer in Heterogeneous Catalysis, MIT Press, Cambridge, MA, 1970, 267 pp. 7. P. Trambouze, H. van Landeghem, J.-P. Wauquier, Chemical Reactors. Design/Engineering/Operation, Editions Technip, Paris, 1988, Ch. 11. 8. G. F. Froment and K. B. Bischoff, Chemical Reactor Analysis and Design, 2nd Ed., Wiley, New York, 1990, 664 pp.

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9.1 Laboratory Catalytic Reactors: Aspects of Catalyst Testing

9. F. Kapteijn, G. B. Marin, J. A. Moulijn, in Catalysis, an Integrated Approach to Homogeneous, Heterogeneous and Industrial Catalysis, J. A. Moulijn, P. W. N. M. van Leeuwen, R. A. van Santen (Eds.), Studies in Surface Science and Catalysis, Vol. 79, Elsevier, Amsterdam, 1993, p. 251. 10. J. C. R. Turner, in Catalysis: Science and Technology, J. R. Anderson, M. Boudart (Eds.), Vol. 1, Springer-Verlag, Berlin, 1981, p. 43. 11. J. M. Berty, Plant Oper. Prog. 1984, 3, 163. 12. K. C. Pratt, in Catalysis: Science and Technology, J. R. Anderson, M. Boudart (Eds.), Vol. 8, Springer-Verlag, Berlin, 1987, p. 174. 13. L. K. Doraiswamy, D. G. Tjabl, Cat. Rev. Sci. Eng. 1974, 10, 177. 14. E. G. Christoffel, Laboratory Studies of Heterogeneous Catalytic Processes, Elsevier, Amsterdam, 1989, 263 pp. 15. F. H. M. Dekker, F. Kapteijn, J. A. Moulijn, A. Bliek, Chem. Eng. Sci. 1995, 50, 3573. 16. K. A. Vonkeman, PhD Thesis, Eindhoven University of Technology, 1990, 179 pp. 17. J. Happel, Isotopic Assessment of Heterogeneous Catalysis, Academic Press, Orlando, FL, 1986, 196 pp. 18. F. H. M. Dekker, J. G. Nazloomian, A. Bliek, F. Kapteijn, J. A. Moulijn, P. L. Mills, J. J. Lerou, Appl. Catal. A 1997, 151, 247. 19. C. Mirodatos, Catal. Today 1991, 9, 83. 20. R. H. Nibbelke, J. Scheerova, M. H. J. M. de Croon, G. B. Marin, J. Catal. 1995, 156, 106. 21. H. A. J. van Dijk, J. H. B. J. Hoebink, J. C. Schouten, Top. Catal. 2003, 26, 111. 22. R. Meijer, F. Kapteijn, J. A. Moulijn, Fuel 1994, 73, 723. 23. J. A. Moulijn, A. Tarfaoui, F. Kapteijn, Catal. Today 1991, 11, 1. 24. H. Redlingsh¨ofer, O. Kr¨ocher, W. B¨ock, K. Huthmacher, G. Emig, Ind. Eng. Chem. Res. 2002, 41, 1445. 25. R. J. Berger, Derivation of a criterion for allowing absence of radial mass transport limitations in catalytic wall reactors, unpublished work, 2004. 26. P. C. Borman, A. N. R. Bos, K. R. Westerterp, AIChE J. 1994, 40, 862. 27. H. P. Calis, PhD Thesis, TU Delft, 1995. 28. G. B. Tatterson, Scaleup and Design of Industrial Mixing Processes, McGraw-Hill, New York, 1994, 312 pp. 29. F. Kapteijn, A. J. van Steen, J. C. Mol, J. Chem. Thermodyn. 1983, 15, 137. 30. M. P. Helmsing, M. Makkee, J. A. Moulijn, Chem. Eng. Sci. 1996, 51, 3039. 31. M. A. den Hollander, M. Makkee, J. A. Moulijn, Appl. Catal. A 1999, 187, 3. 32. D. W. Kraemer, U. Sedran, H. I. d. Lasa, Chem. Eng. Sci. 1990, 45, 2447. 33. S. T. Sie, Chem. Eng. J. 1993, 53, 1. 34. S. T. Sie, Rev. Inst. Fr. P´et. 1991, 46, 501–515. 35. A. W. Patwardhan, J. B. Joshi, Ind. Eng. Chem. Res. 1997, 36, 3904. 36. M. M. P. Zieverink, M. T. Kreutzer, F. Kapteijn, J. A. Moulijn, Ind. Eng. Chem. Res., 2006, 45, 4574. 37. E. Dietrich, C. Mathieu, H. Delmas, J. Jenck, J. Chem. Eng. Sci. 1992, 47, 3597. 38. E. Crezee, B. W. Hoffer, R. J. Berger, M. Makkee, F. Kapteijn, J. A. Moulijn, Appl. Catal. A 2003, 251, 1. 39. B. W. Hoffer, P. H. J. Schoenmakers, P. R. M. Mooijman, G. M. Hamminga, R. J. Berger, A. D. van Langeveld, J. A. Moulijn. Chem. Eng. Sci. 2004, 59, 259.

40. S. Tajik, P. J. van den Berg, J. A. Moulijn, Meas. Sci. Technol. 1990, 1, 815. 41. R. J. Berger, E. H. Stitt, G. B. Marin, F. Kapteijn, J. A. Moulijn, CATTECH 2001, 5, 30. 42. A. C. J. M. van de Riet, H. Vonk, X. Xu, E. Otten, A. Cybulski, A. Stankiewicz, R. K. Edvinsson, J. A. Moulijn, React. Kinet. Catal. Lett. 1997, 60, 339. 43. R. Broucek, Chem. Eng. Sci. 1983, 38, 1349. 44. R. D¨umpelmann and A. Baiker, Chem. Eng. Sci. 1992, 47, 2665. 45. S. Wedel and J. Villadsen, Chem. Eng. Sci. 1983, 38, 1346. 46. H. Gierman, Appl. Catal. 1988, 43, 277. 47. D. E. Mears, Chem. Eng. Sci. 1971, 26, 1361. 48. C. F. Chu, K. M. Ng, AIChE J. 1989, 35, 148. 49. S. S. E. H. Elnashaie, M. E. Abashar, Chem. Eng. Sci. 1990, 45, 2964. 50. R. J. Berger, J. P´erez Ram´ırez, F. Kapteijn, J. A. Moulijn, Appl. Catal. A 2002, 227, 321. 51. V. W. Weekman Jr., AIChE J. 1974, 20, 833. 52. D. E. Mears, J. Catal. 1971, 20, 127. 53. A. A. C. M. Beenackers, W. P. M. van Swaaij, Chem. Eng. Sci. 1993, 48, 3109. 54. A. Cybulski, M. J. van Dalen, J. W. Verkerk, P. J. van den Berg, Chem. Eng. Sci. 1975, 30, 1015. 55. C. N. Satterfield, Heterogeneous Catalysis in Industrial Practice, 2nd Ed., McGraw-Hill, New York, 1992, Ch. 11. 56. R. D. Hawthorn, AIChE Symp. Ser. 1974, 70(137), 428. 57. R. K. Shah, A. L. London, AIChE J. 1976, 22, 344. 58. A. Cybulski, J. A. Moulijn, Catal. Rev. Sci. Eng. 1994, 36, 179. 59. J. K¨arger, D. M. Ruthven, Diffusion in Zeolites and Other Microporous Solids, Wiley, New York, 1992. 60. R. Krishna, Gas Sep. Purif. 1993, 7, 91. 61. P. Schneider, Catal. Rev. Sci. Eng. 1975, 12, 201. 62. J. R. Rostrup-Nielsen, L. J. Christiansen, J.-H. Bak Hansen, Appl. Catal. 1988, 43, 287. 63. D. Kunii, O. Levenspiel, Ind. Eng. Chem. Res. 1991, 30, 136. 64. M. G. Freiwald, W. R. Paterson, Chem. Eng. Sci. 1992, 47, 1545. 65. R. J. Wijngaarden, K. R. Westerterp, Chem. Eng. Sci. 1993, 48, 1273. 66. F. M. Dautzenberg, ACS Symp. Ser. 1989, 411, 99. 67. M. F. M. Post, A. C. van’t Hoog, J. K. Minderhout, S. T. Sie, AIChE J. 1989, 35, 1107. 68. C. A. Bernardo, D. L. Trimm, Carbon 1979, 17, 115. 69. R. M. Koros, E. J. Nowak, Chem. Eng. Sci. 1967, 22, 470. 70. R. J. Madon, M. Boudart, Ind. Eng. Chem. Fundam. 1982, 21, 438. 71. J. A. F. Kunst, A. Cybulski, X. D. Xu, P. J. T. Verheijen, J. A. Moulijn, in Precision Process Technology, Perspectives for Pollution Prevention, M. P. C Weijnen, A. A. H. Drinkenburg (Eds.), Vol. 1, Kluwer, Dordrecht, 1993, p. 197. 72. S. T. Sie, AIChE J. 1996, 42, 3498. 73. J. A. Moulijn, J. P´erez-Ram´ırez, R. J. Berger, G. Hamminga, G. Mul, F. Kapteijn, Catal. Today 2003, 81, 457. 74. R. J. Hendershot, C. M. Snively, J. Lauterbach, Chem. Eur. J. 2005, 11, 806. 75. R. Thomas, J. A. Moulijn, V. H. J. de Beer, J. Medema, J. Mol. Catal. 1980, 8, 161. 76. J. G. Creer, P. Jackson, G. Pandy, G. G. Percival, D. Seddon, Appl. Catal. 1986, 22, 85. 77. T. Zech, J. Klein, S. A. Schunk, T. Johann, F. Sch¨uth, S. Kleditzsch, O. Deutschmann, in High Throughput Analysis: a Tool for Combinatorial Material Science, R. A. Potyrailo, E. J. Amis (Eds.), Kluwer Academic/Plenum Press, New York, 2004, p. 491.

9.2.2 Overall Equipment 78. J. P´erez-Ram´ırez, R. J. Berger, G. Mul, F. Kapteijn, J. A. Moulijn, Catal. Today 2000, 60, 93. 79. J. A. Moulijn, J. P´erez-Ram´ırez, R. J. Berger, G. Hamminga, G. Mul, F. Kapteijn, Catal. Today 2003, 81, 457. 80. F. Kapteijn, L. Singoredjo, N. J. J. Dekker, J. A. Moulijn, Ind. Eng. Chem. Res. 1993, 32, 445. 81. U. Rodemerck, P. Ignaszewski, M. Lucas, P. Claus, M. Baerns, Top. Catal. 2000, 13, 249. 82. M. Lucas, P. Claus, Appl. Catal. A 2003, 254, 35. 83. P. L. Mills, J. F. Nicole, Chem. Eng. Sci. 2004, 59, 5345. 84. P. Claus, D. H¨onicke, T. Zech, Catal. Today 2001, 67, 219. 85. W. Ehrfeld, V. Hessel, H. L¨owe, Microreactors: New Technology for Modern Chemistry, Wiley-VCH, Weinheim, 2000, 288 pp. 86. K. F. Jensen, AIChE J. 1999, 45, 2051. 87. K. F. Jensen, Chem. Eng. Sci. 2001, 56, 293. 88. S. V. Gokhale, R. K. Tayal, V. K. Jayaraman, B. D. Kulkarni, Int. J. Chem. React. Eng. 2005, 3, R2. 89. K. Kusakabe, K. Tokunaga, G. Zhao, K. Sotowa, S. Morooka, J. Chem Eng. Jpn. 2002, 35, 914. 90. D. Tibiletti, E. A. B. de Graaf, S. P. The, G. Rothenberg, D. Farrusseng, C. Mirodatos, J. Catal. 2004, 225, 489. 91. C. Mirodatos, Y. Schuurman, C. Hayaud, A. Holzwarth, D. Farrusseng, T. Richter, European Patent, EP 1 293 772, 2003, assigned to CNSR-AMTEC. 92. www.amtec-chemnitz.de. 93. S. K. Ajmera, C. Delattre, M. A. Schmidt, K. F. Jensen, J. Catal. 2002, 209, 401. 94. S. K. Ajmera, C. Delattre, M. A. Schmidt, K. F. Jensen, Sens. Actuators B 2002, 82, 297. 95. N. Steinfeldt, D. Dropka, D. Wolf, M. Baerns, Trans Inst. Chem. Eng. 2003, 81, 735. 96. J. C. Schouten, E. V. Rebrov, M. H. J. M. de Croon, Chimia 2002, 56, 627. 97. E. V. Rebrov, M. H. J. M. de Croon, J. C. Schouten, Chem. Eng. Res. Des. 2003, 81A, 744. 98. M. J. V. Mies, E. V. Rebrov, M. H. J. M. de Croon, J. C. Schouten, Chem. Eng. J. 2004, 101, 225. 99. E. R. Delsman, M. H. J. M de Croon, G. D. Elzenga, P. D. Cobden, G. J. Kramer and J. C. Schouten, Chem. Eng.Technol. 2005, 28, 367.

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mixtures is usually carried out on a small scale, in a laboratory apparatus accommodating typically 0.1–10 g of catalyst. To ensure that meaningful and reproducible results are obtained in such laboratory-scale units, they must be properly designed and constructed and parts which are prone to malfunctions should be avoided. It is the intention of the present chapter to discuss critical building blocks of laboratory-scale catalytic units which are frequently designed in an inadequate manner. Since the core part of such units, viz. the catalytic reactor, is the subject of Chapter 9.1, this contribution will focus on the peripheral building blocks, i.e. the devices for the preparation of the feed mixtures to be sent to the reactor and the systems downstream of the reactor for transferring product samples to an analytical instrument. 9.2.2

Overall Equipment

All research in the field of heterogeneous catalysis ultimately aims at a more detailed understanding of the various chemical and mass transfer steps and, based on this knowledge, at the design and preparation of improved catalysts. Testing the performance of a solid catalyst in the conversion of gaseous, liquid or supercritical feed

The overall equipment used for catalytic studies is usually classified into (i) batch methods, (ii) semi-batch methods (e.g. a hydrogenation reaction in which hydrogen gas bubbles continuously through a batch reactor containing the liquid substrate and the suspended catalyst), (iii) transient methods (including pulse reactors) and (iv) continuous flow methods. Although all these methods have their specific advantages and disadvantages [1], continuously operated flow-type units strongly prevail in practice. The reasons are obvious: such units are not only well suited to measure the kinetics of the catalytic reaction, but they also enable the experimentalist to detect readily and quantitatively whether the catalyst, while working, preserves a constant activity, deactivates (which happens frequently) or, conversely, becomes more active (which sometimes happens). A rough scheme of a continuously operated flow-type unit for studying a gas-phase reaction on a solid catalyst is depicted in Fig. 1. Arbitrarily, a fixed-bed reactor (see Chapter 10.1) was chosen. In fact, fixed-bed reactors are most popular in heterogeneous catalysis, because they are easy to construct, relatively inexpensive (even if designed for high pressure), robust (since there are no moving parts) and a downscaled image of the most frequently employed reactor type in industrial catalysis. Ideally, the fixed-bed reactor behaves like a plug-flow (or, synonymously, piston-flow) reactor (PFR) with no radial gradients of partial pressure or gas velocity and with a complete absence of axial mixing. In a PFR, the reactant and product partial pressures are thus only a function of the reactor length. Its antipode is the continuous stirred tank reactor (CSTR), which

∗

References see page 2053

9.2

Ancillary Techniques in Laboratory Units for Testing Solid Catalysts .. Jens Weitkamp∗ and Roger Glaser

9.2.1

Introduction

Corresponding author.

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9.2 Ancillary Techniques in Laboratory Units for Testing Solid Catalysts

Solid catalyst Feed mixture A + B ( + C + ...)

Oven

B or B + C + ...

Off - gas stream

Product sample(s)

A Device for generation of feed stream

Product stream Reactor

Sampling device

Analytical instrument(s)

Scheme of a continuously operated flow-type unit with a fixed-bed reactor for studying a gas-phase reaction on a solid catalyst. A is the carrier gas; B and C are solids or liquids at room temperature and vaporizable at elevated temperature.

Fig. 1

also plays an important role in heterogeneous catalysis, especially in kinetic measurements. Ideally, the content of a CSTR is perfectly mixed (homogeneous), i.e. the partial pressures or concentrations of all reactants and products are uniform everywhere and equal to those of the effluent stream. Both the ideal PFR and the ideal CSTR behavior can be approximated to a large extent in real laboratory reactors. The description of the various reactor types for the study of heterogeneously catalyzed reactions is beyond the scope of this chapter. For their pros and cons, their suitability for a rapid catalyst screening and the experimental strategies for collecting kinetic data in these reactors, reference is made to an abundant number of valuable textbooks, monographs and review articles [1–12] (see also Chapters 6.1, 6.3 and 9.1). By contrast, the problems associated with the ancillary building units, e.g. the proper design of devices for the generation of well-defined and stationary feed streams or of suitable devices for product sampling (cf. Fig. 1) have received scant attention in the literature [13, 14]. From a critical evaluation of the pertinent literature, one is led to conclude that these problems tend to be overlooked by too many experimentalists and that unduly simple and inappropriate methods are often employed. In the subsequent sections, selected problems encountered in the authors’ own group and examples for the solution of these problems will be discussed. 9.2.3

Generation of Feed Streams

The device for generating a gaseous feed stream consisting of vapors of one (B) or several (B + C + · · ·)

component(s) in a gas stream A (cf. Fig. 1) must meet a number of requirements: (i) the composition of the mixture must be strictly constant (stationarity requirement) during the whole experiment; (ii) the partial pressure(s) of the higher molecular weight component(s), i.e. B or B + C + · · ·, in the gas mixture to be generated should be variable independently from each other, preferentially over a wide range of compositions; and (iii) the device should be reliable, durable (preferentially without moving parts) and inexpensive. Let us first restrict ourselves to the simpler and more frequently encountered case where the feed mixture consists of vapors of a single component B in the carrier gas A. The optimum device, at both ambient and elevated pressure, will then be a saturator which contains component B in the liquid or solid state (Fig. 2). On its way through the saturator, the carrier gas A is loaded with vapors of B. Since its vapor pressure depends exponentially on temperature, the saturator must be closely thermostated. An externally thermostated water- or oil-bath circulating through a jacket around the saturator is often the best solution. If temperatures above ca. 200 ◦ C are required, a saturator surrounded by a stirred bath or molten salt with efficient temperature control can be used alternatively. In practice, various saturator designs have been employed (Fig. 2). If the carrier gas is simply bubbled through the liquid B (version a), it may happen that the gas/liquid mass and/or heat transfer are insufficient. On their way upwards, the gas bubbles will then be saturated incompletely with vapors of B, and the partial pressure of B in the gas mixture leaving the saturator will be illdefined (generally too low) and irreproducible. To avoid such malfunctions, improved versions of saturators have

9.2.3 Generation of Feed Streams

Version a

Version b

Version c

A + B gas mixture

A + B gas mixture Gas A

Gas A

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Version d

A + B gas mixture Gas A

A + B gas mixture

T2 Gas A

T1 Gas bubbling through liquid Fig. 2

Condensation after supersaturation (T 1 > T 2 )

Artificially lengthened path of gas bubbles

Enhanced mass transfer through inert solid

Design of saturators for the generation of gaseous feed mixtures consisting of a carrier gas A and vapors of another component B.

been designed. For example (Fig. 2, version b), a too high saturator temperature T1 can be deliberately applied, whereupon partial recondensation of B is enforced in a cooler held precisely at the appropriate saturation temperature T2 . Another suitable method (Fig. 2, version c) is to prolong artificially the route of the gas bubbles through the liquid and, hence, the contact time between the two phases. This can be achieved, for instance, by conducting the gas bubbles through a spiral-shaped chimney inside the saturator. Perhaps the best technical solution (Fig. 2, version d) consists in adding a chemically inert solid, porous or not, to liquid B inside the saturator. Normally, this brings about a mass transfer enhancement such that the vapor–liquid equilibrium is safely attained within a bed height of a few millimeters. Suitable porous solids may be impregnated with some 20–30 wt.% of the liquid, and the use of liquid B in such a ‘‘solidified’’ state is particularly advantageous in saturators of high pressure units. It is also obvious that none of the saturators sketched in Fig. 2 is appropriate if the carrier gas A is to be loaded with vapors of more than one component (B + C + · · ·; cf. Fig. 1): In these semi-batch systems, the more volatile component, say B, would be enriched in the gas phase and, for simple mass balance reasons, this would inevitably lead to a gradual depletion of the liquid reservoir with respect to B. In other words, since the partial pressures of B, C, . . . in the gas mixture depend directly on the mole fractions of B, C, . . . in the liquid mixture inside the saturator, the stationarity requirement (see above) would be violated. Therefore, most experimentalists employ either of the methods shown in Fig. 3. Method I provides two or more

saturators in parallel (for the generation of A + B + C and A + B + C + · · · mixtures, respectively), each containing one pure component. The two-component gas streams (A + B or A + C and, if applicable, A + · · ·) leaving these saturators are combined in a mixing chamber to give the desired multicomponent A + B + C (or A + B + C + · · ·) feed stream. According to our experience, method I is satisfactory for units working at atmospheric pressures. Even then, however, it may happen that a flow adjustment through one saturator brings about undesired changes in the flow through the other(s). With modern flow controllers, this problem can be solved, but, at elevated pressures and/or with more than two saturators, the method becomes clumsy and unduly expensive, because of the cost of the flow controllers and the increasing number of thermostats or cryostats required. A further complication may arise when the temperatures of the saturators are largely different. In that case, efficient preheating of the colder gas stream is necessary to avoid condensation of the higher boiling compound in the mixing chamber. In method II, the components B, C, . . . are premixed in the liquid state in exactly the same ratio as desired in the gaseous A + B + C + · · · feed stream. By means of an appropriate pump, typically a pulsation-free piston displacement pump, the premixed liquid is continuously injected into a device for total vaporization, where the vapors are combined with the stream of the carrier gas. The composition of the final A + B + C + · · · gas mixture is determined by the flow-rate of A, the piston displacement velocity of the pump and the composition References see page 2053

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9.2 Ancillary Techniques in Laboratory Units for Testing Solid Catalysts

I Gas A

FIC

A+B+C gas mixture

A+B+C gas mixture

II Gas A

MCh

FIC

B+C

Th / Cr

Vaporizer

Th / Cr

Pump

Conventional methods for the generation of gaseous feed mixtures consisting of a carrier gas A and vapors of two other components B and C. FIC, flow indicator and controller; MCh, mixing chamber; Th/Cr, thermostat or cryostat.

Fig. 3

of the liquid. The sophisticated pumps (strictly pulsationfree operation at low fluxes, typically in the order of 0.1 to several cm3 h−1 ) are nowadays commercially available, even for high-pressure units. The main shortcoming of this method, which is too often overlooked, stems from the difficulty in completely vaporizing the small liquid stream smoothly and pulsation-free, even if the vaporizer is kept at a high temperature. We tested a number of differently designed vaporizers, and the result was very disappointing: It turned out to be extremely difficult to avoid the formation of droplets at the entrance nozzle of the vaporizer. These droplets ultimately break off the nozzle and reach a hot wall, where they vaporize almost instantaneously. The result is a strong and uncontrollable fluctuation of the partial pressures of B, C, . . . in the gas mixture. These fluctuations are usually difficult to detect; in particular, they remain hidden, if an integral, i.e. a time-averaged, sample of the A + B + C + · · · or B + C + · · · mixture is analyzed. In view of all these obvious or hidden drawbacks of methods I and II, a multicomponent saturator has been developed for the generation of gaseous feed streams in laboratory flow-type units [15, 16]. Its salient features (Fig. 4) are a vertical saturator column packed with an inert solid, typically glass beads, and a moving liquid phase (as opposed to the saturators shown in Fig. 2), which trickles downwards, while the carrier gas A is conducted countercurrently. A pump conveys the liquid B + C + · · · mixture from the storage vessel via a heating line inside the thermostated jacket on to a specially designed tray at the top of the column. This tray ensures a uniform distribution of the liquid over the crosssection of the column. In this countercurrent column,

TIR

Gas A

A + B + C + ··· gas mixture

Pump

B + C + ··· Thermostat or cryostat

Collection vessel

Storage vessel

Multicomponent saturator for the generation of gaseous feed mixtures consisting of a carrier gas A and vapors of two or more other components B, C, . . . . TIR, temperature indicator and recorder.

Fig. 4

the vapor–liquid equilibrium is efficiently established. Unvaporized liquid leaves the column at its bottom and accumulates in the collection vessel (this surplus liquid mixture, on adjustment of the concentrations of its

9.2.4 Devices for Product Sampling

constituents, can be re-used in the storage vessel in the subsequent experiment). The gaseous A + B + C + · · · stream at the top of the column, which has a strictly stationary composition [16], is sent to the catalytic reactor. On its way downwards, the liquid phase is, of course, depleted with respect to the more volatile component(s) and enriched in its heavier one(s). At the decisive locus, however, where both phases have their final contact (i.e. the top of the column), the composition of the liquid is obviously stationary. For a desired composition of the gas mixture, the appropriate values for the liquidphase composition and the saturator temperature must be found. This is best done in two successive steps, viz. by phase equilibrium calculations followed by experimental refinement of the calculated values. The multicomponent saturator showed excellent performance, both in a unit for atmospheric pressure [15] and in a high-pressure apparatus [16, 17]. Moreover, it can be utilized for generating a stationary stream of one component B in the carrier gas A for virtually infinite times-on-stream by continuously or periodically adding liquid B to the storage vessel. Thus, the shortcoming of limited hold-up of liquid B in the saturators shown in Fig. 2 is overcome. These saturators can only be used as long as the liquid B is not used up. As another method for generating feed mixtures with defined compositions, the feed components can be formed by a chemical reaction, which starts, e.g., from a liquid substrate (mixture), is highly selective for the desired components and goes to completion. The

Reactor

2049

stoichiometry of the reaction determines the composition of the components in the feed stream. Thus, the direct handling of toxic or highly reactive feed components outside the catalytic testing unit can be avoided. Examples are the thermal generation of oxygen (and water) from aqueous hydrogen peroxide solutions [18] and the catalytic formation of hydrogen (and carbon dioxide) from formic acid, e.g. on a supported Pt catalyst [19]. In the latter case, the carbon dioxide formed can be utilized in its supercritical state as a reaction solvent (see Chapter 8.4). To adjust the hydrogen concentration in the reaction mixture, the concomitant decomposition of ethyl formate supplies carbon dioxide and ethane, which both act as additional reaction solvents and dilute the hydrogen. So far, the discussion of methods for generating welldefined feed mixtures in flow-type units has been mainly restricted to gaseous streams. As a rule, liquid feed streams are much easier to prepare, simply by premixing the reactants in a reservoir and conveying this mixture to the reactor by means of a pump with pulsation-free characteristics. 9.2.4

Devices for Product Sampling

A standard arrangement for sampling gaseous products downstream of the reactor is shown in Fig. 5. In a needle valve (or a similar device), the reactor effluent is depressurized and the flow-rate is controlled. In most References see page 2053

Carrier gas for gas chromatograph Capillary gas chromatograph

Sampling loop

FI Needle valve

Six - port sampling valve

Off - gas stream

Cooling trap e.g.,−196 °C

Standard equipment for sampling gaseous product mixtures in a continuously operated unit downstream of the reactor. FI, flow indicator. The parts within the area circumscribed by the dotted line are heated to avoid condensation of vapors.

Fig. 5

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9.2 Ancillary Techniques in Laboratory Units for Testing Solid Catalysts

cases, the analytical instrument of choice will be a gas chromatograph equipped with a capillary column, because such an instrument often allows a good separation of the products and, if equipped with an appropriate detector, a reliable quantitative analysis of these products. The working principle of gas chromatography, however, is inherently discontinuous with a time required for one analysis between ca. 1 min and several hours, depending on the nature and the number of components to be separated and also the desired resolution of the analytical separation. To link continuous flow-type and discontinuous analytical instruments, a kind of interface is needed, viz. a multiport sampling valve. In its normal position (Fig. 5), it directs the reactor effluent through the sample loop into the off-gas line and, upon actuating the valve, the content of the loop is swept into the gas chromatograph. It is obvious that the samples flushed from the loop into the gas chromatograph correspond to a well-defined time-on-stream of the catalyst, hence they are differential samples. They usually cannot be stored, so no product sample is available for further analysis after the experiment. We recommend routinely collecting an additional sample by condensation of the products in a cooling trap downstream of the six-port sampling valve (Fig. 5). This furnishes a time-averaged, i.e. integral, product sample which can, at least in principle, be easily stored, although it will tend to lose volatile components on being warmed to ambient temperature. Although it may therefore happen that the integral sample is not suitable for quantitative evaluation of the full range of products, it is often very desirable to have such a sample stored for a while, so that additional analytical work can be done, e.g. qualitative product identification by coupled gas chromatography/mass spectrometry.

Obviously, standard equipment as sketched in Fig. 5 enables one to observe the time-on-stream behavior of the catalyst, simply by repeating the product analysis at certain time intervals. This is only true, however, if the time required for product analysis is short compared with the typical time of catalyst deactivation (or, which occurs less frequently, catalyst activation). Numerous other catalytic systems are known in which the catalyst deactivates within minutes or even seconds, while complex product mixtures are formed, the adequate gas chromatographic analysis of which takes several hours. Examples are the catalytic cracking of higher hydrocarbons [20] (see also Chapter 13.5), the alkylation of isobutane with butenes on solid acids [21] (see also Chapter 13.8) and the conversion of methanol to hydrocarbons on acid zeolites [22] (see also Chapter 13.14). The simultaneous occurrence of complex product mixtures and rapid catalyst deactivation requires special sampling techniques, generally referred to as differential or instantaneous sampling, and the possibility of storing the samples without loss of volatile components. Various methods for instantaneous sampling have been developed. In the simplest approach, storable samples can be withdrawn from the product stream by gas syringes. It is evident that this is, in many instances, an unduly simplistic method. It fails, in particular, if the samples contain higher molecular weight components which condense inside the syringe at ambient temperature. The gas syringe method has been significantly improved [23, 24] by placing a heatable sampling column between the syringe and its needle (Fig. 6). The sampling column simply consists of a metal (e.g. steel) tube with a shut-off valve at each end. It can be heated, e.g. by directly passing a low-voltage current (ca. 2 V, 15 A) through the

Gas stream from reactor

Heatable sampling column (e.g., from steel)

Gas syringe

Shut - off valves

Gas stream, to cooling trap Fig. 6

Instantaneous sampling and product storage with a gas syringe and heatable sampling columns (after Ref. [24]).

9.2.4 Devices for Product Sampling

column. At an appropriate position downstream of the reactor, a sampling device is attached to the heated product line. The sampling device is virtually identical with the injection port of a gas chromatograph. Through its septum, the needle connected with the sampling column protrudes directly into the gaseous product stream. At the desired sampling time, both valves of the column are opened and an appropriate amount of product is sucked into the column, which is then shut again. Heating is now stopped and the sample can be stored. For analysis, the column is again heated (so that all product components vaporize) and connected to the injection port of a gas chromatograph, whereupon the product sample is flushed on to the chromatographic column. More details on this technique, including the recommended construction of the shut-off valves, can be found in Ref. [24]. Although the technique is relatively simple, easy to handle and applicable to high-pressure reactions, it has not found widespread application. For its routine application, a large number (in the order of 100) of sampling columns are needed. Another technique, which is much more frequently employed, relies on commercially available multiport valves with a large number of sampling loops (an example is a valve carrying 16 sample loops, with 16 positions and 2 × 16 + 2 = 34 ports, i.e. two for each loop plus a gas entrance and a gas exit port). The simplest arrangement of such a multiport valve downstream of the catalytic reactor is shown in Fig. 7 (for the sake of simplicity, only two sampling loops are drawn). The multiport valve is assisted by an eight-port valve with

two positions. The needle valve is to compensate for the pressure drop inside the multiport valve. During the catalytic experiment, the reactor effluent is passed along the heavy line, through the multiport valve. At the desired time-on-stream, the multiport valve is switched to its next position, whereby a product sample is captured in a loop. After the catalytic experiment, the eight-port valve is actuated, whereupon the carrier gas for the gas chromatograph passes through the multiport valve, and sample by sample can now be analyzed. However, the investment cost for the valves increases sharply if a reasonably large storage capacity for product samples (ca. 100) is desired. Some experimentalists encountered difficulties with the tightness of multiport valves at the elevated temperatures required to prevent condensation of vapors inside the sampling loops. Multiport sampling loops can also be employed for the sequential withdrawal of samples from product streams of several reactors run in parallel for the highthroughput screening of solid catalysts [25] (see also Chapter 9.3). The experimental challenge in the parallel testing in heterogeneous catalysis is similar to that of analyzing complex product mixtures while a catalyst undergoes rapid deactivation: the time required for gas chromatographic analysis with sufficient analytical resolution may be considerably longer than the changes in the product mixtures downstream of the different reactors in a parallel testing unit. Other approaches to References see page 2053

Gas stream from reactor

Needle valve

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Sampling loop

Capillary gas chromatograph

Sampling loop

Off - gas stream Carrier gas for gas chromatograph

Cooling trap

Instantaneous sampling and product storage with a multiport valve. The parts within the area circumscribed by the dotted line are heated to avoid condensation of vapors.

Fig. 7

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9.2 Ancillary Techniques in Laboratory Units for Testing Solid Catalysts

Sampling

Carrier gas for GC

Analysis

Gas stream from reactor To cooling trap To GC (a) Fig. 8

(b)

Instantaneous sampling and product storage by means of glass ampoules. GC, gas chromatograph.

overcoming these experimental difficulties are described in Ref. [25] and in Chapter 9.3. Still another method for instantaneous sampling has been developed which allows a virtually unlimited number of product samples, because extremely cheap, homemade glass ampoules are used for sampling. The glass ampoule technique was first described by Pichler and G¨artner [26] and later refined and systematically employed in catalytic units by Schulz and coworkers [20, 27–29]. Ample experience with the technique also exists in our group [21, 30, 31]. The principle is shown in Fig. 8. The gaseous effluent from the reactor is passed through a heated sampling compartment (Fig. 8a). A long capillary connected with an evacuated glass ampoule protrudes into the compartment through a sealing system which resembles the injection port of a gas chromatograph. At the desired time-on-stream, the tip of the capillary is broken, e.g. by actuating a piston mounted at an appropriate position of the compartment, whereupon the ampoule fills up instantaneously with the product mixture. The ampoule with its capillary is then withdrawn from the compartment, the products inside the ampoule are frozen at liquid nitrogen temperature and the ampoule is quickly sealed. The sample is now captured hermetically inside the ampoule and can be stored until its analysis. A new sampling cycle can then be initiated by insertion of another evacuated ampoule through the sealing system (Fig. 8). A skilled experimentalist is able to repeat the cycles in very short intervals of ca. 10 s, if necessary. For analysis (Fig. 8b), the glass ampoule is simply destroyed inside a heated device connected with a capillary gas chromatograph, whereby the product sample is liberated and swept into the chromatographic column. Very impressive recent examples that demonstrate the potential of the ampoule technique include studies

on the conversion of methanol to hydrocarbons over acidic zeolites [32] and the Fischer–Tropsch synthesis on various iron- or cobalt-based catalysts [29, 33–35]. 9.2.5

Conclusions

Testing the performance of solid catalysts in laboratory units offers numerous pitfalls of which the experimentalist needs to be aware. Some examples of inappropriate designs have been discussed in this chapter, and proven solutions for the key sections of catalytic testing units were described. This discussion was selective rather than exhaustive and focused on two areas, viz. the generation of feed gas streams and methods for product sampling. Other areas were omitted, since they are more intimately related to the catalytic reactor and are, hence, the subject of other chapters of this Handbook. These include, for instance, the various methods for calculating the conversion, selectivities and other kinetic parameters from the analytical raw data and the pros and cons of using an internal analytical standard in the feed stream. Moreover, the increasingly sophisticated equipment developed for high-throughput testing of solid catalysts was not specifically treated in this chapter. These developments will without doubt contribute to the improvement of catalytic testing units in the decades ahead. Leading research groups in the field of heterogeneous catalysis devote much time to the identification and elimination of weak parts in their experimental testing units. Without sophisticated and tailored equipment, certain catalytic reactions cannot be investigated at all. Examples are reactions with a complex product slate and simultaneously rapid catalyst deactivation or where more than one phase, e.g. liquid and gaseous, are involved or are

9.3.1 Introduction

formed as a result of the conversion. The novel catalysts emerging from modern materials science deserve to be tested in units which are free from malfunctions and shortcomings. References 1. J. R. Anderson, K. C. Pratt, Introduction to Characterization and Testing of Catalysts, Academic Press, New York, 1985, p. 257. 2. K. C. Pratt, in Catalysis – Science and Technology, J. R. Anderson, M. Boudart (Eds.), Vol. 8, Springer-Verlag, Berlin, 1987, p. 173. 3. O. Levenspiel, Chemical Reaction Engineering, 3rd Ed., Wiley, New York, 1999, p. 376. 4. J. M. Smith, Chemical Engineering Kinetics, 3rd Ed., McGrawHill, Singapore, 1981, p. 359. 5. L. K. Doraiswamy, D. G. Tajbl, Catal. Rev. Sci. Eng. 1974, 10, 177. 6. V. W. Weekman, AIChE J. 1974, 20, 833. 7. J. M. Berty, Chem. Eng. Prog. 1974, 20(5), 78. 8. C. Gelain, Chim. Ind. G´enie Chim. 1969, 102, 984. 9. A. M. R. Difford, M. S. Spencer, in Standardization of Catalyst Test Methods, C. R. Adams, S. W. Weller (Eds.), AIChE Symposium Series 143, Vol. 70, American Institute of Chemical Engineers, New York, 1974, p. 42. 10. P. A. Ramachandran, R. V. Chaudhari, Chem. Eng. 1980, 87(24), 74. 11. M. Herskowitz, J. M. Smith, AIChE J. 1983, 29, 1. 12. S. T. Sie, in Deactivation and Testing Hydrocarbon-Processing Catalysts, P. O’Connor, T. Takatsuka, G. L. Woolery (Eds.), ACS Symposium Series, Vol. 634, American Chemical Society, Washington, DC, 1996, p. 6. 13. J. Weitkamp, in Innovation in Zeolite Materials Science, P. J. Grobet, W. J. Mortier, E. F. Vansant, G. Schulz-Ekloff (Eds.), Studies in Surface Science and Catalysis, Vol. 37, Elsevier, Amsterdam, 1988, p. 515. 14. P. A. Jacobs, J. A. Martens, in Introduction to Zeolite Science and Practice, H. van Bekkum, E. M. Flanigen, J. C. Jansen (Eds.), Studies in Surface Science and Catalysis, Vol. 58, Elsevier, Amsterdam, 1991, p. 445. 15. J. Weitkamp, H. Dauns, Chem.-Ing.-Tech. 1984, 56, 929. 16. J. Weitkamp, H. Dauns, Appl. Catal. 1988, 38, 167. 17. H. Dauns, S. Ernst, J. Weitkamp, in New Developments in Zeolite Science and Technology, Proceedings of the 7th International Zeolite Conference, Y. Murakami, A. Iijima, J. W. Ward (Eds.), Studies in Surface Science and Catalysis, Vol. 28, Kodansha, Tokyo, Elsevier, Amsterdam, 1986, p. 787. 18. D. Br¨oll, A. Kr¨amer, H. Vogel, I. Lappas, H. Fuess, Chem. Eng. Technol. 2001, 24, 142. 19. J. R. Hyde, M. Poliakoff, Chem. Commun. 2004, 1482. 20. H. Schulz, A. Geertsema, Erd¨ol, Kohle-Erdgas-Petrochem. 1977, 30, 313. 21. J. Weitkamp, in Catalysis by Zeolites, B. Imelik, C. Naccache, Y. Ben Taarit, J. C. Vedrine, G. Coudurier, H. Praliaud (Eds.), Studies in Surface Science and Catalysis, Vol. 5, Elsevier, Amsterdam, 1980, p. 65. 22. H. Schulz, W. B¨ohringer, W. Baumgartner, Z. Siwei, in New Developments in Zeolite Science and Technology, Proceedings of the 7th International Zeolite Conference, Y. Murakami, A. Iijima, J. W. Ward (Eds.), Studies in Surface Science and Catalysis, Vol. 28, Kodansha, Tokyo, Elsevier, Amsterdam, 1986, p. 915.

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23. D. Rodewald, J. Lorenz, H. G. Struppe, DDR Patent WP 123 124, 1976. 24. W. Leipnitz, in Handbuch der Gaschromatographie, E. Leibnitz, H. G. Struppe (Eds.), Akademische Verlagsgesellschaft Geest & Portig, Leipzig, 1984, p. 153. 25. S. A. Schunk, D. Demuth, A. Cross, O. Gerlach, A. Haas, J. Klein, J. M. Newsam, A. Sundermann, W. Stichert, W. Strehlau, U. Vietze, T. Zech, in High-Throughput Screening in Heterogeneous Catalysis, A. Hagemeyer, P. Strasser, A. F. Volpe Jr. (Eds.), Wiley-VCH, Weinheim, 2004, p. 19. 26. H. Pichler, R. G¨artner, Brennst.-Chem. 1962, 43, 336. 27. H. Schulz, A. Geertsema, in Proceedings of the 5th International Conference on Zeolites, L. V. C. Rees (Ed.), Heyden, London, 1980, p. 874. 28. H. Schulz, Erd¨ol, Kohle-Erdgas-Petrochem. 1983, 36, 279. 29. H. Schulz, T. Riedel, G. Schaub, Topics Catal. 2005, 32, 117. 30. J. Weitkamp, in Proceedings of the 5th International Conference on Zeolites, L. V. C. Rees (Ed.), Heyden, London, 1980, p. 858. 31. S. Unverricht, S. Ernst, J. Weitkamp, in Zeolites and Related Microporous Materials: State of the Art 1994, J. Weitkamp, H. G. Karge, H. Pfeifer, W. H¨olderich (Eds.), Studies in Surface Science and Catalysis, Vol. 84, Elsevier, Amsterdam, 1994, p. 1693. 32. H. Schulz, K. Lau, M. Claeys, Appl. Catal. A 1995, 32, 29. 33. H. Schulz, G. Schaub, M. Claeys, T. Riedel, Appl. Catal. A: General 1999, 186, 215. 34. H. Schulz, Topics Catal. 2003, 26, 1. 35. T. Riedel, G. Schaub, Topics Catal. 2003, 26, 145.

9.3

High-Throughput Experimentation in Heterogeneous Catalysis .. Ferdi Schuth∗

9.3.1

Introduction

High-throughput experimentation (HTE) in heterogeneous catalysis is a relatively novel technology suite designed to accelerate the discovery of catalytic materials. The origin of the high-throughput approach is the pharmaceutical industry, where the combination of highly efficient assays and parallelizable methods for synthesis, initially of polypeptides, but subsequently also of other molecular compounds, allowed several thousandfold increases in throughput, albeit often at the expense of information depth [1, 2]. Parallel experimentation had been used for decades in many fields of science, and also in materials science and catalysis such approaches have been reported from time to time. Parallel reactors were in operation as early as the 1920s in Fischer’s References see page 2072 ∗ Corresponding author.

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9.3 High-Throughput Experimentation in Heterogeneous Catalysis

The full power, however, can best be exploited in a concerted, fully integrated high-throughput approach. In this chapter, the different elements of an HTE program in heterogeneous catalysis will be discussed. Several reviews have been published covering different aspects of HTE in catalysis, which may be consulted in addition to this chapter [8–18]. 9.3.2

Typical Program

Photograph of a typical bench in the 1920s with so-called four-tube ovens. The photograph is from 1932. Reproduced from Ref. [3].

Fig. 1

laboratory in M¨ulheim (Fig. 1) – and possibly also elsewhere – in order to increase the rate at which experiments could be conducted [3]. Other parallel reactors were occasionally described in the literature in the following decades [4, 5]. For materials synthesis, Hanak used the so-called compositional spread approach, in which simultaneous evaporation of different materials and deposition of these materials from the gas phase on a substrate resulted in compositional gradients over the substrate [6]. However, HTE in catalysis and materials science as a field only began to form in the middle of the 1990s, starting with a landmark paper on luminescent materials [7]. There are a number of different reasons which may explain the rapid expansion of HTE during that time: (i) probably most importantly, the way to think along the lines of massive parallelization had been paved by the experience of about 10 years in the pharmaceutical industry; (ii) the easy availability of PCs with a requirement for minimal knowledge of computers to use them efficiently made it comparably easy to handle the large amounts of data generated in high-throughput programs, and, in addition, it allowed full integration of the different components in one common platform; (iii) advances in automation technology developed in the pharmaceutical industry, for instance dispensing robots, could easily be leveraged into the field of catalysis. For these reasons, the ideas spread relatively quickly and now there are a number of companies providing HTE services and most major firms active in the field of catalysis are using HTE components in their research programs. Although HTE does not replace conventional tools used in catalysis research, it now seems clear that the components, such as automated catalyst synthesis and parallel reactor technology, are important methods in any catalyst development program.

HTE programs in catalysis can have different objectives: typically, one is interested in finding a new catalyst system for a given reaction, but possibly also the discovery of value-generating pathways starting with a given feedstock could be the target. Finally, the optimization of reaction conditions, such as temperature, exact feed composition, pressure, etc., could be the topic of a high-throughput approach. If the target is a new catalytic system, the prior knowledge, and thus the compositional and the parameter space which needs to be explored, can vary widely. The simplest case is the selection of known, perhaps even commercially available, catalyst systems for a given reaction. This can be important, for instance, in the pharmaceutical industry, where speed in developing a specific synthesis is crucial and the different reaction steps in a synthetic sequence need to be optimized, or in the petroleum refining industry, if a new feedstock is used and the catalyst needs to be adapted to its properties. In these cases, it is very helpful if available commercial catalysts can be rapidly screened in order to find an optimal solution as quickly as possible. The other extreme is the exploration of catalysts for a completely novel transformation for which no precedence exists. If no heuristic knowledge is available, one needs to find optimal ways for the fast exploration of a wide parameter space. Between these two extreme cases, screening of commercially available catalysts for a given transformation and the exploration of many different solids for a reaction for which no precedent is known, there are intermediate situations which in industrial HTE programs are the most frequently encountered cases: There may be a lead for a promising development from previous studies which needs to be systematically optimized, there may be information on successful catalysts for related transformations, there could be a patent which should be circumvented by systematically exploring compositions and conditions which do not fall in the protected domain, and so on. Looking at these various possibilities for the kinds of questions which may arise, it is obvious that the type of program that one would engage is strongly dependent on the nature of the problem to be solved.

9.3.3 Technology Components

Conceptually, one often distinguishes between ‘‘discovery’’ programs (also termed ‘‘stage I’’) and ‘‘optimization’’ programs (‘‘stage II’’), although in practice these approaches cannot always be clearly discriminated. In a discovery program, one typically does not have much information on possible catalyst compositions and reaction conditions and therefore the screening of a wide parameter space is required. This normally means that a high degree of parallelization is needed, the reactors typically have several hundred channels. The depth of information obtained is normally not very high; instead, preliminary information is collected on whether a specific solid seems to be promising, is clearly not promising or may be promising. Obviously, the better a method discriminates between promising and not promising cases, the fewer ambiguous cases are generated and the fewer false positives or false negatives are created, the better such a preliminary screen is. Since the screening method for very high throughput itself is often the most difficult component to develop, the synthesis format is often governed by the requirements of the subsequent catalytic evaluation. If testing is done on flat substrates, evaporation or sol–gel methods are the most suitable synthetic techniques, by which the potential catalysts are directly created on the substrate, which is later exposed to the reagent feed. In contrast, in an optimization (stage II) project, it is mandatory that the data quality is identical with that of a conventional, sequential catalyst development program. This normally means that the catalyst is synthesized in the form of a powder or as pellets and tested in a grain size which ensures suitable fluid dynamics in the reactor. The reactors are typically multichannel reactors with residence times and residence time distributions corresponding to single-tube systems. Since detailed information on activity and selectivity is required in a stage II project, the method of choice for the analysis is mostly gas chromatography for complex product mixtures or IR spectroscopy in reactions involving only relatively simple molecules. A stage I and a stage II project can also be combined in order to increase the efficiency of the methodology. Initially, a stage I project is pursued in order to identify promising leads. These leads are then used as starting points for a stage II project in which the compositional space around the initial leads is more carefully explored. In either case, before the HTE project is started, the design of the compound library is crucial. This is often guided by the intuition and knowledge of the chemist, supplemented by statistical design tools. However, more and more artificial intelligence methods are being developed which support the library design in an efficient manner. These techniques are discussed to some extent in Chapter 2.2 and will also be treated later in this chapter.

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9.3.3

Technology Components 9.3.3.1

Synthesis

9.3.3.1.1 General Remarks In general, synthesis of solid catalysts is a difficult task to be carried out in highthroughput mode. Although several catalyst synthesis methods can be parallelized to some extent, in many high-throughput programs specific synthetic steps are often still being carried out manually and in a sequential fashion. Since the implementation of automated synthesis protocols is a highly time-consuming task, one should always ask oneself, when entering an HT program in heterogeneous catalysis, whether the implementation of an automated synthesis protocol is justified by the savings in time compared with the conventional catalyst synthesis. Typically, the effort is justified when the number of catalysts to be prepared is high, the variability in the synthetic protocol is low and the operations involved are reasonably simple. If the variability of the protocols is too high, then the work needed in programming synthesis robots is often not compensated for by the time savings, although user-friendly interfaces to synthesis robots are available from companies such as Symyx Technologies and hte AG as parts of their software suites. Catalyst synthesis may involve the handling of solids and/or liquids. In general, the processing of liquids is easier, which is mostly due to the very different properties that solids may have. Solids can vary with respect to particle size, stickiness, electrostatic charging, moisture content and many other variables, which make the automated handling of solids difficult. However, if the nature of the solids does not vary much, solids can also be processed by robot systems. For the automated weighing of solids in order to adjust concentrations of different catalyst components, the desired solids weight is mostly only approximated by the automated weighing process but precisely measured, and then solution volumes in the subsequent steps are adjusted according to the real weight of the solid. If the solids themselves cannot be easily processed, handling them as suspensions is an alternative option. However, care has to be taken in this case to avoid settling, for instance by constantly shaking the vessels with the suspended solid. Possible contamination of the synthesis apparatus should also be kept in mind, since this is a twofold concern in HTE synthesis of solid catalysts: First, cleanup of the synthesis vessels and the equipment can be a very time-consuming step which is often forgotten when planning a high-throughput workflow. However, it can become the bottleneck with respect to timing, since the References see page 2072

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9.3 High-Throughput Experimentation in Heterogeneous Catalysis

automation of cleaning procedures is very difficult and one therefore often has to resort to manual operation. Second, the cleaning of HTE equipment has to be done especially carefully, since often only small vessels with a high surface to volume ratio and small amounts of materials are involved. Thus, even traces of contaminants can result in rather high concentrations of impurities in the resulting solids. Ideally, synthetic steps are carried out in one of the standard titerplate formats used in the pharmaceutical industry, i.e. 24-, 96- or 384-well format on their standard footprint. This facilitates the use of commercial instrumentation and software, since many tools have been developed for the needs of HTE in drug discovery. However, often this is not possible, because the solids amounts synthesized in such formats would be too small or the synthetic format is not compatible with the requirements of the subsequently used screening technology. Since the scalability of a catalyst synthesis is an issue in a high-throughput program, preference is given to methods which either seem to be easily scalable from the microgram to the gram and larger scale or to such methods which directly result in milligram to gram quantities of materials. Therefore, a number of other solutions deviating from the standard formats of the pharmaceutical industry have been used in the catalysis sector. Fortunately, most of the available instrumentation is freely programmable and therefore can be adapted to almost any need. Although catalysts are mostly synthesized in individual vessels, such as the wells of titer plates or in beakers, an alternative synthetic protocol is the split-and-pool technique, the principle of which is shown schematically in Fig. 2. Here the materials are synthesized in a highly efficient manner, but one needs an encoding scheme to trace back the course which each synthesized bead has taken through the synthetic sequence. This can be done, for instance, by analyzing the chemical composition with an imaging X-ray fluorescence system, and this has proven to be successful in a number of cases [19–22]. The split-and-pool protocol in combination with an adapted workflow and suitable reactor formats is very powerful in achieving very high throughputs. However, some problems, such as cross-contamination of the different beads in the same vessel, are unique to this approach and have to be carefully considered. Thus, the majority of high-throughput programs use the production of different catalysts in spatially separated domains. 9.3.3.1.2 Robotics and Automation Systems The simplest tools for assisting catalyst synthesis are dispensing robots. Such systems are mostly mounted on a flat-bed system and have single or multiple syringe systems

Pool

Split

Split-and-pool principle for the example of the synthesis of 27 different compounds. In the first step, beads are reacted with a first compound in the three different vessels. Then the contents of the three vessels are combined and split up again into three fractions. Each flask then contains a mixture of all three compounds. Then in each vessel again a reagent is added, different for each flask, resulting altogether in the formation of nine different compounds. The procedure is repeated, so that finally 27 compounds are created in only nine reaction steps. Tracking of the beads by some kind of tagging is necessary to evaluate which pathway the beads have taken through the different flasks.

Fig. 2

for dispensing liquids or – in fortunate cases – slurries. Dispensing systems are supplied by a number of different companies, such as Zinsser [23], Tecan [24], Gilson [25] and others (mention of company names throughout this text does not imply any endorsement or judgement of the quality of the products and services offered). They are distinguished by the size of the working area, the flexibility with respect to syringe volume and syringe needles and the support in programming the robots. All systems are fully software controlled and the dispensing programs can be uploaded either directly from a computer or via disks or memory sticks. Figure 3 shows a typical dispensing robot in operation. Volumes which can reliably be dispensed range from the microliter range to several milliliters with one filling of the syringe. However, the systems may have to be carefully calibrated for each liquid which should be delivered, since the exact amounts which are dispensed depend on a number of factors, such as the viscosity of the solutions, their surface tension, the temperature and others. Less common are systems which are able to deliver highly viscous fluids which can be obtained, for instance,

9.3.3 Technology Components

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freely programmable and can work on footprints as large as 2 × 2 m2 . In such systems, one can integrate precipitation units, impregnation units, filtration or centrifugation stations and so on. However, such highly integrated solutions are rather expensive, require substantial maintenance and also need major programming efforts. Hence they are most useful for industrial environments, where a large number of repeated operations are carried out in well-developed, robust workflows, but less so for syntheses with a high variability in procedures.

Photograph of a dispensing system (Packard) in operation. Courtesy of hte AG.

Fig. 3

in slurry impregnation processes or in zeolite synthesis. Dedicated instruments have been developed which are able to handle such mixtures, such as the system described by Corma’s group suitable for handling zeolite synthesis gels [26] and the multipurpose formulation platform commercialized by hte AG (Fig. 4). However, these are not routine instruments yet and therefore the fully automated handling of highly viscous fluids or fluids with unusual rheological properties still poses problems. There are scattered reports on fully integrated synthesis stations which are mounted on industrial robot platforms [27, 28]. Fully integrated systems can be centered around SCARA (Selective Compliance Assembly Robot Arm) robots or based on a Cartesian coordinate robot platform which moves linearly and at right-angles in three coordinates (called a gantry robot if the horizontal member is supported at both ends). These systems are

Photograph of a detail of an automated formulation workflow suitable for dosing of highly viscous fluids. Courtesy of hte AG.

Fig. 4

9.3.3.1.3 Impregnation Impregnation is one of the most often used methods for the synthesis of supported catalysts. For HT approaches, impregnation has many advantages: Impregnation can start with pre-shaped support materials which have the correct size for optimized fluid dynamics in the reactor. If instead of shaped bodies powders are synthesized, these could create problems in the later catalytic test, such as excessive pressure drop or ill-defined fluid dynamics. This is avoided if shaped bodies are used as the starting materials. Impregnation also typically needs relatively small fluid volumes and, after the impregnation step, solids are obtained which do not need excessive processing, but can often be directly calcined. Moreover, adjusting the concentration of the (supposedly) active materials is simple, since the pore volume of the support is known and thus the loading is determined by the known concentrations of the stock solutions. Automated impregnation processes have therefore been described several times, be it by using simple dispensing systems [29] or completely integrated robotics platforms [27]. One of the greatest problems in the synthesis of impregnated catalysts, whether by conventional means or by HT techniques, is the possible inhomogeneity of the resulting catalysts. In the conventional approach, this can be avoided relatively easily by agitation, for instance by kneading the paste during the impregnation. This is in principle possible also in a parallelized setup, for instance with a magnetic multistirrer plate, but the forces and intensity of agitation often fall behind those achieved when impregnating larger amounts in a single batch. Nevertheless, impregnation processes are possibly the most suitable synthetic method for parallelization. Figure 5 shows selected examples of catalysts which have been prepared by parallelized impregnation processes. As a special case, slurry impregnation of slurries filled with up to 50% of solid has been described [30]. The slurries were then used to coat monolithic materials, socalled miniliths, which are hollow cylindrical extrudates having a porous, structured interior (top row, middle and References see page 2072

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9.3 High-Throughput Experimentation in Heterogeneous Catalysis

be equally suited for precipitation [32]. A centrifugation system has also been described by Bein’s group [33] and is also used at hte AG [34]. An alternative to filtration or centrifugation is the so-called evaporation approach, where the solvent is removed by cryopumping [35]. The workflow described in the latter publication is also capable of pelletizing and sizing the synthesized powders in an automated fashion.

Photographs of different catalysts prepared by automated synthesis. The middle and right photographs in the upper row show the miniliths which can be fitted directly into the channel of a parallel reactor. Courtesy of hte AG.

Fig. 5

right, in Fig. 5). The miniliths can be loaded directly into the channels of adapted high-throughput reactors. 9.3.3.1.4 Precipitation Precipitation is a synthetic procedure which is relatively difficult to implement in a parallelized approach, because typically relatively high solution volumes are needed to synthesize sufficient amounts of solid which can be reliably isolated and catalytically evaluated. In addition, compared with impregnation, where ideally no post-impregnation steps except drying and calcination are required, in the case of precipitated catalysts the filtering or centrifugation step introduces severe complications. Nevertheless, several setups have been described in the literature which can be used for parallel synthesis via precipitation. The degree of parallelization is typically not very high in order to accommodate the high solution volumes. Hoffmann et al. [31] modified a commercial dispensing unit by replacing the needle of the syringe with one of larger diameter. Such larger diameter needles allow the handling of suspensions which can be transferred between containers. In order to avoid settling of the suspension, the beakers have to be constantly shaken, which maintains a homogeneous distribution of the solid. Whether this is possible or not depends on the density of the solid. Dense catalytic materials, which precipitate as relatively large particles, cannot be handled in this way. Separation of the precipitated solid is possible by either filtration or centrifugation, both of which can be parallelized or automated. Filtering modules were described, for instance, by Hoffmann et al. [31] for a precipitation system and by Bein’s group for a hydrothermal reaction setup (see below) which would

9.3.3.1.5 Hydrothermal Synthesis A special case of a precipitation reaction is hydrothermal synthesis, which is the primary method for the synthesis of zeolites and other microporous materials and which is carried out in HTE mode in autoclave blocks or so-called ‘‘multiclaves’’. Hydrothermal synthesis is distinguished from a conventional precipitation reaction by the higher temperatures used and the resulting autogenous pressure, which exceeds atmospheric pressure and thus requires the use of autoclaves. In addition, the times in which a zeolite forms are typically longer than the times for completion of the precipitation of simpler products and often the liquids involved are highly viscous or even a gel is formed. Compared with other synthetic procedures for catalysts, zeolite synthesis thus poses special problems which have been discussed by Newsam et al. [36]. These include:

• inhomogeneities of precursor mixtures • dispensing of microliters of fluids of different viscosities and vapor pressures • mixing of initial solutions or gels • mixing of multiclave well content under reaction conditions • evaporative fluid losses • oxidative degradation, such as of organic additives • contamination from debris, dirt or chemical attack on multiclave components. In spite of these problems, parallelized synthesis of zeolites was one of the first problems tackled in highthroughput synthesis, probably caused by the largely heuristic approach in the discovery of novel zeolites and long synthesis times often required, which made the acceleration of this synthesis route particularly attractive. Akporiaye et al. at Sintef described the first parallelized zeolite synthesis [37], which was carried out in a 100well autoclave. Depending on the conditions, the block consisted simply of Teflon, but a steel-housed version is also described. However, since only 0.5 mL of synthesis solution was present in each well, the pressures could be handled without any metal in this case. The autoclaves were filled with precursor by a commercial dispensing robot. Product isolation and analysis were done manually, which is the greatest disadvantage of the approach. Later publications addressed this issue. Maier and coworkers

9.3.3 Technology Components

synthesized zeolite libraries directly in minute amounts on a flat silicon substrate which could be analyzed without prior isolation using a focusing XRD system [38]. However, for catalytic evaluation in typical zeolitecatalyzed reactions, the sample amounts are probably too small and therefore this method serves more for the screening of a synthesis space. Preparative amounts of materials are thus better prepared in autoclave blocks and solutions for the automated and parallelized recovery of samples, either by centrifugation or filtration, have now been described [32–34]. Figure 6 shows an autoclave block system for the synthesis of larger sample amounts used at hte AG from which samples are isolated by centrifugation. 9.3.3.1.6 Sol–Gel Synthesis The sol–gel synthesis is in principle a very suitable technique for the preparation of solids for stage I screening, since the material can be directly created on a substrate which can later be evaluated as a whole by a primary screening method. It has been used, for instance, by Maier and coworkers for the synthesis of fairly large libraries, which were subsequently evaluated by IR thermography [39, 40], and by Cong et al. [41]. Basically, for this synthesis one needs a matrix which is easy to gelate, such as a silicon alkoxide or a titanium alkoxide, and other elements are added to this matrix precursor. Gelation is induced by admission of water from the air and solvent evaporation, which is easy, since only minute amounts of materials are prepared on the substrates. Sol–gel synthesis is not a generally applicable technique, since due to the requirements of a gellable matrix the choice of materials is limited. On the other hand, within a given matrix material, a high degree

(a)

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of variability is possible due to the diversity of dopant elements. 9.3.3.1.7 Other Techniques for High-throughput Synthesis of Catalytic Materials A variety of other methods have been described which are less generally applicable than those described in the preceding sections and have more limited scope. However, for specialized purposes, such methods can be highly interesting. A range of different methods have been used primarily for the purpose of stage I screening, in which libraries are prepared as thin films on flat substrates. These include evaporation and sputtering techniques using different types of masks [42–44] or ink-jet printing [45]. Although for electrocatalytic evaluation such a synthesis protocol seems scalable, this may pose more problems if the catalysts are to be used in gas-phase reactions and therefore such thin-film libraries are not often used for such purposes. An interesting method relies on the activated carbon route for the production of high surface area materials [46]. According to this route, the pores of activated carbon are impregnated with oxide precursor solutions and then the carbon matrix is removed by calcination. This method is easily parallelized [47] and spherical oxide particles can be produced directly, if suitable spherical carbons are used as the carbon matrix [48].

Characterization Tools High-throughput characterization is one of the least developed technology components in the field. This is primarily 9.3.3.2

References see page 2072

(b)

Autoclave blocks for parallelized hydrothermal synthesis (a) and centrifugation module (b). For centrifugation a filter-paper between two Teflon elements is placed on top of the 16 Teflon inserts and the system is placed upside down in the centrifuge to remove the liquid. Courtesy of hte AG.

Fig. 6

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9.3 High-Throughput Experimentation in Heterogeneous Catalysis

due to the fact that this step is not an absolute necessity in a screening program, if one is interested only in the identification of a well-performing catalyst. Therefore, initially, the development of characterization tools was not pursued with high priority. However, if a deeper understanding is desired and composition-property–activity relationships need to be explored, it is mandatory to collect as much information as possible on the properties of the catalytic materials. A brief survey is given in the following sections on parallelized analytical methods. For phase analysis, X-ray diffraction (XRD) is the method of choice. Increasing the throughput in XRD has become a necessity in connection with the advances in parallelized zeolite synthesis, since the high rate of synthesis shifted the bottleneck towards phase identification. For a number of years, diffractometers equipped with microdiffraction possibilities and x, ysample stages have been available which can be used for high-throughput phase analysis. The key element in these setups is a G¨obel mirror [49], a parabolic multilayer device which allows the generation of a precisely collimated beam which can be focused to spot sizes below 100 µm. This makes it possible to analyze many different spots on a single substrate with reasonable signal-to-noise ratio, as shown in the initial work by Klein et al. [38]. For the determination of the elemental composition of catalysts, X-ray fluorescence (XRF) is a suitable method in many cases. However, sample preparation for dedicated XRF machines is complex, since the sample has to be embedded in a matrix to obtain reliable results. This kind of sample preparation is not possible if high throughput is to be achieved. Therefore, systems have been developed in which a focused X-ray beam is used as in microdiffraction, but instead of collecting the diffracted X-rays, the element-specific fluorescence is detected. Since the matrices and the sample geometries are rather ill-defined in this measurement mode, analysis is more of a semiquantitative than quantitative character. Nevertheless, for screening purposes this is a valuable method. The method has been used in the analysis of the elemental composition of sol–gel-derived materials [40] and in the analysis of the beads obtained from a split-and-pool protocol [20, 21]. Commercial systems are available, as for instance the Eagle II µprobe from Roentgenanalytik [20, 40] and the ThermoNORAN Omicron system [21]. These do not use a G¨obel mirror, but a capillary focusing element. The effect, however, is similar: a spot size on the order of 100 µm can be achieved. Parallelized IR spectroscopy has become possible with the commercial availability of so-called FPA (focal-plane array) detectors. These detectors integrate a large number of individual detector elements in one array, typically 64 × 64 elements, on to which the area which is to be

analyzed is imaged by mirror optics. In combination with an FTIR spectrometer, each of the 4096 elements is capable of recording a full IR spectrum with a time resolution down to the seconds range. The use of such FPA detector setups was pioneered by Lauterbach’s group for the parallel characterization of solid samples and the product gas stream from catalytic reactors [50, 51]. An important improvement, which this group has made, was the introduction of the rapid scan mode instead of the previously employed step scan mode. This improved the time resolution by a large margin [52, 53]. For the study of the interaction of probe molecules, such as CO or pyridine, with solids, cells have been described which allow the parallel study of different samples. We have developed a sorption cell for use in connection with an FPA IR setup for the analysis of eight samples in parallel [54, 55]. This cell is integrated into an imaging mirror system which uses the beam coming out of the external port of an FTIR spectrometer. In order to facilitate sample handling, the samples are placed on a horizontal CaF2 window, and IR light passes through the cell from top to bottom. The whole cell can be evacuated and heated, so that sample activation is possible. Optical spectroscopy could in principle also rely on spatially imaging detectors, but no system for application in the characterization of solid catalysts seems to have been developed so far. However, a setup for time-resolved UV/visible characterization of materials, amongst them also catalytic materials, has been described, which relies on fiber optics in combination with an x, y-sample stage for positioning the sample in the focus of a laser beam [56]. TPD of probe molecules is a method which is often used for the analysis of specific sites of solid catalysts. In a parallelized implementation [57] of this technique, the samples to be studied were placed in a parallel channel reactor body built analogous to that described by Hoffmann et al. [31]. The effluent from this reactor could be directed to a mass spectrometer by a multiport valve. With flushing and analysis times of 8 s per channel, 10 samples can be analyzed in 1 min 20 s, which proved to be appropriate for achieving sufficient resolution in TPD experiments at a heating rate of 10 K min−1 . Finally, a setup for transient kinetic analysis using a parallelized TAP (temporal analysis of products) reactor system has also been described [58]. The actual experiments in this setup are done sequentially on a maximum of 12 samples. Pretreatment times, however, are largely reduced, since all samples are treated simultaneously. Reactors and Analytics In the initial phase of HTE in heterogeneous catalysis, it seemed as if the design and construction of robust and 9.3.3.3

9.3.3 Technology Components

versatile parallel reactors would be one of the greatest challenges, although there was some precedence for mildly parallelized systems in the literature, as discussed in the Introduction to this chapter. However, many of the problems with reactor construction have been solved in the first decade of HTE in heterogeneous catalysis and for most of the challenges there are now solutions at hand, which will be described in more detail in the following. With respect to reactor technology, it is useful to discriminate between stage I and stage II systems, because technologically they are substantially different. In addition, whereas in stage II systems generic analytical tools can be used, because such reactors are basically compatible with any analytical configuration, this is different in stage I systems. Here, analytics and reactor layout are mostly closely correlated, so that one cannot be discussed without the other. In the following, first generic methods suitable for use with high-throughput reactors are discussed, then stage I systems are covered and, finally, stage II reactors giving essentially identical, sometimes even better, data quality as conventional single-tube reactors are introduced. The latter requires that no compromise is made in the precision and reliability of the analytical data and thus the highest standards have to be applied with respect to analytics. 9.3.3.3.1 Generic Analytical Methods In stage II systems, the analytical tools are normally linked to a reactor via multiport valves which select sequentially the different channels of the parallel reactor and direct it to the analytical instrument. There may therefore be a substantial volume between the reactor and the analytical unit which needs to be completely flushed before the next sample is analyzed. This problem is even increased for work at elevated pressure, since then the linear flow velocities are rather low, even if the volumetric flow at standard temperature and pressure is substantial. It is recommended that at least 10 times the volume of the whole system between the reactor exit and analytical instrument has flowed through the setup before injecting the next sample [27]. If the time needed for flushing is longer than the cycle time of the analytical unit, it is therefore not advisable to invest efforts in adding analytical instruments or decreasing the analytical cycle time, but instead the tubing and valve system should be optimized. Here, as in all other components of a high-throughput program, it is essential to analyze the true bottlenecks which limit the throughput.

A Gas Chromatography Gas chromatography (GC) is probably the most versatile method for the analysis of complex mixtures leaving a catalytic reactor and in most cases is therefore the method of choice for

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the analysis of products in a stage II program. Since many heterogeneously catalyzed reactions are gas-phase reactions where the products have sufficient volatility, they can in principle be analyzed by GC on suitable columns using adapted temperature programs. GC is not an inherently fast analytical method. Especially if the separation is not isothermal, the cooling times can increase the overall cycle time substantially. Nevertheless, using optimized systems, cycle times in the minutes range are possible for selected problems. Fast analysis with GC can be achieved by using multiple instruments with one HTE reactor, or at least multiple columns with adapted valve systems, carefully selecting columns and separation programs or multidimensional techniques [27, 59]. Discussing this in more detail would far exceed the scope of this chapter, and readers are referred to specialized publications. Two advanced GC methods deserve to be mentioned here. On the one hand, so-called micro-GC systems offer great opportunities. In these GCs the centerpiece, the column, is miniaturized and this results in short columns and very rapid heating up and cooling down. This substantially decreases analysis times. A drawback of these systems is the limited range of compounds which can routinely be analyzed with micro-GCs due to the restricted choice with respect to the columns. On the other hand, two-dimensional chromatographic methods allow shortening of analysis times, since those sections of the chromatogram that are difficult to resolve and thus would need longer run times are injected into a second column, which allows efficient separation of these compounds. These techniques are also discussed in more detail in specialized publications [59]. An exciting emergent technology for GC, which is not commercially available as yet, is so-called Hadamard transform GC. Combining specially designed injectors with pseudo-random injection sequences and deconvolution methods allows a true multiplexing GC, which reduces cycle times down to the seconds range [60]. Interesting developments in this field can be expected in the next few years. B Mass Spectrometry Mass spectrometry is another generally applicable technology, which can be adapted to almost any high-throughput reactor. However, compared with gas chromatography, mass spectrometry is less easily quantified due to fragmentation and resulting overlap of signals. Also, signal stability is often only limited in mass spectrometry, so that an internal standard is mandatory. For highly precise analyses, as required for a true stage II program, mass spectrometry is therefore less suitable and References see page 2072

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most reports in which mass spectrometry has been used as the analytical method are rather found in the stage I domain. Most mass spectrometric methods are implemented in the scanning mode. This means that either the library is moved in front of a fixed inlet capillary to the mass spectrometer [61] or that the inlet capillary is mounted on an x, y, z-stage and moved into the outlet channels of the parallel reactors or into the vicinity of the catalytically active spot in a library on a substrate [62–65]. Whether the mass spectrometric analysis is more a stage I or stage II method is less dependent on the quality of the mass spectrometer itself, but more on the type of reactor and the mode of reagent feed. If the inlet capillary is moved into the outlet of the channels of well-designed catalytic flow-through reactors, relatively high-quality data can be generated, provided that the mass spectrometer is well calibrated and the product mixture is not too complex. However, this mode of operation is not the prevailing one. In some implementations, the catalysts are not evaluated in flow-through mode. Rather, catalysts are placed on solid substrates and the reagent gases are passed over the surface of the solid, with rather ill-defined residence time distribution. Often two capillaries are used in such systems, either concentrically arranged [61] or bundled together [62]. One of the capillaries feeds the reagents to the surface of the catalyst and the other transfers the gas after reaction into the mass spectrometer. In other setups, the reagents are passed through channels in which the catalyst is located, but the flow patterns and the form in which the catalyst is employed differ strongly from those of more conventional fixed-bed reactors [63–65]. These different types of reactors are discussed in more detail in Section 9.3.3.3.2. C IR Spectroscopy IR spectroscopy is a quantitative method if the system is properly calibrated. It is well suited for the analysis of gas mixtures which are not too complex, because otherwise the strongly overlapping bands make quantification very difficult. Thus, it has been used for the analysis of reaction products from DeNOx reactions and CO oxidation. Two principally different IR analytical methods can be used, non-dispersive IR analysis, where the specific absorption of the target molecule is used to heat the gas in the detector cuvette, which is then converted to an electrical signal, as realized in the URAS instruments [66], or IR spectroscopy, where absorption intensity is recorded against wavenumber. Using nondispersive IR analysis, only one target compound can be quantified in a single channel, whereas IR spectroscopy can also be used to quantify product mixtures, provided that they are not too complex. Nowadays, Fourier transform (FT) IR spectrometers are mostly used for

wavelength-dependent analysis, even if technically FTIR spectrometers do not use dispersive elements. Analysis times, with both non-dispersive and dispersive IR analysis, are normally very short, in the range of seconds. However, since the signal intensity is governed by the Lambert–Beer law, it is dependent on the extinction coefficient of the analyte and the pathlength of the light in the gas mixture. For gases with low extinction coefficients, relatively long pathlengths are necessary. The analysis time per channel is therefore mostly not determined by the analysis time itself, but by the purge times needed to replace completely the gas in the analysis cuvette. For FTIR spectrometers, multipass cuvettes are available which realize long pathlengths of the light in the gas mixture without having an overly long cuvette by using mirrors in the cells. This reduces the length of the cuvette, but still relatively high volumes with correspondingly long purge times are necessary to accommodate the mirrors. Therefore, conventional non-dispersive or dispersive IR instruments are in principle suitable for HTE, but often not the optimal solution due to relatively long cycle times. The situation is different with FPA detector instruments. Since now several thousand IR spectra can be recorded simultaneously, multicuvette setups can be analyzed in truly parallel fashion. Lauterbach pioneered the use of this technique for the analysis of gases from parallel reactors by passing it through multicuvette setups [53, 67, 68]. If the cuvettes are used in transmission mode (Fig. 7a) [30], parallelization degrees higher than 15–20 are difficult to achieve, since the plumbing of the connections to the parallel reactor and within the cuvette becomes more and more difficult. This problem can be solved by using the reflection mode instead of transmission spectroscopy (Fig. 7b) [69]. However, this increases problems in aligning the setup and specific reactor–cuvette combinations are required which are not as generally applicable. Nevertheless, IR analysis in combination with an FPA detector is a highly versatile method for truly parallelized, precise analysis of the products of heterogeneously catalyzed gas-phase reactions. The disadvantage is the relatively high price of such systems and the maintenance efforts required to keep all mirrors properly aligned. Stage I Reactors and Assays Several stage I systems have A Mass Spectrometry been described in which the catalyst is placed on a flat substrate and the gases are passed over its surface [61, 62]. As discussed above, the gas is fed to the active spot via a capillary and the products are fed into the vacuum of a mass spectrometer via a second capillary. The scanning mass spectrometry reactor described by Cong et al. [61] consists of a stainless-steel system in which the substrate 9.3.3.3.2

9.3.3 Technology Components

2063

just large enough to let the feeding and sampling capillary bundle pass through. Also for this type of reactors it has been shown that cross-talk is low.

(a)

CaF2 window

7 × 7 catalyst array Capillary bundle

Gas inlet nozzle (= gold mirror) Exhaust

(b)

(a) 16-channel transmission cuvette for parallelized IR analysis with FPA detector. (b) 49-channel reflection cuvette connected directly to a parallelized reactor. The IR beam enters the cuvettes on the right-hand side after passing a beamsplitter, is reflected at the end of the capillary bundle at the reactor exit and is reflected back through the window and to the beamsplitter.

Fig. 7

wafer, with the catalysts on it, is mounted on an x, y-stage. The probe is moved on the z-axis toward the active spot to be investigated. This spot is heated from the back side using a laser and the temperature is controlled by adjusting the power of the laser. Since only the investigated spot is heated, cross-talk between adjacent catalyst spots is minimized to a large extent. In addition, the majority of the inlet gas flows on to the spot under investigation, and also the sampling is done close to this spot, which further reduces cross-talk. This reactor system has been shown to allow reliable prediction of trends in catalytic behavior. A simplified version of such a reactor has been employed by Maier’s group [62]. Here the whole reactor is heated and the catalysts are placed in small wells which are mostly covered. The hole in the cover is

B Photothermal Methods The scanning mass spectrometry system can also be coupled with a so-called photothermal deflection detection [35, 70]. For this method, the product gas is introduced into a small gas cell. A laser beam tuned to a specific absorption wavelength of the analyte is passed through the cell and due to the absorption of the laser radiation a temperature field is created, the gradient of which depends on the concentration of the analyte. The temperature field creates a lens effect for a perpendicularly passed probe laser, the deflection of which can be very precisely measured. Photothermal deflection is suitable for the detection of very low analyte concentrations. A related method uses the photoacoustic effect [71, 72]. Also here the specific absorption of laser light is used to heat the analyte stream. The laser is used in the pulsed mode and the heating of the gas translates into a pressure pulse, which can be picked up with sensitive microphones. This method can be used in a truly parallelized fashion, since the intensity of the sound signal gives information on the concentration, while the time delay, according to the sound velocity, between the laser pulse and the detection by the microphone is used to analyze the position of the corresponding channel (Fig. 8). A spatial resolution in the centimeter range can be achieved and the results correspond well with those measured with conventional analytical systems. C Electronic Excitation by Laser Light Other laser-based detection methods rely on laser-induced fluorescence (LIF) [73] or resonance-enhanced multiphoton ionization (REMPI) [74]. For REMPI, the target molecule (in the investigated case benzene) is excited to an excited state with a first laser pulse, then a second, tuned pulse is used to ionize the molecule very selectively. The ions produced are then detected with an electrode array placed on top of each outlet channel of the multichannel reactor. The laser beam is passed over an array of outlets from a parallel reactor by means of a mirror arrangement. However, as attractive as the method appears to be, it does not seem to have been used much subsequent to this first publication, probably due to the complexity of the setup in reactions with more diverse product spectra. LIF is in principle a simpler method. A laser beam is passed over the outlet channels of a parallel reactor (in the feasibility study six channels were used [73]). If the analyte (naphthoquinone as an oxidation product of References see page 2072

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Laser

Microphone signal / V

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Time / s Schematics of the photoacoustic detection setup and recorded microphone signal intensity depending on the time lag between laser pulse and detection at the microphone. The signal for the reactor outlets further away from the microphone is weaker due to the longer distance. This effect is corrected for in the calibration procedure. The sound intensity is a measure of the concentration; the time lag allows assignment of the channel.

Fig. 8

naphthalene in Ref. [73]) is present in the product gas, it is excited and returns to the ground state with the emission of visible light, which can be detected with a chargecoupled device (CCD) camera. However, this method also has not found widespread application, probably because difficulties with interference can be expected for complex product mixtures. The case of LIF illustrates one point which is important for successful parallelization: optical detection methods are easily parallelized, because spatially resolving detectors are readily available. Therefore, some of the most interesting methods make use of parallelized visible light or IR radiation detection coupled with suitable reactor setups. D Fluorescence or Color Assays One of the first publications in the field of HTE in catalysis was directed at the discovery of novel anode catalysts for methanol fuel cells [45]. In this reaction, protons are released. Hence an active catalyst can be detected by the production of protons in its vicinity. In the cited study [45], this detection was effected by the luminescence of a dye that is nonluminescent in the unprotonated state. If an array of catalysts on a substrate is evaluated in the presence of the dye, active spots show up due to their luminescence, and can therefore easily be detected at a glance. Other optical methods rely on a color reaction of the educt or the product with a suitable detection dye [48, 75, 76]. Ideally, the dye

can be impregnated on a filter-paper which is placed downstream of the reactor and connected to the outlet channel array via an adaptor. The color reaction between dye and product or educt gives information on the catalyst performance almost instantaneously for all catalysts. The method is suitable for a very high degree of parallelization and a 529-channel reactor based on a ceramic monolith with square channel cross-section has been described [48]. However, quantification would be extremely difficult and has therefore not been attempted. In addition, the method is suitable only for selected reactions where one possible product selectively causes the color reaction and other reagents or products do not interfere with it. A similar principle is used in a microfluidic reactor setup in a slightly different embodiment [35]. Here the product gases, after having contacted the catalyst in 256 CSTR-like chambers, are trapped on an adsorption unit which can as a whole be detached from the reactor and then sprayed with the detection reagent. Readout is possible with a CCD camera. E Thermography The most versatile, truly parallel stage I screening method which again needs adapted reactor setups is IR thermography. This was in fact the method of choice in one of the first papers on parallelized catalyst evaluation [77]. Spatially resolving IR cameras are used, which allow the thermal signature of an array of catalysts to be recorded. If an exothermic reaction is studied, active catalysts are indicated by their higher temperature compared with the others; for endothermic reactions, they are cooler. Taylor and Morken [78] used such a method in a study of heterogenized molecular catalysts in a very simple setup, a beaker in which the acylation of ethanol with acetic anhydride was studied. Since chloroform was the solvent, the polystyrene beads used for immobilization of the molecular catalysts floated on top and thus their thermal signature could easily be recorded. The most powerful implementation of IR thermography was developed by Maier’s group with the introduction of background correction routines which allow temperature differences as small as some 10 mK to be picked up [39, 79]. Figure 9 shows a reactor used in such a system together with a typical thermographic image. In this reactor setup, the catalysts are placed on a slate substrate because this has been found to be optimal with respect to its own emissivity. It is not a flow-through reactor, but a common gas phase is in contact with all catalysts simultaneously, so that no detailed information on activity can be obtained. One should also keep in mind that IR thermography gives only global information on activity, and no selectivity information. For any more complex reaction, it is therefore only useful as a pre-screening

9.3.3 Technology Components

IR-camera

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Activity (all materials)

Reactants

Products x Sampling capillary

y z

Selectivity (only active materials) e.g. MS

Fig. 10 Schematics of the King system. IR thermography is used for pre-screening; active spots are analyzed in more detail with mass spectrometry. Reproduced from Ref. [27].

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Reactor for thermographic imaging of a catalyst library (a) and background-corrected thermal image (b). The bright spots in the thermographic image indicate active catalysts. Reproduced from Ref. [39].

Fig. 9

tool. The need to employ IR-transparent windows also complicates the construction of the reactors. The problem with the lack of selectivity information can be solved by the combination of thermography with a selective analytical method. hte AG has developed the King system [27], in which thermography and mass spectrometry are combined (Fig. 10). First the library is evaluated as a whole by thermography. The thermographic image is analyzed by a software routine and then the inlet capillary of the mass spectrometer is moved to only those spots which appear to be interesting according to their thermal signature. This saves substantial analysis time, because even for a fast method such as mass spectrometry, analysis times are in the minutes range. F Massively Parallelized Microreactors Many of the stage I systems described above are scientifically very interesting, but from a practical point of view either too

complex or too specific for wide-ranging and robust use in an industrial environment. For this, more versatile and generic reactors and methods with fully adapted workflows are needed to be able to achieve the desired very high throughput. Microstructured reactors are systems which allow this high integration [80]. The first systems had been assembled from individual micromachined inlays [81]. However, nowadays the most highly integrated reactors are micromachined on the basis of a standard silicon wafer technology [82, 83] with parallelization ranging up to 625 wells, after earlier generations based on the same principle had 105 or 384 wells [27]. The analytical tool of choice for such microstructured reactors is scanning mass spectrometry in which an inlet capillary mounted on an x, y, z-drive is positioned at the exit of the reactor. Trends in reactivity and selectivity established with such highly integrated reactors correspond well with those measured in conventional setups, so that it is indeed a useful very high-throughput screening tool. However, as innovative and intellectually fascinating as the stage I screening tools are, most industrial projects skip the optional stage I screening stage and enter directly a high-throughput program with a stage II project, since in most cases sufficient prior knowledge has been accumulated (or at least one thinks so) that the more detailed information from a stage II approach is needed. Stage II Reactors A Flow Distribution Various types of stage II reactors are the work-horses of HTE in heterogeneous catalysis. A selection of different types is shown in Fig. 11. The degree of parallelization is typically below 100; mostly 16or 48-channel reactors are used. Most of the described 9.3.3.3.3

References see page 2072

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9.3 High-Throughput Experimentation in Heterogeneous Catalysis

.. Photographs of different stage II HT reactors. From upper left clockwise: 49-channel reactor (Max-Planck-Institut fur Kohlenforschung), 49-channel reactor (hte AG), 16-channel high-pressure reactor (hte AG), 16-channel pretreatment unit (hte AG), .. 48-channel slice reactor (hte AG), 49-channel high-pressure reactor (Max-Planck-Institut fur Kohlenforschung).

Fig. 11

systems are based on the same construction principle, although they differ in detail [27, 29, 31, 64, 84–86]: they have a common gas feed which is distributed to the channels of the high-throughput reactor. At the reactor exit, a multiport valve is used to direct one or more gas exit streams to an analytical instrument, for a stage II screening mostly GC, and the rest is vented. In each of the channels, which have typical diameters of several millimeters, the catalyst of the correct grain size is placed in a fixed bed as in most conventional single-channel laboratory reactors. In most setups, liners or cartridges are used as inserts in the reactor to facilitate handling and to avoid contamination of the whole reactor body. Care has to be taken that linear flow-rates are sufficiently high in order to prevent backmixing in the inlet section in front of each cartridge. The flow distribution between the different channels is in most cases passively controlled by placing identical flow restrictors in each channel, for instance as capillaries in the outlet or as metal frits. Since these elements provide by far the highest flow resistance, small irregularities in the packing of the catalyst bed do not influence the pressure drop much and thus similar flows are ensured. Deviations between different channels can be reduced to substantially below 5% even with passive flow control. A particular problem is caused by non-volume-constant reactions, if the flow restrictors are placed downstream of the catalyst bed. Since the flow through each channel is predominantly governed by the volume passing through the constrictions, according to Hagen–Poiseuille’s law, a non-volume-constant reaction would lead to substantial deviations in flow-rate between the channels. In such cases, active flow control techniques have to be used. One

possibility is to place a mass flow controller in the line to the analytical unit [31]. Since the mass flow does not change in non-volume-constant reactions, this ensures identical flows in each of the channels at least during analysis. Alternatively, one can use separate inlet controls for the channel under investigation on the one hand and all other channels together on the other [87]. If the catalyst performance is not too sensitive to slightly changing flowrates, these measures are sufficient to achieve data quality as in a conventional single flow reactor. In principle, one could also use a system where the inlet flows to all channels are controlled. However, such systems would be very complex and substantially more expensive. In addition, failure probabilities for each part of the system have to be multiplied and the higher the number of sensitive parts integrated in a flow system, the lower the overall availability of the system will be. Multiport valves are available from a number of commercial suppliers. Typically the number of ports ranges up to 16 and valve combinations can be used to collect and select gas streams from more than 16 channels. The problem with these valves is their limited temperature and pressure range. If valves should be operated simultaneously at temperatures exceeding 150 ◦ C and pressures exceeding 2 MPa, one has to use customized valves. A suitable solution is the use of a pressurized valve housing, so that inside the valve the pressure differences against which sealing is necessary are modest [85]. B Heating Different methods are used for heating of such reactors, because a homogeneous thermal profile

9.3.3 Technology Components

is crucial. Reactors can be heated by integrating heating cartridges or heating conductors in the reactor body. Alternatively, the whole reactor body can be placed in an oven. For highly exothermic reactions, even fluidized sandbaths have been used to avoid thermal crosstalk between adjacent channels. Otherwise, if adjacent channels are not thermally decoupled, this could create a problem in reactions with high heat effects, such as oxidation reactions. Simulation can help to arrive at reactor designs which minimize thermal cross-talk [27]. A properly designed heating system can reduce temperature deviations between channels over the whole reactor to around ±1 K. C Possible Problems Although high-throughput reactors for stage II screening are commercially available, from relatively simple general-purpose systems with a low degree of parallelization to highly sophisticated, customized systems providing better performance than typical single flow reactors, one may face severe challenges depending on the kind of application. Challenges can include instability of sealing materials, blank activity of the reactor material, corrosion of the reactors, forming and aging of catalysts on-stream performing three-phase reactions or the formation of high-boiling point liquids from gaseous precursors. a Sealing The choice of sealing material depends on the pressure and/or temperature range to be studied. Polymeric materials, such as Teflon, Kalrez or PEEK (polyether ether ketone), can be used at temperatures up to around 300 ◦ C, with the exact temperature being dependent on the grade of the polymer, the pressure and the time for which the seal is exposed to this temperature. If higher temperatures are required, then graphite or metal seals are necessary, which are, however, more difficult to handle. An alternative to using hightemperature sealing materials is the removal of seals from the hot zone. However, this often results in a more complex layout of the reactor setup. The schematic drawing of a parallelized reactor in Fig. 12 shows the regions in which sealing materials are necessary. b Construction Materials Construction materials are also a crucial issue. Depending on the application, different problems with respect to corrosion and blank activity may arise. If CO is present in the reaction gas at high pressure, so-called metal dusting and subsequent corrosion of the system may occur. Blank activity is a problem in hydrocarbon oxidation reactions at high temperatures. One should bear in mind that alloy components may only segregate to the surface over time, hence the absence of blank activity at the beginning

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Reaction gas inlet

Seal Stainless steel

Catalyst cartridge

Fig. 12 Schematic drawing of a high-throughput reactor for high-pressure reactions as shown in Fig. 11, lower left photograph. Seals are indicated by the hatching.

of an experimental campaign does not mean that also after some time on-stream there is no blank activity. It is therefore always advisable to fill one channel of an HTE reactor with an inert material to detect possible blank activity. In order to have an inherent check on the reliability of a setup, it is equally advisable to have one benchmark catalyst with known activity included in every filling to detect possible problems immediately. Special steel grades can help to decrease the problem with blank activity, but at the expense of more difficult machining and much higher costs. Problems and possible materials solutions are, however, very specific for each catalytic system and a detailed and comprehensive discussion would far exceed the scope of this chapter. Readers are referred to more specialized publications [88]. c Gas/Liquid Separation A particular problem is posed by the simultaneous presence of gas and liquid, as in three-phase reactions or in reactions where liquid products are formed from gases. Such reactions are often encountered in the petroleum refining industry, where typically analytical challenges are also severe. In fortunate cases, the whole system can be heated sufficiently high that no condensation of products occurs in any part of the system. The crucial component is typically the multiport valve, because here the most severe temperature limitations exist. However, even if liquid products form and have to be condensed, solutions are available which allow closing of a full mass balance by analysis of the liquid and the gas phase. Unfortunately, no details of the design of such systems have been published, because these are proprietary solutions [27, 89]. References see page 2072

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d Formation and Aging Formation and aging periods are typical effects observed in heterogeneous catalysis. They may elude detection in both conventional and HT experiments in catalysis. However, since in an HTE experiment the performance is typically not monitored continuously for one channel, the probability of missing such effects is higher. In addition, since the catalysts are normally simultaneously exposed to the reaction gas, but analyzed sequentially, performance is not compared for the same time on-stream, unless very complex valve setups are realized which allow a staged switching of the channels on-stream. Hence it is mandatory that for a multichannel reactor each position is analyzed at least twice. If significant deviations between different analysis times are detected, one has to evaluate the temporal behavior of the corresponding catalyst more carefully. D Liquid-Phase Reactors Parallelized liquid-phase reactors are less problematic than gas-phase reactors. The most significant challenges are avoiding mass transfer limitations, sampling under reaction conditions and – often forgotten – cleaning of the parallelized setups, especially if these are multiautoclave benches. For temperatures in the range up to about 100 ◦ C and atmospheric pressure, one can use glass reactor systems with automated liquid handling and sampling [30, 90]. If higher degrees of parallelization are desired, then 96well plates are suitable solutions. One problem in this case is solvent loss, which may significantly influence the reaction, since the reaction volumes are very small to begin with. Therefore, sealed microtiter plates are mostly used, as described by Desrosiers et al. [91]. For reactions at elevated pressure, however, a common headspace for all wells has to be accepted. For sampling after reaction, an autosampler is used with automated injection into analytical systems. For work more closely resembling conventional liquidphase experiments under elevated pressure, autoclave arrays are suitable [92]. Intensive stirring under reaction is mandatory for gas/solid/liquid reactions and also for this purpose reactor systems have been described which can be operated under non-mass transfer-limited conditions [93]. Informatics Environment A well-developed and powerful informatics environment is essential for a fully integrated high-throughput program. The informatics environment has to serve a number of different purposes: it needs to provide the software support for the library design, interface the different analytical instruments and methods, capture data and store them in a powerful database and support the analysis of the data and data mining. 9.3.3.4

9.3.3.4.1 Data Management Data management for HTE in heterogeneous catalysis is a highly complex task, probably more complex than in the pharmaceutical industry, because the type of information which need to be stored and retrieved is more complex. Nevertheless, one can build upon the expertise which has been gathered in high-throughput programs in drug discovery. A database structure and the interface to the experiments have to fulfill a number of different tasks related to catalyst synthesis, analysis and testing [87, 94]. First, reliable sample tracking is essential, so that all information related to the synthesis, characterization and testing of a sample can unambiguously be related to the sample. With respect to workflow management, this can be facilitated by bar coding individual samples and using bar code readers for data input at each station in which the sample is handled. If libraries are managed on a substrate, for instance a 96-well plate, the library is assigned a barcode and each sample can be uniquely identified by its position in a library. The database has to allow storage of all samplerelated information, i.e. synthesis protocol, stock reagents used for synthesis with lot numbers, supplier name, etc., chemical analysis, spectra used for characterization of the sample, machine IDs on which these spectra have been collected, conversion and selectivities of the catalytic tests, conditions of the test, unit ID on which the test was performed, and many others. Data are retrieved and correlated either by home-made query builders and reporters or the standard query language provided by the supplier of the database. In spite of the complexity of the problem, well-developed software packages are available for managing the data generated in a high-throughput program. In the framework of the EU Combicat project, the database package STOCAT, suitable for an academic environment and consisting of different modules, has been developed [87, 95]. This is integrated in a software environment called OPTICAT, which is accessible on the Internet [96]. Industrial solutions are available which allow full integration of the complete HTE workflow, including interfacing of the different instruments included in the workflow. The system introduced by hte AG is based on a standardized language called hteML which itself is based on XML, one of the standard languages for exchanging data on the Internet. Zech et al. [94] described in detail such a software architecture, together with the challenges and possible solutions. To illustrate the overall task, Fig. 13 shows the architecture of this software system [94]. The centerpiece of the software package is the myhte application server. The application server on the one hand links to the database backbone, and on the other integrates a number of applications and tools, such as visualization, statistical analysis and experimental design, but also hteCapture (capturing and combining data from

hteCapture

Statistical data analysis

Analytical instruments

hteControlTM

hteSetup

HTE equipment

3rd party SW

Design of experiments hteML interface layer

Data management Data retrieval Querying and filtering Visualization Reporting

hteControlTM

myhte application server

hteControlTM

9.3.3 Technology Components

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3rd party instrument

Database server

Fig. 13

Software architecture of an integrated HTE software environment. Courtesy of hte AG.

different sources) and hteSetup (transforming high-level experimental description to actual hardware sequence scripts). Interfacing to the hardware is provided by the hteControl modules, which control the operation of HTE test rigs and instruments, and also that of commercial instruments directly or on top of third-party software. This last point is in fact the most problematic, since only a few instrument manufacturers provide open and well-defined interfaces to their data. Hence substantial efforts have to be invested to integrate commercial instrumentation in an HTE workflow. It may therefore be preferable to work with a less powerful instrument, instead of a more powerful one, which has, however, a closed and proprietary interface. Symyx Technologies also have their own proprietary technology suite, called Renaissance , consisting of many modules. This software environment also allows the capture of the full HTE workflow. Industrial solutions, as well as the STOCAT package, are based on commercial database software from Oracle. With these packages, the problem of managing the high-throughput workflow is in principle solved, although, depending on the overall environment, severe adaptation might be necessary and especially the integration of new analytical and synthetic tools can be time demanding and often not straightforward. A problem still is that no common e-language, protocols and standards exist in the field of HTE in catalysis, which makes combination of different software components from different suppliers difficult. Whether this field will follow the path of many other technology sectors and eventually reach a standardization still remains to be seen.

9.3.3.4.2 Library Design and Data Mining Library design and data mining are closely related, because a high-throughput approach typically follows an iterative procedure, in that first an initial library is designed, possibly with software support. The library is then synthesized and evaluated and then the data are analyzed in order to synthesize an improved new library. Thus, the knowledge created in a high-throughput program is immediately fed back into the program. The topics treated in this section are also to some extent covered from a different angle in Chapter 2.2.

A The Initial Library If the library to be designed is the first one for a given problem in an HT program, the intuition of the chemist is needed to create the initial library. However, even if no previous knowledge is available, statistical methods should be used to cover the available search space as well as possible. The chemist will decide on the variables to be studied, for instance, the nature of the components in the catalyst and their concentration ranges. Then the question is how a maximum of information can be obtained with a given number of experiments. For this purpose, a number of different statistical methods, often called design of experiments (DoE) techniques, are known, which are implemented in many commercial statistical software packages and also in the software packages available for HTE in heterogeneous catalysis. Readers are referred to special publications for more information on these References see page 2072

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methods [97] and some applications of DoE methods in catalysis can also be found [98–101]. It would be very interesting, however, to integrate previous knowledge about catalysis in a softwaresupported design process for initial libraries. This kind of knowledge is present as the ‘‘catalytic knowledge background’’ in the mind of the chemist. Accessing such knowledge in a more structured form would be highly desirable. One source of this information could be the available literature. However, on closely inspecting published data, one realizes that it is by no means structured in a useful form. Catalyst synthesis protocols are very often ill-described, catalytic experiments are carried out under widely differing conditions, analytical data are incomplete, and so on. The information from the literature can therefore hardly be used as the basis for a software-supported library design. In the pharmaceutical industry, a ‘‘virtual’’ screening of molecular compounds is based on so-called descriptors, which encode properties of molecules in computerreadable form. Such descriptors can be molecular mass, dipole moment, the presence of certain functional groups, solvent-accessible surface area and many others [102]. For solids, and especially for catalytic solids, these concepts cannot be applied directly. Thus, an approach has been developed [103–106] in which a set of descriptors is first generated by a high-throughput approach, using a library of diverse solids, the catalytic performance of which is tested in a target reaction. This is done by correlating the performance of the solids in the catalytic reaction with the properties of the solids, such as composition, synthesis protocol, post-treatment method and numerous physicochemical parameters. This generates the descriptor set, which is then the basis of a virtual screening. For this, possible catalysts are generated at random by software and the descriptors for each random ‘‘in silico’’ catalyst are determined. With these descriptors, a promising subset of catalysts is determined, which should optimally cover the diversity space or which have a high probability of good performance in a target reaction. Only this subset has to be synthesized, which substantially reduces the experimental effort. The predictive power of descriptors, which were generated following this approach, substantially exceeds the statistical expectation value. Based only on artificial intelligence methods, the software achieves approximately what an advanced chemistry student would be able to do in designing an initial library. B Library Design and Data Analysis in Iterative Programs Compared with the synthesis of an initial library, where basically no prior information is available, more methods exist for improving on existing expertise and finding the optimum catalyst or at least a local optimum. In

principle, if a first library has been synthesized and evaluated, one could try to fit the generated data points with an analytical function which would correlate the catalyst performance, indicated by some metrics, with the composition of the catalyst, pretreatment conditions, synthesis pathway and so on. However, there are two problems with this approach: First, the types of variables which could influence the catalytic performance can be rather diverse, some numerical and continuous, such as concentration of certain elements, but others may be ordinal and discontinuous [105]. This makes the construction of a fit function extremely difficult. Second, even if such a function could be created, it would be very complex, possibly containing discontinuities, in a high-dimensional space. The probability of finding the right function and, based on it, predicting optimum catalyst/condition points, is rather low. Thus, the typical optimization methods which need an analytical function and then follow, for instance, gradients using adapted algorithms cannot be used. Nevertheless, there are methods which are applicable for the optimization of solid catalysts in iterative programs and two strategies have emerged as the most promising ones, i.e. neural networks and evolutionary algorithms, and in the latter class especially genetic algorithms. A full discussion of these techniques would exceed the scope of this chapter. In the following, only the basic principles of these techniques will be highlighted and references to case studies published in the literature will be given. In addition, readers are referred to Chapter 2.2 and a number of more general publications on the use of artificial intelligence methods in heterogeneous catalysis [87, 95, 107]. a Genetic Algorithms Genetic algorithms are modeled after the evolutionary process in Nature. They were first applied for the optimization of solid catalysts by Baerns’ group [108] and since then a number of times [109–111] for the solution of different catalytic problems. A genetic algorithm approach requires several successive generations of catalysts, each generation typically having a size between 30 and 100 in the studies published so far. The process is illustrated schematically in Fig. 14. Catalysts are encoded in the form of ‘‘genes’’ on a ‘‘chromosome’’, where each ‘‘gene’’ corresponds to some catalyst property. In the simplest case, a bit is set = 1 if an element is present in the formulation and to zero if it is absent (this is the example in Fig. 14). However, the ‘‘genes’’ can be more complex, encoding also concentrations of elements or post-synthesis steps, for instance. After the first generation has been synthesized, ideally using DoE methods to maximize the information obtained with this first library, it is evaluated and the best

9.3.3 Technology Components

1st generation

Ag Au Pt Pd Cu Ni Fe Cr

Mutate

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Cross-over

‘‘Winners’’ selected

2nd generation

Fig. 14

Ag Au Pt Pd Cu Ni Fe Cr

Mutate/cross-over 3rd generation

Schematic representation of a genetic algorithm. For explanation, see text.

catalysts of this first generation form the basis for the next generation. Two operators are used to create the next generation, i.e. mutation, the exchange of one bit of the ‘‘chromosome’’ of a catalyst or a set of bits corresponding to some chemical information (a whole ‘‘gene’’), such as the concentration of an element. The other possibility is crossing over, i.e. the ‘‘genes’’ of two catalysts are cut and the corresponding pieces are exchanged, where either just two ‘‘genes’’ could be exchanged or whole sections of the ‘‘chromosome’’. The library created by these two operations is then synthesized and evaluated and the process is repeated, until a satisfactory solution is achieved. This method can be very useful to reach an optimum quickly; however, there is no guarantee that this is the global optimum and not a local one. A number of parameters are decisive for the success of a genetic algorithm, which are related to the algorithm itself, i.e. the mutation rate, crossing over rate and the fraction of retained individuals from the preceding generation. Equally important is the target function according to which the performance is ranked. In heterogeneous catalysis, the target function typically is multidimensional, i.e. contains information not only on activity, but also on selectivity, specific by-products which need to be avoided, catalyst cost and many others. A good deal of effort should therefore be invested in the construction of a suitable target function [111a]. b Artificial Neural Networks (ANNs) ANNs can be considered as a black-box method for the approximation of a performance function. They mimic neural networks in biological systems in that they contain ‘‘neurons’’ which receive information (input), neurons that process the information (hidden) and neurons which serve for returning information (output). Connections exist between the neurons that determine the path of a signal from the input neurons to the output neurons.

Inputs

Hidden layer

Outputs

Fig. 15 Schematic representation of an artificial neural network. The input layer is connected with the output layer via one – or more – hidden layers.

A neural network typically consists of an input layer, in which the parameters which are thought to be relevant for each catalyst are fed into the ANN, the hidden layers and an output layer, which is the performance of the catalysts. Figure 15 illustrates the structure of an ANN. Various different architectures for neural networks are known [112] and are implemented in commercial software packages, such as Statistica [113]. The most often used neural network for applications in heterogeneous catalysis is the so-called multilayer perceptron. If an ANN is used in a catalyst development program, one first needs to ‘‘train’’ the network. As in a biological network, those connections which provide the correct prediction are amplified and those which give a wrong prediction are weakened. For ‘‘training’’, it is therefore necessary to use real data from a training library for which input and output are known. After verification of the training with a second set of catalysts with known performance, the network is ready to be used. It will now predict with some precision the performance of catalysts which have not References see page 2072

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been evaluated before and it can therefore be used as the basis for in silico screening. However, since the hidden layers are a black box and the multilayer perceptron networks are only forward propagating, i.e. signals only pass from input to output layer, no rules can be derived for synthesizing promising catalysts, since there is no unique pathway from a desired output to the input layer, i.e. the description of the catalysts, through the network. Therefore, various different catalysts have to be evaluated in silico and those for which the network predicts the desired performance are synthesized, to arrive at a more focused, better performing library than the initial one. ANNs are not specific to HTE approaches in heterogeneous catalysis. They have been used before, primarily by Hattori’s group [114–116], but success seems to have been limited due to the rather small database available at that time. This improved with the advent of HTE in catalysis, because now catalytic data could be generated much more quickly. Thus, ANNs have, for instance, been used successfully for the development of catalysts for ethane [117] and propane [118] oxidative dehydrogenation, direct dimethyl ether synthesis [119], methanol synthesis [120], hydroisomerization of paraffins [121] and water gas shift reaction [122]. ANNs were also a crucial component in our approaches to developing a descriptor-based design of solid catalysts [103–106]. c Other Software Tools There are a number of other software tools which can support the high-throughput process in heterogeneous catalysis, not least the visualization tools which are implemented in HTE software suites, but can also be obtained as stand-alone solutions from commercial suppliers, such as Spotfire [123]. The so-called holographic search strategy [124] can also be considered a powerful visualization tool. A more comprehensive approach than those described so far was followed by Caruthers et al. [125]. They have formulated a procedure which integrates kinetic modeling of reactions and HTE and they used the results of HT experiments for the formulation of kinetic models. The models in turn were employed to suggest new sets of experiments targeted at maximizing the information and finally at arriving at a valid kinetic model based description of the catalytic system. This comprehensive approach also makes use of other software-based techniques, such as genetic algorithms. Finally, first attempts are being made to integrate quantum chemical calculations and HTE. By quantum chemical calculations, the performance of possible catalysts is evaluated computationally and only the systems with high predicted performance are synthesized, as shown, for instance, for electrocatalysts [43, 126] and for gas-phase methanation catalysts [127].

Outlook Although this field did not exist at the time when the first edition of this Handbook was published, HTE in catalysis can now be considered as an almost routine technology in catalyst development. It certainly does not fully replace the conventional ways of finding and improving catalysts, but it is a valuable component in many development programs. Many major chemical and catalyst companies either have their own HT programs or have formed strategic alliances with technology companies offering integrated services in this field. These include, in order of their foundation and probably also according to their present market share, Symyx Technologies [128], hte AG [89], Avantium [129] and Accelergy [130], in addition to the companies providing hardware for specific parts of the workflow, as mentioned throughout the text. As the field is still young, future developments are difficult to predict. It appears that the major technology components are in place and that the methodology can now be used in production mode. One can hope that the application of HTE will both increase the rate at which novel catalysts are discovered and further the fundamental understanding of catalysis by providing a much broader and consistent database for the generation of new knowledge. 9.3.3.5

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45. E. Reddington, A. Sapienza, B Guraou, R. Viswanathan, S. Sarangapani, E. S. Smotkin, T. E. Mallouk, Science 1998, 280, 1735. 46. M. Schwickardi, T. Johann, W. Schmidt, F. Sch¨uth, Chem. Mater. 2002, 14, 3913. 47. T. Johann, A. Brenner, O. Busch, F. Marlow, S. Schunk, F. Sch¨uth, Angew. Chem. Int. Ed. 2002, 41, 2966. 48. F. Sch¨uth, L. Baumes, F. Clerc, D. Demuth, D. Farrusseng, J. Llamas-Galilea, C. Klanner, J. Klein, A. Martinez-Joaristi, J. Procelewska, M. Saupe, S. Schunk, M. Schwickardi, W. Strehlau, T. Zech, Catal. Today 2006, 117, 284. 49. M. Schuster, H. G¨obel, J. Phys. D: Appl. Phys. 1995, 28, A270. 50. C. M. Snively, S. Katzenberger and G. Oskarsdottir, J. Lauterbach, Opt. Lett. 1999, 24, 1841. 51. C. M. Snively, G. Oskarsdottir, J. Lauterbach, Catal. Today 2001, 67, 357. 52. C. M. Snively, G. Oskarsdottir, J. Lauterbach, Angew. Chem. Int. Ed. 2001, 40, 3028. 53. R. J. Hendershot, P. T. Fanson, C. M. Snively, J. A. Lauterbach, Angew. Chem. Int. Ed. 2003, 42, 1152. 54. O. Busch, W. Brijoux, S. Thomson, F. Sch¨uth, J. Catal. 2004, 222, 174. 55. P. Kubanek, H.-W. Schmidt, B. Spliethoff, F. Sch¨uth, Microporous Mesoporous Mater. 2005, 77, 89. 56. P. Atienzar, A. Corma, G. Garcia, J. M. Serra, Chemistry - Eur. J. 2004, 10, 6043. 57. H. Wang, Z. Liu, J. Chen, H. Liu, Catal. Commun. 2004, 5, 55. 58. A. van Veen, D. Farrusseng, M. Rebeilleau, T. Decamp, A. Holzwarth, Y. Schuurman, C. Mirodatos, J. Catal. 2003, 216, 135. 59. H. J. Huebschmann, Handbook of GC/MS. Fundamentals and Applications, Wiley-VCH, Weinheim, 2001, 608 pp. 60. O. Trapp, Angew. Chem. Int. Ed. 2007, 46, 5609. 61. P. Cong, R. D. Doolen, Q. Fan, D. M. Giaquinta, S. Guan, E. W. McFarland, D. M. Poojary, K. Self, H. W. Turner, W. H. Weinberg, Angew. Chem. Int. Ed. 1999, 38, 484. 62. M. Orschel, J. Klein, H.-W. Schmidt, W. F. Maier, Angew. Chem. Int. Ed. 1999, 38, 2791. 63. P. Claus, D. H¨onicke, T. Zech, Catal. Today 2001, 67, 319. 64. I. Hahndorf, O. Buyevskaya, M. Langpape, G. Grubert, S. Kolf, E. Guillon, M. Baerns, Chem. Eng. J. 2002, 89, 119. 65. S. Senkan, K. Krantz, S. Ozturk, V. Zengin, I. Onal, Angew. Chem. Int. Ed. 1999, 38, 2794. 66. www.abb.com, accessed 25 January 2007. 67. C. M. Snively, G. Oskarsdottir, J. Lauterbach, J. Comb. Chem. 2000, 2, 243. 68. C. M. Snively, J. Lauterbach, Spectroscopy 2002, 17, 26. 69. P. Kubanek, O. Busch, S. Thomson, H.-W. Schmidt, F. Sch¨uth, J. Comb. Chem. 2004, 6, 420. 70. Y. Liu, P. Cong, R. D. Doolen, H. W. Turner, W. H. Weinberg, Catal. Today 2000, 61, 87. 71. T. Johann, A. Brenner, O. Busch, F. Marlow, S. Schunk, F. Sch¨uth, Angew. Chem. Int. Ed. 2002, 41, 2966. 72. T. Johann, A. Brenner, M. Schwickardi, O. Busch, F. Marlow, S. Schunk, F. Sch¨uth, Catal. Today 2003, 81, 449. 73. H. Su, E. S. Yeung, J. Am. Chem. Soc. 2000, 122, 7422. 74. S. M. Senkan, Nature 1998, 394, 350. 75. O. Busch, C. Hoffmann, T. Johann, H. W. Schmidt, W. Strehlau, F. Sch¨uth, J. Am. Chem. Soc. 2002, 124, 13527. 76. K. Omata, Y. Watanabe, T. Umegaki, M. Hashimoto, M. Yamada, J. Jpn. Pet. Inst. 2003, 46, 328. 77. F. C. Moates, M. Somani, J. Annamalai, J. T. Richardson, D. Luss, R. C. Wilson, Ind. Eng. Chem. Res. 1996, 35, 4801.

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78. S. Taylor, J. Morken, Science 1998, 280, 267. 79. G. Kirsten, W. F. Maier, in High-Throughput Screening in Heterogeneous Catalysis, A. Hagemeyer, P. Strasser, A. F. Volpe (Eds.), Wiley-VCH, Weinheim, 2004, p. 175. 80. A. M¨uller, K. Drese, in High-Throughput Screening in Heterogeneous Catalysis, A. Hagemeyer, P. Strasser, A. F. Volpe (Eds.), Wiley-VCH, Weinheim, 2004, p. 89. 81. P Claus, D. Hoenicke, T. Zech, Catal. Today 2001, 67, 319. 82. T. Zech, G. Bohner, O. Laus, J. Klein and M. Fischer, Rev. Sci. Instrum. 2005, 76, 062215. 83. T. Zech, G. Bohner, J. Klein, Catal. Today 2005, 110, 58. 84. J. Perez-Ramierez, R. J. Berger, G. Mol, F. Kapteijn, J. A. Moulijn, Catal. Today 2000, 60, 93. 85. C. Kiener, M. Kurtz, H. Wilmer, C. Hoffmann, H.-W. Schmidt, J. D. Grunwaldt, M Muhler, F. Sch¨uth, J. Catal. 2003, 216, 110. 86. U. Rodemerck, P. Ignaszewski, M. Lucas, P. Claus, M. Baerns, Chem. Eng. Technol. 2000, 23, 5. 87. D. Farrusseng, C. Mirodatos, in High-Throughput Screening in Heterogeneous Catalysis, A. Hagemeyer, P. Strasser, A. F. Volpe (Eds.), Wiley-VCH, Weinheim, 2004, p. 239. 88. G. Kreysa, M. Sch¨utze (Eds.), Corrosion Handbook, Vols. 1–13, Wiley-VCH, Weinheim, 2004–2009, approx. 6500 pp. 89. www.hte-company.de, accessed 25 January 2007. 90. www.chemspeed.com, accessed 25 January 2007. 91. P. Desrosiers, A. Guram, A. Hagemeyer, B. Jandeleit, D. M. Poojary, H. Turner, H. Weinberg, Catal. Today 2001, 67, 67. 92. M. Lucas, P. Claus, Chem.-Ing.-Tech. 2001, 252, 73. 93. S. Thomson, C. Hoffmann, S. Ruthe, H. W. Schmidt, F. Sch¨uth, Appl. Catal. A 2001, 230, 253. 94. T. Zech, A. Sundermann, R. F¨odisch, M. Saupe, Rev. Sci. Instrum. 2005, 76, 062203. 95. D. Farrusseng, L. Baumes, C. Mirodatos, in High-Throughput Analysis – A Tool for Combinatorial Materials Science, E. Amis, R. Potyrailo (Eds.), Kluwer Academic/Plenum Press, New York, 2003, p. 551. 96. OPTICAT, http://eric.univ-lyon2.fr/∼fclerc/opticatUK.htm, accessed 25 January 2007. 97. J. Cawse (Ed.), Experimental Design for Combinatorial and High-Throughput Materials Development, Wiley, Hoboken, NJ, 2003, 318 pp. 98. J. P. Pirard, B. Kalitvenzeff, Ind. Eng. Chem. Fundam. 1978, 17, 11. 99. M. Iborra, J. F. Izquierdo, F. Cunill, J. Tejero, Ind. Eng. Chem. Res. 1992, 31, 1840. 100. P. Rao, S. Divakar, World J. Microbiol. Biotech. 2002, 18, 341. 101. R. J. Hendershot, W. B. Rogers, C. A. Snively, B. A. Ogunnaike, J. Lauterbach, Catal. Today 2004, 98, 375. 102. E. M. Gordon, J. F. Kerwin (Eds.), Combinatorial Chemistry and Molecular Diversity in Drug Discovery, Wiley, New York, 1998, 516 pp. 103. C. Klanner, D. Farrusseng, L. Baumes, C. Mirodatos, F. Sch¨uth, QSAR Comb. Sci. 2003, 22, 729.

104. C. Klanner, D. Farrusseng, L. Baumes, M. Lengliz, C. Mirodatos, F. Sch¨uth, Angew. Chem. Int. Ed. 2004, 43, 5347. 105. J. Procelewska, J. L. Galilea, F. Clerc, D. Farrusseng, F. Sch¨uth, Comb. Chem. High-Throughput Screen 2007, 10, 37. 106. D. Farrusseng, C. Klanner, L. Baumes, M. Lengliz, C. Mirodatos, F. Sch¨uth, QSAR Comb. Sci. 2005, 24, 78. 107. M. Holena, in High-Throughput Screening in Heterogeneous Catalysis, A. Hagemeyer, P. Strasser, A. F. Volpe (Eds.), Wiley-VCH, Weinheim, 2004, p. 153. 108. D. Wolf, O. V. Buyevskaya, M. Baerns, Appl. Catal. A 2000, 200, 63. 109. O. V. Buyevskaya, A. Br¨uckner, E. V. Kondratenko, D. Wolf, M. Baerns, Catal. Today 2001, 67, 369. 110. A. Holzwarth, P. Denton, H. Zanthoff, C. Mirodatos, Catal. Today 2001, 67, 309. 111. J. S. Paul, R. Janssens, J. F. M. Denayer, G. V. Baron, P. A. Jacobs, J. Comb. Chem. 2005, 7, 407; (a) O. C. Gobin, A. Martinez-Joaristi, F. Sch¨uth, J. Catal. 2007, 252, 205. 112. M. T. Hagan, H. Demuth, M. Beale, Neural Network Design, PWS Publishing, Boston, 1996, 628 pp. 113. www. statsoft. com, accessed 25 January 2007. 114. T. Hattori, S. Kito, Catal. Today 1995, 23, 347. 115. S. Kito, T. Hattori, Y. Murakami, Appl. Catal. 1994, 114, L173. 116. S. Kito, T. Hattori, Y. Murakami, Ind. Eng. Chem. Res. 1992, 31, 979. 117. A. Corma, J. M. Serra, E. Argente, V. Botti, S. Valero, Chem. Phys. Chem. 2002, 3, 939. 118. U. Rodemerck, M. Baerns, M. Holena, D. Wolf, Appl. Catal. A 2004, 223, 168. 119. K. Omata, S. M. Hashimoto, G. Ishiguro, Y. Watanabe, T. Umegaki, M. Yamada, Ind. Eng. Chem. Res. 2006, 45, 4905. 120. Y. Watanabe, T. Umegaki, S. M. Hashimoto, K. Omata, M. Yamada, Catal. Today 2004, 89, 455. 121. J. M. Serra, A. Corma, E. Argente, S. Valero, V. Botti, Appl. Catal. A 2003, 254, 133. 122. L. Baumes, D. Farrusseng, M. Lengliz, C. Mirodatos, QSAR Comb. Sci. 2004, 23, 467. 123. www.spotfire.com, accessed 25 January 2007. 124. A. Tompos, J. L. Margitfalvi, E. Tfirst, L. Veqvari, M. A. Jaloull, H. A. Khalfalla, M. A. Elgarni, Appl. Catal. A 2005, 285, 65. 125. J. M. Caruthers, J. A. Lauterbach, K. T. Thomson, V. Venkatasubramanian, C. M. Snively, A. Bhan, S. Katare, G. Oskarsdottir, J. Catal. 2003, 216, 98. 126. J. Greeley, T. F. Jaramillo, J. Bonde, I. B. Chorkendorff, J. K Norskov, Nature Mater. 2006, 5, 909. 127. M. P. Andersson, T. Bligaard, A. Kustov, K. E. Larsen, J. Greeley, T. Johannessen, C. H. Christensen, J. K. Norskov, J. Catal. 2006, 239, 501. 128. www.symyx.com, accessed 25 January 2007. 129. www.avantium.nl, accessed 25 January 2007. 130. www.accelergy.com, accessed 25 January 2007.

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10

Reaction Engineering 10.1

The easiest, and by far the most common, way to carry out a heterogeneously catalyzed gas phase reaction is by passing the gas mixture over a fixed catalyst. The arrangement of the fixed catalyst is generally called a fixed-bed, and the respective reactor a fixed-bed reactor. The simplest type of fixed catalyst bed is a random packing of catalyst particles in a tube. Different particle shapes are in use such as spheres, cylinders, rings, flat disc pellets or crushed material of a certain sieve fraction. Mean particle diameters range from 2 to 10 mm, the minimum diameter being limited primarily by pressure drop considerations, and the maximum diameter by the specific outer surface area for mass and heat transfer. Much effort has been devoted to the question of replacing the random packing by a regular catalyst arrangement to reduce the pressure drop of the flowing gas and/or to increase the heat transfer to a cooling or heating system. Pressure drop considerations led to the introduction of monolith catalysts with parallel channels in catalytic gas purification applications (automotive exhaust and flue gas treatment). The increase of heat transfer to the tube wall remains the subject of ongoing research. The influence of catalyst shape on fixed-bed reactor performance will be discussed in Section 10.1.2. The different types of fixed-bed reactor applied in industry are detailed in Section 10.1.3. The reaction chambers used to remove nitrogen oxides from power station flue gases constitute the largest type of fixed-bed reactor with regards to reactor volume and throughput, while automobile exhaust purification represents by far the most widely employed application of fixed-bed

reactors. Within the chemical industry, fixed-bed reactors are used as standard in heterogeneously catalyzed gasphase reactions, and only if special conditions such as rapid catalyst deactivation or operation in the explosive regime have been taken into account, is the alternative of fluidized-bed operation considered (see Chapter 10.2). With regards to reactor application and construction, it is convenient to differentiate between fixed-bed reactors for adiabatic operation and those for non-adiabatic operation. Since temperature control is one of the most important methods to influence a chemical reaction, adiabatic reactors are used only when the heat of reaction is low, or where there is only one major reaction pathway; in these cases no adverse effects are expected on selectivity or yield due to the adiabatic temperature development. The characteristic feature of an adiabatic reactor is that the catalyst is present in the form of a uniform fixed bed (Fig. 1a) which is surrounded by an outer, insulating jacket. Adiabatic reactor designs are outlined in Section 10.1.3.1, and multi-stage adiabatic reactors with intermediate heat or reaction component addition or removal in Section 10.1.3.2. Reactions with a large heat of reaction, as well as those that are very temperature-sensitive, are carried out in reactors in which indirect heat exchange occurs with a heat transfer medium that is circulated through the fixed bed. As in most cases the task of the heat-transfer cycle is to maintain the temperature in the fixed bed within a specific range, this concept is frequently described as an ‘‘isothermal fixed-bed reactor’’. The most common arrangement is the multitubular fixed-bed reactor, in which the catalyst is arranged in tubes, around which the heat carrier circulates externally (Fig. 1b). These reactors are considered in Section 10.1.3.3. Within a production plant the reactor may justifiably be regarded as the central item of apparatus. However, compared to the remaining parts of the plant for preparing the feed and for separating and processing the products, it is often by no means the largest and most cost- intensive

∗

References see page 2105

Catalytic Fixed-Bed Reactors Gerhart Eigenberger∗

10.1.1

Introduction

Corresponding author.

Handbook of Heterogeneous Catalysis, 2nd Ed. .. .. Edited by G. Ertl, H. Knozinger, F. Schuth, and J. Weitkamp Copyright  2008 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim ISBN: 978-3-527-31241-2

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10.1 Catalytic Fixed-Bed Reactors

Feed gas

Feed gas

Heat carrier Catalyst bed

Heat carrier

(a)

(b)

Basic types of catalytic fixed-bed reactors. (a) Adiabatic fixed-bed reactor; (b) multitubular fixed-bed reactor.

Fig. 1

component. In many cases, the achievable conversion in the reactor is limited for thermodynamic (equilibrium) and kinetic reasons (selectivity). It is then usual to separate the material discharged from the reactor into products and unreacted feed components, which are recycled to the feedstock (see Fig. 2). This recycling procedure involves costs for: • product separation • recycle compression

• repeated heating and cooling of the circulating material to the reaction temperature and back to the temperature of the separating device • loss of product resulting from the need to remove part of the recycle as a bleed stream to limit the accumulation of inert substances or byproducts in the recycle loop. In order to minimize these costs it is therefore necessary to maximize the conversion in the reactor and to avoid inert substances in the reaction mixture as far as possible. With irreversible reactions (e.g., partial oxidations) the trend is therefore towards a highly concentrated, approximately stoichiometric feed composition, which may occasionally be in the explosive range. The resulting problems are addressed in Sections 10.1.3.3.2 and 10.1.5. A reaction process as depicted in Fig. 2 follows the classical Unit Operation Concept of Chemical Engineering, according to which a clear division into educts preparation (mixing and preheating), chemical reaction, product separation, and cooling shall be achieved in different units. A more recent development under the heading of ‘‘integrated’’ or ‘‘multifunctional’’ reactor concepts contrasts the unit operation concept. Its aim is to provide optimal reaction conditions in the reaction unit by incorporating optimal heat and reaction component addition or extraction at the reaction site. This trend and its main concepts are reviewed in Section 10.1.4. Feed stream

d

b

Bleed stream

a

e

c

Product

Reaction cycle for synthesis reactions with incomplete conversion. (a) Fixed-bed reactor; (b) feed preheater; (c) exit cooler; (d) recirculation compressor; (e) separation device.

Fig. 2

10.1.2 Catalyst Shapes for Fixed-Bed Reactors

Since catalytic fixed-bed reactors are often operated with flammable or potentially explosive gas mixtures at high temperatures and elevated pressures, issues of operational safety are of major concern. These are considered in Section 10.1.5 with respect to parametric sensitivity, runaway, and general safety. Throughout this chapter, model simulation results will be used to explain and discuss specific points, but the underlying models and their numerical simulation will not be addressed. Instead, the reader is referred to Chapters 6.1, 6.3, and 6.6 of this Handbook, and for further reference to standard text books and monographs [1–4]. 10.1.2

Catalyst Shapes for Fixed-Bed Reactors

The essential part of a fixed-bed reactor is the catalyst, where the reaction takes place. Seen from the flowing gas, the following steps of the overall reaction can be distinguished (see also Chapters 6.1 to 6.3): • diffusion of the reactants from the flowing gas through the outer gas–particle boundary layer, macropores, and micropores • chemisorption on active sites • surface reactions • desorption of the products • back-diffusion of the products into the flowing gas. As most reactions take place with a considerable heat of reaction, a corresponding heat transport is superimposed on the mass transport. The control of the microkinetics, consisting of micropore diffusion, chemisorption, surface reaction, and desorption, is the specific task of the catalyst developer. Once the catalyst is specified together with its microkinetic properties, reaction conditions (feedstock concentrations, pressure, temperature, and residence time) can be found that lead to optimum yields. The reaction engineer must determine these conditions and ensure that they are maintained in an industrial reactor. This is not a simple straightforward procedure, but requires multiple iteration loops in order to ensure that the optimal catalyst is available for the optimal reactor in an optimally designed process chain of feed preparation, reaction and product separation. In the case of selectivity-sensitive multistep reactions, any deviation from the optimum reaction conditions inevitably leads to a decrease in yield. Such deviations may be the result of a non-uniform residence time distribution due to flow dispersion and flow bypass phenomena in the fixed bed, as well as of deviations from uniform reaction conditions in the catalyst as a

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result of mass-transport resistance in the particles and the outer boundary layer [5]. The influence of mass-transport resistance in the particles can only be excluded if the reaction rate is substantially lower than the mass transfer velocity. This leads to the need for good external mass transfer (i.e., a sufficiently rapid flow rate in the packed bed), as well as for short diffusion paths and large transport pores in catalyst particles. In the case of exothermic or endothermic reactions, the local reaction rate must be controlled by the packedbed temperature. Thus, temperature control plays a predominant role in selective reaction control in general, and in particular in the case of exothermic multistep reactions. Under non-adiabatic conditions, catalysts must therefore be assembled in the fixed bed in such a way that good heat transport to the heat-transfer medium is ensured (see Sections 10.1.2.4 and 10.1.3.3). A further requirement placed on the catalyst is a low pressure loss. This applies particularly if the conversion in a single pass through the reactor is low, so that a large amount of gas has to be recirculated. It is of prime importance for off-gas purification, in which large off-gas streams must be treated with minimal additional cost (Section 10.1.2.3). Finally, the catalyst should be available in a sufficiently high concentration in order to keep the construction volume of the reactor low. The decisive parameters here are the specific external catalyst surface ap (area of catalyst surface per volume of reactor) for reactions controlled by external mass transfer, and the catalyst fraction (1 − ε), where ε (volume of free gas space per total reactor volume) is the void fraction of the fixed bed. The above requirements are to some extent contradictory, and this has led to the proposition of a large number of different catalyst shapes and arrangements. However, only a few of these have proved really effective in practical operation. Suitable catalyst forms and arrangements include random packings of spheres, solid cylinders, and hollow cylinders, as well as uniformly structured catalyst packings in the form of monoliths with parallel channels, parallel stacked plates, and crossed, corrugated-plate packets (Fig. 3). Random Packings Industrial fixed-bed reactors are generally operated with a mass flux of Gz ≥ 1 kg gas per square meter of reactor cross-section per second. This loading produces a sufficiently strong turbulence in random packings. As a result, the external gas-catalyst mass transport resistance is small compared to the transport resistance in the catalyst pores. 10.1.2.1

References see page 2105

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10.1 Catalytic Fixed-Bed Reactors

(a)

(b)

j

2a (c)

Common shapes of monolith catalysts. (a) Square-channel monolith; (b) parallel-plate monolith; (c) corrugated-plate packing.

However, the thermal conductivity of the catalyst matrix is usually larger than that of the gas. Hence, the external gas-catalyst heat transport resistance exceeds the thermal conduction resistance in the catalyst particles. The temperature and concentration profiles established in a spherical catalyst are illustrated for a partial oxidation reaction in Fig. 4. The profiles can be calculated from the model equations given in Chapter 6.3. An essential precondition for the establishment of the above-mentioned profiles is that the catalyst particle is uniformly exposed over its entire surface to a flow with uniform temperature and concentration. This is, of course, never the case in random packings. Figure 5 provides an image of the local mass transfer distribution around cylindrical or ring-shaped particles in a random packing. A test reaction producing dark deposits has been used. The intensity of the dark coloration is directly proportional to the local reaction rate of the surface reaction. As the test reaction is masstransfer controlled, the coloration is also proportional to the local mass transfer, and, with the analogy of mass and heat transfer, also to the local heat transfer [6]. Figure 5 shows that the local conditions in random packings are much more complex than assumed in current single-pellet models. However, as a fixed-bed

Reaction: A + O2 B B + x O2 n CO2 + m H2O Catalyst pellet

T CA

T, Ci

Fig. 3

CB C CO2 Boundary layer Pellet radius

Temperature and concentration profiles for a partial oxidation reaction in a spherical catalyst pellet.

Fig. 4

10.1.2 Catalyst Shapes for Fixed-Bed Reactors

cm

1

2

wo = 0,4 m /s

3

4

5

6

7

8

9

Rep = 1000

cm

1

2

3

wo = 0,4 m /s

4

5

6

7

8

2079

9

Rep = 1000

Visualization of local mass-transfer distribution at the surface of individual cylindrical or ring-shaped catalyst pellets in a fixed-bed packing [10]. The darkness is proportional to the local mass transfer.

Fig. 5

reactor contains a large number of particles at the same cross-section, Fig. 4 can be considered a mean value approximation of a process that varies greatly as regards detail. For the same reason, it is inappropriate to compare model predictions of fixed-bed reactor behavior with a few local temperature or concentration measurements. Indeed, a certain degree of averaging is always necessary in the measurement procedure. With the advancement of modeling and numerical simulation it now becomes possible to model the above details of transport and reaction in a section of a random packing [7]. This will help to derive even better mean value approximations for mass and heat transport parameters or global reaction rates, to be used in the more coarse reactor design and operation models. The literature contains a number of correlation equations for the mean gas-catalyst mass transfer and heat transfer as a function of gas properties, catalyst geometry, and flow conditions [8, 9]. In practice, gas-catalyst mass transfer resistance is usually not dominating for catalyst packings and can be neglected. Nonetheless, the safest simulation results can be expected if a quasi-homogeneous model is used, the (macro-) kinetic parameters of which are adjusted to single-tube experiments with the final catalyst shape, preparation and tube dimension.

Monoliths Monolith catalysts are used primarily in environmental catalysis where large gas streams must be processed with a low pressure drop (see Chapters 11.2, 11.3, and 11.6). In contrast to random packings, the external heat and mass transfer in monolith catalysts is much more uniform; in fact, unlike random packings, it may become a limiting factor at high reaction rates. This applies in particular to channel-type monoliths with narrow parallel channels, where the flow is generally laminar under industrial operating conditions. Examples include monolith catalysts with a square-channel cross-section for automobile exhaust purification, and also for the removal of nitrogen oxides from flue gases. Figure 6 shows the results of a visualization of local mass transfer by using the same technique as in Fig. 5 [10]. The marked decrease in coloration in the flow direction is mainly the result of increasing educt consumption at the reactive wall; that is, the build up of the so-called ‘‘laminar boundary layer’’. The depletion is particularly pronounced in the corners of the monolith (Fig. 6, lower part), as two reaction surfaces meet here. The more acute their enclosed angle, the greater is the depletion in the corner region and the smaller the contribution of the wall surface to further heat transfer and mass transfer. The efficiency of channel 10.1.2.2

References see page 2105

2080

10.1 Catalytic Fixed-Bed Reactors

DG 1000 / 3 / 0

QK 1000 / 3

Visualization of mass transfer in monolith structures (flow from left to right). Top: Corrugated packing (as in Fig. 3c); bottom: square-duct monolith [10].

Fig. 6

monoliths of equal cross-sectional area but different shape therefore decreases in the sequence: circle, hexagon, rectangle, triangle. This is illustrated in Table 1, which lists the asymptotic dimensionless mass-transfer and heat-transfer coefficients for tubes of the above crosssections [9]. For more accurate calculations, the radially and axially variable velocity, concentration, and the actual reaction rate profiles must also be taken into account [11]. In addition to monoliths with straight parallel channels, a large number of regularly structured packings have been proposed in the past. Among these, corrugated plate packings (see Fig. 3c) have obtained some industrial importance; such a design is also frequently used in parallel plate heat exchangers, as column packings, and static mixers (e.g., Sulzer type SMV). Asymptotic dimensionless laminar heat-transfer coefficients Nu = αW .dh .λ−1 G (constant wall conditions) and Fanning friction factor f for pressure drop calculation (p = 2f ηzL .vG .dh−1 ) for ducts of different cross-section

Tab. 1

Geometry a

Nu

f

dh

2.47

13.33

√ 2a/ 3

2.98

14.23

a

3.34

15.05

√ 2 3a

3.66

16.00

a

7.54

24.00

2a

a a a

a a

In contrast to the monolith with straight parallel channels, the flow conditions in corrugated-plate packings may be turbulent under typical industrial velocities. The uniformity of the mass-transfer distribution depends on geometric parameters (wave form, amplitude, wavelength, angle of incidence) [12]. The transfer coefficients are considerably higher than those of laminarly traversed channel monoliths (see the difference in coloration between the corrugated packing and the square-duct monolith in Fig. 6 for similar flow conditions), but the pressure loss is also increased. These structures offer considerable advantages for convective heat transport transverse to the flow direction and for transverse mixing; these are discussed in greater detail in the following sections. Gas Flow and Pressure Drop in Fixed Beds Conventional industrial catalysts differ considerably as regards pressure loss. For example, for equal mean dimensions and the same void fraction ε, random packings generally have a considerably higher pressure loss than monolith structures, and among these, corrugated structures have a higher pressure loss than monoliths with straight parallel channels. External catalyst mass and heat transfer are strongly correlated with the pressure loss which means, that the benefits of a low pressure drop are (partly) compensated by increased heat and mass-transfer resistance. The catalyst geometry for a given process is therefore chosen to combine the required mass and heat transport with the lowest pressure loss. The Hagen–Poiseuille equation [Eq. (1)] is used to calculate the pressure loss p in laminarly traversed 10.1.2.3

10.1.2 Catalyst Shapes for Fixed-Bed Reactors

Thus, the pressure drop depends very strongly on the pellet diameter dp and on the void fraction ε of the packing. From the standard forms for packedbed catalysts, hollow cylinders of thin wall thickness (ε ≈ 0.6–0.8) are therefore preferred over spheres (ε ≈ 0.37–0.4) and solid cylinders (ε ≈ 0.35). The pressure loss coefficient ζ in Eq. (2) can be determined for typical packing forms, for example, according to Ref. [13]. A comparison of different catalyst packings with respect to pressure drop over bed length, p/L, is contained in Fig. 7. The flow into a fixed-bed reactor is generally achieved by means of a feed pipe and a distribution hood (see Fig. 1). It is essential for a good operation that the gas flow is distributed evenly over the whole fixed bed or tube bundle. A uniform flow distribution can be easily obtained if the pressure drop in the fixed bed is substantially larger than in the entrance and exit hoods. This is not the case

monolith channels, where η is the dynamic viscosity of the fluid, zL the channel length, vG the channel velocity and dh the channel hydraulic diameter: 32 ηzL vG dh2

(1)

The pressure loss of randomly packed tubes can be described either by means of a pressure loss coefficient ζ and the pressure drop equation p = ζ

ρG 2 v 2 G

(2)

where ρG is the fluid density or by the Ergun equation: 2 p = f1 vG + f2 vG

(3)

where f1 = 150η

(1 − ε)2 (1 − ε) ; f2 = 1.75ρG 2 2 ε dp εdp 3000

(4)

References see page 2105

0.9

ap

2500

e

1500 1000

0.3

500 0

a

b

c

3.0

300

aw

2.5

250

lr

1.0 0.5 0

aw, W m−2 K−1

1.5

d

e

f

g

h

i

0

50 40

∆p/L

200

2.0

lr / W· m−1·K−1

0.6

e

ap / m2 ·m–3

2000

30 150 20 100 10

50 0

∆p/L / mbar .m−1

p =

2081

0 a

b

c

d

e

f

g

h

i

Specific outer surface area ap , void fraction ε, wall heat-transfer coefficient αw , effective radial heat conductivity λr , and pressure drop p, for different packings in a tube of 50 mm internal diameter (i.d.) and an air mass flux Gz = 1 kg m−2 s−1 at ambient conditions [23]. (a) Glass spheres, dp = 5 mm; (b) glass spheres, dp = 10 mm; (c) Raschig rings, ceramic, do = 10 mm, di = 6 mm, l = 11 mm; (d) Raschig rings, stainless steel, do = 10 mm, di = 9 mm, l = 11 mm; (e) hollow ceramic cylinders of irregular length, do = 9 mm, di = 4.5 mm, l = 15 mm; (f) full ceramic cylinders of irregular length, do = 5 mm, l = 11 mm; (g) crossed corrugated metal-plate packing (Sulzer Katapak ), wide channels; (h) crossed corrugated metal-plate packing, narrow channels; (i) automotive exhaust monolith, Cordierite, square channels, d = 0.9 mm. Fig. 7

2082

10.1 Catalytic Fixed-Bed Reactors

with monolith reactors with laminarly traversed parallel channels. They need special care in the design of the hoods or must be equipped with static mixers to avoid non-uniform flow patterns [14]. A related problem occurs in multitubular fixed-bed reactors with random packings if different tubes have different pressure drops due to non-uniform packing or different batches of catalysts. The resulting differences in throughput may cause the danger of reaction run-away in tubes of higher pressure drop and low throughput [15]. Great care must therefore be taken to adjust the pressure drop of all tubes of a multitubular reactor for highly exothermic reactions to the same level. A particular case of flow non-uniformity is inherent to all randomly packed tubes. It is due to the fact that the catalyst pellets make only point contact with the wall, whereas they overlap and cross over one another inside the packing, thereby reducing the free volume and hence the velocity. The conditions for a spherical packing are illustrated schematically in Fig. 8. The radially varying empty space distribution and velocity distribution has been taken into account in detailed reactor calculations, as well as in the determination of accurate heat transport parameters [16–18]. This will be addressed in the next section. Standard catalysts have a sufficiently dense structure, and all convective transport is directed around the pellets or separated through the walls of the monolith. In a sequence of publications [19, 20], macroporous catalyst pellets with high permeability of the reacting gas have also been proposed. Here, part of the diffusional resistance is overcome by convection through the pellets, which can increase conversion and selectivity.

Void fraction

0.5

0

Radius 5.0 4.0

Interstitial velocity

Heat Transfer in Catalyst-Filled Tubes With non-adiabatic reaction control, heat must be exchanged laterally through the fixed bed with heatexchange surfaces. At the usual mass flow rates of Gz ≥ 1 kg m−2 s−1 , this heat transport takes place mainly by convection – that is, the fixed bed must allow for a cross-flow transverse to the main flow direction. In catalyst pellet packings a transverse-flow component is established automatically as a result of the non-uniform arrangement and the twisted flow around the pellets. Hollow and full cylinders with a length-to-diameter ratio of 1 to 3 are particularly effective in this respect. Despite the fact that radial heat transport takes place mainly by convection, it is formally described by means of a so-called effective radial thermal conductivity λR , transverse to the flow direction. In addition, the heat transport at the boundary between the fixed bed and the heat-exchanging walls is also decisive for the heat exchange. The latter heat transport is generally

1.0

10.1.2.4

3.0 2.0 1.0 0

Radius Radial distribution of void fraction and axial (interstitial) flow velocity in a tube packed with spheres.

Fig. 8

described by a wall heat-transfer coefficient, αw . This lumps the complex interplay between convective flow at the tube wall and conduction transport by (point) contact between the fixed bed and the heat-exchange surface. Heat transport in packed tubes has been investigated and discussed in detail [8, 16–18]. In Ref. [21], it is shown that a plug-flow model with axial dispersion using adequate λR - and αW -values is usually able to capture the detailed behavior (with radially non-uniform flow) even of

10.1.3 Types of Fixed-Bed Reactor

strongly exothermic reactions with reasonable accuracy. Only close to run-away conditions does the more detailed model predict a somewhat earlier onset of run-away. However, considering the uncertainties involved in the determination of transport and kinetic parameters, these differences can be considered negligible. Monolith structures with straight parallel channels are not well suited for non-adiabatic reaction control due to the missing lateral cross-flow. An increase of the lateral heat transport via conduction in the monolith structure is beneficial [22], but its success is hampered by the insufficient heat contact between a monolith structure and the tube wall, due to unavoidable manufacturing tolerances.

2083

In addition, the flow behavior in tubes of circular crosssection is rather non-uniform over the circumference, why it is advantageous to arrange short packing sections in series, each section being displaced by 90◦ . The heat-transport parameters in Fig. 7 were determined for structures arranged in this way. A general problem in the use of monolith structures in reactor tubes is incomplete wall contact. As individual tubes of multitubular reactors always have a diameter tolerance of ca. 1 mm and interlocking of the structure with the tube wall must be avoided, the bypass of flow at the wall is even greater than with random packings. 10.1.3

Types of Fixed-Bed Reactor Comparison of Different Catalyst Shapes A comparison of different catalyst shapes is given in Fig. 7, where the two heat-transfer parameters, λr and αw , plus the external catalyst surface ap , the bed void fraction ε and the pressure drop p/L are given for a selection of different random and regular catalyst packings in a tube of 50 mm internal diameter and a mass flow velocity of Gz = 1 kg m−2 s−1 [23]. The dimensions of the irregular packing bodies were chosen so that their specific external surface area ap is ca. 500 m2 m−3 . Under these conditions, hollow, thinwalled cylinders have clear advantages over other packing forms, exhibiting the lowest pressure loss and the highest thermal conductivity. Only as regards wall heat transfer are they inferior to spheres or full cylinders. However, good wall heat transfer is apparently less decisive from a reaction engineering viewpoint than good radial thermal conductivity, as the former can be compensated by an appropriate temperature profile of the heat-transfer medium, whereas the radial thermal conductivity directly influences the uniformity of the reaction conditions over the tube cross-section. On the other hand, solid cylinders with a large length-to-diameter ratio have good heattransport values, but at the cost of a very high pressure loss. Monolith structures can have very high specific surfaces combined with a very low pressure loss. Monoliths with straight, parallel channels (as are used for automobile exhaust control) have only very poor radial heattransport properties. Crossed corrugated structures are considerably more favorable for isothermal reaction control. They have a very high radial thermal conductivity which is almost independent of the specific surface area; the latter can be varied over a wide range by the geometry of the corrugations. It must be remembered that with crossed corrugated structures, convective radial heat transport occurs only in one plane perpendicular to the main flow direction. 10.1.2.5

Adiabatic Reactors Adiabatic fixed-bed reactors constitute the oldest configuration of this reactor type. In the simplest case, they consist of a cylindrical jacket in which the catalyst is loosely packed on a screen support and is traversed in axial direction (Fig. 9a). To avoid catalyst abrasion by partial fluidization, random catalyst packings should always be traversed from top to bottom. If this is not possible, for example because of periodic flow reversal (see Section 10.1.4.1.1), special measures must be taken to stabilize the packing. If fixed beds composed of monolith catalyst sections are used, the flow direction can be chosen freely. It can be shown [14] that the requirement for a low pressure loss leads to a fixed bed of large diameter and low height (Fig. 9b). Such an arrangement (‘‘disk concept’’) is used particularly when very short residence times, followed by direct quenching of the reaction, are required. Examples include ammonia oxidation in nitric acid production and oxidative dehydrogenation on silver catalysts (e.g., synthesis of formaldehyde by dehydrogenation of methanol). In the first case, the ‘‘fixed bed’’ consists of several layers of platinum wire gauze, and in the second case of a porous silver layer several centimeters in height. The bed diameters can be up to several meters. Flat-bed reactors with a residence time in the order of milliseconds have recently achieved considerable attention for partial oxidations of hydrocarbons to synthesis gas, olefins, or oxygenates (this subject is treated in detail in Chapter 10.5). On account of the difficulties involved with obtaining uniform flow as well as for structural reasons, the disk concept is limited to relatively small catalyst volumes. The radial flow concept (Fig. 9c) is used when large amounts of catalyst are involved. The catalyst packing is accommodated in the space between two concentric screens or perforated plate cylinders, and is traversed in radial 10.1.3.1

References see page 2105

2084

10.1 Catalytic Fixed-Bed Reactors

Perforated wall

Unperforated wall

Axial radial gas flow zone Catalyst Quench

(a)

(b) Radial gas flow zone

Fig. 10

(c)

Main design concepts for adiabatic reactors. (a) Adiabatic packed-bed reactor; (b) disk reactor; (c) radial-flow reactor.

Fig. 9

direction, either from the inside to the outside or from the outside to the inside. This design is particularly suitable for large catalyst volumes as well as for operation at elevated pressure, since at moderate reactor diameters the catalyst volume can be varied over a wide range by altering the reactor length, without affecting the flow-through length of the packing and hence the pressure drop. A critical feature of packed radial-flow reactors is the shape of the upper bed closure. A simple horizontal covering is not practicable as a gap is formed due to the unavoidable settling of the packing through which unreacted gas can pass. The arrangement shown in Fig. 10 has proved effective as it produces mixed axial and radial flow through the bed in the upper bed closure. The required geometric shape must be determined by simulation of the local two-dimensional (axisymmetric) flow through the packing [24].

Upper bed closure in a radial-flow reactor [24].

Monolith catalysts with straight, parallel channels are particularly well suited for axial-flow adiabatic reactors as they allow for a low pressure drop and provide a high specific outer surface area. Since monolith catalysts are usually produced with a rectangular crosssection, the fixed bed is constructed by arranging these individual monoliths in rows in the form of a large ‘‘box’’. Conventional DENOX reactors for removing NOx from power station flue gases are therefore designed as rectangular chambers (Fig. 11). The catalyst is often arranged in the form of several layers in series, the spaces between the individual layers permitting crossmixing, so that the influence of non-uniform flow as well as any possible local blockage of the next layer can be compensated to some extent. Reference was been made in Section 10.1.2.3 to the importance of uniform flow into and through adiabatic fixedbed reactors. This is not easy to achieve, particularly with low pressure-loss monolith reactors, and requires a careful design of the inflow hood. On account of the low pressure loss, unfavorable flow conditions in the outflow hood may also affect the flow behavior through the catalyst bed. This is of particular importance for the design and operation of automotive exhaust catalysts, the inlet or outlet pipes of which often induce a swirl component in the flow, which leads to a non-uniform flow through the monolith [14].

10.1.3 Types of Fixed-Bed Reactor

2085

purification, in which small amounts of noxious components are converted. The chambers used to remove NOx from power station flue gases, with a catalyst volume of more than 1000 m3 , are the largest industrial adiabatic reactors, and the exhaust catalyst for internal combustion engines, with a catalyst volume of ca. 1 L, the smallest. Typical applications in the chemical industry include the methanation of traces of CO and CO2 in NH3 synthesis gas, as well as the hydrogenation of small amounts of unsaturated compounds in hydrocarbon streams. The latter case requires accurate monitoring and regulation when hydrogen is in excess, in order to prevent complete methanation due to an uncontrolled temperature run-away.

Structure Flue gas + NH3 in

Catalyst element

Multi-Stage Reactors In the majority of fixed-bed reactors for industrial synthesis reactions, direct or indirect supply or removal of heat in the catalyst bed is necessary to adapt the temperature profile over the flow path as far as possible to the requirements of an optimal reaction pathway. Here, a clear historical trend can be observed. This started with the adiabatic reactor (Fig. 12a) which, on account of the adiabatic temperature change, could only be operated to give a limited conversion. Higher conversions were achieved at the same mean temperature level when several adiabatic stages were introduced, with intermediate heating or cooling after each stage. The simplest form involves injecting hot or cold gas between the stages (Fig. 12b). For a constant tube diameter, the main disadvantages of this temperature control strategy are cross-sectional loading, which increases from stage to stage, and the mixing of 10.1.3.2

Catalyst layers Catalyst module Flue gas out

Fig. 11 Reaction chamber for removal of nitrogen oxides from power station flue gases [25].

Purely adiabatic fixed-bed reactors are used mainly for reactions with a small heat of reaction or for catalytic combustion (see Chapter 11.6). A typical example is gas

References see page 2105

Feed gas

Interstage gas feed

(a)

(b)

Interstage heat exchangers

(c)

Development of fixed-bed reactors. (a) Single-bed adiabatic packed-bed reactor; (b) adiabatic reactor with interstage gas feed (ICI concept); (c) multi-bed adiabatic fixed-bed reactor with interstage heat exchange.

Fig. 12

2086

10.1 Catalytic Fixed-Bed Reactors

hot and cold streams, which is energetically unfavorable. The composition is changed by injection, which can have a positive or negative effect on the desired reaction. The next step in development was the replacement of injection cooling by interstage heat exchangers, through which the required or released heat of reaction is supplied or removed (Fig. 12c). More stringent requirements led to the development of fixed-bed reactors where the heatexchange surface is integrated in the fixed bed. This type of reactor will be considered in the next section. Today, adiabatic multi-stage fixed-bed reactors with intermediate cooling or heating are used particularly when the reaction proceeds selectively to a single product but is limited by the equilibrium conditions. Intermediate cooling or heating is used to change the gas temperature towards higher equilibrium conversion. Typical examples include the synthesis of ammonia, sulfur trioxide, and methanol. In these exothermic reactions the equilibrium conversion to the target product decreases with increasing temperature (see Fig. 13a), while the reaction rate always increases with temperature. For a given conversion, X, a temperature can therefore be found at which the reaction rate to the target product becomes a maximum. This temperature must be below the equilibrium temperature, but at the same time not so low that the reaction becomes too slow for kinetic reasons. The points obtained in this way can be joined to form an optimal reaction rate curve (Fig. 13b). In the case of an adiabatic reaction the temperature increases linearly with the achieved conversion X according to the equation Tad =

(−hR ) c0 X ρG cpG

(5)

where hR is the reaction enthalpy, c0 the feed concentration and cpG the heat capacity of the reaction

mixture. Each adiabatic reaction pathway therefore lies on a straight line of gradient T /X (see Fig. 13a). A reasonable reaction pathway for a multi-stage adiabatic reaction can therefore be composed of the straight lines of an adiabatic bed followed by the vertical lines of indirect intermediate cooling (Fig. 13c). The kinetically optimum reaction pathway with the smallest required catalyst volume results when the trajectory follows, in a large number of small steps, the line of maximum reaction rate. In practice, the apparatus and equipment expenditure involved in a large number of stages must be weighed up against the savings in catalyst. Conventional multi-stage reactors for this class of reaction therefore have three to five stages. Similar considerations lead to the design of multi-stage adiabatic reactors for endothermic reactions, where the required heat of reaction after each adiabatic stage is again provided by interstage heat exchangers or (as in styrene synthesis) by the addition of a hot gas (e.g., steam). Figure 14 shows the layout of an ammonia synthesis reactor designed on this basis. For structural reasons, the heat exchanger between the inflow and outflow is incorporated in the lowest part of the pressure casing. The reaction gas then flows upward in the annular gap between the pressure casing and the fixed beds. There, it is further heated and at the same time protects the pressurebearing structural components against excessively high fixed-bed temperatures. The three adiabatic fixed beds are traversed from top to bottom, part of the heat of reaction being utilized to generate steam in the two intermediate heat exchangers. To start up the cold reactor, hot gas must be added to the uppermost bed, for example through an external start-up preheater. Although industrial adiabatic multi-stage reactors often differ in many details from Fig. 14, they are of a comparable basic design. For example, radially instead of axially traversed beds can achieve a smaller pressure

2

Locus of optimal reaction rate

3

Feed temperature

Be d

T

Adiabatic reaction path

T

T

1

Equlibrium curve

∆X1 (a)

Conversion, X

(b)

X

(c)

∆X2

∆X3

X

Fig. 13 Equilibrium limitation of exothermic reactions. (a) Equilibrium conversion as a function of temperature; (b) optimum reaction rate curve; (c) improvement of conversion by using multiple adiabatic beds with interstage cooling.

10.1.3 Types of Fixed-Bed Reactor

2087

400 °C, 300 bar, 2% NH3 Steam

Stage 1 Stage 1

Water Stage 2

Steam

Stage 2

Water Stage 3 Stage 3 Fresh gas N2 + 3H2 (+CH4) 20% NH3 (N2, H2, CH4)

Circulating gas Off–gas Exit

Condensation coolers Fig. 14

Feed

NH3, liquid

Schematic of a multistage reactor for ammonia synthesis [23].

loss with a more favorable structural arrangement; heat exchange with the cold feedstock can be effected by heatexchange surfaces integrated in the first catalyst bed; alternatively, a cold gas quench can be used to achieve additional temperature regulation. An interesting novel extension for endothermic reactions concerns oxidative dehydrogenation reactions, where the required temperature increase after each adiabatic stage is provided by an intermediate catalytic combustion of part of the hydrogen produced [26]. State of the art adiabatic multi-stage reactors may thus become quite complex, and the question must be answered as to whether a multitubular design (see next section) might be the more favorable alternative. If the heat exchange is integrated in the fixed bed, the required heat-exchange surface area is smaller than in the case of free gas flow due to the enhanced bed-towall heat transfer of the catalyst packing. Furthermore, it does not induce additional pressure loss, and the optimum reaction rate curve (see Fig. 13b) can be better approximated by controlling the cooling temperature profile (see Section 10.1.3.3.4) than by a stepwise temperature reduction. On the other hand, a multi-stage arrangement may be considered for structural, kinetic or economic reasons in the following cases: 1. If, in the case of large single-train plants, subdivision into several individual units is necessary for reasons of transport or construction.

2. If a catalyst must be replaced in individual stages at different times on account of different catalyst deactivation. 3. If a step-wise addition of reactant has kinetic advantages compared to the total addition to the feed (here, a suitably designed intermediate heat exchanger ensures a uniform distribution and mixing within the reaction gas stream). 4. If the intermediate stages are used to extract a limiting product in the case of equilibrium-limited reactions; an example is the intermediate absorption of SO3 before the last stage of the SO3 synthesis. 5. With reaction temperatures above 300 ◦ C, intermediate cooling can still be performed directly with boiling water, whereas in a fixed bed a high-temperature heattransfer medium must be used as coolant (see next section). Cooled or Heated Multitubular Fixed-Bed Reactors The development of reactors in which the heat-exchange surfaces are integrated in the fixed bed occurred in parallel with the development of multi-stage adiabatic reactors with intermediate heating or cooling. The main aim is to supply or remove the heat of reaction as close as possible to the reaction site. The multitubular fixed-bed reactor (see Fig. 1b) constitutes the oldest and still predominant representative of this class. The catalyst is packed into the individual tubes 10.1.3.3

References see page 2105

2088

10.1 Catalytic Fixed-Bed Reactors

Steam

T,TC Coolant loop I

Z

Feed water

T TCI

Coolant loop II

(a) Fig. 15

(b)

Circulation turbine

Steam generator

TCII

(c)

Heat-transfer medium flow control in tube-bundle fixed-bed reactors. (a) Cross-flow; (b) parallel flow; (c) multiple cooling sections.

of the tube bundle. The heat-transfer medium is circulated around the tube bundle and through an external heat exchanger, in which the heat of reaction is supplied or removed (Fig. 15). Whereas with endothermic reactions, circulating gas is also in use as heat-transfer medium, for strongly exothermic reactions exclusively liquid or boiling heat-transfer media are used. Only in this way can the catalyst temperature be held in the narrow temperature range necessary for selective reaction control (e.g., in the case of partial oxidations). Initially, the integration of heat exchange in the fixed bed was utilized to ensure a reaction being conducted as isothermal as possible. Therefore, reactors of the type shown in Fig. 15a and b are also commonly referred to as ‘‘isothermal reactors’’. They are comprised of reaction tubes of 20 to 80 mm internal diameter with 0.5 to 15 m length, and a carefully designed flow control of the liquid heat-transfer medium, with largely constant heat-transfer conditions and maximum temperature changes of the heat-transfer medium throughout the tube bundle of a few K. More recent concepts are aimed at establishing a freely selectable optimum temperature profile over the tube length. This requires complex heat-transfer medium control with several sections and temperature levels (Fig. 15c). Strongly exothermic synthesis reactions such as partial oxidations can only be carried out in fixed-bed reactors if the catalyst temperature is controlled in a narrow optimal window. This can be achieved if the following requirements are fulfilled: • Large heat-transfer areas must be available per catalyst volume (but a small tube diameter requires small pellet diameters which increase pressure drop).

• The temperature of the heat-transfer medium must be close to the desired catalyst temperature to ensure an effective catalyst temperature control. • A sufficiently large mass flow velocity of the reaction gases is generally necessary to ensure good heat transport from the packing to the tube walls. Although called ‘‘isothermal’’, the temperature profiles of a liquid-cooled or heated packed-bed reactor with a strongly exothermic or endothermic reaction are far from being isothermal. Figure 16 shows simulated temperature and conversion profiles of the phthalic anhydride synthesis in a wall-cooled tube of 25 mm diameter at typical synthesis conditions (as specified in the figure legend). Both the temperature profiles in the tube center and close to the tube wall are given for 387 ◦ C wall temperature, considering plug flow (dotted lines) or a radially varying void fraction and flow velocity profile as shown in Fig. 8 (solid lines). The comparison shows that the plug-flow assumption is well justified as the conversion and yield profiles are almost identical and only the temperature in the tube center is a little higher in the more detailed model [21]. Comparably changing temperature profiles have been obtained for strongly endothermic reactions such as methane steam reforming, where the fixed-bed temperature drops rapidly at the entrance of the active catalyst [27, 28]. Truly isothermal conditions could only be achieved if the catalyst were to be placed directly at the reactor tube wall of constant temperature. Such a ‘‘wall-coated’’ catalytic reactor is a concept which has been discussed for some time, but hardly transferred into industrial application. The main reasons are problems with in-situ uniform wall coating and catalyst replacement in large-scale reactors. Recently, however, ‘‘micro-reactors’’ using the design of parallel-plate heat exchangers with a channel width

10.1.3 Types of Fixed-Bed Reactor

420

1.0 X 0.8

T / °C

410 Tube center

Φ

0.4

390

0.2

Tube wall 380

X

Φ 0.6

400

0

1

2089

2

3

4

0.0

0

1

2

3

4

z/ m

z/m

Fig. 16 Temperature (T), conversion (X) and yield profile () for the partial oxidation of o-xylene to phthalic anhydride in a tube of 25 mm internal diameter (i.d.) at 387 ◦ C coolant temperature, p = 105 Pa (1 bar), feed weight fraction o-xylene 3.4%, mass flux 1 kg m−2 s−1 . After Ref. [21]. Solid lines: calculated with a 2-D model assuming the actual flow and void fraction profile, comparable to Fig. 8. Broken lines: calculated with a two-phase, 1-D plug- flow model.

in the order of 1 mm or less have become available (see Chapter 10.8). In these reactors, the catalyst is generally coated at the walls, which allows for an excellent control of its temperature via the heat transfer medium and makes the conduction of highly exothermic reactions with undiluted feed possible. However, similar to tubular reactors a uniform catalyst deposition on the channel walls and the possibility to replace the catalyst without dismantling the reactor are among the unsolved problems. 10.1.3.3.1 Heat-Transfer Media The above-mentioned requirement that the catalyst temperature should be close to the coolant temperature to prevent thermal run-away of exothermic reactions presupposes an assortment of heattransfer media, which cover the whole temperature range of interest for gas-phase reactions in fixed-bed reactors. It is reasonable to distinguish between gaseous, liquid, and vaporizing heat-transfer media. Vaporizing heat-transfer media are used exclusively to remove heat from exothermic reactions. Whereas formerly petroleum fractions such as kerosene (e.g., in ethylene oxide synthesis) have been used more widely, they are now largely replaced by pressurized boiling water, as water is not flammable, has a higher heat of vaporization, and the steam produced can be used directly in other parts of the plant. Depending on the saturated vapor pressure, the temperature range from 100 to 310 ◦ C (107 Pa; 100 bar) can be covered with boiling water. In this range it is the preferred heat-transfer medium if an isothermal cooling temperature is required. A standard design of a boiling-water-cooled multitubular fixed-bed reactor is shown in Fig. 17, and the details will be discussed in the next section. Locally variable cooling temperature profiles can be established most easily with liquid heat-transfer media that do not vaporize in the intended operating range. In order to avoid cavitation, pressurized water should

Feed gas

Saturated steam

Steam drum Feed water

Fig. 17

Multitubular reactor with boiling-water cooling [23].

be used only up to ca. 220 ◦ C; heat-transfer oils cover the temperature range up to 300 ◦ C, but above this temperature salt melts are now used exclusively [23]. Salt melts cover a larger temperature range and have the particular advantage over oils that they are incombustible. The potential danger of a relatively large amount of hot salt melt obviously exists, but is reliably dealt with by experienced reactor construction companies. Special nitrate melts (HITEC ) can be used in the temperature range 200 to 500 ◦ C. Gradual decomposition begins above this temperature, and can accelerate violently above References see page 2105

2090

10.1 Catalytic Fixed-Bed Reactors

600 ◦ C. The access of organic components to the melt (nitrate decomposition) and of water (steam explosion) must be excluded [29]. For the temperature range of 400 to 800 ◦ C salt melts based on carbonates have been developed. In this case, it is not so much the thermal stability of the molten salt but rather the corrosion of the reactor materials that presents problems. Gases are the only heat-transfer media usable over the entire temperature range, but because of their low density they have an unfavorable heat-transport behavior. Gas cooling of strongly exothermic reactions is not possible for industrial reactors, and even not advisable for small-diameter reactors used for kinetic experiments. In addition, to achieve the same heat transfer at the same temperature difference, the necessary pumping energy would exceed the respective value of liquids by a factor of 100 to 200 [23, 30]. Gases are therefore used exclusively as flue gases to supply heat at high temperatures (as discussed in Section 10.1.3.3.3). 10.1.3.3.2 Design Concepts for Exothermic Reactions As with adiabatic reactors, a principle task of reactor development is to produce uniform reaction conditions over the whole reactor, and to maintain such conditions during the entire operating time. This involves the conditions in the catalyst packing (residence time, catalyst concentration and activity, heat transport) and in the heat-transfer medium circuit (throughput, temperature, heat transfer). When liquid or vaporizing heat-transfer media are used, the heat- transfer coefficients at the heat transfer medium side are usually one order of magnitude greater than those on the catalyst side, which facilitates the task as regards the heat-transfer medium. On the other hand, with strongly exothermic, selectivity-sensitive reactions (especially partial oxidations), a temperature constancy of the heat-transfer medium of ca. 1 K is often required. This leads to a high energy requirement for circulation and necessitates careful design and control of the heat-transfer medium circuit. Tubular reactors of the types shown in Fig. 15 have been in use for the longest time, and hence have been furthest developed. In the case of a vaporizing heat-transfer medium an arrangement as shown in Fig. 17 is generally chosen, in which the liquid heat-transfer medium surrounds the stationary tube bundle. The rising vapor bubbles escape through the ascending pipe into a vapor drum, where they are separated from the liquid. A circulation flow is established due to the difference in density in the downpipe (pure liquid) and in the reactor jacket, which means that circulating pumps are generally not required. The heat-transfer medium temperature is regulated via the vapor pressure through the valve in the steam drum.

Apart from the types of multitubular reactor shown in Figs. 15 and 17, other multitubular reactors are sometimes used in which the catalyst bed is arranged around the tubes and the heat-transfer medium flows through the tubes (Fig. 18b). An interesting development has been introduced by Linde [31] (Fig. 18a); here, the tube bundle is composed of counter-wound spirals in which upwardly flowing water is evaporated. The tubes run into a vertical vapor drum located at the reactor head. The tube bundle is connected to a central down-pipe at the bottom such that, as in the arrangement in Fig. 17, a natural circulation of the evaporating water is established. Advantages in construction and in the heat transfer from the reaction gas to the tubes of the bundle are claimed for this design [32]. Circulation systems with parallel and crossed cocurrent or countercurrent flow of the heat-transfer medium (see Fig. 15) are commonly employed as liquid heat-transfer media. The main part of the heat-transfer medium is generally circulated with a high-capacity pump in order to achieve uniform heat-exchange conditions. A partial stream is passed through a heat exchanger for supplying or removing the heat of reaction. The desired heat-transfer medium temperature is attained by regulating this partial stream. With exothermic reactions, the heat exchanger is normally a steam generator which produces saturated steam at a pressure corresponding to a boiling point of 30–80 K below the minimum cooling temperature. Apparatus construction companies specializing in these reactors have developed a detailed and comprehensive know-how as regards flow control of the heat-transfer medium [33, 34]. This concerns the uniform supply and removal of the heat-transfer medium, which generally takes place via external annular channels, as well as the flow control within the reactor. Some publications illustrate the major differences in the behavior of different tube sections that can arise due to an inadequate design and layout of the heat-transfer medium circuit [35–37]. Design Concepts for Endothermic Reactions Endothermic reactions are not in the danger of thermal run-away, and therefore a close coupling of the reaction temperature to the temperature of a heat-transfer medium is not essential. Here, the limiting points are usually the transfer of the necessary amount of heat into the catalyst bed, and control of the high-temperature residence time in order to limit side reactions. Moderately endothermic reactions such as dehydrogenations are often carried out in multi-stage adiabatic reactors with interstage heating. Typical examples are reactors for styrene synthesis through ethylbenzene dehydrogenation, where superheated steam is added to the reaction mixture between the (radial flow) adiabatic stages. 10.1.3.3.3

10.1.3 Types of Fixed-Bed Reactor

2091

Steam

HP-steam Recycle water

2000 °C

Gas in

Air Syngas Gas

850 °C

Water

Fig. 19 General concept of a top-fired reformer unit for syngas-production. In spite of a 2000 ◦ C flame temperature, the catalyst temperature at the exit of the reformer tubes reaches only 850 ◦ C.

Gas out (a)

Feed water

(b)

Fig. 18 (a) Design concept of the Linde isothermal reactor for methanol synthesis; (b) section through the tube bundle surrounded by catalyst pellets (from Ref. [32]).

The alternative for more strongly endothermic reactions is the multitubular reactor which is either heated directly or indirectly by burners. A typical example is methane steam reforming, where the reformer tubes of about 50 mm internal diameter and 10 m length are arranged along the walls of large boxes, fired by burners at the ceiling, the bottom or/and the sidewalls (Fig. 19). At the required catalyst temperatures, the reaction is rapid and essentially heat-transfer-controlled. This means that the whole tube length is needed to transfer the heat from the hot burner gases into the packed catalyst bed. In the general design shown in Fig. 19, only about 50% of the heat of combustion is transferred into the endothermic reforming reaction. The remainder leaves the reformer box as hot flue gas, requiring a complex heat-exchanger network for heat recovery. Such a design is therefore only economical in combination with a larger chemical complex where the excess heat can be utilized. For decentralized or remote locations, novel concepts with improved heat integration have been developed. An example is the Haldor Topsøe heat-integrated reformer shown in Fig. 20 [38]. As shown in the insert, the reformer tubes consist of a tube-in-tube arrangement with the catalyst in the annulus in good heat contact to the hot flue gases, and the inner tube empty to prevent back-reaction at decreasing temperatures and recover heat from the product gas. A similar tube-in-tube arrangement is used in the ICI gas-heated reformer (GHR) concept [39], which provides an even better heat integration (Fig. 21). Here, References see page 2105

2092

10.1 Catalytic Fixed-Bed Reactors

Air / O2

Syngas CH4, steam

Flue gas

Off

GHR

Syngas

ATR

CH4

Feed

Steam Syngas

800–950 °C Fig. 21 Gas-heated reformer (GHR) of ICI [39], thermally coupled with an autothermal reformer (ATR).

Fig. 20 Heat-integrated reformer concept for methane steam reforming of Haldor Topsøe. The insert shows the tube-in-tube arrangement with catalyst in the outer section and countercurrent heat exchange of the entering and leaving reformer gases [38].

the hot gas leaving the secondary autothermal reformer (ATR) is used to heat the primary steam reformer tubes. The limiting factor for the above concepts is the temperature stability of the reformer tubes used. This limits the maximum reforming temperature to about 850 ◦ C in the catalyst tubes, leading to about 90% equilibrium conversion for methane steam reforming

at 2 MPa (20 bar). Higher temperatures would require ceramic tube materials, but to date the only large-scale synthesis using a ceramic tube multitubular reactor seems to be the Degussa BMA process for HCN synthesis from methane and ammonia [40]. A reactor sketch with the reactor temperature profile and the two main reactions is provided in Fig. 22. This reactor also seems to be the only large-scale multitubular reactor where the catalyst is coated at the inner tube wall. To prevent ammonia decomposition to nitrogen, a water quench cools the product temperature of about 1250 ◦ C down to about 200 ◦ C. The design and the safe operation of such a reactor,

HCN, H2, CH4, NH3 Sintered corundum pipe

Cooling water Cooling water

x /m

2

Fuel 1300 CH4, NH3 NH3 + CH4 2NH3

T / °C

0 100

CH4 + NH3

1250 °C

2 cm

HCN + 3 H2

15 µm Catalyst layer 60% Pt a = 100 m2 / m3 32% Al 8% Ru h = 170 (450) W / m2 − K

HCN + 3H2 ; ∆H = 252 kJ mol−1 N2 + 3H2 ; ∆H = 105 kJ mol−1

Sketch of the Degussa BMA-reactor with approximate temperature profiles for HCN synthesis in a multitubular arrangement of sintered corundum tubes (2 cm i.d., 2.5 m length) with Pt/Ru catalyst coated on the inner tube walls [40].

Fig. 22

10.1.3 Types of Fixed-Bed Reactor

in addition to the manufacture of the ceramic tubes used, cannot yet be considered a standard technology due to the risk of tube failure, but it points at the direction in which short-contact time, high-temperature synthesis in tubular reactors is seeming to develop.

420

2093

Reactor temperature Cooling temperature Reference temperature

380

Vs = −0.14 m ⋅s−1

340

Reaction Control by Controlling the HeatThe conversion and Transfer Medium Temperature selectivity of the reaction can be decisively influenced by the design and the operation of the heat-transfer circuit. The most obvious – though technically most complex – solution is to arrange different heat-transfer circuits in order to achieve a stepwise approximation of an optimum temperature profile. The purposeful utilization of the temperature change of the heat-transfer medium flowing through the reactor is technically simpler, and will be discussed here in connection with cocurrent or countercurrent cooling of a fixed-bed reactor with an exothermic reaction. Figure 23 illustrates temperature profiles for three different ways of controlling the cooling stream in a partial oxidation reaction. If the coolant is circulated so rapidly that its temperature in the reactor hardly changes, then its flow direction is irrelevant and a temperature profile with a pronounced temperature maximum may be established, similar to the example discussed in Fig. 16. This situation is shown in Fig. 23a. If the coolant is circulated in cocurrent and its velocity is chosen so that it becomes noticeably hotter over its path, an almost isothermal temperature behavior can be achieved (Fig. 23b). This is because the reactive gas at the inlet is in contact with the coldest coolant, and the cooling temperature rises in step with the consumption of the reactants, such that the reaction rate remains virtually constant over a fairly long section [41–43]. Until now, the stabilizing effect of cocurrent cooling has barely been exploited in industrial reactors. This may be due to the fear that, at the required low flow velocity (in the example of Fig. 20b, vS = 0.01 m s−1 ), heat transfer will be inadequate and natural convection will occur in the cooling jacket. However, vS describes the mean coolant velocity parallel to the tube axis. With a cross-cocurrent flow of the coolant, as in Fig. 15a, the actual flow velocity may in fact be substantially larger, depending on the number of deflections, and the aforementioned problems do not arise. Compared to cocurrent flow, countercurrent flow has a markedly destabilizing effect for an irreversible exothermic reaction at low flow velocities (Fig. 23c). As the incoming reaction mixture in this case is in contact with the warm coolant outflow, the maximum temperature rises to much higher values. Countercurrent cooling can even lead to the occurrence of multiple steady 10.1.3.3.4

300 (a) 420

Vs = + 0.01 m ⋅s−1

T /°C

380

340

300 (b)

420

Vs = − 0.01 m ⋅s−1

380

340

300 0.6

1.0

(c)

1.4

1.8

2.2

2.6

3.0

Z/m

Fig. 23 Influence of the coolant flow direction and flow velocity vs on the reaction temperature profile [23]. (a) Isothermal; (b) cocurrent flow; (c) countercurrent flow.

states, and in general favors the run-away of a strongly exothermic, irreversible reaction [44, 45]. In contrast to pure heat exchange without a reaction, countercurrent heat transfer in reactors involving exothermic reactions should therefore be chosen only in particular cases. The temperature control of an exothermic equilibriumcontrolled reaction can constitute such a case. As illustrated in Fig. 13b, the optimum temperature profile should in this case decrease with increasing conversion – that is, along the tube length. On account of the equilibrium inhibition of the reaction, it is not possible for the reaction to run away in the front region. References see page 2105

2094

10.1 Catalytic Fixed-Bed Reactors

50

700

40

660

T 30

wt%

20

580

T / °C

620 EB

St 10 0

540

0

1

2

(a)

3

Catalyst

50

700 660

EB

T

30 20

580 St

10 0

540

0

1

2

(b)

3

4

ZL / m 50

700 EB

40

660 620

T 20

580 St

10 0 (c)

0

1

T /°C

wt%

30

540

2

3

Reaction Control through Catalyst Activity Profile A further possibility of influencing the course of the reaction is to use catalysts of different activities over the reactor length. Particularly with strongly exothermic reactions that are liable to run-away (e.g., partial oxidations), a less active catalyst is occasionally used in the front part of the reactor in order to avoid a toohigh maximum temperature. Figure 25a illustrates the use of two catalysts of differing activity in series. The resulting temperature profiles have a typical double-hump shape [46, 47], but this can be avoided if an activity profile is established by using a continuously varying mixture of catalysts with different activities (Fig. 25b and c) [48]. In these cases, the fully active entry region (relative activity 1) is designed such that the temperature rises to a prespecified maximum value. In order to maintain the temperature at this value, the activity in the following region is sharply decreased and then raised to a relative activity of 1 as the reaction rate falls due to depletion of the reactants. A smooth temperature profile can be achieved even if the optimized activity profile in Fig. 25b is only approximated for experimental verification (Fig. 25c). The control of the maximum temperature by using locally differing catalyst activities presents problems if the main reaction zone moves into the region of high catalyst activity due to changes in the operating conditions. For example, in the case of Figs. 25b and c, a decrease in the throughput may already result in reaction run-away in the short, fully active front region. This can be avoided by reducing the activity of this zone. Catalyst deactivation occurring during the operation may have more severe effects. If, for example, the catalyst is poisoned in a front migrating from the entrance to the rear, the main reaction zone finally migrates to the highly active rear catalyst region, which may lead 10.1.3.3.5

620

T /°C

40

wt%

4

ZL / m

With adiabatic and with isothermal reaction control, styrene formation decreases with increasing tube length, whereas it remains roughly constant with countercurrent flow of the heating medium. A significant advantage of non-isothermal control of the heat-transfer medium in cocurrent or countercurrent flow is the saving in circulation energy, as much smaller heat transfer-medium streams must be circulated (see the differences in coolant flow velocity vS between case a and b in Fig. 23). Overall, the combination of several heat-transfer medium circuits (Fig. 15c) and the purposeful utilization of the temperature change of the heat-transfer medium in the reactor offer a wide range of possibilities to establish optimum temperature profiles for a given reaction. This is particularly important for counteracting any changes in activity by influencing the coolant temperature profiles.

4

ZL / m

Fig. 24 Influence of heating strategy on the temperature and concentration profiles in styrene synthesis [23]. (a) Adiabatic; (b) isothermal; (c) countercurrent operation; EB = ethylbenzene; St = styrene.

Countercurrent flow of the heat-transfer medium is also advantageous for endothermic equilibrium reactions. Figure 24 shows the calculated temperature and concentration profiles with different heating conditions in the synthesis of styrene (dehydrogenation of ethylbenzene).

10.1.4 Multifunctional Fixed-Bed Reactors

100 80

690 T0 = TC = 642 K

660

T0 = TC = 633 K

0

75

60

120

40

80

T − T 0/ °C

T0 = TC = 647 K

720

T − T 0/ °C

Temperature/K

750

630

2095

20 0

150 225 300 375

40 0

(a)

1.0

0.8

0.8

1.0

0.6

0.6

0.6 0.4

Reactor length

(b)

Activity

1.0

Activity

Activity

ZL / cm

0.4 0

Reactor length

(c)

0.2

Reactor length

Influence of activity profiles on the temperature profile of a strongly exothermic reaction [23]. (a) Catalyst with 66% activity in the front section and 100% activity in the back section [45]; (b) linear (broken line) and optimal catalyst activity distribution (solid line) for limiting the maximum temperature to 370 ◦ C (simulation result for T 0 = 300 ◦ C); (c) experimental verification of (b) [48].

Fig. 25

to excessively high temperatures that can no longer be controlled. Rapid catalyst deactivation may also lead to a potentially dangerous transient run-away (this will be addressed in Section 10.1.5.2). In general, influencing the reaction course via control of the cooling stream is more flexible than incorporating catalysts of different activities. Particularly with nonisothermal control of the heat-transfer media, the reaction course can be influenced over a wide range by means of the inflow temperature of the heat-transfer medium as well as by its volumetric flow rate [49]. 10.1.4

Multifunctional Fixed-Bed Reactors

The Unit Operation Concept of Chemical Engineering asks for a clear separation of educt preparation, feed preheating, chemical reaction, down-cooling and product separation in different units. The major advantage of this concept is the maximum degree of freedom in the combination of different unit operations. It became clear, however, that better conversion and selectivity can be achieved if educts are added in the required amounts, and intermediates or products are immediately withdrawn from the place of reaction. In addition, the energy requirements can be minimized if an efficient heat exchange between the cold feeds and the hot effluents is incorporated into the reactor. Novel designs to improve the interaction of transport and reaction have therefore attracted considerable interest in recent years; these can be placed under the heading ‘‘multifunctional reactors’’ [26, 50–55].

It should be noted, however, that any integration of multiple functions into one reaction unit leads to a close coupling of the process steps involved, which generally reduces the freedom in the design and operation parameters considerably. Process integration will therefore lead to problem-specific solutions. Three general alternatives are presently under discussion to achieve this goal, heat-integrated reactors, membrane reactors and sorptive reactors. Heat-Integrated Reactors In heat-integrated reactors the hot effluent of a fixed-bed reactor is used to heat up the cold feed to the ignition temperature of the catalytic reaction. If the reaction system is all together moderately exothermic, an ‘‘autothermal’’ operation results where no other addition or removal of heat is necessary. Autothermal operation always requires a special start-up procedure to raise the catalyst temperature above the ignition temperature of the respective reactions. This means that an autothermal reactor always operates in a region of multiple steady states in the ignited state [54]. 10.1.4.1

10.1.4.1.1 Heat-Integrated Reactors for Exothermic Reactions The conventional autothermal reactor design consists of an adiabatic packed-bed reactor coupled with a countercurrent heat exchanger (Fig. 26a). As mentioned above, the reaction must be started with help from a separate preheater through which the catalyst bed temperature is raised above the ignition temperature of the reaction. During operation, control measures must References see page 2105

2096

10.1 Catalytic Fixed-Bed Reactors

z

∆Tad

2∆Tad

Conventional design T

T ign

(a)

2∆Tad

z

∆Tad

Countercurrent fixed-bed reactor T (b)

T ign

Fig. 26 Autothermal fixed-bed reactors with recuperative heat exchange. Basic design and typical temperature profiles [60]. (a) Conventional design with separate heat exchanger; (b) countercurrent fixed-bed reactor.

be taken to prevent the reaction from extinction if, for example, the feed to the reactor is too lean. Because the hot effluent is used to heat up the cold feed, the maximum temperatures achieved under autothermal operation may exceed many-fold the adiabatic temperature rise [see Eq. (5)]. To prevent overheating, autothermal operation is limited to reactions with a maximal adiabatic temperature rise Tad in the order of 300–400 K (depending upon the temperature stability of the catalyst and the reactor construction). It is best suited for reactions with a low overall exothermicity. The catalytic combustion of traces of organic compounds or catalytic hydrogenations are typical examples. With conventional tube and shell design for the countercurrent heat exchanger, the design of Fig. 26a is only suited for an adiabatic temperature rise exceeding about 150 K. To reduce this limit, new designs with better heat-transfer characteristics have been conceived [54].

Figure 26b illustrates the concept of the countercurrent fixed-bed reactor where the catalyst is placed inside and outside of a tube bundle that forms a countercurrent heat exchanger for the reacting gas. Alternatively, a parallel-plate design can be used with the catalyst deposited at the plate surface or between the plates. One advantage of this design is the improved heat transfer caused by the presence of the catalyst packing. Another advantage stems from the integration of heat exchange and reaction within the reactor as compared to its separation in Fig. 26a. This can be seen from the temperature profiles for a single irreversible reaction in Fig. 26 (right-hand side). As long as the temperature is below the ignition temperature of the catalytic reaction, the slope of the temperature profiles in the heat-exchange section is the same in both designs (provided that the heat-transfer parameters are identical and the heat capacity flux is the same in both directions). As soon as the temperature of the incoming gas exceeds the ignition temperature, T ign , the reaction will start immediately in the countercurrent reactor. If the activation energy is sufficiently high, the temperature will increase rapidly and the reaction will be completed after a short time. In a limiting case, the heat exchange over the reaction front can be neglected and the maximum temperature reached can be estimated from the ignition temperature plus the adiabatic temperature rise, Tad . This is of course only a rough estimate as the ignition temperature is not a properly defined quantity, nor is the influence of the heat exchange over the length of the reaction front really negligible. The estimate is sufficient, however, to obtain a reasonable picture of the main differences between the two types of autothermal reactor. As can be seen in Fig. 26, a doubling of the adiabatic temperature rise leads to a doubling of the total temperature rise in case (a) whereas a substantially lower total temperature increase results in case (b). The countercurrent fixed-bed reactor is obviously self-adaptive in that it establishes the necessary temperature level for the respective reaction, whereas in the standard type the temperature level is directly proportional to the slope of the temperature front in the heat-exchanger section, dT /dz [see Eq. (6)], which in turn is proportional to the adiabatic temperature rise Tad [54, 56]. dT = dz

Tad  1 λs (1 − ε) 2 + Gz cpG Gz cpG αaP


(6)

Here, λS is the axial bed conductivity, ε is the bed void fraction, Gz is the axial mass flux density (mass per reactor cross-section and time), α is the pellet heat-transfer coefficient, and ap is the specific outer pellet surface area (area per reactor volume). A completely different design of an autothermal reactor has been proposed and developed by Matros and

2097

Z

t

10.1.4 Multifunctional Fixed-Bed Reactors

τ

T°

C°

T°

T′

c,T

(a)

(b)

∆Tad

(c)

Reverse-flow reactor with direct (regenerative) heat exchange for an irreversible reaction [14]. (a) Basic arrangement; (b) local concentration and temperature profiles prior to flow reversal in the periodically steady state; (c) variation of outlet temperature with time in the periodically steady state.

Fig. 27

coworkers [57–59]. A sketch of the basic concept, showing an adiabatic fixed-bed reactor with feed/exit tubes and two valves for periodic flow reversal, is illustrated in Fig. 27a. Before starting the operation, the fixed bed must again be heated above the ignition temperature of the catalytic reaction. If the reactor is then fed with cold feed from one side, the cold feed gas is heated up by the hot catalyst on one side and cooled down by the cold catalyst on the other side. Two temperature fronts develop and move through the packed bed. In the first front, the gas is heated up and reacts, but after a certain time period the direction of flow is reversed via the valves and the temperature fronts are pushed back. In the cyclic steady state a hot plateau in the middle of the packed bed thus moves up and down, while the exit and inlet sections of the packed bed serve as regenerative heat exchangers. As the reactor is adiabatic, the heat of reaction can only be removed with the leaving gas, and the gas exit temperature shows a saw-tooth-like variation in time (Fig. 27c). From an overall energy balance, it is clear that the integral of the exit temperature over time must exceed the feed temperature just by the adiabatic temperature rise Tad . It is interesting to note that, in the limit of rapid flow reversal, the reverse-flow reactor and the countercurrent fixed-bed reactor show completely similar temperature and conversion profiles [56]. This can be understood with the help of Fig. 28. With rapid flow reversal, the catalyst temperature will remain constant due to the large heat

capacity of the packing, while the gas temperature will be below the catalyst temperature in the respective feed section and above in the exit section (Fig. 28b). This behavior is completely similar to that of a countercurrent fixed-bed reactor, where the catalyst is placed at the separating walls between the up- and down-flowing gas (Fig. 28c). It only has to be considered that, instead of pushing the reacting gas for a short period in one direction and for another period in the other direction, half of the mass flow will now go permanently in direction one and the second half in the other direction. As it is much easier to solve the steady-state model of the countercurrent fixed-bed reactor than the dynamic model of the reverseflow reactor, the analogy can be used for rapid design calculations for the reverse-flow reactor [14, 54]. Moreover, the simple design approximation using the slope of the temperature front [Eq. (6)] and the maximum temperature estimated from the ignition temperature and the adiabatic temperature rise (Fig. 28d) helps to understand the basic features of both reactors [56]. The obvious advantage of the reverse-flow reactor over the countercurrent fixed-bed reactor is the simple form of the adiabatic packed bed and the excellent efficiency of the regenerative heat exchange. Reactions with an adiabatic temperature rise as low as 10 K can be run autothermally at maximum temperatures exceeding 500 ◦ C in a properly designed fixed bed. It is therefore presently a favorite design for the catalytic oxidation References see page 2105

2098

10.1 Catalytic Fixed-Bed Reactors

(a) Feed

Catalyst or wall

Exit

(a)

T

∆Tad

z

(b)

(c) ∆Tad

Feed

Ti

(b)

C°

c,T

∆T

Exit Purge

Fig. 29 Alternative arrangements for autothermal reactor design with periodic flow reversal. (a) Radial-flow concept [14]; (b) multiple-bed arrangement with bed purge prior to flow reversal [60].

T (d)

z

Fig. 28 Equivalence of reactor operation with periodical flow reversal and countercurrent heat exchange for a weakly exothermic irreversible reaction. (a) Fixed-bed reactor with periodic flow reversal; (b) temperature profiles with rapid flow reversal; (c) countercurrent reactor with catalyst at the wall; (d) schematic concentration and temperature profiles in both reactors [14].

of traces of combustible components in exhaust air. However, its prime disadvantage is the unsteady mode of operation and the need to switch large gas streams periodically. In the basic design of Fig. 28, unreacted gas in the entrance hood and the preheating section of the packed bed is flushed into the exit with every flow reversal. To avoid this breakthrough, most commercial air purification units use a three-bed design (Fig. 29b), where one bed is purged with clean air prior to flow reversal. In addition to the standard design, a number of modifications have been proposed and applied (see also

Chapter 10.4). As explained in Section 10.1.3.1, large adiabatic packed beds are preferably designed as radialflow reactors to avoid excessive pressure drop with increasing bed height. In the radial-flow design depicted in Fig. 29a, the hot region in the middle insulates itself against heat losses. This design can easily be extended to a three-bed arrangement. Designs with rotating fixed beds similar to the Ljungstr¨om heat exchanger design have also been proposed and tested [14]; this design offers the possibility of a continuous, valveless operation at the expense of rotating seals at the cold end of the rotating fixed bed. Typical applications for exhaust air purification are characterized by rapid concentration changes. An efficient control strategy must prevent the reaction from extinction during times of low concentrations, and the catalyst from overheating and sintering during times of rich feed. Several possibilities to achieve this goal have been discussed in Ref. [60]. Aside from air purification, autothermal operation has been proposed and demonstrated for a number of exothermic synthesis reactions with equilibrium

2099

10.1.4 Multifunctional Fixed-Bed Reactors

10.1.4.1.2 Heat-Integrated Reactors for Endothermic Reactions Interesting aspects of autothermal reactor operation arise if an endothermic synthesis reaction is closely coupled with an exothermic auxiliary reaction in such a way that the combination is weakly exothermic. Under these conditions, three different cases can then be distinguished, as discussed in Ref. [54]. If the endothermic and exothermic reactions take place simultaneously in the catalyst bed, as in the case of oxidative hydrogenation of hydrocarbons, the same concepts as discussed in the previous section can be applied. An early example with regenerative heat exchange and periodic flow reversal has been given by Amoco for the oxidative steam reforming of methane [62]. If the endothermic reaction has to be run separated from the exothermic one, a close thermal coupling between both reactions should be achieved, where the heat of the exothermic reaction must be taken up by the endothermic reaction as soon as it is set free. In addition, all hot product gases should be used to heat up all cold feeds. Clearly, the ‘‘heat-integrated’’ concepts discussed in Figs. 20 and 21 fulfill this requirement only to a limited extent, as the heat of the hot flue gases is not used completely to heat up the cold feeds. In the context of decentralized hydrogen production for fuel cells, several new concepts have been developed. The most obvious design is an arrangement of parallel-flow channels, where in every second channel the endothermic reaction and in every other channel the exothermic reaction takes place. In order to incorporate an efficient heat exchange between the hot products with the cold feeds, a countercurrent flow should be established while the catalyst in both channels is confined to the hot central region (Fig. 30a). Unfortunately, such a design has not been successful. An explanation is given in Fig. 30b, showing the simplified temperature profiles in both channels [27]. Here, the exothermic (combustion) reaction has been considered to be instantaneous (which is a reasonable

Exo feed Endo feed

Catalyst

(a) exo Tmax

2000

1500

T/ °C

limitations such as methanol, ammonia and sulfur trioxide synthesis (see Chapter 10.4 and references therein). An advantage has been seen in the fact that the decreasing temperature profile in the exit section of the reactor leads to an increased equilibrium conversion over that obtained adiabatically [39, 54]. However, a twostage adiabatic reactor with interstage cooling will still lead to higher conversion [61], with the additional advantage that the heat of reaction after each stage can be used for the additional production of steam (as in Fig. 14). Purely autothermal concepts for the above-mentioned bulk chemicals will therefore probably be restricted to sites where simple reactor operation without waste heat utilization is required.

exo ∆T ad

1000

500

0.1

0.2

0.3

0.4

z /m

(b) Endo feed

Air Burner gas

(c)

Endo feed Catalyst (d)

Exo feed

Fig. 30 Coupling options for endothermic and exothermic reactions [27]. (a) Pure countercurrent coupling; (b) schematic temperature profiles for the countercurrent coupling of an instantaneous exothermic (combustion) reaction with a constant-rate endothermic reaction with equal heat capacity flows; (c) coupling with distributed burner gas side feed; (d) coupling with cocurrent reaction section.

approximation at the high temperatures in the reactor center), whereas a constant reaction rate has been assumed for the endothermic reaction. In order to achieve an efficient heat exchange in the heat-exchange sections, the heat capacity flux in both reaction channels should be equal. If the endothermic reaction has an adiabatic temperature drop of 1000 ◦ C (as in case of methane steam reforming), the adiabatic temperature rise of the exothermic reaction exo must be at least about 1100 ◦ C to compensate for Tad heat losses. With an instantaneous exothermic reaction, this would lead to intolerable maximum temperatures, as shown in Fig. 30b. In fact, all pure countercurrent concepts for the coupling of endo- and exothermic reactions have shown the danger of an almost adiabatic run-away of the exothermic reaction, as the main reaction zones of both reactions have an inherent tendency to separate [27]. References see page 2105

2100

10.1 Catalytic Fixed-Bed Reactors

Two alternatives to circumvent the above difficulties are shown in Figs. 30c and d. In case c, one educt of the exothermic reaction (e.g., the burner gas) is added at several locations along the length of the active reaction zone. This enforces a heat input, which is more uniform over the length of the reaction zone. In case d, the heatexchange sections remain countercurrently, whereas the flow in the reaction zone is cocurrent. This leads to a more uniform overlapping of both reaction zones. Depending on the intensity of the heat contact between both channels, and on the reaction kinetics, the exothermic reaction may still develop a hot spot or the endothermic reaction may quench the exothermic reaction, leading to a cold spot, although the reactor behavior is generally more uniform and robust [28]. The direct coupling of endo- and exothermic reactions has also been addressed, applying micro-reactor technology (see Chapter 10.8). Here, an excellent heat transfer between the endothermic and exothermic reactions can be achieved if the catalyst is deposited on opposite sides of the channel walls. In addition, homogeneous ignition of the combustion reaction can be prevented if the channel dimensions are in the sub-millimeter range (≈200 µm). Wall heat conduction assures an almost uniform temperature profile as long as the reactor dimensions are in the low-centimeter range. However, wall heat conduction prevents an efficient heat recovery from the hot effluents to heat up the cold feeds if the heat-exchange sections are not extended to several tens of centimeters in length [28]. Sorptive and Membrane Reactor Concepts Just as heat can be added to or removed from the reaction site either continuously through the wall in recuperative heat exchange or discontinuously in direct, regenerative heat exchange, specific reaction components can also be provided or removed selectively by two different processes. In membrane reactors the addition or extraction takes place continuously through a semi-permeable membrane which can be immersed into the catalyst bed. This leads to the concept of membrane reactors (this is treated in detail in Chapter 10.7). For industrial applications, one major disadvantage of catalytic membrane reactors is the fact that no convincing large-scale concepts have yet been proposed. This concerns both the implementation of large membrane areas necessary for the production of bulk chemicals within a chemical reactor, and its combination with devices for the addition or removal of the required heat of reaction. Membrane reactor concepts are therefore presently limited to laboratory-scale investigations. An alternative is the sorption reactor concept, where specific reaction components are deposited at sorption sites inside the catalyst bed, which must periodically be 10.1.4.2

filled-up or emptied (i.e., regenerated). As between the continuous recuperative and the periodic regenerative heat exchange (see Section 10.1.4.1.1), there is a general analogy between sorptive and membrane methods for the limiting case of short regeneration periods [63]. The basic difference is that membrane methods allow for a continuous addition or removal of components through a semi-permeable wall, whereas sorptive methods are essentially dynamic; that is, they provide a limited storage capacity for some components if a suitable adsorbent is placed near the site of the catalytic reaction. Sorptive methods are briefly discussed in Chapter 10.4, but for a more detailed discussion of the current state, see Ref. [64]. 10.1.5

Safety Issues of Fixed-Bed Reactors

Fixed-bed reactors belong to the largest and most widely used reactor group in the chemical industry. As large amounts of combustible or potentially decomposable gases are processed in these reactors – usually at high temperatures and elevated pressures – safety issues are important for their design and operation. Since the early 1970s, such issues have been treated mainly under the heading ‘‘thermal run-away’’ in numerous publications which might give the impression that fixed-bed reactors constitute a particularly critical reactor type, with a large potential risk. In fact, the opposite is true. Compared to a liquid-phase reactor of the same size, a fixed-bed reactor with a gas-phase reaction contains a mass of reactants several orders of magnitude smaller. The danger of a rapid pressure increase due to evaporation does not exist and, due to the reduced mass at comparable volume, the danger of decomposition of reactants, accumulated in the reactor, is considerably smaller than in liquidphase reactors. The catalyst mass additionally damps the uncontrolled temperature rise in fixed-bed reactors. Parametric Sensitivity and Run-Away Nevertheless, instabilities can arise in fixed-bed reactors, particularly with strongly exothermic reactions, and can lead to excess temperatures that may damage the catalyst and the reactor construction materials. Several causes for this behavior have been identified [23], but colloquially they are all considered under the heading ‘‘run-away’’. In fixed-bed reactors, run-away usually occurs under operating conditions of high parametric sensitivity, where small changes in the operating parameters can lead to large changes in the maximum temperature and yield. One main reason is the exponential dependence of the reaction rate on temperature (Arrhenius law). Figure 31 shows calculated temperature profiles for a partial oxidation reaction in a wall-cooled, fixed-bed reactor tube 10.1.5.1

10.1.5 Safety Issues of Fixed-Bed Reactors

2101

1000

1000

344 °C 800

500 °C

T / °C

T / °C

800

600

600 343 °C

400 °C

320 °C

400

400

300 °C

300 °C 200 (a)

0

0.5

1.0

200

1.5

0

(b)

Reactor length / m 1400

80

0.5

1.5

Reactor length / m 1100 Main + consecutive reaction

1200

70

1.0

900 1000

50 600 40

Tmax / °C

800

Tmax / °C

Yield / %

60

Main reaction

700 500

400

No reaction 300

30 300 (c)

320

340

200

Tc / °C

200 (d)

300

400

500

600

Tc / °C

Fig. 31 Parametric sensitivity of a partial oxidation reaction in a fixed-bed reactor of typical dimensions as a function of the coolant temperature Tc with T(z = 0) = Tc [23]. (a) Temperature profile over reaction length (main reaction only); (b) temperature profile including total oxidation as side reaction; (c) maximum temperature Tmax and yields as a function of coolant temperature Tc in case (b); (d) Tmax as a function of Tc for both cases.

of typical dimensions. In Fig. 31a, only the main reaction is considered, while in Fig. 31b the total combustion to CO2 and water is additionally taken into account. Both cases are almost identical up to coolant temperatures of 330 ◦ C. As considerably more heat is liberated in the total combustion than in the desired main reaction, the sensitivity is increased substantially as soon as the ignition temperature of the second reaction is exceeded. As a measure of the parametric sensitivity, Fig. 31c shows the change in maximum temperature via the cooling temperature for case b. The sensitivity is only moderate at low cooling temperatures, whereas above Tc = 343 ◦ C small changes in TC – and also in other parameters such as throughput, feed concentration, or pressure – lead to large changes in reactor behavior. In

fact, under the conditions considered, a new steady-state is approached if the coolant temperature is raised above Tc = 344 ◦ C (Fig. 31d). This means that the maximum temperature remains at an upper value, even if the coolant temperature is decreased below Tc = 340 ◦ C. Due to the unavoidable differences between individual tubes, multitubular reactors cannot be operated in the range of high parametric sensitivity. In the case considered, the cooling temperature must be kept below ca. 340 ◦ C, and the tubes made longer to give a good conversion. This example emphasizes the requirement discussed in Section 10.1.3.3 for keeping the conditions in the tubes of the tube bundle and in the cooling circuit as References see page 2105

2102

10.1 Catalytic Fixed-Bed Reactors

uniform as possible to avoid premature run-away reaction in individual tubes. Published reports include numerous run-away criteria with which operating ranges of high parametric sensitivity can be precalculated for known reaction kinetics [65, 66]. In practice, these parameters are of only limited importance because they rarely take into account the peculiarities of individual cases. Sensitive reactions such as partial oxidations and partial hydrogenations are, therefore, generally tested in single-tube reactors of the same dimensions as those in the subsequent multitubular reactor. This allows the range of parametric sensitivity to be determined directly. Recalculation of the results for other tube diameters is only possible to a limited extent due to uncertainties in the quantification of the heat-transfer parameters (see Section 10.1.2.4). A fixed-bed reactor can enter the region of high parametric sensitivity through changes in the catalyst properties or operating conditions. Rapid changes in feed temperature, feed concentration or throughput may induce migrating reaction zones which can lead to transient excess temperatures, a phenomenon known as ‘‘wrong-way behavior’’ (see Section 10.1.5.2 and Chapter 10.4). Under run-away conditions, initially a few particularly sensitive tubes of the bundle will run-away – that is, the reaction changes, perhaps from a selective partial oxidation to a total combustion – and the temperature rises rapidly. In a multitubular reactor with thousands of tubes, not every tube can be equipped with temperatureprofile measurements; it is therefore likely that this run-away will remain undetected, especially if it involves only a few tubes. Although temperatures >1000 ◦ C can often be reached in the catalyst during such run-aways, there is no safety risk, provided that the tube is surrounded by a liquid heat-transfer medium. Because of the good heat transfer to the fluid, the tube temperature remains close to that of the heat-transfer medium, and melting of the tube does not occur even if the catalyst pellets are destroyed by the high temperatures. The most certain method of detecting a run-away is on-line analysis of a product formed in the run-away reaction. For example, CO2 can be monitored in the off-gas during the run-away-sensitive synthesis of ethylene oxide. If its concentration increases above a specified limit, the reactor must be shut down, purged with nitrogen, and for a certain period cooled to a lower temperature before operation is recommenced. Wrong-Way Behavior and Transient Run-Away If in an adiabatic fixed bed (without reaction) the feed temperature is suddenly increased or decreased, a dispersive temperature front develops with a front velocity 10.1.5.2

wT ερG cpG vZ ρS cS

wT =

(7)

where ερG cP G and ρS cpS are the heat capacities of gas and solid (packing), respectively. If in addition an irreversible exothermic reaction takes place, a self-sharpening reaction front is generated, moving with velocity wR , which is in general different from wT . This case is shown in Fig. 32, where a reaction front has been started by preheating an adiabatic catalytic fixed-bed reactor locally above the ignition temperature of the reaction. The cold feed is then heated up by the hot packing so that the reaction can proceed and the reaction front moves downstream with wR . Simultaneously, the heat of the reaction zone is carried ahead in a thermal front with wT , and the heat liberated by reaction heats up an increasingly larger portion of the bed between the two moving fronts. From a simple heat balance between the heat of reaction and the heat accumulated in the hot portion of the bed (the shaded area in Fig. 32), it follows εvz ρG cpG Tad = (wT − wR )ρS cS T

(8)

Substituting Eq. (7) leads to T =

wT Tad w T − wR

(9)

This shows that the total temperature rise T in the reaction front may substantially exceed the adiabatic temperature increase, a result already observed in connection with the reverse-flow reactor (Figs. 27 and 28). There, it was shown that the total temperature rise T can be approximated by T ≥ Tign + Tad , leading to wR ≥ wT

Tign Tign + Tad

(10)

A similar moving reaction front occurs if the flow velocity in a fixed-bed reactor is suddenly increased above the T T max ∆Tad

T ign

T0

wR

TS

wT ∆T

TG z

Fig. 32 Development of a reaction front for an irreversible exothermic reaction in a locally preheated, sufficiently long adiabatic fixed-bed reactor.

10.1.5 Safety Issues of Fixed-Bed Reactors

blow-out velocity of the reaction, or the feed temperature is suddenly decreased substantially below the ignition temperature. For the (temporal) maximum temperature increase, resulting from a feed temperature drop, the term ‘‘wrong-way behavior’’ has been established [67]. From Eq. (10) it follows that, in a moving reaction front, usually wR < wT . This means that the heat of reaction is distributed over an extended portion of the bed (see Fig. 32), and hence the maximum temperature rise T is bounded. A notable exception occurs if the reaction front velocity wR is not freely established but is forced to increase. This may happen if the catalyst is rapidly deactivated, for example by a catalyst poison in the feed [68]. The deactivation front then pushes the reaction front ahead. A critical situation obviously occurs if the deactivation front velocity equals the thermal front velocity [wR = wT in Eq. (9)]. Then, the heat of the reaction can no longer be distributed over an extended bed portion, as in Fig. 32, but rather accumulates in a narrow region, resulting in a constantly increasing maximum temperature. In this case, the simplified Eq. (9) predicts in the long term an infinite temperature increase. Although it is rather unlikely that a poison front velocity coincides with the temperature front velocity, the situation is different if the deactivation is a result of too-high catalyst temperatures. Then, a vicious circle can develop where the high catalyst temperature in the reaction front deactivates the catalyst, which induces a reaction front movement through which the catalyst temperature is further increased. Such a situation has been observed experimentally [69, 70], and is shown as a simulation result in Fig. 33. The specific example was the oxidation 2500 2000

800

1500

700

T /K 1000 600

t = 500s

500 0

0.5 Z /L

1

Fig. 33 Run-away of CO-oxidation on a Ni catalyst due to rapid catalyst deactivation [69].

2103

of traces of CO in an industrial adiabatic fixed-bed reactor, using a Ni catalyst with a too-low thermal stability [69]. The danger of such a situation is that the steadystate operation of a weakly exothermic reaction with an adiabatic temperature rise below 100 K seems to require no special precautions. Therefore, the gradual development of the high-temperature front may remain unnoticed as long as it is inside the packed bed. The more dramatic is the sudden substantial exit temperature rise, as soon as the front leaves the catalyst bed. Decoking of Fixed-Bed Reactors So far, only gas–gas reactions catalyzed by a fixedbed catalyst have been considered. Under certain circumstances, however, gas–solid reactions also take place in catalytic fixed-bed reactors. A typical example is the periodic burn-off of the coke which is formed during reforming or dehydrogenation reactions, and tends to block the catalyst surface. Another example of particular importance is the regeneration of Diesel soot filters. In both cases, the fixed-bed entrance is periodically heated up above the coke or soot ignition temperature, and a gas flow containing oxygen is introduced. The developing temperature fronts can be treated similarly to the previous section [71, 72]. The main difference is that now a gas component AG (oxygen) is reacting with a solid reactant BS , which is deposited at the fixed-bed surface. In the reaction front the initial coke loading qB0 is completely burned off, leading to the simple stoichiometric relationship: 10.1.5.3

0 wR · qB0 = vG · cA

(11)

0 is the oxygen feed concentration. Hence, the where cA reaction front velocity wR is now completely specified by the gas feed velocity vG , the coke loading, and the oxygen feed concentration. If the coke loading is low and the oxygen concentration high (Fig. 34, left), wR will be large and exceed the thermal front velocity. In this case the reaction front will precede the thermal front (Fig. 35, left). If the coke loading is high and the oxygen concentration low (Fig. 34, right), the thermal front will precede the reaction front (Fig. 35, right). In both cases, the heat of coke combustion will be dispersed over a larger portion of the fixed bed (the shaded areas in Fig. 35) and the maximum temperature rise TF is limited. Again, a simple heat balance between the heat released and stored in the fixed bed gives an approximation for TF (see the equations in Fig. 35). A singularity with (almost) unlimited temperature rise can be expected if the reaction front velocity approaches the thermal front velocity. To avoid this situation,

References see page 2105

2104

10.1 Catalytic Fixed-Bed Reactors

is comparatively low. Leaving out the peculiarities of individual cases, the following safety risks can be assumed for fixed-bed reactors:

G

A

c 0A

• Leaks which result in the release of large amounts of gas or vapor and the formation of explosive clouds • Leaks resulting in the release of large amounts of liquid heat-transfer media (oils, salt melts) • Occurrence of ignitable or decomposable gas mixtures in the reactor • Melting of the reactor due to a run-away reaction.

products

wR

wR

VZ

Reaction front movement with velocity wR during a gas-solid (decoking) reaction in a fixed-bed reactor.

Fig. 34

decoking of industrial fixed-bed reactors is usually started with very low oxygen concentrations. Then, w T  w R . Following the equation in Fig. 35 (right), the maximum temperature rise can now be controlled through the adiabatic temperature rise with respect to the oxygen 0 , irrespective of the initial coke loading. concentration cA Diesel soot filters, on the other hand, are usually regenerated with Diesel exhaust containing about 10 vol.% oxygen. With this high oxygen concentration the regeneration temperature can only be controlled by limiting the soot loading qB0 (left side of Fig. 35). As on-line monitoring of actual soot loading is far from trivial, soot filter regeneration is usually initiated in short intervals to safely avoid excess temperatures during regeneration. Other Safety Aspects Because of the small mass storage capacity compared to liquid-phase reactors, the danger of a sudden reaction of accumulated reactants in gas-phase, fixed-bed reactors 10.1.5.4

Temperature in °C

q 0B low, c A0 high

q 0B limiting

q 0B high, c A0 low

1000 800 600 400

S ∆T ad

WT

∆TF

200 0 • (WR − WT) × rcP ∆TF = ∆TF =

The safety aspects of liquid heat-transfer media have been discussed in Section 10.1.3.3.1. Ignitable gas mixtures can arise particularly during partial oxidation reactions, and they are especially critical where large, packing-free volumes are present. This is the case in the inflow and outflow hoods of the reactor, whereas in the reactor tubes the catalyst packing dampens the propagation of a flame front due to its heat capacity. The complete avoidance of an ignitable mixture is generally not possible in partial oxidations because, at least during mixing of the gas streams prior to the reactor, the ignition limit is exceeded locally. However, the operation of fixedbed reactors with ignitable mixtures has generally been avoided, either by dilution with inert gas or by operating in the non-stoichiometric range. The former requires additional efforts for heating, cooling, and separation of the inert gas, while the latter gives only low conversions of the reactants in a single pass. New developments in partial oxidation therefore aim for stoichiometric operation in the ignitable range [34]. Referring to multitubular fixed-bed reactors of conventional dimensions, a prerequisite for this is a pressure-resistant construction with check valves

WR q B0 (−∆hR)

WR WR − WT

×

q 0B (−∆hR) rcP S ∆ Tad

WR

Temperature in °C

AG + BS

BS

c 0A limiting

1000 800

WT

600 400

WR

200 0 • (WR − WT) × rcP ∆TF = VZ c A0 (−∆hR)

∆TF =

WT

c 0A(−∆hR)

WT − WR

e(rcP)G ∆ TaGd

Temperature profiles during coke burn-off in fixed-bed reactors [72]. Left side: low coke loading q0B , high oxygen concentration cA0 : preceding reaction front. Right side: high coke loading q0B , low oxygen concentration cA0 : trailing reaction front (wT , thermal front velocity; wR , reaction (combustion) front velocity).

Fig. 35

References

and flame barriers, so that a possible ignition is confined to the interior of the reactor. Melting of reactor tubes during run-away reaction is only to be feared in multitubular reactors if the respective tube is not surrounded by liquid heat-transfer medium. Thus, appropriate design must ensure that running dry of reactor tubes cannot occur. In the case of corrosive reaction gases, provision for the detection of leaks caused by corrosion must be made, particularly when pressurized or boiling water is used as coolant. A new option to cope with some of the above-mentioned safety concerns are ‘‘micro-reactors’’ (see Chapter 10.8). It has been shown [73, 74] that extremely exothermic reactions with potentially explosive feed composition, such as catalytic hydrogen oxidation with stoichiometric air feed, can be carried out without thermal run-away in micro-reactors if the cross-sectional dimensions of the reaction channels are sufficiently small (the order of 200 µm). In addition, the micro-mixing of reactants prior to reaction or controlled addition of reactants with immediate mixing becomes possible. This opens the possibility of carrying out reactions which have so far been impossible because of a lack of sufficient temperature and residence time control. The challenges to be solved are the distribution and collection of the feed streams/the effluents to/from a large number of micro-reactors, even if the volumetric productivity of a single micro-reactor can exceed the productivity of conventional large-scale reactors by orders of magnitude. 10.1.6

Conclusions

The different industrially established fixed-bed reactor configurations of the adiabatic, the multi-stage and the multitubular reactor types represent a mature and versatile class of reactors for heterogeneously catalyzed gas-phase reactions. Companies specialized in their construction offer a wide variety of different designs to meet specific requirements of operation and temperature control, including operation above the explosion limit of the gas mixture. Only in the case of rapid catalyst deactivation are other reactor concepts such as fluidizedbed reactors (see Chapter 10.2) mandatory. In addition to the well-established fixed-bed reactor configurations mentioned above, new concepts are being discussed where specific features such as distributed educt addition and selective product removal are integrated in the fixed bed. From these ‘‘multifunctional reactors’’, autothermal reactor concepts where the heat exchange between the cold feed and the hot effluent is integrated in the fixed bed have already been established in industrial practice.

2105

‘‘Sorptive reactor concepts’’, where periodic operation is used to temporarily store or remove educts or products in the fixed bed, have been occasionally applied in industrial practice, whereas membrane reactor concepts with permselective inert or catalytically active microporous membranes are still at the laboratory or pilot plant stages. They promise the highest potential for a further improvement of catalytic reactor technology, and present the biggest challenges. With micro-reactors, the possibility has been introduced selectively to carry out catalytic reactions of very high exothermicity, even in the potentially explosive range, due to strict temperature and residence time controls. These may be considered an interesting option to extend Chemical Reaction Engineering into areas of operation which, to date, have been inaccessible. At present it remains open, however, as to whether this option will be attractive for more than a small number of industrially relevant reactions. References 1. G. F. Froment, K. B. Bischoff, Chemical Reactor Analysis and Design, John Wiley, New York, 1990, 664 pp. 2. K. R. Westerterp, W. P. M. van Swaaij, A. A. C. M. Beenackers, Chemical Reactor Design and Operation, John Wiley, Chichester, 1987, 767 pp. 3. M. Baerns, H. Hofmann, A. Renken, Chemische Reaktionstechnik, 3rd Ed., Georg Thieme Verlag, Stuttgart, 1999, 428 pp. 4. J. B. Rawlings, J. G. Ekerdt, Chemical Reactor Analysis and Design Fundamentals, Nob Hill Publishers, Madison, 2002, 609 pp. 5. L. Riekert, Appl. Catal. 1985, 15, 89–102. 6. V. Kottke, H. Blenke, Verfahrenstechnik (Mainz) 1982, 16, 504–509. 7. M. Nijemeisland, A. G. Dixon, AIChE J. 2004, 50, 906–921. 8. Verein Deutscher Ingenieure, VDI-W¨armeatlas, 7th Ed., Section Gh, VDI-Verlag, D¨usseldorf, 1994. 9. R. K. Shah, A. L. London, Advances in Heat Transfer, Supplement 1, Academic Press, New York, 1978. 10. G. Gaiser, V. Kottke, Chem. Ing. Tech. 1989, 61, 729–731. 11. K. Ramanathan, V. Balakotaiah, D. H. West, Chem. Eng. Sci. 2003, 58, 1381–1405. 12. G. Gaiser, V. Kottke, Chem. Eng. Technol. 1989, 12, 400–405. 13. Verein Deutscher Ingenieure, VDI-W¨armeatlas, 7th Ed., Section Le, VDI-Verlag, D¨usseldorf, 1994. 14. G. Eigenberger, U. Nieken, Chem. Ing. Tech. 1991, 63, 781–791; G. Eigenberger, U. Nieken, Int. Chem. Eng. 1993, 34, 4–16. 15. K. R. Westerterp, K. J. Ptasinski, Chem. Eng. Sci. 1984, 39, 245–252. 16. D. Vortmeyer, R. P. Winter, Chem. Ing. Tech. 1983, 55, 312–313. 17. M. Winterberg, E. Tsotsas, Chem. Eng. Sci. 2000, 55, 5937–5943. 18. O. Bey, G. Eigenberger, Int. J. Therm. Sci. 2001, 40, 152–164. 19. A. E. Rodrigues, R. M. Quinta Ferreira, AIChE Symp. Ser. 1988, 84, 80–87. 20. A. E. Rodrigues, R. M. Quinta Ferreira, Chem. Eng. Sci. 1990, 45, 2653–2660.

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55. F. M. Dautzenberg, P. J. Angevine, Catal. Today 2004, 93–95, 3–16 56. U. Nieken, G. Kolios, G. Eigenberger, AIChE J. 1995, 41, 1915–1925. 57. G. K. Boreskov, G. A. Bunimovich, Y. S. Matros, A. A. Ivanov, Kinetika i Kataliz 1982, 23, 402–406; G. K. Boreskov, G. A. Bunimovich, Y. S. Matros, A. A. Ivanov, Int. Chem. Eng. 1982, 22, 335–342. 58. Y. S. Matros, Chem. Eng. Sci. 1990, 45, 2097–2102. 59. Y. S. Matros, G. A. Bunimovich, Catal. Rev. -Sci. Eng. 1996, 38, 1–68. 60. U. Nieken, G. Kolios, G. Eigenberger, Chem. Eng. Sci. 1994, 49, 5507–5518. 61. B. Young, D. Hildebrandt, D. Glasser, Chem. Eng. Sci. 1992, 47, 1825–1837. 62. R. F. Blanks, T. S. Wittrig, D. A. Peterson, Chem. Eng. Sci. 1990, 45, 2407–2413. 63. A. Gorbach, G. Eigenberger, G. Kolios, Ind. Eng. Chem. Res. 2005, 44, 2369–2381. 64. D. W. Agar, in Integrated Chemical Processes–Synthesis, Operation, Analysis, and Control, K. Sundmacher, A. Kienle, A. Seidel-Morgenstern (Eds.), Wiley-VCH, Weinheim, 2005, p. 203–232. 65. G. F. Froment, Front. Chem. React. Eng. 1984, 1, 12–38. 66. M. Morbidelli, A. Varma, AIChE J. 1986, 32, 297–306; A. Varma, M. Morbidelli, Chem. Eng. Sci. 1991, 46, 3330–3332. 67. V. Pinjala, Y. C. Chen, D. Luss, AIChE J. 1988, 34, 1663–1672. 68. G. Eigenberger, Chem. Ing. Tech. 1978, 50, 924–933; G. Eigenberger, Int. Chem. Eng. 1981, 21, 17–28. 69. E. Blaum, Chem. Eng. Sci. 1974, 29, 2263–2277. 70. G. Emig, H. Hofmann, U. Hoffmann, U. Fiand, Chem. Eng. Sci. 1980, 35, 249–257. 71. G. Kolios, A. Gritsch, B. Gl¨ockler, G. Eigenberger, in Integrated Chemical Processes–Synthesis, Operation, Analysis, and Control, K. Sundmacher, A. Kienle, A. Seidel-Morgenstern (Eds.), Wiley-VCH, Weinheim, 2005, p. 3–43. 72. A. Salden, Adsorption/Incineration Processes for Waste Gas Purification, Logos Verlag, Berlin, 2002, 165 pp. 73. G. Veser, Chem. Eng. Sci. 2001, 56, 1265–1273. 74. D. G. Norton, D. G. Vlachos, Chem. Eng. Sci. 2003, 58, 4871–4882.

10.2

Fluidized-Bed Reactors1 Joachim Werther∗

10.2.1

Introduction The Fluidization Principle In fluidization, an initially stationary bed of solid particles is brought to a ‘‘fluidized’’ state by an upward stream of gas or liquid as soon as the volume flow rate of the fluid 10.2.1.1

1 A list of abbreviations/acronyms used in the text is provided at the end of the chapter. ∗ Corresponding author.

10.2.1 Introduction

exceeds a certain limiting value V˙mf (where mf denotes minimum fluidization). In the fluidized bed, the particles are held suspended by the fluid stream; the pressure drop pfb of the fluid on passing through the fluidized bed is equal to the weight of the solids minus the buoyancy, divided by the cross-sectional area At , of the fluidized-bed vessel (Fig. 1): pfb =

At · H · (1 − ε) · (ρs − ρ f ) · g At

granulation, to many heterogeneous catalytic gas-phase reactions as well as non-catalytic reactions. In this chapter, we survey the fluid-mechanical principles of fluidization technology, gas and solid mixing, gas–solid contact in the fluidized bed, typical industrial applications, and approaches to the modeling of fluidizedbed catalytic reactors. Further information is provided in textbooks [2] and monographs [3–8], while summary treatments are provided in Refs. [9–14]. Other useful literature includes reports of the Engineering Foundation Conferences on Fluidization [15–17], and the Circulating Fluidized-Bed Conferences [18–20].

(1)

In Eq. (1), the porosity ε of the fluidized bed is the void volume of the fluidized bed (volume in interstices between grains, not including any pore volume in the interior of the particles) divided by the total bed volume; ρs is the solids apparent density; and H is the height of the fluidized bed. In many respects, the fluidized bed behaves as a liquid. The bed can be stirred; objects of greater specific gravity sink, whereas those of lower specific gravity float; if the vessel is tilted, the bed surface resumes a horizontal position; if two adjacent fluidized beds with different bed heights are connected to each other, the heights become equal; and the fluidized bed flows out like a liquid through a lateral opening. Particularly advantageous features of the fluidized bed for use as a reactor are excellent gas–solid contact in the bed, good gas–particle heat and mass transfer, and high bed–wall and bed–internals heattransfer coefficients. The fluidization principle was first used on an industrial scale in 1922 for the gasification of fine-grained coal [1]. Since then, fluidized beds have been applied in many industrially important processes. The present spectrum of applications extends from a number of physical processes, such as cooling–heating, drying, sublimation–desublimation, adsorption–desorption, coating, and

Forms of Fluidized Beds As the volume flow rate V˙ or the superficial velocity u = V˙ /At , of the fluid increases beyond the value V˙mf or umf (Fig. 2a) corresponding to the minimum fluidization point, one of two things happens: (i) in fluidization with a liquid, the bed begins to expand uniformly; or (ii) in fluidization with a gas (a process which is of greater industrial importance and is discussed almost exclusively in the following sections), virtually solids-free gas bubbles begin to form (Fig. 2b). The local mean bubble size increases rapidly with increasing height above the grid because of coalescence of the bubbles. If the bed vessel is sufficiently narrow and high, the bubbles ultimately fill the entire crosssection and pass through the bed as a series of gas slugs (Fig. 2c). As the gas velocity increases further, increasing amounts of solids are carried out of the bed, the original, sharply defined surface of the bed disappears, and the solids concentration begins to decrease continuously with increasing height. In order to achieve steady-state 10.2.1.2

References see page 2129

∆P

∆Pfb

∆P

0 0

Vmf

V Packed bed

Fig. 1

Pressure drop in flow through packed and fluidized beds.

2107

V Fluidized bed

2108

10.2 Fluidized-Bed Reactors

u = umf (a) Fig. 2

u > umf (b)

u > umf (c)

u >> umf (d)

u >> umf (e)

Forms of gas-solids fluidized beds. See text for details.

operation of such a turbulent fluidized bed (Fig. 2d), solids entrained in the fluidizing gas must be collected and returned to the bed. The simplest way to do this is with a cyclone integrated into the bed vessel and a standpipe dipping into the bed. A further increase in gas velocity finally leads to the circulating fluidized bed (Fig. 2e), which is characterized by a much lower average solids concentration than the previous systems. The high solids entrainment requires an efficient external solids recycle system with a specially designed pressure seal (shown as a siphon in Fig. 2e). Advantages and Disadvantages of the Fluidized-Bed Reactor The major advantages of the (gas–solid) fluidized-bed as a reaction system include: 10.2.1.3

• Easy handling and transport of solids due to liquid-like behavior of the fluidized-bed • Uniform temperature distribution due to intensive solids mixing (no hot spots even with strongly exothermic reactions) • Large solid–gas exchange area by virtue of small solids grain size • High heat-transfer coefficients between bed and immersed heating or cooling surfaces • Uniform (solid) product in batchwise process because of intensive solids mixing. Set against these advantages are the following disadvantages: • Expensive solids separation or gas purification equipment required because of solids entrainment by fluidizing gas • As a consequence of a high solids mixing rate, backmixing of gas and a resultant lower conversion

• In catalytic reactions, undesired bypass or broadening of residence-time distribution for reaction gas due to bubble development • Erosion of internals and attrition of solids (especially significant with catalysts), resulting from high solids velocities • Possibility of defluidization due to agglomeration of solids • Gas–solid countercurrent motion possible only in multistage equipment • Difficulty in scaling-up. A comparison between fluidized-bed reactors with alternative gas–solid reaction systems (i.e., fixed-bed, moving-bed) and entrained-flow reactors is provided in Table 1. 10.2.2

Fluid-Mechanical Principles Minimum Fluidization Velocity The minimum fluidization point, which marks the boundary between the fixed- and the fluidized-bed conditions, can be determined by measuring the pressure drop p across the bed as a function of volume flow rate V˙ (see Fig. 1). Measurements should always be performed with decreasing gas velocity, by starting in the fluidized condition. Only for very narrow particle-size distributions, however, does a sharply defined minimum fluidization point occur. The broad size distributions commonly encountered in practice exhibit a blurred range; conventionally, the minimum fluidization point is defined as the intersection of the extrapolated fixed-bed characteristic with the line of constant bed pressure drop typical of the fluidized bed (Fig. 1). 10.2.2.1

10.2.2 Fluid-Mechanical Principles

Tab. 1

2109

Comparison of gas–solid reaction systems [2, 13]

Characteristics Suitability for heterogeneous catalytic gas-phase reactions

Temperature distribution

Heat supply and removal, heat exchange

Particle size

Fixed bed Only for catalyst that is deactivated very slowly

Moving bed

Fluidized bed

Entrained flow

Can also be used with catalyst that is rapidly deactivated

Catalyst attrition Catalyst attrition may be critical, depending on operating conditions negligible Plug flow of gas ensures high gas conversion Backmixing of gas due to Gas in virtually plug flow; mixing motion of solids high conversion possible and bubble-gas bypass lead to lower conversion Danger of hot spots Temperature gradients can High solids mixing ensures Temperature gradients can with exothermic be held within limits by uniform temperature be held within limits by reactions virtue of high solids distribution in bed; high solids circulation circulation and high gas temperature control by throughput heat exchangers immersed in bed or by admission and removal of solids Poor heat exchange; Poor heat exchange; due to Very efficient heat Properties intermediate heat transport limits high heat capacity of exchange; good heat between fluidized bed scale-up solids transport of large transport by solids and moving bed quantities of heat by way of circulating solids Large pellets Medium size (≈2–6 mm) Broad particle-size Fine (≈0.02–0.5 mm) with (≈8–20 mm), as and uniform; no fines distribution narrow particle-size uniform as possible; (≈0.02–6 mm); high distribution no fines fines content acceptable

The measurement technique already contains the possibility of calculating the minimum fluidization velocity umf : The pressure drop in flow through the polydisperse fixed bed at the point u = umf , given for example by the Ergun relationship [21], is set equal to the fluidized-bed pressure drop given by Eq. (1). From the Ergun relationship: p (1 − ε)2 1−ε η · u + 0.29 · Sv · 3 · ρ f u2 = 4.17 · Sv2 · h ε3 ε it follows umf = 7.14(1 − εmf )ν · Sv   3 εmf (ρ 1 − ρ ) · g s f ·  1 + 0.067 · · 3 − 1 (1 − εmf )2 ρ f ν2 Sv (2) Accordingly, to calculate umf , the characteristics of the gas (ρ f , η), the density ρs of the particles, the porosity εmf of the bed at minimum fluidization, and the volume-specific surface area Sv of the solids must be known. The specific

surface area defined by Sv =

surface area of all particles in the bed volume of all particles in the bed

(this takes into account only the external area, which governs hydraulic resistance, not the pore surface area as in porous catalysts) cannot be determined very exactly in practice. Hence, umf should not be calculated on the basis of the measured particle-size distribution of a representative sample of the bed solids; instead, it is better measured directly. Equation (2) can be employed advantageously to calculate umf in an industrial-scale process on the basis of minimum fluidization velocities measured in the laboratory under ambient conditions [22]. An equation from Wen and Yu [23] can be used for approximate calculations: 

Remf = 33.7 1 + 3.6 × 10−5 · Ar − 1 (3) where Remf =

umf dp ν

References see page 2129

(4)

2110

10.2 Fluidized-Bed Reactors

Ar =

gdp3 ν2

·

ρs − ρ f ρf

(5)

In the case of a particle-size distribution, the Sauter diameter calculated from the mass-density distribution q3 (d) of the particle diameters ds =  d max dmin

1 d −1

(6)

· q3 (d) · dd

should be used for the characteristic particle diameter dp . It should be noted that both the Ergun approach and the Wen and Yu simplification have been confirmed experimentally over a wide range of parameters. Recently, Vogt et al. [24] found that Eqs. (2) and (3) are also applicable to high-pressure fluidized beds with the fluid being under supercritical conditions. Fluidization Properties of Typical Bed Solids In fluidization with gases, solids display characteristic differences in behavior that can also affect the operating characteristics of fluidized-bed reactors. Geldart has proposed an empirically based classification of solids into four groups (A to D) by fluidization behavior [25]. The parameters employed are those crucial for fluidization properties: the mean particle diameter (dp ) and the density difference (ρs –ρ f ) between solid and fluid. Figure 3 shows the Geldart diagram with the interclass boundaries theoretically established by Molerus [26]. Solids of group C are very fine-grained, cohesive powders (e.g., flour, fines from cyclones and electrostatic filters) that virtually cannot be fluidized without fluidization aids. The adhesion forces between particles are stronger than the forces that the fluid can exert on the particles. Gas flow through the bed forms channels extending from the grid to the top of the bed, and the 10.2.2.2

pressure drop across the bed is lower than the value from Eq. (1). Fluidization properties can be improved by the use of mechanical equipment (agitators, vibrators) or flowability additives, such as Aerosil. Solids of group A have small particle diameters (ca. 0.1 mm) or low bulk densities; this class includes catalysts used, for example, in the fluidized-bed catalytic cracker. As the gas velocity u increases beyond the minimum fluidization point, the bed of such a solid first expands uniformly until bubble formation sets in at u = umb > umf . The bubbles grow by coalescence but break up again after passing a certain size. At a considerable height above the gas distributor grid, a dynamic equilibrium is formed between bubble growth and breakup. If the gas flow is cut off abruptly, the gas storage capacity of the fluidized suspension causes the bed to collapse rather slowly. Group B solids have moderate particle sizes and densities. Typical representatives of this group are sands with mean particle diameters between 0.06 mm and 0.5 mm. Bubble formation begins immediately above the minimum fluidization point. The bubbles grow by coalescence, and growth is not limited by bubble splitting. When the gas flow is cut off abruptly, the bed collapses quickly. Group D includes solids with large particle diameters or high bulk densities; examples are sands with average particle diameters greater than 0.5 mm. Bubbles begin to form just above the minimum fluidization point, but the character of bubble flow is markedly different from that in group B solids: group D solids are characterized by the formation of ‘‘slow’’ bubbles. On sudden stoppage of the gas flow, the bed also collapses suddenly. Gas Distribution The gas distribution devices must satisfy the following requirements: 10.2.2.3

10

(rs - rf) / (g·cm−3)

D

1 A

B

C 0.1 10

100

1000

dp / µm Geldart diagram (boundaries according to Molerus [26]). See text for details.

Fig. 3

• Ensure uniform fluidization over the entire crosssection of the bed (especially important for shallow beds) • Provide complete fluidization of the bed without dead spots where, for example, deposits can form • Maintain a constant pressure drop over long operation periods (outlet holes must not become clogged) Often, the gas distributor design must also prevent solids from raining through the grid both during operation and after the bed has been shut off. Porous plates of glass, ceramics, metal, or plastic are commonly used as gas distributors in laboratory apparatus; a variety of designs are used in pilot plant and full-scale fluidized-bed reactors

10.2.2 Fluid-Mechanical Principles

following correlations for vertical gas jets [31]:  

0.2
 u20 L ρ f d0 0.3  1.3 = 5.2 − 1 d0 ρs dp gd0

2111

(7)

and for horizontal jets [32]: 0.4

  ρ0 u20 L ρ f 0.2 dp 0.2 = 5.25 − 4.5 d0 (1 − ε)ρs gdp ρs d0 (a)

(b)

(c)

Industrial gas distributors. (a) Perforated plate; (b) nozzle plate; (c) bubble-cap plate.

Fig. 4

(see Fig. 4). Many more designs can be found, for example, in Ref. [2]. The principal requirement – that of a uniform distribution of fluidizing gas over the bed cross-section – can be met if the pressure drop pd across the gas distribution grid is large enough. Suggested values for the ratio pd : pfb are 0.1 to 0.3 (with a minimum pd of 3.5 kPa) [27], 0.2–0.4 [28], >0.15 [2], and >0.3 [29]. In the operation of fluidized-bed reactors, the quadratic response (pd ∝ u2 ) of industrial gas distributor designs must be borne in mind, because even if the fluidization velocity is lowered only slightly, an unacceptably low pressure drop across the gas distributor may occur. Industrial experience with different distributor designs, practical design rules and a discussion of distributorrelated problems, as for example weepage into the windbox and erosion by grid jets and at grid nozzles has been compiled in Ref. [30]. Gas Jets in Fluidized Beds Gas jets can form at the outlet openings of industrial gas distributors, and also where gaseous reactants are admitted directly into the fluidized bed. A knowledge of the geometry of such jets, in particular the depth of penetration, is important for the implementation of chemical operations in fluidized-bed reactors, and not just from the standpoint of reaction engineering. It is also vital for reasons of design: the strongly erosive action of these jets means that internals, such as heat-exchanger tubes, must not be located within their range. The literature contains many empirical correlations for estimating the mean depth of jet penetration [2–4]; these must, however, be used with care and, whenever possible, only within the range of parameter values for which they were derived. By way of example, Merry provides the 10.2.2.4

(8) Here, d0 is the diameter of the outlet opening, u0 is the outflow velocity, and ρ0 is the density of the jet gas. Donald et al. [33] investigated the penetration of single and multiple horizontal jets into a fluidized bed of FFC catalyst. These authors found that Eq. (8) was best able to predict the jet penetration for the range of conditions investigated (orifice diameter 4.4 to 9.2 mm, jet velocity 50 to 170 m s−1 ). A jet is affected more by another jet below it than by a neighboring jet at the same horizontal level or by one above it. Bubble Development For many applications – especially physical operations and non-catalytic reactions – the state of a fluidized bed can adequately be described in terms of a single quantity averaged over the entire bed, such as the mean bed porosity, ε. In contrast, the design of fluidized-bed catalytic reactors requires that local fluid-flow conditions are also taken into account. The local fluid mechanics of gas–solid fluidized beds are determined by the existence of bubbles, which influence the performance of fluidized-bed equipment in several ways: the stirring action and convective solids transport by the rising bubbles are helpful; the resulting intensive solids motion produces a uniform temperature throughout the fluidized bed and rapid heat transfer between the bed and the heating or cooling tubes submerged in it. The bubbles and the motion of solids that they cause, however, also have some drawbacks: attrition of solid particles, erosion of internals, and increased solids entrainment by bubbles bursting at the bed surface. The existence of bubbles is particularly detrimental in the case of a heterogeneous catalytic gas-phase reaction, because the bypass of reactant gas in the bubble phase limits the conversion achieved in the fluidized bed. The ultimate cause of bubble formation is the universal tendency of gas–solid flows to segregate. Many studies on the theory of stability [3, 4] have shown that disturbances induced in an initially homogeneous gas–solid suspension do not decay but always lead to the formation of voids. 10.2.2.5

References see page 2129

2112

10.2 Fluidized-Bed Reactors

The bubbles formed in this way exhibit a characteristic flow pattern, the basic properties of which can be calculated with the model of Davidson and Harrison [34]. Figure 5 shows the streamlines of the gas flow relative to a bubble rising in a fluidized bed at minimum fluidization conditions (ε = εmf ). The characteristic parameter is the ratio of the upward velocity ub of the bubble to the interstitial velocity umf /εmf of the gas in the suspension surrounding the bubble. A value of this ratio > 1 is typical for solids of Geldart groups A and B. The gas rising in the bubble flows downward again in a thin layer of suspension (‘‘cloud’’) surrounding the bubble. An important point for heterogeneous catalytic gas-phase reactions is that the presence of a boundary between bubble gas and suspension gas leads to the existence of two distinct phases (bubble phase and suspension phase) with drastically different gas–solid contact. If the velocity ratio is below 1, some of the gas in the suspension phase undergoes short-circuit flow through the bubble, while only part of the bubble gas recirculates through the suspension. This type of flow is typical for fluidized beds of coarse particles (Geldart group D). Under the real operating conditions of a fluidized-bed reactor, a number of interacting bubbles occur in the interior of the fluidized bed. As a rule, the interaction leads to coalescence. As detailed studies have shown, this process is quite different from that between gas bubbles in liquids because of the absence of surface-tension effects in the fluidized bed [35, 36]. The effects of both bubble growth and breakup are described in the case of Geldart group A and B solids by [37]: d dv = dh




2εb 9π

1 3

−

dv 3λub

(9)

ub umf / emf

Fig. 5

with the following boundary condition at h = h0 :  1/3 0.008 × εb for a porous plate  

dv0 0.2 = V˙02  m for an industrial gas distributor 1.3 g (10) The local volume fraction of bubble gas εb is given by εb = V˙b /ub

(11)

and the visible bubble flow V˙b is V˙b ≈ 0.8(u − umf )

(12)

The upward velocity ub of bubbles depends not only on the bubble size but also on the diameter dt of the fluidized bed: (13) ub = V˙b + 0.71 × α × gdv where α=



3.2 · dt0.33

0.05 ≤ dt ≤ 1 m, Geldart group A

2.0 · dt0.5

0.1 ≤ dt ≤ 1 m, Geldart group B

(14) Outside these limits, α is taken as constant. The differential equation [Eq. (9)] describes not only bubble growth by coalescence but also the breakup of bubbles (second term on the right hand). The crucial parameter here is the mean bubble lifetime λ: λ ≈ 280 ×

umf g

(15)

In practice, bubble growth is limited not only by the splitting mechanism based on the particle-size

1

10.2.2 Fluid-Mechanical Principles

distribution of the bed solids, but also by internals (screens, tube bundles, and the like) that causes bubbles to break up. Computational techniques for estimating this process are given in Refs. [38–40]. An example of measured and calculated bubble-growth curve is presented in Fig. 16. Elutriation When bubbles burst at the surface of the fluidized bed, solid material carried along in their wake is ejected into the freeboard space above the bed. The solids are classified in the freeboard; particles with a settling velocity ut greater than the gas velocity fall back into the bed, whereas particles with ut < u are elutriated by the gas stream. As a result, both the volume concentration of solids cv and the mass flow rate of entrained solids in the freeboard show a characteristic exponential decay (Fig. 7). With 10.2.2.6

0.05

dv / m

0.02

umf = 0.0023 m·s−1; u = 0.1 m·s−1

0.01

dp = 61×10−6 m; 0

0.1

0.2

dt = 1.0 m

0.3

0.4

(16)

Equation (16) was, however, derived for a bench-scale unit and may not scale to plant-size equipment. The mass flow rate Gs of entrained solids per unit area leaving the fluidized-bed reactor is the sum of contributions from the entrainable particle size fractions

0.03

0

increasing height above the bed surface, the transport disengaging height (TDH) is finally reached. Here, the increased local gas velocities due to bubble eruptions have decayed, and the gas stream contains only particles with ut < u. When the TDH can be reached in a fluidizedbed reactor, this is associated with minimum entrained mass flow rates and solids concentrations, and hence with minimum loading on downstream dust collection equipment. Design of the dust collection system requires a knowledge of the entrained mass flow rate Gs , and the particle-size distribution of the entrained solids. For the design of the fluidized-bed reactor, the distribution cv (h) of the solids volume concentration as a function of height in the freeboard must be known. For solids of Geldart group A, the TDH can be estimated with the diagram shown in Fig. 8 [41]. The following relationship is given for the TDH of Geldart group B solids as a function of the size dv of bubbles bursting at the bed surface [42]: T DH = 18.2 × dv

0.04

2113

10 0.5

h/m 7.5 3 1.5

Bubble growth in a fluidized bed of fine particles (Geldart group A; data points from Ref. [40]; calculation from Ref. [37]).

Fig. 6

0.6 0.3

1.0

5

TDH / m

0.1

Hf

75

0.0

TDH 5m

h

0.1 0.05

H

d t= 0.05

0.02

0.1

1 u –umb

0

Cv

Schematic drawing of fluidized bed and freeboard. TDH, transport disengaging height.

5

/ m·s−1

Estimation of transport disengaging height (TDH), according to Ref. [41]. umb = fluidization velocity at which bubble development begins.

Fig. 8

Fig. 7

References see page 2129

2114

10.2 Fluidized-Bed Reactors

(ut < u):  Gs = xi × χi∗

(17)

i

Here, xi is the mass fraction of particle-size fraction i in the bed material and χi∗ is the elutriation rate constant for this fraction. The literature contains a number of empirical correlations for estimating χi∗ [2–4]. More physically based are the elutriation models of Wen and Chen [43] and of Kunii and Levenspiel [2, 44], which enable not only calculation of the exiting mass flow rate but also estimation of the concentration versus height cv (h) in the freeboard. The most recent literature survey on the factors affecting elutriation and the available modeling tools is given in Ref. [45].

Upper part of riser Annulus

Core

Suspension

Circulating Fluidized Beds Whereas for bubbling fluidized beds the solids hold-up in the upper part of the reactor and the entrainment of catalyst are often negligible, these features become most important in the case of circulating fluidized beds. These systems are operated at gas velocities above the terminal settling velocity, ut , of a major fraction or even all of the catalyst particles used (1 ms−1
u
20 ms−1 ). As a result, the solids mass flow rates to be externally recirculated are high, up to values of more than 1000 kg m−2 s−1 . The flow structure within circulating fluidized beds is very complex and exhibits axial as well as horizontal non-uniformities, as shown in Fig. 9. Unless the solids hold-up is very low and the gas velocity very high, a dense (c¯v  0.15) bottom zone near the gas distributor exists, where bubble-like voids have been observed surrounded by a dense suspension. In the upper part of the riser the solids concentration is lower than near the distributor (c¯v
0.05). Although a number of modeling approaches for prediction of the axial solids–concentration profile have been developed [46–52], there is no generally accepted method available. Experimental studies [49–55] have shown that the circulating fluidized bed exhibits characteristic horizontal profiles in the upper part of the riser, with the concentration cv,wall near the vessel wall being always significantly higher than the value c¯v averaged over the vessel cross-section; for example, cv,wall = 2.3 · c¯v [56]. Local measurements of the solids concentration and solids velocity show that upward-flowing regions of low solids concentration and downward-flowing aggregates of high solids concentration alternate in time at every point inside the fluidized bed, with downward-moving aggregates (strands, clusters) predominating near the wall and upward-moving regions of low suspension concentration predominating in the central zone. 10.2.2.7

Bottom zone

Bubbles

Schematic diagram of flow structure in a circulating fluidized bed.

Fig. 9

Gas

Solids at rest

Fluidized Gas solids Gas (a) Fig. 10

Gas (b)

Design options for solids recycle. (a) Siphon; (b) L-valve.

Solids carried over with the fluidized gas are generally collected in cyclones. In the case of bubbling beds, the solids can easily be returned to the bed through the standpipe of the cyclone, which dips directly into the bed. Due to the large amounts of circulating solids, circulating fluidized beds require very large cyclones arranged beside and outside the bed, with special valves needed to connect the standpipe to the bed vessel. Figure 10 shows two design options, the siphon and the L-valve. Within the

10.2.2 Fluid-Mechanical Principles

siphon the solids are fluidized (i.e., enabled to flow back into the reactor). In the L-valve design, the mass flow rate of the solids can be controlled by varying the gas supplied to the standpipe. Because the solids path does not contain any sort of mechanical closure, the characteristic pressure distribution plotted in Fig. 11 is obtained. The distribution of solids between the fluidized bed and the recycle line is directly related to this pressure distribution. Operating properties differ from one recycle design to another [57]. Catalyst Attrition The attrition of solid particles is an unavoidable consequence of the intensive solids motion resulting 10.2.2.8

3

4 3

P0 4

h

a 5

b

5

2 1

2 1

0

P − P0

Fig. 11 Pressure distribution in solids recycle system of a circulating fluidized bed. (a) Fluidized bed; (b) return leg (the numbers indicate the location of pressure measurements).

+

q3

from the presence of bubbles in the fluidized bed. The attrition problem is especially critical in processes where the bed material needs to remain unaltered for the longest possible time, as in fluidized-bed reactors for heterogeneous catalytic gas-phase reactions. Catalyst attrition is important in the economics of such processes, and may even become the critical factor. Catalyst attrition occurs normally in fluidized-bed reactors as surface abrasion (Fig. 12), which means that surface asperities are abraded and edges of the catalyst particles are rounded off. However, fragmentation may also play a role, especially for some fresh catalyst particles which, upon entering the reactor, may simply be crushed into pieces. If in an industrial process extraordinarily high catalyst losses are observed it is advisable to examine catalyst samples under the scanning electron microscope. If the sample contains many fragments this could be an indication of a poor design (e.g., too-high velocities at the cyclone inlet or at the distributor). When designing catalytic fluidized-bed processes, the attrition performance of candidate catalysts should be tested under standardized conditions. This test can be performed in a small laboratory apparatus; it consists essentially of an extended fluidization test in which the mass of solids carried out of the bed is recorded as a function of time. Figure 13 presents a typical test result: during the first hours of testing, an increased attrition rate is observed which is due to breaking off of weak agglomerates and sharp edges of the catalyst particles. Only after a relatively long operating period is a quasisteady state attained. The attrition rate ratt in this steady References see page 2129

After abrasion Initial distribution

Abrasion

dp Fragmentation After fragmentation

q3

Initial distribution

dp Fig. 12

2115

Attrition modes and their effects on the particle-size distribution (q3 = mass density distribution of particle sizes, dp ).

2116

10.2 Fluidized-Bed Reactors

Mass elutriated matt / kg attrition rate ratt / kg·s−1

∆t

∆matt

matt

matt (t)

ratt

Time t / s

V1 m b (t )

Fig. 13

Result of an attrition test.

state can be defined as: ratt =

1 matt × mb t

(18)

where matt is the attrited mass and mb the bed mass. Usually ratt is expressed as percentage per day; for relatively attrition-resistant, fluidized-bed catalysts, it is of the order of 0.2% per day [5]. Many standard test apparatuses have been proposed for comparative attrition tests [58–60], but all such equipment has been suitable only for comparative studies of different catalysts under consideration for the same process. The attrition measured in large-scale equipment can be far different from the values measured in a test apparatus. A number of sources can be identified for catalyst attrition in industrial fluidized-bed reactors: • Jet attrition at gas distribution grid openings and nozzles where gaseous reactants are admitted to the bed • Bubble attrition in the bed due to solids motion caused by bubbles • Attrition in cyclones • Attrition in pneumatic conveyor lines, such as those between reactor and regenerator beds. Empirical correlations are available for the attriting action of a gas jet in the fluidized bed [61] and for the size reduction effect of solids motion in the bed [62, 63]. Werther and Xi [64, 65] employed the laboratory apparatus shown schematically in Fig. 14, which enables separate study of the attrition due to jets from nozzles of various diameters and that due to bubbles. The jet attrition-related mass production of fines per unit time for a gas distributor with a number no of orifices from mother particles with the diameter dp,i which are

V2 Fig. 14

Experimental apparatus for attrition test.

present in the catalyst inventory with a mass fraction Q3i is proportional to the particle size dpi , the mass fraction Q3i , the density ρo of the gas issuing from the orifice, the square of the orifice diameter do , and to the cube of the jet exit velocity uo [66, 67]: m ˙ att,jet,i = Cj · no · dpi · Q3i · ρ0 · d02 · u30

(19)

Attrition due to the bubble-induced solids movement is given by [65, 67]: m ˙ att, bubble, i = Cb · dpi · Q3i · mb · (u − umf )3

(20)

where mb denotes the bed mass which contains bubbles (i.e., which is located outside the jet-dominated grid region). Equation (20) denotes also the mass production of attrited fines which is due to the size fraction dpi in the bed. The stress on the catalyst particles will be different in contact with a gas jet, in the bulk of the bubbling fluidized bed, and during its passage through a cyclone. Recent investigations of the cyclone-induced catalyst attrition [68, 69] have shown that the mass flow of attrited fines which is produced by attrition inside the cyclone when a solids mass flow m ˙ c · Q3ci of particles of the size fraction dpi enters the cyclone is given by u2 m ˙ att,c,i = Cc · m ˙ c · Q3ci · dpi · √ c µc

(21)

10.2.3 Gas Mixing in Fluidized-Bed Reactors

where uc is the gas velocity at the cyclone inlet and µc is the solids loading of the incoming gas flow, µc =

10000.0

(22)

matt (J) matt,∞

where ρc is the density of the inflowing gas and Ac is the inlet cross-sectional area of the cyclone. The above Eqs. (19–22) describe the catalyst attrition under conditions of steady state – that is, when the particles are more or less rounded off (see Fig. 11). In order to describe also the initial breakage and attrition of fresh catalyst particles it is necessary to follow the fate of the particles upon their introduction into the reactor which is possible with population balance models (cf. Section 10.2.7.3). Klett et al. [70] have defined a stress history parameter ϑ,  t  for jet–induced attrition  ∗    tj   t for in–bed attrition ϑ= (23)  tb∗    np    ∗ for attrition in cyclones np where the definition of the characteristic parameters tj∗ , tb∗ and n∗p may be taken from Fig. 15. np denotes the number of passages of a given particle through the cyclone. tb and tj indicate the time period during which the particle is subjected to bubble and jet stress, respectively. If it is assumed that the effects of the different stress mechanisms on the catalyst particles are additive, then a uniform treatment of the overall stress history for all three attrition mechanisms is given by  m ˙ att (ϑ) 1.1 · ϑ b ϑ ≤ 1.1−1/b = (24) 1 ϑ > 1.1−1/b m ˙ att,∞ The parameter ϑ is characteristic of a given catalyst. Figure 16 shows measurements with an FCC catalyst [71]

Bubble-induced attrition Attrition in cyclones Jet-induced attrition

b = −1.16

1000.0

m ˙c ρc · uc · Ac

2117

100.0 10.0 1.0 0.1 0.001

0.01

0.1

1

10

J/− Fig. 16 Dimensionless attrition rate of an FCC catalyst as function of stress history. See text for details.

which lead to b = −1.16. Equation (24) allows the description of the stress history-dependent attrition, and can be used for the simulation of fluidized-bed reactors (cf. Section 10.2.7.3). A variety of approaches exists for reducing attrition in industrial fluidized-bed reactors. The jet attrition action can be controlled with special gas distributor designs [5, 72], for example by the use of bubble caps (see Fig. 4) such that gas jets do not issue directly into the bed at high velocity. Attrition due to bubbles can be lowered by limiting bubble growth [avoiding high gas velocities and large bed heights; use of fine catalysts with low umf , as implied by Eqs. (9) and (15)]. Attrition in cyclones can be prevented, in the simplest case, by replacing the cyclones with devices such as filters. Attrition can also be minimized by cutting back the load on the cyclone, for example, by placing the cyclones above the TDH. Relatively high catalyst attrition also occurs in circulating fluidized beds where very large quantities of solids must be collected in the cyclones. 10.2.3

Gas Mixing in Fluidized-Bed Reactors matt

The mixing and residence-time distribution of the gas are particularly important for catalytic reactions, but are also significant for gas–solid reactions when gaseous reactants are to be converted to the greatest possible extent in fluidized beds. Gas mixing is closely linked to the motion and mixing of the solids in the bed.

1.1·matt, ∞ matt, ∞

Gas Mixing in Bubbling Fluidized Beds If the flow and mixing of gas in the bubbling fluidized bed are described by a simple one-phase dispersion model, the coefficients Dgv and Dgh of gas dispersion in the vertical and horizontal directions follow trends similar to those of 10.2.3.1

t *b, t *j, n*p

t b, t j, np

Fig. 15 Dependence of attrition on time and number of passages through a cyclone, respectively. See text for details.

References see page 2129

2118

10.2 Fluidized-Bed Reactors

Dgv·u −1/m

10

1 0.5

0.1

0.05 0.1

0.5

1

5

dt /m Fig. 17 Vertical gas dispersion in a fluidized bed of solids of Geldart group A (measurements by various authors [2]).

the solids dispersion coefficients. By way of example, Fig. 17 shows the effect of fluidized-bed diameter dt on vertical gas dispersion. The increase in dispersion coefficient with vessel diameter might be attributable to the formation of large-scale solids circulation patterns, which becomes more marked in larger equipment. The coefficients of horizontal gas dispersion are a factor of 10 to 100 lower than those of vertical gas dispersion. Many studies show that a single-phase dispersion model gives only a rough description of gas mixing in bubbling fluidized beds (e.g. Ref. [73]). A more exact description comes from models that take account of local flow conditions in the bed, especially the presence of bubbles (see Section 10.2.7). A recent study has shed light on the mechanisms of gas mixing in the splash zone above a bubbling fluidized bed. Solimene et al. [74] found that isolated bubbles gave rise, upon bursting, to toroidal vortex rings which grew by entrainment of gas from the mainstream. However, with multiple bubbles bursting at the bed surface strong interaction between the bubble eruptions was observed which indicates the effectiveness of the splash zone for gas and gas–solid mixing, and thus also for the performance of the fluidized-bed reactor in the case of heterogeneously catalyzed reactions. Gas Mixing in Circulating Fluidized Beds Few detailed studies of gas mixing in circulating fluidized beds have been published [53, 75–80]. The bubbles in a bubbling fluidized bed influence the gas residence-time distribution and mixing directly through the bypass action of the bubble-gas flow and gas exchange between the bubbles and the surrounding suspension phase, and also indirectly through the solids motion that they induce. In the circulating fluidized bed, 10.2.3.2

however, the gas-mixing properties are controlled by segregation due to the formation of solid aggregates (strands, clusters) and the rapid downward movement of solids strands predominantly near the wall. Brereton et al. [76], for example, show that a single-phase dispersion model cannot describe the tracer gas residence-time distributions that they measured. Instead, these authors propose a two-phase model featuring exchange between a wall zone with stagnant gas and a core zone with plug flow. For the case of horizontal gas mixing, Werther and coworkers [79, 80] have shown that, for the bed solids they used (quartz sand, dp = 0.13 mm, Geldart group B), horizontal gas mixing in the top part of the circulating fluidized bed in the core zone can be described by the model for gas dispersion in turbulent single-phase flow [81]. The Peclet number P er,c =

uc × 2R ∗ Dr,c

(25)

has a value of 465, which is in fairly good agreement with values measured in single-phase flows (when defined in terms of the superficial velocity in the core zone uc , the radius of the core zone R ∗ and the horizontal dispersion coefficient in the core zone Dr,c ) [82]. This value is independent of the solids circulation rate Gs . The circulating fluidized bed thus exhibits no especially intensive horizontal gas mixing, at least not in the upper section where solids concentrations are relatively low. 10.2.4

Gas–Solid Separation

The fluidizing gas inevitably carries fine catalyst particles by entrainment to the reactor exit, and it is not only for environmental reasons (i.e., to minimize emissions) that the solids must be separated from the gas. Such separation may also be necessary to stop the main reaction and to avoid unwanted side or consecutive reactions, or to protect following process steps or machines from particleladen streams. In fluidized-bed technology the cyclone is mostly used for this purpose. Knowlton [83] has provided an excellent survey on the state of the art of cyclone design and application in fluidized-bed reactors. It must be noted that the cyclone should not be considered as a separate apparatus following the fluidized bed but rather be seen as an integral part of the fluidbed process. The reason is that not only in circulating fluidized beds, but also in bubbling or turbulent fluidized beds, the catalyst particles recovered in the cyclone are recycled to the fluidized bed. The collection efficiency of the cyclone will thus be responsible for maintaining the particle size distribution in the bed inventory, which in

10.2.6 Industrial Applications

turn determines the fluidized dynamics and the chemical performance of the bed as a reactor. The interrelation between fluidized bed and cyclone is discussed in Section 10.2.7.3. The influence of cyclone performance on the overall process performance is increasingly considered. For example, Pulupula et al. [84] investigated the role of cyclones in the regenerator system of a commercial FCC unit. Arnold et al. [85] were able to trace the deterioration of plant performance in the ALMA maleic process back to problems with cyclone efficiency. A change of cyclone design improved the particle size distribution of the bed inventory and, consequently, bed hydrodynamics and chemical conversion. Smit et al. [86] reported on cyclone performance in the turbulent fluidized-bed Synthol reactors for Fischer–Tropsch synthesis. Carbon deposition on the catalyst particles was seen to influence the bed hydrodynamics which in turn, via the elutriation mechanism, influenced cyclone performance.

2119

smaller hot catalyst particle was recently investigated in a three-dimensional direct numerical simulation to analyze droplet–particle collisions in the Leidenfrost regime [95]. The calculations were carried out for conditions prevailing near the feed nozzle in an FCC riser. Vapor layer pressure induced by evaporation and the droplet surface tension are the driving forces for droplet recoiling and rebounding. The contact time for an FCC particle and an oil droplet was found to be about 140 µs. 10.2.6

Industrial Applications

In this section the industrial uses of fluidized-bed reactors for heterogeneous catalytic gas-phase reactions and the polymerization of alkenes are described. The most important applications are listed, and some typical examples analyzed in more detail. Complete descriptions of industrial uses of the fluidized-bed reactor can be found in Refs. [2, 6, 13, 14].

10.2.5

Injection of Liquid Reactants into Fluidized Beds

The injection of reactants in the liquid form into the bed is an essential part already of the first fluidized-bed catalytic process. In the FCC process (cf. Section 10.2.6 and Chapter 13.5), a heavy oil fraction is injected at the base of the reactor and evaporated in contact with the hot catalyst particles. The direct heat transfer is very efficient and avoids the need for a separate evaporator for the feed. The cooling action of the evaporating reactant is a further advantage in the case of an exothermal reaction. Liquid feed injection is therefore practiced not only in the FCC process but also, for example, in the syntheses of aniline [87], caprolactam [88] and melamine [89] and in BP Chemicals’ INOVENE process [90] for the gas-phase production of low-density polyethylene. Despite its industrial significance, our knowledge of the mechanisms of liquid mixing and evaporation in the fluidized bed is relatively scarce. Investigations with nonvaporizing horizontal gas–liquid spray jets have shown that, with a proper design of the injection nozzle, it is possible to penetrate over several decimeters into the bed before the jet breaks up [91, 92]. On the other hand it was found that, under vaporizing conditions for atomizer nozzles with spray angles between 20 and 120 degrees, the injected liquid wetted the bed particles and subsequently evaporated from the surface of the particles while the particles were mixing in the bulk of the bed [93, 94]. This latter mechanism helps to transport the reactant away from the location of the nozzle and thus contributes to an equalization of the feed distribution inside the reactor. The special case that a large oil droplet impinges on a

Heterogeneous Catalytic Gas-Phase Reactions The fluidized-bed reactor offers the following principal advantages over the fixed-bed reactor for heterogeneous catalytic gas-phase reactions: 10.2.6.1

• High-temperature homogeneity, even with strongly exothermic reactions • Easy solids handling, permitting continuous withdrawal of spent catalyst and addition of fresh if the catalyst rapidly loses its activity • Ability to operate in the explosion range, provided that the reactants are not mixed until they are inlet to the fluidized bed. This is because the high heat capacity of the bed solids, together with intensive solids mixing, prevents the propagation of explosions. The ease of solids handling was the basic reason for the success of catalytic cracking of long-chain hydrocarbons in the fluidized bed. The cracking reaction is endothermic and involves the deposition of coke on the catalyst surface, which quickly renders the catalyst inactive. Accordingly, the catalyst must be discharged continuously from the reactor and regenerated in an air-fluidized regenerator bed, where its coke loading is lowered from 1–2 wt.% to 0.4–0.8 wt.%. The combustion in this bed simultaneously furnishes the heat required for the cracking reactor; the catalyst acts as a heat carrier. The temperature in the regenerator is 650–750 ◦ C and in the reactor 480–540 ◦ C [2]. In a stripper, steam is admitted to remove hydrocarbons adhering to the catalyst before it is forwarded to the regenerator. References see page 2129

2120

10.2 Fluidized-Bed Reactors

With the advent of high-activity zeolite catalysts during the 1960s, the bubbling fluidized bed, operated at gas velocities between 0.31 and 0.76 ms−1 [2], was replaced by the riser cracker (Fig. 18), in which the oil fed in at the bottom of the riser (c) is vaporized in contact with the hot catalyst and the mixture of oil vapors and cracking gas transports the catalyst up through the riser. In the reactor bed (a), solids are collected before passing through the stripper (b) to the regenerator (f). By virtue of the short contact time of the order of a few seconds and the narrow gas residence–time distribution, the high activity of the zeolite catalyst is optimally utilized and a higher gasoline yield is achieved [2, 6]. The crucial factor in the successful use of the fluidizedbed reactor for the synthesis of acrylonitrile by the ammoxidation of propene (Sohio process) was reliable control of this strongly exothermic reaction (Hr = −515 kJ mol−1 acrylonitrile):

Product gas

Steam a Water b

NH3, C3H6

C3 H6 + NH3 + 32 O2 −−−→ C3 H3 N + 3H2 O The reaction is carried out at a bed temperature of 400–500 ◦ C and with gas contact times of 1 to 15 s [96] or 5 to 20 s [2]. A schematic representation of the reactor is shown in Fig. 19. Air is fed to the bottom of the fluidizedbed vessel, while the reactants, ammonia and propene, are fed in through a separate distributor (b). Catalyst regeneration by carbon burnoff occurs in the space between the air distributor and the feed-gas distributor. The heat of reaction is removed by bundles of vertical

Product

a

Flue gas

f

b

e c

d Air Feed oil Fig. 18 Riser cracking process (UOP system) [2]. (a) Reactor; (b) stripper; (c) riser; (d) slide valve; (e) air grid; (f) regenerator.

Air

Synthesis of acrylonitrile (Sohio process) [2]. (a) Cooler with internals; (b) distributor.

Fig. 19

tubes (a) inside the bed (horizontal tubes are used in other designs [97]). The Fischer–Tropsch synthesis of hydrocarbons is used on a large scale for fuel production in South Africa [98–100]. Synthesis gas generated from coal in Lurgi fixed-bed gasifiers enters the Synthol reactor (Fig. 20), where it is reacted over an iron catalyst at ≈340 ◦ C. The reactor works on the principle of the circulating fluidized bed. The mean porosity in the riser is 85%, and the gas velocity varies between 3 and 12 m s−1 [2]. Reaction heat is removed by way of heatexchanger tube bundles placed inside the riser. However, experience has shown that this reactor is costly, relatively expensive to operate and maintain, and the scale-up to the size of the reactors in operation is probably close to the maximum achievable for 350 ◦ C and 25 × 105 Pa (25 bar) pressure operation. Therefore, the 16 circulating fluidized-bed reactors operating at Sasol’s Secunda site were, during the 1990s, replaced by eight turbulent fluidized-bed reactors of 10.7 m diameter, each of which achieves a higher per pass syngas conversion [101]. Different process routes have been developed for the synthesis of maleic anhydride. The Mitsubishi process [102, 103] uses the naphtha cracker C4 fraction, while the ALMA process uses n-butane as the feedstock [85, 104, 105]. A more recent development is the DuPont process,

10.2.6 Industrial Applications

2121

Product recovery

f

Gas out

Spent air d

Reactor

a

Seal Inert

e

Regenerator

Seal Butane Inert

d

Air

Recycle gas b

Fig. 21

DuPont’s maleic anhydride process.

Polymerization of Alkenes The gas-phase polymerization of ethylene in the fluidized bed was developed by Union Carbide (Unipol process [119]; see Fig. 22). For the production of linear low-density polyethylene (LLDPE), the reaction gas (ethylene with propene, 1-butene, and 1-hexene) fluidizes the bed at 75–100 ◦ C and ≈2 × 106 Pa. Extremely finegrained catalyst is metered into the bed, and polymerization then occurs on the catalyst surface to yield a granular product with a diameter ranging from 0.25 mm to 1 mm. As the ethylene conversion is comparatively low (i.e., ca. 2% per pass), the reaction gas is recycled. The heat of reaction is removed by cooling the recirculating gas. The catalysts used have such a high activity that more than 105 parts by volume of polymer can be produced per unit weight of active substance in the catalyst [2]. Because of the high degree of catalyst dilution in the granular polymer, the catalyst need not be removed from the product. Mitsui Petrochemical Industries has developed a process for the gas-phase fluidized-bed polymerization of propene; a plant using this process came on stream in 1984 [120]. The Unipol-Shell process was jointly developed by Union Carbide and Shell and commissioned in 1986. Burdett et al. [121] have provided a broad overview on this still-developing technology which contains many challenges for the engineer. One of the biggest problems is the stickiness of the particles under the operating conditions of the process, and this has often led to particle sintering with subsequent defluidization of the bed. Seville et al. [122] monitored the motion of particles in a scaled polymer reactor and studied the sintering 10.2.6.2

Reactant gas

c

Fischer–Tropsch synthesis in the Synthol reactor [2, 23]. (a) Hopper; (b) standpipe; (c) riser; (d) cooler (coil); (e) reactor; (f) gooseneck.

Fig. 20

which is also based on n-butane but uses a circulating fluidized bed as a reactor. This process, which is shown schematically in Fig. 21, is based on a vanadium phosphorus oxide (VPO) catalyst which oxidizes n-butane to maleic anhydride by a redox mechanism on its surface layers. In the riser, n-butane is selectively oxidized by the oxidized catalyst which thus loses the oxide. The spent catalyst is then reoxidized in the fluidized-bed regenerator [106, 107]. Since 1996, a commercial plant is operating in Asturias, Spain [108]. Other catalytic reactions carried out in fluidized-bed reactors include (partly on a commercial scale, partly on a pilot-plant scale) the oxidation of naphthalene to phthalic anhydride [2, 6, 109], the ammoxidation of isobutane to methacrylonitrile [2], the reaction of acetylene with acetic acid to vinyl acetate [2], the oxychlorination of ethylene to 1,2-dichloroethane [2, 6, 110, 111], the chlorination of methane [2], the reaction of phenol with methanol to cresol and 2,6-xylenol [2, 112], the reaction of methanol to gasoline [113, 114], the synthesis of phthalonitrile by ammoxidation of o-xylene [115], the synthesis of aniline by gas-phase hydrogenation of nitrobenzene [116], and the low-pressure synthesis of melamine from urea [117]. An excellent overview on the various fluidized-bed catalytic processes was recently provided by Jazayeri [118].

References see page 2129

2122

10.2 Fluidized-Bed Reactors

a d

b c e

Ethylene Comonomer

Before a reactor model found in the literature can be applied to a given problem, the designer must determine whether numerical values are available for all model parameters – that is, whether the model is appropriate for design calculations, or is a ‘‘learning model’’ [126] in which the numerical values of important parameters can be determined only after the model is adapted to actual test results. Reaction kinetics may be determined in a fixed-bed reactor, provided that the measurements are performed under conditions comparable to those that prevail in the fluidized-bed reactor (e.g., the same solids composition and particle-size distribution, and the same activity state) [127]. However, the kinetic parameters can also be determined directly by measurements in a bench-scale fluidized-bed apparatus [128]. Bubbling Fluidized-Bed Reactors By far the majority of fluidized-bed reactor models described in the literature deals with reactions in bubbling fluidized beds [2, 5, 6, 124, 125]. For a specific application, modeling depends on the bubble flow regime. For slow-bubble systems (see Fig. 5, left), the short-circuit flow of gas through the bubbles must be taken into account [129]. For fast-bubble systems (Fig. 5, right), the species must be balanced separately in the bubble and suspension phases. If models from the literature are employed, those devised in the past, when adequate computing hardware was not available, often sought to obtain closed analytical expressions for the degree of conversion of a single reaction (usually taken as first order). The simplifying assumption of a single ‘‘effective’’ bubble size for the entire fluidized bed was therefore made [2], or the mass transfer area between the bubble and suspension phases was taken as uniformly distributed over the height of the bed [130]. Today, in view of the computing power available at the PC level, the recommended procedure is to start from local masstransfer relations, to write balance equations for the differential volume element of the reactor, and then numerically to integrate. Figure 23 presents a model used for a constant-volume reaction [128, 131]. Here, the simplifying assumption is that flow through the suspension phase is at the minimum fluidization velocity, umf . For a heterogeneous catalytic gas-phase reaction, the material balances for species i in the unsteady-state cases are for the bubble phase, 10.2.7.1

Granular polyethylene

Gas-phase polymerization of ethylene (Unipol process) [2]. (a) Compressor; (b) cooler; (c) catalyst feed hopper; (d) reactor; (e) separator.

Fig. 22

kinetics in order to determine a safe operating window. Cai and Burdett [123] developed a model of single-particle polymerization in the fluidized bed to simulate the particle growth and particle temperature evolution with the residence time of a catalyst particle in the reactor. 10.2.7

Modeling of Fluidized-Bed Reactors

Exhaustive literature surveys of the modeling of fluidizedbed reactors are available [2, 5, 6, 124, 125], and the details of many models have been reported which are classified in the cited references under various schemes. The available information can be summed up as follows: No generally accepted model of the fluidized-bed reactor exists; instead, many models have been proposed on the basis of more-or-less extensive experimental findings for various applications. Any fluidized-bed reactor model can be broken down into separate components that describe, with varying degrees of accuracy, the hydrodynamics (depending on solid properties, operating conditions, and geometry), gas–solid contact, and reaction kinetics. The essential point is that the reactor geometry effect, which is important for scale-up (see Section 10.2.8), manifests itself in the flow conditions and must therefore be included in the hydrodynamic part of the model.

εb

δCbi δCbi = −[u − umf (1 − εb )] · δt δh − kG,i · a · (Cbi − Cdi )

and, for the suspension phase

(26)

10.2.7 Modeling of Fluidized-Bed Reactors

]u –umf(1−eb)] At (Cb + (∂Cb / ∂h)dh)

umf(1−eb) At(Cd + (∂Cd /∂h)dh)

Bubble phase

Suspension phase

Cb

kG,i aAt dh (Cb − Cd)

Cd

(Plug flow)

(Plug flow)

[u –umf (1−eb)] AtCb

umf (1−eb) AtCd

dh

Fig. 23 Two-phase model of the fluidized-bed reactor operating in the bubbling-bed mode.

(1 − εb )[εmf + (1 − εmf )εi ] = −umf (1 − εb ) ·

δCdi δt

+ (1 − εb ) · (1 − εmf )ρs

M 

vij rj

(27)

j =1

In Eqs. (26) and (27) the following simplifying assumptions have been made: • Plug flow through the suspension phase at an interstitial velocity (umf /εmf ) • Bubble-phase in plug flow, bubbles are solids-free • Reaction in suspension phase only • Constant-volume reaction (Ref. [128] shows how to handle a change in number of moles) • Sorption effects neglected (see Ref. [131] for handling sorption). Here, εi is the porosity of the catalyst particles, a is the local mass-transfer area per unit of fluidized-bed volume, which for spherical bubbles can be calculated as a=

6εb dv

The freeboard space above the bubbling fluidized bed must be considered in the reactor model if the entrainment rate is high and the reactions in the freeboard are not quenched, for example, by cooling. Most fluidized-bed models include concentration profiles only for the vertical direction. This one-dimensional modeling is acceptable when the reactants are admitted uniformly over the bed cross-section. If, however, reactants are metered into the bed at individual feed points, then three-dimensional modeling may become necessary [133–135]. As a rule, the modeling of solids behavior in fluidizedbed reactors is based on that in stirred tanks, and temperature homogeneity is a virtually fundamental property of these systems. Circulating Fluidized-Bed Reactors Since the number of different applications of circulating fluidized beds for catalytic reactions on the commercial scale is still small, the development of models for these systems is at an early stage. However, throughout the past decade many promising approaches have been developed, the most important of which are described in an overview by Grace [136]. In the early days of circulating fluidized-bed reactor modeling, a very low axial dispersion and a laterally uniform flow structure was believed to characterize these systems; thus, simple plug-flow models were used [137]. This approach was found to oversimplify the behavior of circulating fluidized-bed reactors, because a significant amount of axial dispersion was observed. As a result, the plug-flow model has often been modified by adding a dispersion term to the balance equations. Axial dispersion coefficients have been determined by many authors, who measured the residence time distribution of tracer gases [138, 139]. Typical values of Peclet numbers, Peax , found in these studies are of the order of 10. By means of a model reaction it has recently been proved that, in many cases, circulating fluidized-bed reactors cannot be characterized solely by considering mixing phenomena [140]. Instead, the presence of mass-transfer limitations and bypassing was found to have a significant influence. In analogy to low-velocity fluidized beds, a detailed description of the local flow structure within the reaction volume must serve as a basis for appropriate reactor modeling. The highly non-uniform flow structure of circulating fluidized beds described in Section 10.2.2.7 has led to reactor models which separately deal with different axial zones. The bottom zone – if existing under the given operating conditions – can be described by models, the 10.2.7.2

δCdi + kG,i · a · (Cbi − Cdi ) δh

(28)

rj is the rate of partial reaction j per unit mass of catalyst, and νij is the stoichiometric number of species i in reaction j. The relationship  umf 4Di εmf ub (29) kG,i = + 3 πdv proposed by Sit and Grace [132] has proved useful for describing the mass-transfer coefficient kG,i associated with component i in mass transfer between the bubble and suspension phases, where Di is the molecular diffusion coefficient of species i.

2123

References see page 2129

2124

10.2 Fluidized-Bed Reactors

basic approaches of which were originally developed for modeling of bubbling fluidized beds as presented in Section 10.2.7.1 [141]. Modeling of the upper section of the circulating fluidized bed is in most cases based on a proper description of the heterogeneous core-annulus flow structure [142–145]. These state-of-the-art models are one-dimensional, and define two phases or zones which are present at every axial location: • A dense phase or annulus zone; this has high solids concentration, and gas stagnant or moving downwards • A dilute phase or core zone; this has low solids concentration, gas flowing quickly upward. Similar to the situation in bubbling fluidized beds, the two phases exchange gas with each other and are modeled by separate equations which are obtained from mass balances for each component in each phase. Just as for this model, all similar approaches found in the literature suffer from the problem that not all fluid mechanical variables can be precalculated on the basis of the operating conditions. Instead, reasonable estimations or measurements in cold flow models are used to obtain numerical values for many variables. It is, therefore, absolutely necessary to refer to the original literature in order to perform quantitative calculations. A common feature of all models for the upper part of circulating fluidized beds is the description of the mass exchange between the dense phase and the dilute phase. Analogously to low-velocity fluidized beds, the product of the local mass-transfer area, a, and the masstransfer coefficient, k, may be used for this purpose. Many different methods to determine values for these important variables have been reported, such as tracer-gas backmixing experiments [141], non-steady-state tracer gas experiments [146], model reactions [144], and theoretical calculations [143]. Similar to the bubbling fluidized-bed reactor, the solids behavior of the circulating fluidized-bed reactor can usually be described as completely mixed. This does not hold for riser reactors with very high gas velocities, as they are used in many FCC-units (u > 10 m s−1 ). Here, better modeling results will be obtained by assuming dispersed plug flow of solids [139]. As well as for bubbling fluidized beds, it can be assumed that circulating fluidized beds exhibit a high degree of temperature homogeneity even in the case of highly exothermic reactions. New Developments in Fluidized-Bed Reactor Modeling The models described above follow the ‘‘classical’’ chemical engineering approach, which replaces the complex particle–fluid interaction in the fluidized bed 10.2.7.3

by idealized configurations (plug flow, stirred tank, either overall valid or in regions) with mixing and mass-transfer coefficients describing the transport of matter. However, during recent years there has been a strong tendency also to model the fluid mechanics of fluidized-bed reactors from first principles. The problem of modeling in this area is that the particle–particle and particle–fluid interactions need to be considered on the particle scale, whereas the reactor performance must be described on a much larger scale which, typically, is of the order of several meters. This leads to computational difficulties and problems with available computing capacities. At present, there is no generally accepted computational fluid dynamics (CFD) model of the fluidized-bed reactor available, although rapid progress is currently being made in this area [147, 148]. One promising approach appears to be a multi-scale modeling strategy [149], where the idea is essentially that fundamental models which take into account the relevant details of fluid–particle (lattice Boltzmann model) and particle–particle (discrete particle model) interactions, are used to develop closure laws to feed continuum models which can then be used to simulate the flow structures on a larger scale. Figure 24 illustrates this approach which, ultimately, leads to the discrete bubble model that should be applicable to the large industrial scale of the bubbling fluidizedbed reactor. Although the multiscale methodology [150] requires further development, it provides the chance to arrive at more realistic fluidized-bed reactor models in the not-too-distant future. Another line of development which should be mentioned here is the modeling of fluidized-bed reactor systems. Whilst, previously, the isolated fluidized bed was modeled, the current focus is on the coupling between the fluidized bed and the cyclone for catalyst recovery and recycle [151], or even on the coupling between two fluid-bed reactors [152] – for example, reactor–regenerator systems as they are used in the FCC or the maleic anhydride processes. As an example, Fig. 25 shows a fluidized bed coupled with a cyclone, and its translation into the model system. Attrition leads to a loss of material from the system which, after some time, requires the addition of fresh catalyst (Fig. 26). A population balance model which considers the changes in the particle size intervals of the catalyst allows the change in the catalyst inventory to be followed with time. It can be seen that it takes several weeks for the system to reach a quasi-steady state. As a consequence of attrition and incomplete separation in the cyclone, the mean particle diameter in the bed increases with time, which leads to larger bubbles and a reduced area of mass transfer between bubbles and the surrounding suspension in the bed.

10.2.7 Modeling of Fluidized-Bed Reactors

2125

Larger geometry

Lattice Boltzmann model

Discrete particle model

Fluid – particle interaction

Particle – particle interaction

Continuum model

Particel – particle interaction; Bubble behavior

Discrete bubble model

Large-scale motion industrial size

Larger scale phenomena

Fig. 24

The multi-scale approach for computational fluid dynamics (CFD) modeling of fluidized-bed reactors (after Ref. [149]).

Product + catalyst

Gaseous product

Freeboard Transport (no reaction)

Cyclone Separation (no reaction)

Solids loss flux

Cylone attrition

Freeboard Fluidized bed

Cyclone

Bubbling bed (reaction)

Fresh catalyst

Catalyst feed

Bubble region Standpipe Jet region

Catalyst discharge

Reactant gas

Fig. 25

Return line Bubble attrition

Transport (no reaction)

Distributor region (reaction) Jet attrition

Reactant gas

Fluidized-bed reactor model system [152].

As a further consequence, the conversion rate of a simple first-order reaction falls off with time. Finally, Fig. 27 shows that improvements in the efficiency of the solids recovery system are able to increase the conversion

rate, and this again is in agreement with large-scale industrial experience [153, 154]. References see page 2129

2126

10.2 Fluidized-Bed Reactors

100

Sauter diameter /µm

Catalyst inventory / kg

15 000

14 500

14 000

13 500

13 000

80 60 40 20 0

0

100

200

300

400

500

600

Time / h

(a)

0

100

200

300

400

500

600

400

500

600

Time/h

(b) 1.0 0.9

Conversion X /-

Mass flow/kg·s−1

0.01 1E-3 Loss flux 1E-4

Bubble attrition Cyclone attrition

1E-5

0.8 0.7 0.6 0.5

Grid jet attrition 1E-6

0.4 0

100

200

300

500

600

Time / h

(c)

Fig. 26

400

(d)

100

200

300

Time/h

Reactor behavior as a function of operating time [152].

450 to 600 mm, which should allow a reliable scale-up [5]. Full-scale fluidized-bed reactors have diameters up to ≈10 m. As equipment size increases, characteristic changes take place in the flow regimes that can decisively affect reactor performance. These changes result either directly from the geometry, or indirectly from design changes made as the unit is enlarged. In particular, experience has shown that the following factors affect the performance of bubbling fluidized beds during scale-up [72, 155].

1.0 3-stage separation 2-stage separation High-efficiency cyclone

Conversion /-

0

3 -stage separation with multi-cyclone as primary stage

0.8

Multi-cyclone

Standard cycone 0.6

Fig. 27 Influence of the Sauter diameter on the chemical conversion of a simple first-order reaction [152].

Bed Diameter According to Eq. (13), the mean upward bubble velocity increases as the bed diameter dt increases. As a result, the bubbles have a shorter residence time in the bed and the exchange area between the bubble and suspension phases is smaller, so conversion is reduced [156].

10.2.8

10.2.8.2

10.2.8.1 50

60

70

80

90

Sauter diameter in bed/µm

Scale-Up

Typical diameters of bench-scale, fluidized-bed reactors are approximately 30 to 60 mm, and of pilot-scale units

Grid Design In the laboratory, porous plates are the preferred type of gas distributor because of the ease of working with them. Gas distribution becomes worse when these are replaced by industrial distributor designs. Thus, the exchange area

10.2.9 Conclusions

between the bubble and suspension phases is reduced, again with consequently lower conversion [156, 157]. Internals Whereas the laboratory fluidized bed is generally operated with no internals, plant equipment often must contain bundles of heat-exchanger tubes. Screens, baffles, or similar internals are frequently used to redisperse the bubble gas in industrial reactors. The mass-transfer area is thus increased relative to the fluidized bed without internals; the extra area can be utilized to partially offset the conversion-reducing effects of bed diameter and gas distributor [158]. 10.2.8.3

Catalyst Particle-Size Distribution Bubble growth is influenced by the proportion of fines in the particle-size distribution of the bed (usually measured as the weight fraction below 0.044 mm) or by the mean grain size dp [via umf ; Eq. (9)]. If the content of fines increases, the bubbles collapse sooner and the equilibrium bubble size becomes smaller, with a resultant greater bubble-suspension mass-transfer area. This effect generally is fully developed only in the plant-scale reactor, where bubbles can grow without the hindrance of vessel walls. Thus, in principle, the performance of catalytic fluidized-bed reactors can be controlled by modifying the catalyst particle-size distribution [153, 154]. The recommended content by weight of fines (8 mm) sizes [5]. Moreover, the gas bubbles show a distinct axial size distribution in the column, with small bubbles being present near the sparger up to the equilibrium height, above which large bubbles predominate. Above this equilibrium height, churn-turbulent flow is completely developed and the bubbles do not grow further in size due to coalescence. 10.3.3.3

Backmixing Backmixing in a slurry reactor is usually detrimental to reaction rate and product selectivity [70]. For a gas–liquid–solid multiphase system, separate residence time distribution (RTD) measurements are needed to evaluate the mixing characteristics for each phase, and many such methods are available for use in complex multiphase systems. Slurry reactors deviate considerably from the ideal cases of plug flow and perfect mixing, respectively. These deviations may be the result of non-uniform velocity profiles, short circuiting, bypassing and channeling, velocity fluctuations, reactor shape and internals, and backflow of fluid due to velocity differences between phases and agitation [70]. The extent of backmixing in each phase (gas, liquid and solid) is, in general, different 10.3.3.4

2137

10.3.3 Hydrodynamics of Slurry Reactors

liquid in the slurry by uG Dc 13F rG = EL (1 + 8F rG0.85 )

P eL =

(4)

where uG F rG = √ gDc

(5)

The data in Fig. 2 show that, in the case of coal and dried mineral ash, Eq. (4) gives reasonable predictions for small particles sizes. The behavior of suspended solid particles and their backmixing have been extensively studied also by Imafuku et al. [111]. The most important conclusion drawn from these studies was that, for small particle sizes in small-diameter columns, the longitudinal dispersion coefficient for the solids, ES , is the same as for the liquid and is given by Eq. (4), when replacing EL by ES . For large columns and particles of different diameter, Kato et al. [110] suggested a modification of Eq. (4): P eS =

13F rG (1 + N Repb F rG−0.8 ) uG Dc = ES (1 + 8F rG0.85 )

(6)

dp vt ρL µL

(7)

where Rep =

The authors reported that at Rep = 0.3 to 2.5, N = 0.009 and b = 1, whereas at Rep = 2.5 to 640, N = 0.023 and

0.045 0.04

×

×

0.035 0.03

EL / m2 s−1

and should be considered separately. It is essential to identify correctly the nature of the gas–liquid flow and to consider the role of the solids phase. For example, in a mechanically agitated slurry reactor, gas-phase mixing might approach plug flow, whereas, in the absence of solids, the gas phase is found to be backmixed [92, 93]. The backmixing characteristics in a slurry reactor are evaluated from the RTD of a tracer injected in the respective phase. The tracer is injected at one location or more in the system, and its concentration as a function of time at one or more downstream positions subsequently detected. The basic requirements for tracer experiments are discussed by Shah et al. [70]. Concentrations of a solid tracer can be measured using a capacitance probe. In principle, however, for solid (and sometimes also gas) phases a suitable radioactive tracer is convenient to use. If the proper safety precautions are enforced, the use of a radioactive tracer has one distinct advantage over other tracers in that the tracer-detecting devices can be placed externally (non-invasive technique). Such radioactive measurements have been performed extensively at Washington University in St. Louis, USA [2, 94–105]. In the past, the axial dispersion model (ADM) has been the most widely used technique to correlate RTDs in multiphase systems. The ADM characterizes the backmixing by a simple one-dimensional Fick’s lawtype equation in which the constant of proportionality is commonly known as the axial dispersion coefficient, E. The latter is expressed in dimensionless form as the Peclet (or Bodenstein) number (Pe = uG Lc /E). In a slurry reactor, Lc may be the column length or diameter, the bubble diameter, or the diameter of the solids. The value of the Peclet number characterizes the degree of backmixing; for example, if Pe = 0 then backmixing is complete, but for Pe = ∞ plug flow prevails. For a slurry reactor, the backmixing in each phase is considered separately because a considerably different degree of backmixing can exist in each phase. The standard ADM is a onedimensional model which neglects radial dispersion as well as non-uniform velocity distributions. The theoretical aspects of three-phase slurry reactors have been analyzed by Govindarao [106]. Ostergaard and Michelsen [107] showed a method for evaluating the tailing problem with the RTD in a three-phase reactor, while Cova [108] measured the steady-state concentration distribution of solids in a three-phase reactor. Farkas and Leblond [109] used such data to calculate the backmixing coefficient of the solids, although the results were not correlated to gas and liquid velocities. The effects of suspended solid particles on liquid backmixing were examined by Kato et al. [110], who correlated the longitudinal dispersion coefficient for the

× ×

0.025 0.02

× 10 µm × 30 µm 70 µm Eq. (4)

0.015 0.01 0.005 0

0

0.05

0.1

0.15

0.2

0.25

0.3

uG / m s−1 Air–slurry axial dispersion coefficient data compared with theoretical predictions [109] at low slurry flow rates (superficial velocity of slurry phase = 0.034 m s−1 , 11 wt.%, Dc = 0.152 m, water-coal and dried mineral ash, ρp = 1300 kg m−3 ). (Adapted from Ref. [41].) Fig. 2

References see page 2152

10.3 Slurry Reactors

b = 0. Kato et al. [110] found an exponential variation of the solid particles concentration with height; this effect may be significant in columns with large bed aspect ratios. Farkas and Leblond [109] discussed the effect of solid distribution on the reaction rate in the SBC. Kara et al. [41] reported that the air–slurry dispersion coefficient depends on gas and slurry velocities and the particle size, and it is relatively independent of the solids concentration. It is commonly assumed that the liquid in the SBC is perfectly mixed, and radioactive tracer investigations confirm close to complete backmixing [2]. Numerous correlations for the liquid axial dispersion coefficient have been reported, but unfortunately there is wide scatter in their predictions [43], caused largely by the unsuitability of ADM in describing liquid backmixing with a single parameter. Extensive studies utilizing computer-aided radioactive particle tracking (CARPT) [2, 112] have revealed that, at sufficiently high uG values and in large aspect ratio columns, in a time-averaged sense a large-scale liquid circulation cell occupies most of the column height, with liquid ascending along the central core region and descending along the annular region between the core and the walls. While a single onedimensional axial liquid velocity profile exists in this large recirculation cell, with negligible radial and azimuthal liquid velocities, two- and three-dimensional velocity profiles are evident in the distributor and free surface (disengagement) region. The height of the entry and the disengagement zone, respectively, is approximately equal to the column diameter in the churn-turbulent flow regime. Liquid recirculation is driven by non-uniform radial εG profiles, which (in a time-averaged sense) have been shown to approach a parabolic shape for churn-turbulent flow. Superimposed on this recirculation are turbulent fluctuations in the axial, radial and azimuthal directions due to eddies that are induced by the wakes of the rising gas bubbles. Liquid mixing is therefore primarily due to convective liquid recirculation, driven by the existing non-uniform radial εG profile and turbulent dispersion due to bubble wake interactions and turbulent eddies [2]. Often, it is assumed that the gas phase moves in plug flow. This may be justified for the SBC, but in the case of mechanically agitated slurry reactors the gas-mixing pattern is uncertain. Niiyama and Smith [93] analyzed the transient response of a slurry adsorber, and showed the mixing pattern of the gas to be unimportant for sparingly soluble and highly soluble gases. Similar effects were predicted by Goto and Smith [113]. Thus, the influence of the gas-phase mixing is likely to be important only for gases in the intermediate solubility range, and when gas–liquid mass transfer is important.

Gas Holdup Correct design and scale-up of SBCs depend largely on the accurate prediction of the gas holdup, εG . The latter is a key parameter to characterize the macroscopic hydrodynamics of SBCs, and many reports have focused on gas holdup in slurry reactors [32, 36, 38, 39, 42, 44, 74, 114–152]. In addition to the classical manometric methods to measure gas holdup, γ -ray densitometry [132], ultrasonics [144, 150] and a dual electroresistivity probe [127] have been used. Gas holdup εG generally increases with an increase in gas velocity, uG , with the effect being more pronounced in a dispersed bubble regime than in a churn-turbulent regime. The distributor design affects εG significantly only at low uG values. The addition of particles into a bubble column usually leads to a larger bubble size, and thus gas holdup is reduced (Fig. 3). According to Fan et al. [153], the particle size effect on gas holdup can be neglected in the particle size range of 44 to 254 µm. It is noteworthy that the maximum in the εG curve for pure organic liquids vanishes as the slurry concentration increases; this sharp maximum denotes a shift in the regime from homogeneous bubbly to churn-turbulent flow. Deckwer et al. [30] identified a minimal effect of pressure on εG in an FT SBC with a porous plate distributor (P = 0.4–1.1 MPa; T = 143–260 ◦ C; uG < 0.035 m s−1 ). The experimental data of Kojima et al. [141] indicated that εG increases with pressure, although no pressure effect was observed at 30 wt.% solids concentration (P = 0.1–1.1 MPa; uG = 0.02–0.09 m s−1 , single-orifice distributor). Inga [154] measured εG in SBCs for FT 10.3.3.5

0.6 0.55 0.5

Total gas holdup eG

2138

0.45 0.4 0.35

+ + + + + + + fs = 0.0 + + + + + fs = 0.05 + + fs = 0.10 +++ + + ++ + fs = 0.16 + + + ++ fs = 0.36 + + +

0.3 0.25 0.2 0.15 0.1 0.05 0

0

0.05 0.1 0.15 0.2 0.25 0.3 0.35 0.4 0.45 0.5

Superficial gas velocity uG / m s−1 Effect of particle concentration on the gas holdup in a (air/paraffin/silica particles, dp = 38 µm) 0.1-m diameter column equipped with sintered plate gas distributor. (Adapted from Ref. [44].)

Fig. 3

10.3.3 Hydrodynamics of Slurry Reactors

synthesis at pressures up to 0.72 MPa, and observed a significant pressure effect whereby the increased gas holdup was directly related to a smaller bubble size and, to a lesser extent, to a reduced bubble rise velocity at higher pressures. The main reason for the bubble size reduction may be attributed to variations in the physical properties of the gas and liquid with pressure [153]. Knowledge of both bubble flow characteristics and bubble size is essential for the fundamental understanding of gas holdup behavior in SBCs. Luo et al. [145] applied the particle image velocimetry (PIV) system in order to examine flow characteristics and to substantiate the applicability of the dynamic gas disengagement (DGD) technique to measure bubble-size distribution in the SBC. Here, bubble size and distribution in a high-pressure SBC were measured directly using a fiber-optic probe. The gas holdup εG in the SBC was measured using the DGD technique when, at steady-state operation, the gas inlet and outlet are simultaneously shut off to maintain a constant system pressure. The bubbles disengage from the column and the particles gradually settle. A differential pressure transducer was used to monitor variations in pressure drop. Luo et al. [145] reported that the presence of particles reduces εG at both P = 0.1 and 5.6 MPa, while Barghi et al. [73] also found εG -values to be lower in the SBC compared to the bubble column. The lower εG in the SBC can be attributed to an increase in apparent suspension viscosity [155]. At ambient pressure, εG in the bubble column was twice that in a slurry of 0.19 solids volume fraction, over the entire uG range [145]. In contrast, at the pressure of 5.6 MPa, the effect of the solids concentration on εG is relatively small at uG ≥ 0.25 m s−1 . Luo et al. [145] found that elevated pressures led to higher εG under all experimental conditions, but that the pressure effect was more significant at low pressures and at higher solids concentrations. The same group [145] developed an empirical correlation to account for εG in high-pressure SBCs operated in the coalesced bubble (churn-turbulent) regime:

α
 u4G ρG ρG β 2.9 σg ρsl εG =  (8)   4.1 0.054 1 − εG cosh Mosl where Mosl is the modified Morton number for the slurry phase: (ξ µL )4 g ρsl σ 3

(9)

0.008 α = 0.21Mosl

(9a)

Mosl = and

−0.011 β = 0.096Mosl

2139 (9b)

The effective slurry density is given by ρsl = φS ρS + φL ρL

(10)

where φS and φL are the volume fractions in the gas-free suspension. In Eq. (9), ξ is a correction factor which accounts for the effect of particles on slurry viscosity:  ln ξ = 4.6φS 5.7φS0.58    × sinh −0.71 exp(−5.8 φS ) ln Mo0.22 + 1 (11) where Mo =

gµ4L ρL σ 3

(12)

When the solids volume fraction φS approaches zero, this correlation reduces to the form for bubble columns. Since εG is basically dictated by the size and rise velocity of bubbles, the large bubbles play a key role in determining εG . Through the bubble wake attraction, the gas holdups in SBCs are closely associated with the size and number of the large bubbles. A fundamental explanation of the pressure effect on εG should be based on a full understanding of the variation of bubble size, especially of the large-bubble size, with the pressure. The increase in εG with pressure is a result of decreases in both bubble size and bubble rise velocity. De Swart et al. [76] also showed that, as the slurry concentration increases, the holdup of small bubbles decreases. At the column axis, the small bubbles move with similar velocity as the large bubbles in the same area. In the wall region or the vortical flow region, the downward liquid motion significantly affects the small bubbles. The liquid circulation becomes more pronounced at higher uG values, and as uG increases the bubble–bubble interactions begin to play an important role. When the DGD technique is applied to SBCs, during the disengagement of large bubbles the small bubbles are significantly affected not only by the wakes of large bubbles but also by the downward motion of liquid. During the disengagement of large bubbles, various assumptions about the disengagement of small bubbles have been used. Their impact on the results of the DGD technique has been discussed in detail by Jordan et al. [156]. Krishna et al. [74] studied the influence of particle concentration on the hydrodynamics of SBCs operating References see page 2152

2140

10.3 Slurry Reactors

in the heterogeneous flow regime. Paraffin oil was used as the liquid phase and slurry concentrations were up to 36 vol.%. The authors identified the following two phases: (i) a ‘‘dilute phase’’ consisting of fast-rising large bubbles that pass through the column virtually in plug flow; and (ii) a ‘‘dense phase’’ consisting of the liquid phase along with solid particles and entrained small bubbles. Krishna et al. [74] used a DGD technique in the heterogeneous flow regime in order to determine the gas voidage in the dilute (εb ) and dense (εdf ) phases. Experimental data show that increasing the solid concentration decreases εG significantly, but the influence on εb is small. It is the εdf value that decreases significantly due to the enhanced coalescence of small bubbles as a result of the introduction of particles. εdf is practically independent of Dc but the εb -value decreases with column diameter due to wall effects on the rise velocity of the large bubbles. The entering gas in the churn-turbulent regime is split into two parts: a portion of the gas (udf ) rises through the column in the form of ‘‘small’’ bubbles, whereas the remainder (uG –udf ) rises through the column in the form of ‘‘large’’ bubbles. The superficial gas velocity through the dense phase udf is determined from the slope of the disengagement curve for the dense-phase gas. The gas holdup of the large bubbles can be estimated as follows [74]: εb =

udf = vsmall εdf

(17)

where vsmall and εdf at the regime transition point are obtained from Eqs. (15) and (16). For the dilute-phase holdup (large bubbles), Krishna et al. [74] suggest the following dimensional equation:
−0.18  uG  udf 0.58 Dc (18) − εb = α m ms−1 ms−1 with α = 0.268 for φS < 0.16 and α = 0.3 for φS > 0.16. The total gas holdup in the heterogeneous flow regime can then be calculated as follows: εG = εb + εdf (1 − εb ) (heterogeneous flow regime) (19)

Bubble Rise Velocity Differences in fluidizing media, pressure and temperature may lead to different bubble rise characteristics. According to Li and Prakash [158], the rise velocities of both small and large bubbles increase with increasing slurry concentrations, indicating an increase in average bubble size in the dispersion. The bubble rise velocity ub in high-pressure slurries can be calculated by a correlation developed by Luo et al. [159], the simplified version, for a large bubble, is  gdb ρsl − ρG 2.8σ + (20) ub = ρsl db 2 ρsl 10.3.3.6

(14) where vsmall,0 is the rise velocity of the small-bubble population in the solids-free paraffin oil. This trend indicates an increase in the size of the small bubbles due to coalescence. The total gas holdup in the homogeneous flow regime can be estimated from vsmall

The dense-phase gas holdup for the gas–liquid system εdf ,0 can be estimated from the correlation for the gas voidage at the regime transition point, as suggested by Reilly et al. [157]. The superficial gas velocity at the regime transition point can be estimated from

(13)

vsmall = vsmall,0 (1 + 8.4φS ) ; vsmall,0 = 0.095 m s−1

uG

(16)

with the dense-phase gas holdup εdf from Eq. (16).

εG − εdf 1 − εdf

The rise velocity of the small bubble swarm in the homogeneous regime is also affected by the solids loading. Krishna et al. [74] found that the rise velocity of the small bubbles increases linearly with the particles volume fraction:

εG =

volume fraction in the suspension:
 0.7 εdf = εdf ,0 1 − φS ; εdf ,0 = 0.27 εdf ,0

(homogeneous flow regime)

(15)

The phenomenon is observable also in the heterogeneous regime, prevailing at uG > udf , where εdf is significantly reduced with increasing solids concentration. Krishna et al. [74] argue that εdf is approximately constant at uG ≥ 0.1 m s−1 (churn-turbulent regime). The εdf value was empirically fitted as a function of the solids

The large bubbles have a dominant effect on ub in SBCs due to bubble-wake attraction. It is found that, under all experimental conditions, the bubble size at elevated pressures is significantly smaller than at ambient pressure. The presence of particles in liquids leads to a larger bubble size, especially at ambient pressure. The increased liquid circulations cause the bubble swarm velocity to increase with increasing uG [160]. As a result of the increased liquid circulations, the bubbles tend to rise faster in the central core. Fan et al. [153] reported that the bubble rise velocity decreases with an increase in pressure

10.3.4 Mass Transfer in Slurry Reactors

or gas density. According to these authors, the decrease in bubble rise velocity may occur due to corresponding variations of gas and liquid properties with pressure. They also stated that the surface tension decreases and liquid viscosity increases with increasing pressure. Moreover, their simulations have shown that the elevated pressure causes the bubble shape to become more flat. The same group [153] concluded that the bubble was more stable when the gas density was higher, and that the bubble rise velocity was reduced with a decrease in temperature. Krishna et al. [44] proposed the rise velocity of the swarm of large bubbles vb to be estimated as follows: vb =

(uG − udf ) εb

(21)

Equation (21) is valid when the E¨otv¨os number is >40. Considering Eq. (18), the large bubble swarm velocity therefore increases with increasing Dc . Krishna et al. [44] developed a model to describe the scale dependence of vb : ! (22) vb = 0.71 gdb,large (SF )(AF ) where SF and AF are the scale correction factor and the acceleration correction factor, respectively. SF accounts for the influence of the column diameter and is expressed as follows: SF = 1 for

db,large

< 0.125

 SF = 0.496

(23b) Dc db,large

for

db,large Dc

10.3.4

Mass Transfer in Slurry Reactors

In a slurry reactor, a number of transport steps must occur. For the reactant in the gas phase, the steps are: • transport of the reactant from the bulk gas phase to the gas–liquid interface • transport from gas–liquid interface to the bulk liquid • transport from bulk liquid to the catalyst surface • intra-particle diffusion in the pores of the catalyst • adsorption of the reactant in the gas phase on the active sites of the catalyst • surface reaction of the gas-phase reactant to yield products. Additional steps such as the desorption of products and transport of products from the catalyst to bulk liquid and, for volatile products, bulk gas, may also become ratelimiting. While in the liquid within the porous particle mass transport and reaction may occur in parallel, as long as the reaction occurs mainly in the bulk rather than only in the liquid film (Fig. 4), the overall rate can still be described by resistances in series. For a pseudo-first-order reaction in the gas-phase reactant A:  r = (kG a)−1 + (HA kL a)−1 + (HA kS aS )−1 + (HA k1 ηεS )−1

(23a)

Dc 
db,large −db,large for 0.125 < < 0.6 SF = 1.13 exp Dc Dc

> 0.6

db,large = 0.069(uG − udf )0.376

−1

cA,G

(26)

where HA is the ratio of the equilibrium concentrations of reactant A in the liquid and gas phase; for the catalyst effectiveness factor ηc of a porous catalyst see Section 10.3.4.7. A detailed review of the literature on

(23c) aS

a

Catalyst

Liquid

Gas

AF accounts for the increase in vb over that of a single, isolated bubble due to wake interactions. Krishna et al. [44] give the following relations (for Tellus oil): AF = 2.25 + 4.09(uG − udf )

2141

cG

(24) (25)

The model agrees well with εb data for paraffin oil slurries. The large bubbles play an important role in dictating the macroscopic hydrodynamics of SBCs due to their large volume, high rise velocity, and the wakes associated with the large bubbles. According to Fan et al. [153], in the churn-turbulent regime more than 70% of the small bubbles are entrained by the wakes of large bubbles, and consequently have a velocity close to that of large bubbles.

cL

c

kG Fig. 4

kL

kS

k1η

Mass transfer resistances in a catalytic slurry reactor.

References see page 2152

2142

10.3 Slurry Reactors

mass transfer in three-phase systems up to 1995 was provided by Schumpe and Nigam [161].

1.4

Volumetric Liquid-Side Mass-Transfer Coefficient (kL a) Gas–liquid mass transfer is often the rate-controlling step, and in some cases it can be enhanced by the presence of solids. Different cases of enhancement can be distinguished (see Ref. [162]):

1

1.2

• enhanced mass transfer by physical adsorption onto small particles • enhanced mass transfer due to inert particles • enhanced mass transfer due to reactive particles • enhanced mass transfer due to catalytic particles. The first two effects are described by a physical enhancement factor, Ephys , and the latter two by a chemical enhancement factor, Echem . Physical and chemical enhancement effects may occur in parallel and thus their distinction might be difficult. It may also be problematic to separate the individual effects on the specific gas–liquid interfacial area a and the liquid-side mass-transfer coefficient kL . Although some interesting enhancement effects have been identified, in concentrated slurries (as used in most industrial applications) the presence of particles will usually be detrimental to mass transfer as the high slurry viscosity induces bubble coalescence. Although extensive investigations have been carried out on the influence of solids on kL a, there is still no general relationship describing the influence of all types of particles at any weight fraction in any liquid. If the density difference between the solids and liquids is small, or if the liquid viscosity is high, the slurry behaves as a pseudo-homogeneous phase and relations for kL a as a function of the effective suspension viscosity can be applied successfully. ¨ and 10.3.4.1.1 kL a in the Slurry Bubble Column Ozturk Schumpe [163] measured kL a in a 0.095 m inner diameter (i.d.) bubble column absorbing oxygen in organic liquids containing up to 40% of various polymer particles and aluminum oxide (Fig. 5). With the exception of highdensity particles in a low-viscosity liquid (alumina in ligroin), all results could be uniformly correlated, with a mean error of 7.7%:
 µeff −0.42 kL a = (27) (kL a)0 µ0 The correlation covers uG values up to 0.08 m s−1 and an effective viscosity range of 0.54 to 100 mPa s.

k La / (k La)0

10.3.4.1

0.8 0.6 0.4 0.2

Al2O3 /tetralin Al2O3 /ligroin PVC/ligroin PE-1/ ligroin PE-2 /ligroin PE-3 /Exsol

0 0.001

0.01

0.1

1

fS Variation of kL a in slurry bubble columns compared to kL a0 in the particle-free organic liquid (Al2 O3 : dp = 22 µm; polyvinyl chloride, PVC: dp = 100 µm; polyethylene, PE-1: dp = 369 µm; PE-2: dp = 46 µm, PE-3: dp = 45 µm). (Adapted from Ref. [163].) Fig. 5

For non-Newtonian suspensions, the effective slurry viscosity µeff can be estimated from the Ostwald–de Waele equation: n−1 µeff = k γ˙eff

(28)

For bubble columns, the effective shear rate is a function of superficial gas velocity: γ˙eff = C uG

(29)

where C = 2800 m−1 [164]. There is an excellent agreement between the above result, obtained in organic media, and the results of Schumpe et al. [33] both for water and aqueous salt solutions obtained in the same bubble column applying kieselguhr (7 µm), alumina (8 µm) and activated carbon (5 µm). Provided that µeff > 2µL , it was found that
 µeff −0.39 kL a = (30) (kL a)0 µ0 Equation (30) is valid for uG < 0.08 m s−1 and 1 < µeff < 100 mPa s. Fine high-density solids in low-viscous liquids such as alumina in ligroin or kieselguhr in water (at low concentration) yield slightly higher kL a values than predicted by Eqs. (27) or (30), possibly by increasing kL [33]. However, non-wettable particles may reduce kL a substantially more than predicted by Eqs. (27) and (30). For non-wettable polypropylene particles in aqueous solutions of carboxymethyl cellulose, the extra decrease could be described by assuming surface blocking of the gas–liquid interface by a Langmuir–Hinshelwood-type of adsorption of the non-wettable particles [165].

10.3.4 Mass Transfer in Slurry Reactors

Nigam and Schumpe [166] found that the relationship developed by Nguyen-tien et al. [167] for three-phase fluidized beds of glass spheres 
kL a φS (31) = 1− (kL a)0 0.58 is also applicable to SBCs with non-uniform distribution of the particles (5 < dp < 200 µm). However, the solids fraction at the bottom (where intense coalescence occurs) should be substituted into Eq. (31). The solids fraction could be measured or calculated with the sedimentation–dispersion model [110]. Sauer and Hempel [168] examined aqueous slurries of a wide variety of low-density particles (1020 < ρp < 1381 kg m−3 ) with relatively large diameters (0.4 < dp < 2.9 mm) at up to 20 vol.% solids fraction. In addition, two solids with a small particle diameter were included (110 µm PVC, ρp = 1376 kg m−3 ; and 200 µm sand, ρp = 2780 kg m−3 ). These authors applied a sintered plate (3-µm pores) and a perforated plate gas sparger (1-mm holes) in a bubble column of 0.15 m diameter, and were able to correlate all results by using the concept of an effective slurry viscosity by the empirical correlation:

 n1 νsl 0.5 uG kL a =C guG (νsl guG )0.25 n2
n3
c¯s νsl × (32) ν¯ eff ,rad cs0 where

µsl ρsl  µsl = µL 1 + 2.5φS + 10.05φS2 νsl =

+2.73 × 10−3 exp(16.6φS )

ν¯ eff ,rad

= 0.011Dc gDc




uG νsl g

(32a)

1/8 (32c)

(c¯s /cs0 ) is the normalized solids distribution; that is, the ratio of the mean solids concentration to that just above the sparger: [1 − exp(−P eS )] c¯s = cs0 P eS

kL a in the Stirred Tank With regards to the influence of solids on kL a in stirred-tank reactors, some very comprehensive studies have been published. Oguz et al. [172] used several dynamic techniques to measure kL a in various slurries in a baffled 0.145 mdiameter stirred-tank reactor in which water and three organic liquids (1-butanol, 1-tetradecene, and 1,2,4trimethylbenzene) were used as liquid phases. The following particles were applied: sea sand, kieselguhr, Al2 O3 , Fe2 O3 , TiO2 , ZnO, CaCO3 , and BaSO4 . The observed effects of the particles on kL a are remarkably complex and seem to be confusingly different for different slurry systems. However, if the increase in apparent slurry viscosity due to the presence of the particles exceeds a factor of 1.3 relative to the solids-free liquid, then all data can be consistently correlated by a single function of power input, gas-sparging rate, apparent slurry viscosity, surface tension, and liquid diffusivity in a similar manner to that proposed earlier by Yagi and Yoshida [173] for solids-free systems. The results of Oguz et al. [172] are as follows: √ kL adI2 1.5 0.5 = 0.162 ReST Scsl F rG0.19 GF −0.6 A0.09 σH2 0 /σ DL 10.3.4.1.2

(34) ReST is the impeller-based Reynolds number: ReST =

(33)

where PeS can be estimated from Eq. (6) derived by Kato et al. [110]. Sauer and Hempel [168] succeeded in correlating their results by applying different sets of numerical values for the constant C and the exponents n1 to n3 for the two gas spargers. Koide et al. [32] and Sada et al. [169] also examined the influence of solids on kL a. Some studies have also been reported on the influence of solids on kL a and εG in the draft tube slurry reactor [170, 171].

ωdI2 ρeff µeff

(34a)

Scsl is the slurry Schmidt number: Scsl =

(32b)

2143

µeff ρeff DL

(34b)

GF is the gas flow number: GF =

σ µeff uG

(34c)

A is the aeration number: A=

ωdI uG

(34d)

As also found in bubble columns, the concept of the effective viscosity fails for relatively high-density particles in low-viscosity liquids. The effective viscosity µeff is estimated from the Ostwald–de Waele equation [Eq. (28)] with the shear rate γ˙eff obtained from the relationship of Metzner and Otto [174]. Schmitz et al. [175] and Kojima et al. [176] also studied the influence of solids on kL a in stirred tanks. Albal et al. [177] examined the influence of pressure on kL a in stirred-tank slurry systems, but no influence of References see page 2152

2144

10.3 Slurry Reactors

pressure on kL a was observed in the range from 2 to 9 MPa. These authors also provided a comprehensive review of the literature on kL a in stirred tanks. Iwanaka et al. [178] measured the influence of solids on both kL a and kL in a slurry reactor with vibrational agitation, while Steinberg and Sch¨oo¨ n [179] described experimental techniques suitable for kL a measurements with catalytic reactions. Lindner et al. [180] found a strong increase in specific interfacial area due to activated carbon particles in concentrated salt solutions encountered in hydroxylamine production (HPO process of DSM). A particle layer on the bubbles may prevent bubble coalescence and increase the specific interfacial area (this effect is not observed in water [33]). Gas–Liquid Specific Interfacial Area (a) Especially for fast reactions, where enhancement at the gas–liquid interface occurs, knowledge of the true gas–liquid specific contact area, a, is desired rather than knowledge of the product kL a only. An overview of measurement techniques is provided by Beenackers and van Swaaij [181]. Recently, a new ultrasonic technique was introduced by Chang and Morala [182]; for an evaluation of the chemical method, see Refs. [183] and [184]. The effects of suspended solids on the interfacial area can be rather diverse (Fig. 6). Pandit and Joshi [185] proposed that four different regions be distinguished, with the first two regions being of primary importance to slurry reactors. The first region is observable with small particles, typically below 100 µm, at low solids holdup ( τy

(44)

γ˙ = 0 for τ ≤ τy

(45)

β

= τy /τ ≤ 0.7, Kawase and Moo-Young [196] For derived:
 eρL 0.25 8 kL = √ (1 − β  )1.45 (46) 15 π µB For bubble columns, the specific energy dissipation e is expressed as guG . For Newtonian liquids β  = 0, and Eq. (46) reduces to
 eµB 0.25 kL = 0.3Sc−0.5 (47) ρL Equation (47) was tested in a bubble column with various aqueous solutions of carboxypolymethylene, simulating slurries with τy up to 0.7 Pa and Bingham viscosity as high as 32 × 10−3 Pa s. When the slurry can be considered as a pseudohomogeneous phase, the effect of solids on kL is relatively well comprehended. Less clear is the situation with small fractions of heavy particles, which is typical for many catalytic slurry reactions. A moderate increase of kL by up to 30% seems possible, depending on factors that are not yet well understood. Gas-Side Mass-Transfer Coefficient (kG ) Despite the availability of adequate experimental techniques [181], no specific information is available on kG a in slurry reactors [197, 198]. Therefore, correlations for gas–liquid systems should be used whenever significant gas-side resistance is encountered. 10.3.4.4

Liquid-Solid Mass-Transfer Coefficient (kS ) For reviews on measurement techniques for kS , see Refs. [40] and [199]. A new method applicable under reactive conditions without requiring a priori knowledge of the kinetics was suggested by Gholap et al. [200]. Relations for the mass-transfer coefficient kS around the solid particles are usually presented in the form of 10.3.4.5

Sh =

n1 Scn2 2 + C Rem

(48)

where kS dp DL µL Sc = ρL DL

Sh =

In early attempts, the Reynolds number was often based on the hindered particle settling velocity – that is, the actual terminal particle slip velocity in the slurry, υp : Rem =

ρL υp dp µL

(48c)

There are two difficulties in this approach: first, the slip velocity must be related to the global system parameters; and second, the concept becomes inapplicable for density differences approaching zero, as is often the case for bioreactors. An illustration of the difficulties involved in the slip velocity concept is the study on kS in SBCs performed by Jadhav and Pangarkar [201, 202]. The second approach is based on the idea that kS is influenced by the local turbulence around the particles. Based on Kolmogoroff’s theory of local isotropic turbulence, Re is defined by means of the velocity of the critical eddies responsible for most of the energy dissipation. For solid particles much larger than the Kolmogoroff scale of these eddies, this leads to Re ∝

e dp4

0.33 (49)

νL3

In Eq. (49), e is the energy dissipation rate per unit mass of liquid. For bubble columns, the power input is expressed as guG . S¨anger and Deckwer [203] developed the following correlation:

4/3

e1/3 ds ρL Sh = 2 + 0.545 µL

0.792 Sc1/3

(50)

An interesting investigation in this area focuses on the influence of liquid, gas and solid properties on kS and on scale-up rules. Lazaridis [204] measured the influence of dp , uG , e, σ, εs and L on kS . The author’s equation can also be applied to non-Newtonian liquids. Particularly valuable are the studies of Jadhav and Pangarkar [202], who tested their relation in three columns of 0.1, 0.2, and 0.4 m diameter. Although these authors also used the concept of the specific energy dissipation rate, their approach is somewhat different, defining the Reynolds number as Re =

vL dp νL

(51)

with the characteristic turbulence velocity vL defined as [185]:

(48a)

 "
#$0.33 (ρp − ρL ) vL = 0.43 gDc uG − εG υb∞ − εs υt ρL

(48b)

(52)

10.3.4 Mass Transfer in Slurry Reactors

Except for concentrated slurries, the last term of Eq. (52) can be neglected, so that % &0.33 vL = 0.43 gDc [uG − εG υb∞ ]

(53)

In fact, vL is proportional to the liquid circulation velocity. The relation of Jadhav and Pangarkar [202] is less easily applicable than the relations of Jadhav and Pangarkar [201], S¨anger and Deckwer [203], Lazaridis [204], and Sano et al. [205]. The reason is that knowledge of εG and υb∞ is required. For the latter, Jadhav and Pangarkar [201] suggested using the relations of Grace et al. [206] for so-called ‘‘clean’’ systems, and of Zuber and Findley [207] for other systems. The approach based on the energy dissipation rate as outlined above is not limited to a particular type of slurry reactor. Therefore, equations of type (48–50) have also been proposed for stirred-tank reactors. For e we must then use the total energy dissipation rate originating from both gas and power inputs via the stirrer. Hence, e = eG + e w

(54)

Ww e = guG + ρL VL

(55)

Various correlations were proposed by Sano et al. [205], Marrone and Kirwan [208], and Asai et al. [209], and these agree well with each other [162]. It is noteworthy that the relation of Sano et al. [205] contains a correction factor for asphericity of the particles. Additional correlations for ks can be found in Refs. [210–215]. Enhancement of Gas–Liquid Mass Transfer Particles may enhance the gas absorption rate by various mechanisms [161, 162]: 10.3.4.6

• higher surface renewal rate due to inert particles (cf. Section 10.3.4.5) • higher interfacial area due to coalescence hindrance (cf. Section 10.3.4.4) • mass transfer enhancement by physical adsorption on small particles • mass transfer enhancement by catalytic reaction on the particles • mass transfer enhancement by reactive particles (soluble or insoluble). The first three effects could be described by a physical enhancement factor, Ephys , and the latter two by a chemical enhancement factor, Echem . Physical and chemical enhancement effects may occur in parallel, and thus their distinction might be problematic.

2147

Adsorptive Transport (‘‘Shuttle Mechanism’’) Enhancement of gas absorption in slurries in the presence of small particles was first observed by Kars et al. [216] and Alper et al. [217], followed by several other groups [194, 218–223]. The effect is explained by gas being adsorbed onto the particles in the film and released in the bulk of the liquid (termed the ‘‘shuttle mechanism’’ or ‘‘grazing particles’’). These observations have been performed with activated carbon slurries in stirred cells with a flat interface where the interfacial area is well known and the mass transfer rate is relatively low (in bubble columns with activated carbon slurries no or only marginal enhancement is observed [192, 220]). The enhancement factors obtained for a typical case of oxygen or carbon dioxide absorption into a carbon slurry [217] can only be explained by a much higher particle concentration in the mass-transfer zone than in the bulk of the slurry. A higher concentration might be caused by the poor wetting properties of activated carbon. Following Schumpe et al. [165], Vinke [224] proposed a type of Langmuir adsorption isotherm to describe the gas–liquid surface fraction covered by particles. Alper et al. [217, 222] proved that the enhancement reaches a stable value for relatively low solids loadings. To explain this, Sada and Kumazawa [219] and Alper and Deckwer [225] assumed the existence of a particle-free layer at the interface with a thickness of one particle diameter, dp . This appears controversial, however, due to the necessity of higher particle concentration as compared to the bulk. These phenomena may occur in parallel to chemical reactions on a carbon-supported catalyst. Activated carbon supports are used in many catalytic slurry systems, and similar phenomena have been observed for other solids [220–222]. An industrially important example of non-catalytic particles is the absorption of carbon dioxide in aqueous amine solutions. The addition of fine activated carbon particles may enhance the absorption rate ninefold in a stirred cell [223]. The explanation of the shuttle mechanism and its effects were recently questioned [226]. In a stirred cell with a flat interface, similarly high mass-transfer rates as with activated carbon could be induced by adding other slightly hydrophobic solids (e.g., graphite) with practically no adsorption capacity for the gas. Moreover, in a carefully cleaned system, the same absorption rate was reached without activated carbon and, upon addition of activated carbon, there was no enhancement (Fig. 7). Kaya and Schumpe [226] concluded that the solids actually adsorb surface-active contaminants (caused by fingerprints or dust) that otherwise would induce surface rigidity. During CO2 absorption into monoethanol amine solutions, the 10.3.4.6.1

References see page 2152

2148

10.3 Slurry Reactors

104 kL / m s−1

1.5

1.0

Contaminated system Clean system 0.5

0

1

2

cs / kg m−3 Effect of activated carbon particles on the oxygen transfer coefficient in a stirred vessel with flat surface when carefully cleaned and when contaminated by dust (open cover). (Adapted from Ref. [226].)

Fig. 7

adsorption of surfactants onto activated carbon was found to enable surface convection and thus lead to very high absorption rates [227]. Based on these recent findings, the adsorptive transport mechanism seems questionable. 10.3.4.6.2 Absorption with Heterogeneous Catalytic Reaction Enhancement of gas absorption due to a heterogeneous catalytic reaction may occur if the catalyst particles have a diameter dp which is much smaller than δL . The reaction within the film can often be considered as quasi-homogeneous and be described by the well-known relations for gas absorption with homogeneous chemical reactions [228]. A complication may occur when particles are attracted to the gas–liquid interface, as has been observed for carbon–water slurries [224, 229, 230] and aqueous sulfur slurries [231]. Enhanced gas absorption will then occur at much lower bulk loadings of catalyst particles. If the concentration of solids at the gas interface is high, the enhanced mass transfer is very sensitive to the actual geometrical position of the particles [232]. More complex systems have been analyzed, on a theoretical basis, by Sada et al. [233]. Chemical-enhanced gas absorption with fine catalyst particles in slurries has been observed for several chemical systems [194, 216–223, 234]. It is difficult to distinguish between enhancement by the heterogeneous catalytic reaction and either gas adsorption or surfactant adsorption (cf. Section 10.3.4.6.1) as the cause of enhancement. With many catalytic reactions, a leveling-off of the enhancement factor takes place at a certain catalyst loading. This can be caused by a full surface coverage [225] or by a limitation of the pseudo-homogeneous approach

when the effective film thickness is decreased to the size of the particle. When a fast reaction occurs in the film only, the catalyst in the bulk is useless. It may be attractive to use a catalyst of slightly reduced wettability so that it sticks to the bubbles, and just such a situation is encountered in the HPO process of the Dutch State Mines for hydroxylamine production where Pd on activated carbon is used in a concentrated salt solution. Here, an important additional effect is that coalescence hindrance increases the interfacial area [180]. The idea of bubble-adhering catalysts was discussed enthusiastically in the previous edition of this Handbook [8], but has not found further application. 10.3.4.6.3 Reactive Particles Doraiswamy and Sharma [235] have reviewed the systems with reactive particles. Often, no chemical enhancement occurs because product layer diffusion limits the conversion rate. However, the latter situation will not arise if the products rapidly dissolve in the liquid. If the particles are very small they may disappear during the contact time at the gas–liquid interface, and a penetration-type model should be used to estimate the absorption rates. Another case is the conversion of small hydridable metal particles in slurry [236–238]. As hydride formation is very rapid, the external mass transfer of hydrogen to the particles is rate-limiting. During the contact time, the particles may become hydrogen-saturated, thus limiting the enhancement. A special case here is the situation when the particles are soluble and mainly serve to provide the liquidphase reactant for the gas. This system has been studied exhaustively [181, 239], and the resultant absorption enhancement was modeled by Uchida et al. [240, 241].

Optimum Catalyst Size When the catalyst particles are large and the reaction rate is high, the concentration of the gaseous reactant will decrease towards the center of particle. The effect on the overall reaction rate can be described with the effectiveness factor η: " # 1 1 η= coth(3) − (56)  3 10.3.4.7

where  is the generalized Thiele modulus. For spherical particles and a first-order reaction: " # R k1 0.5 (57) = 3 De Chaudhari and Ramachandran [1] summarized the equations to be used for more complex kinetics. Any pore-diffusion resistance could be eliminated in slurry reactors by further decreasing the particle size.

10.3.5 Heat Transfer in Slurry Reactors

However, there are limitations caused by difficulties in particle separation and particle agglomeration. Small particles may form clusters with large effective particle sizes and correspondingly poor mass-transfer characteristics. In a detailed analysis of this discrepancy, Gamwo et al. [242] have identified the optimum particle size for the case of methanol synthesis in the SBC to be 60 µm. 10.3.5

Heat Transfer in Slurry Reactors

The heat-transfer behavior strongly depends on hydrodynamics, and is associated with both the macroscopic flow structures and microscopic flow characteristics. For gas–liquid and gas–liquid–solid systems, it has been reported that the heat-transfer coefficient h increases with an increase in gas velocity, but levels off in the uG range of 0.1–0.2 m s−1 [30, 243, 244]. The heat-transfer coefficient is always lower at the wall than in the center of the column, where a larger bubble size prevails [245]. A higher thermal conductivity, as well as heat capacity of the liquid, augment heat transfer, whereas h decreases with an increase in liquid viscosity. Therefore, an increase in temperature augments significantly the heat-transfer coefficient [244] as the liquid viscosity decreases with increasing temperature. Different trends have been reported for the effect of suspended particles [244, 245], and details published up to 1995 have been reviewed by Suh and Deckwer [246]. The effect depends on particle size: larger particles tend to increase the heat-transfer coefficient, whereas very fine particles may lead to a high effective slurry viscosity and thus increase the viscous sublayer at the surface. The data in Fig. 8 illustrate the case of fine particles.

2149

Deckwer et al. [30] measured the heat-transfer coefficient at an immersed heater under conditions simulating the FT slurry process (P = 0.1–1.0 MPa, T = 523–573 K). These authors used hydrocarbon liquid and Al2 O3 powder (dp ≤ 5 µm) at solids concentrations up to 16 wt.%, and the increase in solids content led to a higher heat-transfer coefficient. On the basis of the surface renewal model, and also the Kolmogorov theory of isotropic turbulence, a correlation was proposed to predict the heat-transfer coefficient in SBCs. The original relationship developed by Deckwer [247] was used but, instead of the liquid properties, the effective slurry properties were considered: 
 1/2 uG g 1/2 h = 0.1 cp,sl ρsl ksl (58) µsl The slurry density ρsl can be estimated from Eq. (10), whereas the heat capacity of the suspension can be averaged based on the solids weight fraction ωS : cp,sl = cp,L (1 − ωS ) + cp,S ωS

(59)

The thermal conductivity of the suspension can be calculated as follows: ksl = kL

2kL + kS − 2φS (kL − kS ) 2kL + kS + φS (kL − kS )

(60)

The slurry viscosity was estimated from Einstein’s relation. As a result of the good agreement of Eq. (58) with the authors’ own results and literature data for particles up to 0.1 mm diameter, it is concluded that these suspensions can be considered as a pseudo-homogeneous References see page 2152

Mean heat transfer coefficient h / kW m−2 K−1

10 9 8 7 6 5

uG /m s−1 0.06 0.11 0.17 0.24 0.35

4 3 2 1 0

0

10

20

30

40

50

Slurry concentration / vol. % Time-averaged heat-transfer coefficient (measured at the column axis) as a function of the slurry concentration (Dc = 0.28 m, air-water-glass beads (dp = 35 µm), perforated gas distributor). (Adapted from Ref. [245].)

Fig. 8

2150

10.3 Slurry Reactors

phase. Equation (58) is frequently used, although the effect of pressure is not accounted for as it was found to be negligibly small. Other physical properties, which are less affected by pressure, include liquid density, liquid thermal conductivity, and liquid heat capacity [153]. Yang et al. [244] studied heat transfer in an SBC with up to 35 vol.% glass beads (dp = 53 µm) at pressures up to 4.2 MPa; the addition of particles to the liquid phase also enhanced the heat transfer appreciably. As the pressure was increased, the heat-transfer coefficient in the SBC decreased substantially, a behavior which the authors attributed to variations of the physical properties of the liquid phase, bubble size, and gas holdup with the pressure [244]. As the bubble wake size is proportional to the bubble size, a larger bubble would have a larger wake and stronger vortices associated with the wake, thereby enhancing the rate of heat transfer. Since the bubble size decreases as a function of pressure, the wake contribution to the heat transfer is reduced and thus the pressure would have a negative effect on the heat-transfer coefficient. The effect of solids concentration on the heat-transfer coefficient is more pronounced at ambient pressure than at high pressure because the bubble size is relatively large and the effect of solids concentration on the bubble size is stronger at ambient pressure [145]. Yang et al. [244] proposed the following empirical correlation for the heat transfer coefficient in an SBC at elevated pressures: "
#−0.022 εG Stsl = 0.037 (Resl F rG Pr1.87 ) (61) sl 1 − εG where Stsl = Resl = F rG = P rsl =

h ρsl cp,sl uG uG db ρsl µsl u2G gDc cp,sl µsl ksl

(61a) (61b) (61c) (61d)

The term εG /(1 − εG ) can be estimated from Eq. (8). The slurry properties were calculated from Eqs. (59) and (60), but the suspension viscosity was estimated after Luo et al. [159]:    µsl = exp   µL

 KφS
 φS  1− φS,c

Luo et al. [159] presented the following formulas for K and φS,c : K = {3.1 − 1.4 tanh[0.3(10 − 102 vt )]}/ψ 2

φS,c = {1.3 − 0.1 tanh[0.5(10 − 10 vt )]}εS0

(62a) (62b)

where ψ is the particle sphericity, εs0 is the solids holdup at packed state, and vt is the particle terminal velocity in liquid. A comparison between the experimental data and the predictions of Eq. (61) shows that the deviation of the predictions is within ±20%, and the average deviation is 6.9%. Yang et al. [244] demonstrated that Eq. (58) cannot predict the trend of the pressure effect on the heat-transfer coefficient for systems in which pressure has a significant effect on hydrodynamics, arguing that Eq. (61) can predict reasonably the pressure effect in such systems. 10.3.6

Concluding Remarks

Over the years, reviews of reports on slurry reactors have shown that the most intensely studied (in terms of number of publications) topics are as follows: • • • • •

hydrodynamics: 29% mass transfer: 23% syntheses in slurry reactors: 20% (FT: 16%) mathematical modeling: 12% (CFD: 3%) heat transfer: 7%.

Although numbers of publications centered on numerical simulations of slurry reactors based on CFD codes are expected to increase in the near future, the dynamics of slurry reactors has received scant attention [105, 248]. During the past few years some interesting modifications of slurry reactors have been developed, including mini slurry reactors [249], and boiling [250, 251] and vibrating [252, 253] alternatives, the further exploration of which may prove valuable. List of Symbols and Subscripts Symbols

a a0

(62)

as

where K is a proportionality constant and φS,c is the critical solids volume fraction in the slurry.

A

specific gas–liquid interfacial area, m−1 specific gas–liquid interfacial area in the solids-free system, m−1 specific internal surface area of porous solids (per unit particle volume), m−1 aeration number in a stirred tank, Eq. (34d), dimensionless

List of Symbols and Subscripts

AF b cA,G cp cS ( cS ) cS0 db db,large dI dp ds D Dc De e(e) eG ew E Ephys (Echem ) Fr g GF h HA k k1 kG kG a kL kL a kS ksl L Lc

acceleration factor, Eq. (24), dimensionless exponent, Eq. (6), dimensionless concentration of species A in the gas phase, mol m−3 heat capacity, J kg−1 K−1 (mean) solids concentration, kg m−3 solids concentration at the bottom, kg m−3 bubble diameter, m large bubble diameter, m impeller diameter, m particle diameter, m Sauter-mean bubble diameter, m molecular diffusivity, m2 s−1 column diameter, m effective diffusion coefficient in the pores of the catalyst, m2 s−1 (average) power input per unit mass of liquid, W kg−1 average specific power input due to the gas, W kg−1 average specific power input by stirring, W kg−1 axial dispersion coefficient, m2 s−1 physical (chemical) enhancement factor, dimensionless Froude number, Eq. (5) or Eq. (61c), dimensionless acceleration due to gravity, m s−2 gas flow number, Eq. (34c), dimensionless heat transfer coefficient, W m−2 K−1 Ostwald solubility coefficient for A in a liquid, dimensionless fluid consistency index, Pa sn pseudo first-order rate constant, s−1 gas-side mass-transfer coefficient, m s−1 volumetric gas-side mass-transfer coefficient, s−1 liquid-side mass-transfer coefficient, m s−1 volumetric liquid-side mass-transfer coefficient, s−1 liquid to catalyst particle mass-transfer coefficient, m s−1 thermal conductivity of liquid–solid suspension, Eq. (60), W m−1 K−1 total height of slurry above the gas distributor, m characteristic length, m

Mo n N P Pe Pr sl

r R Re Rem Rep Resl

ReST Re Sc Scsl SF Sh Stsl T tc ub udf uG uG,min vb vb∞ vL vp

2151

Morton number, Eqs. (9, 12), dimensionless flow behavior index, dimensionless constant, Eq. (6), dimensionless pressure, Pa Peclet number, dimensionless Prandtl number based on slurry properties, Eq. (61d), dimensionless reaction rate per unit volume of slurry, mol m−3 s−1 radius of the catalyst particles, m Reynolds number, Eq. (49), dimensionless modified Reynolds number, Eq. (48c), dimensionless particle Reynolds number, Eq. (7), dimensionless Reynolds number based on slurry properties, Eq. (61b), dimensionless impeller-based Reynolds number, Eq. (34a), dimensionless modified Reynolds number, Eq. (51), dimensionless Schmidt number, Eq. (48b), dimensionless Schmidt number of the slurry phase, Eq. (34b), dimensionless scale correction factor, Eqs. (23a–c), dimensionless Sherwood number, Eq. (48a), dimensionless Stanton number based on slurry properties, Eq. (61a), dimensionless temperature, K contact time in penetration model, s bubble rise velocity, Eq. (20), m s−1 superficial gas velocity through the dense phase, m s−1 superficial gas velocity, m s−1 minimum uG for complete suspension of solid particles, m s−1 rise velocity of the large bubble swarm, Eqs. (21, 22), m s−1 terminal rise velocity of a single bubble, m s−1 characteristic turbulence velocity, Eq. (52), m s−1 actual terminal particle slip velocity in the slurry, m s−1

References see page 2152

2152 vsmall vsmall,0

vt VL Vr W (Ww )

10.3 Slurry Reactors

rise velocity of small bubbles, Eq. (14), m s−1 rise velocity of the small-bubble population in the solids-free suspension, m s−1 terminal settling velocity of a single particle, m s−1 volume of liquid phase, m3 volume of reactor, m3 power input (by stirring), W

Greek Letters α proportionality constant in Eq. (18), dimensionless dimensionless β  τy /τ (effective) shear rate, s−1 γ˙ (γ˙eff ) δL liquid-film thickness for mass transfer, m εb gas voidage in the dilute phase (large bubbles), dimensionless gas voidage in the dense phase εdf (suspension + small bubbles), dimensionless εdf ,0 gas voidage in the dense phase for the gas-liquid system, dimensionless total gas holdup, dimensionless εG εS solids holdup, dimensionless solids holdup at packed state, εS0 dimensionless volume of solids per unit εS cross-sectional area of the column, Eq. (2), mm η catalytic effectiveness factor, Eq. (56), dimensionless µ dynamic viscosity, Pa s µB Bingham viscosity, Pa s effective dynamic viscosity of slurry, µeff Eq. (28), Pa s ν kinematic viscosity, m2 s−1 ξ correction factor, Eq. (11), dimensionless ρ density, kg m−3 σ surface tension, N m−1 τ shear stress, Pa τy yield stress, Pa volume fraction of the liquid in the φL gas-free suspension, dimensionless φS (φS,c ) (critical) volume fraction of the solids in the gas-free suspension, dimensionless  Thiele modulus, Eq. (57), dimensionless ψ particle sphericity, dimensionless

ω(ωmin ) ωS

(minimum) stirring speed, s−1 mass of catalyst per mass of liquid, dimensionless

Subscripts 0 without solids G gas L liquid p particle S solid sl suspension Acknowledgments

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10.4

Unsteady-State Reactor Operation Grigorii A. Bunimovich, Peter L. Silverston, and Yurii S. Matros∗

10.4.1

Introduction

When unsteady-state reactor operation is discussed, it is often in conjunction with catalytic processes characterized by rapid catalyst deactivation, such as catalytic cracking (see Chapter 13.5) or the dehydrogenation of alkanes (see Chapter 14.6). In these processes, a sequence of reaction and regeneration steps is unavoidable and so the catalyst, at least, is always in an unsteady state. Unsteady-state ∗

Corresponding author.

reactor operation, however, is encountered much more widely. Many catalytic reactions are performed in batch reactors; these are always operated in transient mode. Millions of vehicular catalytic converters (see Chapter 11.2) always operate in an unsteady state because of varying engine exhaust conditions. In this chapter, none of the above processes will be considered; rather, we will restrict our discussion to a forced unsteady-state operation for continuous rather than batch processes that have industrial importance. Catalyst life in these processes may be as long as several years. The traditional operation of the processes considered here is steady state; indeed, automatic control systems are used to eliminate any fluctuations in performance. The goal of traditional design is to obtain the optimal steady-state operating conditions for maximum capacity consistent with the required product quality. However, reactor performance obtained under optimal steady-state operating conditions may not represent a limit for a catalytic reaction system. Performance can be improved further using a forced unsteady-state operation which adds new useful process features and allows a better exploitation of the non-linear properties inherent in a catalytic reaction system. The benefits are generated by either or both of [1]: • The dynamic properties of the catalyst: Unsteady conditions in the fluid phase can invoke changes of state, composition, and structure of the catalyst which, under certain conditions, may result in improved selectivity and/or activity compared to steady-state operation. • The dynamic characteristics of an entire reactor system: Modulation of inputs can lead to optimum temperature and composition distributions in the reactor that are not possible in any steady-state regime. Under unsteadystate conditions, a bed of catalyst can acquire unique functions that not only accelerate a chemical reaction but also store energy and mass. Various schemes of forced unsteady-state reactor operation are shown in Figs. 1–4. An unsteady or transient state in a reactor can be created by modulating the inlet composition, flow, pressure or temperature by using switching valves (Fig. 1b). As a rule, a simple stepwise periodic variation (Fig. 1a) is preferable compared to other types of inlet parameter forcing (e.g., a sinusoidal variation). A step variation often is simpler to implement and it affects a system more strongly than a smooth type of control. Manipulated variables for modulation are amplitude (A), time of cycle (tc ), split of inlet signal (S), and phase angle when more than one input is forced. One widely applied technique for fixed-bed reactors is that of periodic flow reversal (Fig. 2a). Four switching

10.4.1 Introduction

tc Feed 1

Stc Control function

Valve 1

A

Average control function

Valve 2 Time

Fixed bed reactor

Feed 2 (b)

(a)

(a) Step-wise variation of inlet parameters. (b) Fixed-bed reactor with on-off inlet valves to provide periodic switching between two different feed compositions to implement step-wise control. See text for details.

Fig. 1

valves may be used to provide this operation. In a rotary reactor (Fig. 2b) with a plate-type catalyst bed, the bed position continuously changes relative to the direction of flow, thereby achieving a behavior similar to that in a flow-reversal reactor. The reactor schemes shown in Fig. 2a and b provide for a continuous migration of a temperature or a reaction zone along the catalyst bed. The direction of this migration is changed periodically. Another configuration (Fig. 2c) provides the possibility of moving this zone in one direction by using a system with two catalyst beds. Instead of valve switching, the desired unsteady state of the catalyst can be obtained by catalyst circulation inside the fluidized-bed reactor (Fig. 3b). Circulation can

1

be also obtained using a diagram (Fig. 3a) combining fluidized-bed and riser reactors. The latter arrangement allows separate feeding of two reactants, for example, hydrocarbons and oxygen in partial oxidation and a stripping inert gas. Another group of forced unsteady-state processes combines chemical reaction with separation of products. For these processes, a tubular chromatographic reactor can be used with a continuous flow of a carrier gas and periodic pulses of reactants. This is rather inefficient, so a reactor with a circulated moving bed (Fig. 4a) is a better choice. A more practical system places the catalyst in an annular vessel that revolves past fixed feed and withdrawal ports (Fig. 4b). Circulation of the catalyst can be avoided also by periodic changing of feed and product ports in a reactor with several fixed beds (Fig. 4c), an arrangement referred to as a ‘‘simulated moving-bed reactor’’. Reaction and separation can be performed in a mixed bed of catalyst and adsorbent with multiple ports at the bed entrance and using a cycle of periodic pressure changes. This operation is an adaptation of the widely used pressure-swing adsorption. A typical pressure cycle and a schematic of the reactor-adsorber are illustrated schematically in Fig. 5. The accurate design of unsteady-state catalytic processes requires knowledge of catalyst behavior and reaction kinetics under unsteady-state conditions, processes of unsteady-state intraparticle and interface mass and heat References see page 2172

2

1

2

2157

3

4 4 O C C O

C O O C

O C C O

5

C O O C

Valve position

Valve position

3 1 2 3 4

tc Time

1 2 3 4 5 6

O C C O C O

6 C O O C O C

O C C O C O

C O O C O C

tc (a)

(b)

(c)

Time

Schemes of fixed-bed reactors operated under forced unsteady-state conditions. (a) Reverse-flow reactor; (b) rotary reactor; (c) reactor system with periodic changes between the inlet and outlet ports in two catalyst beds. The tables show the positions of the switching valves during two successive switching cycles. C: valve closed, O: valve open.

Fig. 2

2158

10.4 Unsteady-State Reactor Operation

Regenerator

Stripper

Air

Riser

(a)

Inert gas

Hydrocarbon

(b)

Fluidized-bed reactors for partial oxidation processes. (a) The DuPont system for maleic anhydride production consisting of a riser reactor, a fluidized-bed regenerator and a stripper. (b) Fluidized-bed reactor with internal catalyst circulation.

Fig. 3

Feed, A

Strongly adsorbed product (Extract) B, S

Reactant feed A, S

3

5 Direction of fluid flow and port switching

Section 3

Carrier make-up (Eluent) S

Section 1

1

(b)

Raffinate C, S

Eluent S

4

2

Feed A, S

(a)

Section 4

Raffinate C, S

Carrier recycle

Solid recycle

Lightly adsorbed product (Raffinate) C, S

Carrier Fluid (Eluent) S Rotation

6

8

7

Extract B, S

Section 2

Extract B, S

(c)

Chromatographic reactors. (a) Continuous countercurrent moving-bed reactor; (b) continuous rotating annular-bed chromatographic reactor; (c) simulated moving-bed reactor with four catalyst sections.

Fig. 4

transfer, and flow patterns or mixing behavior inside the reactor vessel. New approaches for reactor modeling and optimization are necessary. These topics form a broad area of research that has been continuously developed during the past four decades. 10.4.2

Dynamic Kinetic Model

Traditional kinetic models of catalytic reactions assume time-invariant concentrations of intermediate species on

the catalyst surface. This assumption often is not valid for unsteady-state conditions characterized by continuous changes in a fluid phase composition and temperature. Besides, the catalyst itself can interact with the reaction mixture and can undergo significant changes, influenced by changing conditions in the fluid phase. Such a modification of the catalyst can be considered as a side process in relation to steps of a reaction mechanism in a catalytic cycle. Side processes comprise both changes in chemical composition of the catalyst and reconstruction of its surface structure. At steady state, reactions involving

10.4.3 General Approaches to Reactor Modeling

Feed

Pressure

Feed Delay

Delivery Exhaust Exhaust

Time Pressure swing adsorptive reactor. The left part of the figure indicates pressure changes during the cycle.

Fig. 5

surface intermediates proceed at fixed rates while the side processes affecting the catalyst state are in equilibrium and are thus invariant. The development of unsteady-state kinetic models requires the application of dynamic relaxation methods (see Chapter 5.2.6). More reliable models can be achieved by coupling these transient techniques with direct determination of intermediate products on the surface of the catalyst [2, 3]. Microkinetic models based on surface-science studies can be supplemented by a proper description of catalyst state transformations. Unfortunately, however, there is as yet no unified or standard method suitable for the discrimination of unsteady-state or dynamic kinetic models. Several examples of unsteady-state kinetic models are listed in Table 1. These models are represented in a form of rate dependencies for catalytic reaction steps and side processes. Parameters of the models, such as rate constants and activation energies, are given in the references (see Table 1), and were determined mainly from experimental data using transient response techniques. For one of the examples, namely CO oxidation over a supported platinum catalyst, the kinetic theory of gases was applied for estimating adsorption constants. In the simplest case (Example 1), the model assumes an ideal catalyst surface and describes the rates of the catalytic steps on the basis of the law of mass action. A special definition of vacant active sites on the catalyst surface is used, as in Example 3 for the rate of CO adsorption. A linear effect of surface coverage on the activation energy of the reaction steps (Example 2) can be introduced if necessary. In Example 3, dramatic changes of oxygen adsorption and surface reaction rates depending on carbon monoxide coverage result from the adsorbate-induced phase reconstruction of the platinum surface. A parameter (θCO ), which is a simplified form for describing the influence of catalyst phase changes, was necessary for quantitative explanation of the various transient and steady-state kinetic experiments performed. Example 4 represents the simplest model capable of describing the side process of catalyst bulk change. The oxidation or the reduction of the vanadium oxide catalyst

2159

is represented as a reaction between surface oxygen and a bulk-phase oxygen vacancy. The same approach was used for the description of the side process of V5+ reduction to V4+ that proceeds during SO2 oxidation (see Table 1, Example 5, reaction 4). The second side process (Example 5, reaction 5), the saturation of vanadium complexes by pyrosulfate anions, is assumed to be at equilibrium. Each model includes parameters characterizing the total or balance concentration of intermediate species, CAS , in the catalyst volume. Some models require several balances. For example, the balance of bulk or lattice oxygen (CAS,V ) must be taken into consideration along with the balance of surface species for the oxide catalyst (Table 1, Example 4). The balance itself can change during the course of reaction due to various structural transformations of the catalyst surface. The ratio of CAS to the rate of catalytic reaction under steady-state conditions (r) provides a rough estimate of the reaction time scale. For typical heterogeneous catalytic processes for the production of bulk chemicals and petrochemicals, this value is estimated to be between 10−2 and 101 s (Fig. 6). Changes of reaction rate caused by the side processes of catalyst modification can take considerably longer, but this may be attributed to the larger amount of material in the catalyst bulk phase that can be involved in side processes, as well as to a slow rate of these processes in comparison with the steps of a catalytic cycle that occur on the surface. 10.4.3

General Approaches to Reactor Modeling

As a rule, mathematical models of unsteady-state processes cannot be formulated by simply adding time derivatives to the equations describing steady-state behavior of the reaction system. This is because dynamic kinetic models (see above) should be used, and because the rates of the heat and mass transfer processes must be considered. For example, the modeling of unsteady-state processes in a fixed-bed reactor must account for heat and mass transfer between the catalyst surface and the gas phase, though for steady-state operation these factors often may be neglected. The development of an adequate model for unsteadystate reaction processes utilizes two general procedures: (i) decomposition of the complex system into elements of different scales; and (ii) analysis of the time scales of essential processes in the reactor. Commonly used levels of decomposition are catalyst surface, single catalyst pellet, catalyst bed, reactor (including heat-exchangers, mixing and distributing devices), and the entire process chain with the catalytic reactor as one of its elements. Each References see page 2172

Reference

[4]

[5]

[6]

1

2

3

Dynamics kinetic models

Example

Tab. 1

CO oxidation over supported Pt

Addition of acetic acid to ethylene over silica-alumina impregnated by sulfuric acid

N2 O decomposition over NiO on SiO2

Process

3

3

∗ CO2 + ∗ S −− −− −− S CO2

4

∗ −−− CO∗ + O∗ − −− CO2 + 2 Side process (support surface):

∗ CO + ∗ −− −− −− CO 2 ∗ ∗ O2 + 2 −− −− −− 2O

1

Catalytic cycle (Pt crystallite surface):

∗ −−− A∗ + E∗ − −− AE + 2

∗ −−− E + ∗ − −− E 2 ∗ A + ∗ −− −− −− A

1

Reaction: E A CH3 COOH + C2 H4 −−−→ AE CH3 COOC2 H5 Catalytic cycle:

∗ −−− 2O∗ − −− O2 + 2

2

N2 O + ∗ −−−→ N2 + O∗

1

Scheme of reaction mechanism

y

(qco)L

y2 (qco)H

y1

" # 1 − NCO θCO r1 = CAS k1 CCO (1 − θCO − θO ) − −k−1 θCO ; 1 − θCO r2 = CASLPt k2 CO2 (1 − θCO − θo )2 ; r3 = k3 CAS LPt θCO θO ; r4 = CS k4 CCO2 (1 − ϕCO2 ) − k−4 ϕCO2 ;

1/2  RT E−1i k1 = 100 L−1 k−1 = A−1 exp − ; Pt SCO ; 2π MCO RT

1/2  RT E3 k2 = 100 ψL−2 k3 = A3 ψ exp − ; Pt SO2 ; 2π Mo2 RT

1/2  RT E−4 L−1 S ; k = A exp − ; k4 = 100 CO2 −4 −4 S 2π MCO2 RT θCO , θO = platinum surface coverage by CO and oxygen; SCO , SO2 , SCO2 : sticking coefficients; MO2 , MCO : molecular weights; Ai = pre-exponential factors; CAS = adsorption capacity of Pt surface per unit of catalyst volume; Lpt = adsorption capacity of Pt surface (mol cm−2 Pt ); Ls = surface CO2 adsorption capacity on the support (mol cm−2 ); φCO2 = support coverage by CO2 ; CS = CO2 adsorption capacity on the support per unit of catalyst volume (mol cm−3 ); NCO = factor limiting CO adsorption;  = factor, reflecting reversible transition between two surface phases of Pt. Function  is determined from the graph:

2 k θ ∗θ ∗; r3 = CAS 3 E A θE∗ , θA∗ , θmax = concentrations of adsorbed ethylene, acetic acid and maximum concentration of acetic acid in the liquid phase (mol (mol H2 SO4 )−1 ); CAS = molar concentration of catalyst, sulfuric acid, in the catalyst pellet (mol cm−3 ); b = Elovich constant.

) * r1 = CAS k1 CN2 O (1 − θ); r2 = C2AS k2 θ 2 − k−2 CO2 (1 − θ)2 . θ = surface coverage by oxygen; CAS = total concentration of surface species (mol cm−3 ). 

r1 = CAS k1 e−bθA∗ CE − k−1 e−bθA∗ θE∗ ;  r2 = CAS k2 (θmax − θA∗ ) CA − k−2 θA∗ ;

Rate equations

2160 10.4 Unsteady-State Reactor Operation

[7]

[3]

4

5

SO2 oxidation over vanadium catalyst: SO2 + 1/2 O2 = SO3

2-Butene oxidation over supported vanadium oxide

2

1

3

1

5

2− 5+ (V5+ −− −− (V2 )j S2 O7 2 )j + S2 O7 −− (V52 )j : binuclear vanadium complexes 5+ 5+ V5+ 2 OSO3 , V2 O, and V2 O2 .

−−− 4+ + SO3 V5+ 2 SO3 −−− 2V

4

Side processes:

5+ V5+ −− −− V2 O2 + SO3 2 SO3 + O2 −−

3

−−− 5+ V5+ 2 O + SO2 −−− V2 SO3

2

−−− 5+ V5+ 2 O2 + SO2 −−− V2 O + SO3

Catalytic cycle:

−−− S−O + L− − −− S− + L−O S− and S−O: reduced and oxygen covered surface site; L− and L−O: lattice oxygen and oxygen vacancy in the catalyst lattice.

Side process:

−−− S− + 1/2O2 − −− S−O

−−− C4 H8 + nS−O − −− Products + nS−

Catalytic cycle:

Reaction: C4 H8 + n/2 O2 −−−→ Products

    r1 = CAS θa k1 C1L θ1 − k−1 C2L θ2 ; r2 = CAS θa k2 C1L θ2 − k−2 θ3 ;     L L L r3 = CAS θa k3 C2 θ3 − k−3 C3 θ1 ; r4 = CAS θa k4 θ3 − k−4 C3 θ4 ; 4 1 θa = ; CLi = Hi pi ; i = 1, 2, 3; j=1 θj = 1; L 1 + K5 C3 CL1 , CL2 and CL3 are SO2 , O2 and SO3 concentrations in liquid phase (mol cm−3 ); pi = partial pressure of gaseous components (kPa); Hi = Henry’s constants (mol cm−3 kPa−1 ); θ1 , θ2 , θ3 , and θ4 are dimensionless concentrations of binuclear vanadium species in the melt; K5 = equilibrium constant of reaction 5 (cm3 mol−1 ); CAS = concentration of binuclear vanadium complexes in the melt (mol cm−3 ).

r1 = nk1 CAS CB (1 − θ); r2 = k2 CAS CO2 θ; r3 = CAS,V CAS [k3 θ(1 − θV ) − k−3 θV (1 − θ)] ki = Ai exp(−Ei /RT), i = 1, 2, 3, −3; r2 − r1 = r3 ; θ = surface coverage by oxygen; θV = concentration of lattice oxygen; CB and CO2 = butene and oxygen concentration in gas phase (mol cm−3 ); n = number of oxygen atoms consumed by every butane molecule in reaction 1; CAS = concentration of surface active sites (mol cm−3 ); CAS,V = concentration of oxygen vacancies in the lattice (mol cm−3 ).

1/2

(θCO )L and (θCO )H are critical CO coverages corresponding to changes of surface phase.

10.4.3 General Approaches to Reactor Modeling

2161

2162

10.4 Unsteady-State Reactor Operation

Hydrodynamics

Processes

Chemical reactions Mass transfer Heat transfer Activity changes 10−3 10−2 10−1

1

10

102

103

104

105

Time scale, s Time scales for various processes in a fixed-bed reactor. (From Ref. [9].)

Fig. 6

lower element is a component of the higher level, but each can be studied separately. The application of this procedure is illustrated by the analysis of unsteady-state processes in a single catalyst pellet [8, 9]. Separate consideration of this element allows estimation of the ranges of parameters in which certain heat- and mass-transfer steps can be neglected and the mathematical model thus simplified. These criteria for simplification, derived assuming a steady-state reaction rate (r), are listed in Table 2. The time scale technique employs estimation of the particular time scales or characteristic times (τi ) for each significant process in every element of the reaction system. The characteristic time for heat/mass transfer can be estimated as a ratio of the heat/mass capacity of an element to the rate of the respective transfer process. In addition, a time scale or response time of the whole system or element can be estimated. Estimates of characteristic times for important processes in a fixed-bed reactor [9] are given in Fig. 6. These

characteristic times can differ substantially from the time used to characterize process performance (process time) or the cycle time for forced unsteady-state operation. In this way, it is possible to determine very slowly changing variables – the changes of which may perhaps be neglected – and rapidly changing variables that can be assumed to be in a quasi-steady state. A condition of quasi-steady state means that the capacitances or inertial terms included in the model equations for those variables can be neglected for a major fraction of the time considered. For example, the assumption of an invariant concentration distribution inside a catalyst particle is valid for the following relationships between the process time and characteristic times of heat transfer (τpT ) and mass transfer (τpm ) inside the pellet: τpT ≥ 40 τpm η

and t > τpm

where τpT = ((cs ρs + εp cg ρg )/λef )Rp2 and τpm = 2 ), c is the heat capacity of catalyst pellet, (εp Rp2 /Def s (J kg−1 K−1 ); ρs is the density of the catalyst (kg m−3 ), and εp is the void fraction of the catalyst pellet; the other symbols are defined in Table 2. 10.4.4

Analysis and Optimization of Cyclic Processes

A set of the possible unsteady or cyclic states of the dynamic system includes a narrower set of steady states. Therefore, the maximum value of an object or performance function (J max ) obtained under optimal unsteady-state conditions may not be lower than that

Criteria allowing for simplification of the mathematical model of unsteady-state processes in a catalyst pellet

Tab. 2

Dimensionless group ratioa  ≤ 0.5 and Bim ≤ 0.3 √ T η ≤ 0.2 and BiT ≤ 0.1 Bim ≥ 30 and Bim ≥ 20

Simplification of mathematical model to be allowed. Diffusion limitations have no importance. No substantial temperature gradients inside the pellet. Limitation to mass-transfer between catalyst particle and bulk of gas phase has no importance.

= (kg Rp /Def );  = Rp (r/Def C); T =  2 θad Le−1 ; Le = (λef /Def cg ρg ); BiT = (hf Rp /λef ); θad = ((−H) · C · E/cg ρg RT 2 ); η = (3Bim (ψcth ψ − 1)/ψ 2 (Bim − 1 + ψcth ψ)); Bim and BiT = Biot numbers for mass and heat transfer;  and T = Thiele modulus; Le = Lewis number; θad = dimensionless adiabatic temperature rise; η = effectiveness factor; kg = mass-transfer coefficient (m s−1 ); Rp = radius of catalyst pellet (m); Def = effective diffusion coefficient (m s−2 ); r = rate of reaction (mol m−3 s−1 ); C = concentration of reactant (mol m−3 ); λef = coefficient of effective heat conductivity inside the catalyst particle (J m−1 s−1 K−1 ); cg = gas heat capacity (J kg−1 K−1 ); ρg = gas density (kg m−3 ); hf = coefficient of heat transfer between catalyst surface and gas phase (J m−2 K−1 s−1 ); −H = reaction enthalpy (J mol−1 ); E = activation energy (J mol−1 ); R = universal gas constant (J mol−1 K−1 ); T = temperature (K). a Bi

m

(1)

10.4.4 Analysis and Optimization of Cyclic Processes

for optimal steady-state (Jsmax ), or J max ≥ Jsmax . However, efficient or proper periodic operation includes only those which provide higher performance in comparison with the optimal steady state, J max > Jsmax . The goal of a mathematical investigation is to establish whether or not a maximum steady-state performance can be improved using a forced unsteady-state operation. Then, if the answer is positive, the optimal operation should be discoverable. The discovery involves cost-analysis of the optimal unsteady-state system in comparison with the steady-state system. For the class of systems described by ordinary differential equations the general problem includes: • Equations for state variables: x˙ = f (x, u)

(2)

where x = [x1 (t), x2 (t), . . . , xn (t)] is a vector of states that are continuous time-dependent functions. • Definition for forcing control variables: u(t) = [u1 (t), u2 (t), . . . , um (t)]

(3)

where u(t) is a vector of piecewise continuous functions determined over the time interval [0, tc ]. • Periodicity constraints for control functions u(0) = u(tc ) (periodicity constraints for state variables x(0) = x(tc ) are also usually assumed). • Definition of the objective function: + 1 tc fo (x, u) dt (4) J= tc 0 where fo (x, u) is an instantaneous process performance index. • Constraints on instantaneous values of control variables: ≤ uj (t) ≤ umax umin j j ,

j = 1, . . . , L

(5)

or on their averaged values. For steady state, solution of the system is determined from the equation 0 = f (xs , us ), and the objective function becomes Js = fo (xs , us ). Sufficient conditions for optimality of forced unsteadystate operation, which provides J > Js , can be determined by the analysis of two limiting types of periodic control [10]. The first limiting type is a so-called quasisteady operation, which corresponds to a very long cycle duration compared to the process response time, τ . In this case the description of the process dynamics is reduced to the equations x(t) = h(u(t)), where h is defined as a solution of the equation describing a steady-state system 0 = f (h(u(t)), u(t)). The second limiting type of

2163

operation, a so-called relaxed operation, corresponds to a very small cycle time compared to the process response time (tc τ ). The description of the system is changed to: + 1 tc x˙r = f (xr , u(θ)) dθ (6) tc 0 where xr is a vector of relaxed states. If this reduced system does not exhibit oscillatory behavior, the steadystate solution + 1 tc f (xrs , u(θ)) dθ (7) 0= tc 0 determines a vector of relaxed steady states, xrs . These states are assumed to be asymptotic values of state variables, x, at a very small cycle duration (tc → 0). If periodic control exhibits a non-linear effect on the state variables, the relaxed steady-state process indices can differ substantially from those for the steady state, and a positive or a negative effect from periodic control may be produced. A general approach to the analysis of low-amplitude periodic operation based on a so-called ‘‘-criterion’’ is described in Ref. [11]. The shape of the optimal control function can be found numerically using an algorithm proposed by Horn and Lin [12]. In Refs. [9, 13], this technique was extended to the simultaneous optimization of forcing function shape and cycle period. The technique is based on the periodic solution of the original system of state variables coupled with the solution of equations for adjoined variables [λ1 , λ2 , . . . , λn ]. These adjoined equations are: λ˙ i = −

∂H (x, u) ∂xi

(8)

with the boundary conditions: λi (0) = λi (tc ), where  H (x, u) = fo (x, u) + tc nj λj fj (x, u). At each step of an iterative solution, small increments in the control functions and cycle time are generated in accordance with the equations: uk+1 = ukl + εl l

∂H ∂ul

(9)

and + tck+1 = tck + εt

1 0

  n   λj fj (x, u) dτ

(10)

j

where k is the iteration index, and εl [l = 1, 2, . . . , m] and εt are small positive numbers. The calculation for periodic References see page 2172

10.4 Unsteady-State Reactor Operation

processes is performed using the Newton–Raphson procedure described in Ref. [12]. For the analysis of distributed-parameter systems such as, for example, a tubular fixed-bed reactor, numerical simulation of periodic operation at various values of the control parameters is typically employed. Asymptotic models for quasi-steady and relaxed steady states are valuable instruments for substantial simplification of the original distributed-parameter system. A method allowing for numerical optimization of the shape of inlet perturbation is described in Ref. [9]. 10.4.5

Time average reaction rate / µmol s−1 g−1

2164

6 2

5 4 1

3 2 1

0.0

Inlet Composition and Temperature Modulation

0.2

0.4

0.6

0.8

1.0

Mean mole fraction of hydrogen (balance N2)

A large number of experimental studies have demonstrated performance improvement under composition modulation for various heterogeneous catalytic processes, as reviewed in Ref. [14]. The list of these processes includes sulfur dioxide oxidation over a vanadium catalyst, ammonia synthesis over a promoted iron catalyst, hydrogen sulfide oxidation by SO2 on a bauxite catalysts, Fischer–Tropsch synthesis over ruthenium and cobalt catalysts, CO oxidation on different metallic and oxide catalysts, methanol synthesis on a Cu/ZnO catalyst, and dehydrogenation or partial oxidation of hydrocarbons over various oxide catalysts. A typical experimental system used an isothermal fixed-bed reactor with timer-operated valves to switch the inlet gas composition. In order to prove that periodically forcing concentration is superior to steady-state operation, the experiments usually included measurement of the steady-state reaction rate in the entire region of possible inlet compositions. This permits a comparison to be made between the cycle average rate for the forced unsteady-state process and a maximum steadystate rate obtained at optimal experimental conditions. A typical comparison of cycle average and steady-state rate data for a reaction with two reactants is shown in Fig. 7. Early attempts to carry out temperature modulations (for a review, see Ref. [16]) were generally unsuccessful because of the high thermal inertia and low surface-tovolume ratio in conventional reactors that do not allow the generation of fast temperature modulations necessary for maintaining a catalyst at unsteady state. Recently developed micro-reactors [17] allow periodic temperature changes to be made of more than 100 K in the second to sub-second range. A recent study [17] of temperature modulation for CO oxidation as a model reaction over a micro-channel supported platinum catalyst has shown a considerably higher yield compared to steady state. Two major factors – the non-linearity of chemical reaction rates and complexity of the reaction systems – are generally considered to be responsible for conversion

Steady-state (curve 1) and cycle average reaction rate (curve 2) versus feed composition, experimental data for ammonia synthesis reaction over a conventional iron catalyst. (After Ref. [15].)

Fig. 7

or selectivity improvement under forced unsteady-state operation [18]. Non-linearity of the reaction rate is the evident reason for improvement of the simple irreversible reaction An → B. Low-frequency temperature modulation around the average value T¯ increases the reaction rate compared to a steady-state calculated at an average temperature. A positive effect can also be obtained due to concentration variation if n > 1 and the time average concentration is restricted. Systems with two parallel or consecutive reactions can give rise to a selectivity enhancement at high frequencies of oscillation of temperature if the activation energy of the reaction to the desirable product is higher than that for the other reaction. For example [19], with high-frequency temperature modulation, the reaction pairs: (A → B, A → C), (A → B, B → C) and (A → B, A + B → C) will show a higher selectivity to product B than at the optimal steady state, if E1 > E2 , where E1 and E2 are activation energies of the first and second reactions in each of the pairs. For a simple Langmuir–Hinshelwood reaction mechanism [20], concentration forcing of one of the adsorbing reactants or a temperature modulation may lead to a maximum or minimum in reaction rate, depending on the modulation frequency. For concentration forcing, this resonant behavior is observed in the case of high total surface coverage, and the maximum cycle average reaction rate is not higher than that achieved at the optimal steady state. Temperature forcing may give rise to a resonant behavior for both high and low surface coverage, and the maximum rate is higher than the maximum steady state rate [20]. A selectivity increase in a complex catalytic process can be expected if: (i) competition between two gaseous

10.4.6 Partial Oxidation in Fluidized-Bed and Riser Reactors

species for an empty active site exists; and (ii) the characteristic time for changing concentration of an intermediate participating in the formation of the desired product is lower than that for an intermediate leading to an unwanted product. For example [9, 13] reactions A + B → C, A + B → D proceed in accordance with the following simple reaction mechanism: A + [Z] ← → [AZ]; B + [AZ] → C + [Z]; B + [Z] ← → [BZ], and A + [BZ] → D + [Z]. If the concentration of the intermediate [BZ] changes more slowly than the concentration of [AZ], then periodic variation of the concentration of component A increases the selectivity to product C compared to the optimal steady state. Variation at an intermediate frequency is preferred. The same reaction system shows an even larger selectivity improvement with the simultaneous variation of the inlet concentrations of both A and B. The reaction rate can be increased by concentration cycling in systems with substantial inhibition by the adsorbed product. The reaction of ethylene addition to acetic acid (see Table 1, Example 2) is one of the earliest examples of this. A rate increase was achieved with high-frequency modulation of acetic acid concentration. The cycle average concentration of the adsorbed intermediate A∗ (see Table 1) is lower than the steady-state concentration. This facilitates ethylene adsorption and leads to a favorable distribution between the two adsorbed species which substantially increases the reaction rate. Product inhibition can also be avoided through periodic purging of the catalyst with an inert gas. For example, in the oxidation of SO2 over a vanadium catalyst (Table 1, Example 5), periodic purging of the catalyst by a large amount of an inert gas or air allowed achieving an aboveequilibrium SO2 conversion [21]. The explanation of this surprising result is that the purge desorbs the SO3 product from the vanadia surface, thereby shifting the actual equilibrium state. Processes exhibiting multiplicity and/or self-sustained oscillations can be improved by imposing a forced perturbation of the inlet composition. For example, the oxidation of CO (Table 1, Example 3) exhibits high- and low-rate steady states in the region of intermediate CO coverage. Simultaneous modulation of the CO and oxygen inlet concentrations shifts the catalyst to a state that corresponds to the higher-rate branch in the multiplicity region [6]. This is a further example of the use of cycling to avoid reactant inhibition. It is rate inhibition by adsorbed CO that is responsible for multiplicity. Probably the most dramatic example of composition cycling correcting an inhibition problem is the Stopeffect [22]. Periodic interruption in reactant feed to a system in which one of the reaction products is adsorbed increases remarkably the rate of the catalytic reaction. For example, for the dehydration of alcohols or deamination of

2165

primary amines on acid–base catalysts [22], the optimal operation corresponds to a short cycle time in which the fraction of time for interruption of a reactant feed depends on the inlet concentration of that reactant. For the reaction to proceed, two adjacent unoccupied sites are needed. Stopping the addition of an inlet reactant to the system allows adjacent sites to become available. Some examples considered in Ref. [14] show that the reaction rate or selectivity increase from concentration cycling can be induced by a favorable change in catalyst states. For example, in the partial oxidation of propene over a copper molybdate catalyst, improvement in selectivity to the oxygenates arose through recrystallization and restructuring of the catalyst surface induced by the concentration variation. Changes in the molten vanadia phase due to cycling appears to provide an explanation for the significant increase in the rate of SO2 oxidation over a conventional potassium-promoted vanadia catalyst at low levels of SO2 conversion. Cycling between H2 and N2 over an iron catalyst greatly increases the rate of ammonia formation at low conversion that appears to result from dissolution of these reactants in the iron phase during cycling. Nevertheless, in most systems studied, it is the disruption of reactant inhibition by composition cycling that leads to the performance improvements. Composition modulation appears to be not always attractive for single reactions where the objective is to increase conversion. It is simpler, and perhaps even less costly, to increase conversion by employing more catalyst – that is, by increasing the reactor size. Adding more catalyst is not a solution when selectivity needs to be increased in a multiple reaction system; concentration cycling deserves consideration in such situations. 10.4.6

Partial Oxidation in Fluidized-Bed and Riser Reactors

Exposing the catalyst to time-varying fluid phase compositions can be achieved by moving the catalyst constantly between zones of different composition and temperature in one reactor, or between several reactors where each reactor has a different fluid-phase composition and temperature. Catalyst movement is most readily carried out using very fine, fluidizable catalyst particles. Many composition modulation experiments have demonstrated that selectivity in dehydrogenation or partial oxidation reactions can be substantially increased by bang-bang switching between a hydrocarbon and an oxidant, usually air [14]. Here, a redox mechanism is involved. The general redox mechanism of metal oxidecatalyzed oxidation of hydrocarbons involves two steps: (i) reduction of the surface layers by hydrocarbons; and References see page 2172

2166

10.4 Unsteady-State Reactor Operation

(ii) their re-oxidation by oxygen. While these two steps occur simultaneously in a reactor with the catalyst working under steady state, they can be carried out in two separate reaction zones in a reactor with catalyst circulation [23]. A hydrocarbon is fed into the first zone, where the desired oxygenate is formed by reaction with the oxidized catalyst. In the second zone, gas-phase oxygen re-oxidizes the catalyst. Clearly, the residence time of the catalyst in the first zone should be short enough to prevent the formation of an inactive reduced catalyst. If only the catalyst surface layers participate, the time for the catalyst reduction will be only a few seconds. Early attempts to implement catalyst circulation for partial oxidation processes were undertaken during the late 1960s [24], although problems with catalyst attrition and the transport of large amounts of solid per unit of production hindered the realization at that time. A more recent development is the DuPont Riser reactor for n-butane oxidation to maleic anhydride over a vanadyl pyrophosphate catalyst [23]. In the DuPont riser reactor (see Fig. 3a), the catalyst is exposed briefly (for ca. 10–30 s) to the hydrocarbon feed in the riser, where spherical catalyst particles with a dimension of about 100 µm are transported by a butane-inert mixture at a velocity of about 0.5 m s−1 . Reoxidation of the catalyst is carried out in the fluidizedbed regeneration zone. An important advantage of the recirculating catalyst process in comparison with a traditional fixed-bed or fluidized-bed reactor system is the possibility of dramatically increasing the concentration of n-butane in the reactor feed. The feeding of oxygen and n-butane into separate vessels avoids the explosion limit restriction on n-butane concentration. DuPont’s major problem with the process was the development of an attrition-resistant catalyst. Process selectivity from pilot unit data was 80–85% [23], which is 5–10% higher than for a conventional, wall-cooled, multi-tubular reactor. This selectivity gain is attributed to suppression in the riser of the highly active surface − oxygen species O− 2 and O , which are assumed to cause side reactions resulting in carbon oxides. Steam stripping of the catalyst after the fluidized reoxidation reactor also improved selectivity. Evidently, an accurate accounting of the factors determining selectivity in the riser reactor system can be obtained from a dynamic kinetic model. If the characteristic time for catalyst reduction or reoxidation is higher than the characteristic time for particle dispersion in a fluidized bed, the catalyst will adjust to the average conditions in the reactor volume rather than to the local temperature and composition at every point of the reactor. This situation can be used for improving the selectivity of the partial oxidation process. For example, the oxidation of o-xylene to phthalic anhydride over a V2 O5 −TiO2 catalyst occurs at higher selectivity for high

o-xylene conversion, if the catalyst state corresponds to a low conversion of o-xylene [25]. Therefore, stirring of the catalyst in a fluidized bed can result in selectivity improvement, and this effect is even greater in a fluidized bed that includes a special low-volume fixed packing that promotes mass transfer between the low- and highdensity phases of the bed. In such a reactor, a dropping temperature profile with height can be created that increases the number of catalyst particles with high concentrations of surface intermediates that lead to the formation of phthalic anhydride. 10.4.7

Dynamic Phenomena in a Fixed-Bed Reactor

Recently, several forced unsteady-state processes have been already commercialized, based on a dynamic phenomenon associated with an exothermic reaction in a fixed bed of catalyst. This phenomenon is referred to in the literature as ‘‘wrong-way’’ behavior of a fixed-bed reactor [26]. Substantial differences in the characteristic times of heat and mass transfer in a packed-bed reactor result in a surprising rise in temperature inside the reactor, after a sudden reduction of inlet temperature or after an increase in the flow rate. This unsteady-state phenomenon occurs because a reaction mixture with a higher concentration of reagents travels faster than the following low-temperature wave, and so reaches the downstream section of the bed that retains heat and a higher temperature due to high thermal capacity of the catalyst. At a large difference between the inlet and outlet temperatures, and substantial bed length, the transient temperature and conversion profiles do not change and move through the bed with essentially constant velocity (Fig. 8). These transient patterns, which are referred to as ‘‘traveling waves’’ or ‘‘creeping fronts’’, have been widely studied [27]. Estimates of the creeping front parameters can be derived from asymptotic models assuming unchanged temperature and concentration profiles moving through an infinitely long bed at a velocity uF : ξ = l − uF t, where ξ is the creeping front coordinate, l is the reactor length, and t is time. At constant velocity the energy balance for the catalyst bed may be written as: 
u 1 Tad xe (TF ) (11) uF = 1− εg Tad xe (TF ) γ TF − Tinp 1− γ TF − Tinp where u is the linear velocity of the reaction mixture (m s−1 ), Tad = ((−H )Co /cg ρg ) is the adiabatic temperature rise (K), γ = (1 − ε)(1 − εp )(cs ρs /cg ρg ) + εg ; εg = εp (1 − ε) + ε; ε, εp , and εg are the packed-bed porosity, the void fraction in the catalyst particle, and the total fraction of gas volume in the catalyst bed respectively, TF

10.4.8 Forced Unsteady-State Processes in Fixed-Bed Reactors

Ze = (TF − Tinp − Tad xe (TF )/RTF2 )E is the dimensionless Zeldovich number characterizing the temperature rise. The parameter ηI (TF ) is a weight-average effectiveness factor along the heat front coordinate:

600 1 500 3

4

5

400

300

200

ln[shψ(TF )] − ln ψ(TF ) (14) ψ 2 (TF ) where ψ(T ) = 3(Vp /Sp ) (k(T )/xe (T )/Def ) is a Thieletype modulus for a reversible reaction, and Vp (m3 ) and Sp (m2 ) are the volume and external surface areas, respectively, of the catalyst particle. An estimate of λ˜ ef is: ηI = 6

100 90 80 70 60 50 40 30 20

Conversion / %

Temperature / °C

2

1′ 2′

3′

5′

4′

100

0 0.0

0.2

0.4

0.6

0.8

uF (ucg ρg )2 γ u hf aV
#2 " Vp ucp ρg Sp 3 uF + γ 5 u (1 − ε)λp

λ˜ ef = λs,ef +

1.0

Catalyst bed depth / dimensionless Creeping fronts in a fixed catalyst bed. Curves 1 to 5 are temperature profiles, curves 1 to 5 are conversion profiles at different moments of time.

Fig. 8

is the maximum temperature (K), xe (TF ) is the equilibrium conversion at maximum temperature, and uF is the velocity (m · s−1 ) of the creeping profile. In many practical cases (εg /γ ) ≈ 1 · 10−3 , and the second term of the denominator on the right side of Eq. (11) can be eliminated. The creeping heat front can move not only in the direction coinciding with the flow direction (Fig. 8) but also in the opposite direction. In the latter case, the reaction heat will accumulate in the catalyst bed and the outlet (or maximum) temperature will be less than its adiabatic value: Tad xe (TF ) > TF − Tinp at uF < 0. A stagnant wave with uF = 0 is also possible. An analytical theory of the phenomenon is developed in Refs. [9, 28]. An estimate of the maximum temperature in a creeping front with a first-order reaction is:  u2 cg ρg 1E Tad (12) TF = xe (TF ) m R k(TF )ηI (TF )(1 − e−Ze ) λ˜ ef while the width of the zone, where the major temperature increase occurs is given by: L =

2167

λ˜ ef TF − Tinp ucg ρg Tad xe (TF )

(13)

where m ≈ 1 is a numeric coefficient; k(TF ) is the reaction rate constant (s−1 ) at the temperature TF , ηI (TF ) is a dimensionless parameter dependent on the temperature and diffusion coefficient inside the catalyst pellet, λ˜ ef is the effective thermal conductivity of the packed bed (J m−1 s−1 K−1 ), E is the activation energy (J · mol−1 ), R is the universal gas constant (J mol−1 K−1 ), and

(15)

where λs,ef is an effective axial thermal conductivity of the packed-bed solid phase (J m−1 s−1 K−1 ), and aV is a specific surface of the catalyst pellet per unit of packed-bed volume (m−1 ). Other symbols are listed in Table 2. Dependencies of the heat-transfer parameters hf , λs,ef on linear velocity, particle dimension, packed bed porosity and physical properties of the fluid can be found elsewhere [8, 29]. Equations (11) to (14) together reflect the main qualitative features of the creeping front that were demonstrated in experimental and numerical studies. An important point for practical applications is the prediction by this theory of the possibility of controlling the maximum temperature in a reaction front by variation of design parameters such as the linear velocity and size of the catalyst pellet. 10.4.8

Forced Unsteady-State Processes in Fixed-Bed Reactors Reverse-Flow Operation This is a simple technique for that allows an exothermic process continuous migration of the reaction front through a fixed catalyst bed. Usually, after several flow reversals, the periodic changes in temperature and other process parameters become stable and reproducible, and the reactor attains a cyclic or periodic steady state. Figure 9 shows the typical profiles observed during one half of a cycle for a reversible exothermic reaction. During the second half of the cycle, the profiles are symmetric to those in Fig. 9, and migrate in the other direction. In the reverse-flow reactor the catalyst not only accelerates the chemical reaction rate but also serves 10.4.8.1

References see page 2172

2168

10.4 Unsteady-State Reactor Operation

600 1

500

3 4 5

300

100 90 80 70 60 50 40 30 20

1′

Conversion / %

Temperature / °C

2 400

200 2′ 3′ 4′

5′

100

0 0.0

0.2

0.4

0.6

0.8

1.0

Catalyst bed depth / dimensionless

Temperature (curves 1 to 5) and conversion (curves 1 to 5 ) profiles in the reactor with periodic flow reversal for an exothermic reversible reaction.

Fig. 9

as a heat-exchange and heat-accumulation medium. A high-temperature reaction zone is efficiently trapped in the middle of the catalyst bed (Fig. 9), between two cold boundary zones. Due to regenerative heat transfer the reactor can maintain autothermal operation; that is, operation without external heat even for weakly exothermic processes (treatment of dilute gases). Another important feature of a reverse-flow reactor is a gradual decrease in temperature towards the outlet of a bed; this allows for higher conversion in an adiabatic reactor than under steady state for exothermic equilibriumlimited reactions such as SO2 oxidation or ammonia synthesis. Conventional unidirectional flow operation with an exothermic reaction provides a continuous temperature rise along the adiabatic catalyst bed. Estimates of creeping front parameters given by Eqs. (11–15) can be used with a long catalyst bed operated at a low flow-reversal frequency. Another simplified model describes relaxed steady-state operation with intermediate flow-reversal frequencies. The cycle time for such frequencies is much shorter than the time needed for the reaction zone to travel through the catalyst bed, but it is also much longer than the residence time of the gas flow in the reactor. This situation is possible for processes carried out under atmospheric pressure when the densities of the catalyst and gas phases differ substantially. The temperature profiles for relaxed steadystate operation exhibit very small oscillations around a certain cycle average profile while the concentration profiles are quasi-steady relative to the temperature profile

and flow direction. The process model can be reduced to a system of ordinary differential equations [9, 30–32] which, as shown in Ref. [31], appears to be similar to a model of a countercurrent fixed-bed reactor where the catalyst is deposited over the separation wall. The theory of reverse-flow reactors involves the stability problem of the high-temperature zone inside the catalyst bed when a low-temperature gas enters. It has been shown (e.g., see Refs. [9, 30–32]) that periodic flow reversal results in the multiplicity of temperature and conversion profiles. In particular, for adiabatic wall-insulated reactors, three possible temperature profiles were computed for a single exothermic reaction. Two extreme profiles – an upper one, characterized by a high maximum temperature and high conversion; and a lower one having zero conversion – are stable, while an intermediate profile is unstable. A reactor operating with the high-temperature profile can be extinguished by some perturbation of the operating parameters, such as an increase in flow rate, a decrease in inlet concentration, or an increase in the duration of the cycle. For multiple reactions more than three steady states are possible. The wall-cooled reactor can exhibit a more complex behavior (see Ref. [33]), including aperiodic and chaotic states when quick periodic changes caused by flow reversals superimpose with slow modulations. The theoretical analysis helped to determine possible ways for controlling and optimizing a reverse-flow reactor. For example, higher yields for reversible exothermic reactions correspond to a larger residence time along with a decreasing temperature profile. This can be obtained using larger catalyst pellets, a lower linear velocity, or a lower adiabatic temperature rise. However, these changes do not improve stability and require a longer catalyst bed [9, 32]. Possible applications of reverse-flow reactors are reviewed elsewhere [9, 32]. Hundreds of large-scale reverse-flow reactors (also known as regenerative catalytic oxidizers) operate in industry for the catalytic incineration of volatile organic compounds (VOCs). In most cases these systems contain ceramic or other inert packing material at the ends of a catalyst bed [9, 32]. The temperature after the inert packing can be estimated using the following expression:

ucg ρg inr Tout ≈ Tinp + Tad xout 1 + (16) Linr 2λ˜ inr ef

where the index ‘‘inr’’ relates to the inert packing, Linr (m) is the length of one of the boundary beds, and λ˜ inr ef (J m−1 s−1 K−1 ) is the effective heat conductivity of the inert packing. Preheating the inlet gas in the inert packing beds enables a substantial reduction to be made in the catalyst bed length. In addition, various recycle and purge

10.4.9 Flow Modulation in Three-Phase Reactors

techniques have been developed to prevent a back-flash of non-reacted gas from the void fraction of the inert bed and spaces before the catalyst bed at the moment of flow reversal. In contrast to conventional steady-state operated catalytic incineration systems, the reverse-flow reactor can operate at a low concentration of VOCs and low inlet gas temperature, without external preheating of the gas mixture. The minimum adiabatic temperature rise from VOC oxidation (Tad ) allowing this autothermal operation is as low as 15 to 25 ◦ C [9, 32]. The high energy recovery achieved in reverse-flow operations has been applied to weakly exothermic processes other than VOC oxidation; examples include the selective reduction of low-concentrated NOx [34] and the removal of CO and methane after natural gas-fired diesel engines [35]. When the temperature rise due to the exothermic reaction is moderately high, all (or a fraction) of the hot gas from the middle of the catalyst bed can be directed to a waste heat boiler or another heat recovery unit for energy utilization [9, 32, 36]. Such a technique can be practical, for example, in the removal and utilization of methane in coal mine ventilation air [36]. One promising application of reverse-flow reactors is to combine exothermic and endothermic reactions. For example, steam reforming may be combined with partial oxidation of methane; methanol reforming with its partial oxidation and dehydrogenation of ethylbenzene with oxidative dehydrogenation [32, 37]. Extensive experimental and theoretical studies (for a review, see Ref. [37]) led to the development of design fundamentals for a variety of process operating modes. These included the simultaneous continuous feeding of a mixture of reactants for both the endo- and exothermic reactions and alternating feeding of different reactants during different fractions of a flow reversal cycle in order to separate the endo- and exothermic processes in time and space. The high energy-efficiency of a reverse-flow reactor makes it possible to maintain gross heat intake only slightly higher than the energy consumption for the endothermic process, and thereby to reduce energy consumption for the desired endothermic process. Several industrial catalytic reactors with periodical flow reversal are used in non-ferrous metallurgy for the treatment of exhaust gases with SO2 concentrations ranging from 1 to 4.5% [9, 32, 38]. The inlet temperature of these exhaust gases is typically 40–60 ◦ C, while the minimum temperature needed to initiate SO2 oxidation is 370–400 ◦ C. Utilization of these gases by traditional methods of sulfuric acid production relies on a multi-bed catalytic reactor with feed gas preheating and interstage cooling. Reverse-flow operation substantially simplifies the reactor by decreasing the amount of catalyst required, and lowering the bed pressure drop. The simplest flow diagram (see Fig. 2a) is used for treating exhausts

2169

with SO2 concentrations below 2–3%. At higher SO2 concentrations, reactors with one or two intermediate heat exchangers for heat removal are recommended [9, 32, 38]. The flow reversal reactor can be applied for other reversible processes such as the synthesis to methanol [9, 39] and of ammonia [9, 40], and sulfur production over a bauxite catalyst using the Claus process [9, 41]. In the latter application, sulfur condensing at low temperature blocks the active catalyst surface, eventually stopping the reaction. In a reverse-flow reactor there is periodic evaporation of condensed sulfur from the outlying parts of the catalyst bed. Although it is difficult to remove all of the sulfur which has condensed within the catalyst pellets, after some time a balance is attained between the amounts of sulfur which condense and evaporate. Migration of a Reaction Zone through a Reactor Network A continuous movement of reaction zones through a network of two fixed-bed reactors can be created using a flow diagram as shown in Fig. 2c. Generally, this unit may include more than two catalyst beds. In these arrangements, the thermal wave travels continuously through a series of packed beds in one direction, as in a closed ring. Similar to the reverse-flow operation, the process is autothermal for weakly exothermic processes [9, 42] such as the oxidation of dilute VOCs. However, the system is more complex and requires larger amounts of catalyst than the simple reverse-flow reactor. 10.4.8.2

Inlet Temperature Modulation The modulation of temperature to the inlet of an adiabatic fixed-bed reactor provides a conversion improvement for reversible exothermic reactions [9]. A cycle average inlet temperature under modulation can be substantially lower than the inlet temperature in a steady-state operation. This leads to a lower outlet temperature and a higher equilibrium conversion for a reversible reaction. A better performance is achieved if the temperature oscillations attenuate on passage through the catalyst bed [9]. The process requires a more complex flow diagram and control than the reverse-flow operation. 10.4.8.3

10.4.9

Flow Modulation in Three-Phase Reactors

The liquid-flow modulation in cocurrent trickle-bed reactors has been explored in two frequency ranges [43]. Fast modulation at frequencies above 1 Hz can extend the natural highly intensive slug or pulsing flow regime to the References see page 2172

2170

10.4 Unsteady-State Reactor Operation

region of low liquid and gas flow rates normally associated with a less-intense trickling regime. Advantageously, the flow modulation at low mean liquid velocity does not incur a high pressure drop typical for the natural slug regime. The results of a few laboratory studies [43] performed to date have indicated that this unsteady-state operation can improve performance for liquid mass transfer-limited processes. Low-frequency modulations at frequencies below 0.01 Hz are usually undertaken by the periodically interrupting liquid flow in a cocurrently operating trickle-bed reactor. This unsteady-state operation is applied for lowintensity trickle-bed processes limited by gas-phase transfer [43]. Flow interruption either eliminates or severely reduces the liquid film surrounding a particle, thereby enhancing mass transfer to the particle external surface. Stopping and re-starting the liquid flow improves the catalyst irrigation and raises the volume mean reaction rate. Finally, flow interruption results in higher operating temperatures for exothermic reactions because this retards cooling of the catalyst by the liquid stream. It is important that temperature run-away does not occur, although the mean trickle-bed operating temperature will rise. All of these effects contributed to a remarkable fourfold increase in the rate of cumene formation in the hydrogenation of α-methyl styrene over a supported palladium catalyst [44]. Low-frequency, on-off modulation has been applied to the oxidation of dissolved phenol, the hydrogenation of dicyclopentadiene and cyclohexene, and to the disproportionation of cyclohexene over a Pd catalyst, as well as the adsorption and oxidation of SO2 over a carbon catalyst [45]. In all of these applications, improvements in reaction rate and selectivity over the steady-state operation were observed, although these were heavily dependent on the cycle time, cycle split, and other parameters [46]. 10.4.10

Periodically Operated Adsorptive Reactors

A combination of reaction and adsorption units into a single adsorptive reactor may be advantageous for equilibrium-limited and selective chemical processes [47–57]. Adsorptive reactors usually employ mixed beds of catalyst and adsorbent, although occasionally a catalyst exhibiting a strong adsorption of one of the products can be used alone. Adsorption is an unsteady-state process, and a catalyst in a mixture with the adsorbent operates under unsteady-state conditions. If the adsorption capacity of a particular adsorbent is limited, a cascade of a catalyst bed followed by an adsorbent bed may be used. Graded beds can be applied, in which the adsorbent to catalyst ratio changes moving downward through the reactor.

Chromatographic Reactors The chromatographic reactor is the simplest adsorptive reactor, and was first described during the early 1960s [48]. It consists of a fixed-bed column which is fed continuously with a carrier gas into which reactant pulses are periodically injected. Adsorption with a continuous flow of a carrier gas separates the reactant and products one from another. For an equilibrium-limited reaction such as A ← → B + C, if one of the reaction products (B or C) is adsorbed more strongly than the other, then separation of the products prevents the reverse reaction. Consequently, the reaction can be forced to completion. The frequent injection of a pulse leads to the faster-moving, more weakly adsorbed product overtaking the more slowly moving product, thus destroying the separation. This frequency limitation reduces the capacity or throughput of a chromatographic reactor. One solution to the throughput problem is the continuous countercurrent, moving-bed chromatographic reactor [49] (see Fig. 4a). This design employs simultaneous countercurrent flow of the solid catalyst–adsorbent mixture and the fluid phase. Several attempts to build and operate a moving bed enjoyed limited success [47], mainly due to problems associated with uneven of the solid flow and attrition. An alternate technique of moving the solid phase is to place the mixture of catalyst and adsorbent in an annular bed, and to rotate this bed past multiple annular inlets for carrier fluid and a single inlet for reactants. At the bottom of the bed there are discharges for the reaction products and multiple take-off points for the waste carrier fluid, which can be recycled to the reactor inlet (see Fig 4b). This design, known as a continuous rotating chromatographic reactor, is well developed and used commercially in biochemical applications [47, 50]. The throughput advantages of the continuous countercurrent moving-bed chromatographic reactor may be obtained without the problems caused by transporting solids through using a simulated countercurrent movingbed chromatographic reactor (SMBR) [47, 50, 51], an unsteady-state operation. The operation of a moving bed is simulated by sequentially changing the function of each port from inlet to outlet and back again. The arrangement of the ports at any point in time is as shown in Fig. 4c; namely, that the reactant enters the system between the extract port that carries away the weakly adsorbed product and the raffinate port that recovers the strongly adsorbed product. The carrier gas or liquid, often called an ‘‘eluent’’, enters above the raffinate port. Changing sequentially the function of the ports between the beds simulates the countercurrent movement of solid and fluid phases in a discrete manner. The four functions – reactant feed, extract and raffinate recovery, and eluent 10.4.10.1

10.4.10 Periodically Operated Adsorptive Reactors

feed – divide an SMBR into four sections, each of which contains at least one bed of mixed catalyst and adsorbent, though frequently more beds are used. The ports switch function after a switching time, tswitch . Experimental investigations with SMBRs have been conducted for liquid-phase processes [47, 50] such as acetic acid esterification with methanol, ethanol or β-phenethyl alcohol (using an ion-exchange resin as a catalyst), the synthesis of methyl tert-butyl ether (also using an ion-exchange resin), and the synthesis of bisphenol-A from acetone and phenol. Enzymecatalyzed reactions, such as sucrose inversion and hydrolysis of lactose, have also been studied. For strongly equilibrium-limited esterification reactions, the use of SMBRs results in complete conversion of the initial reactant (acetic acid) and the acquisition at the raffinate port of a desirable product (ester) which is completely separated from another product (water). Examples of studies with gas-phase reactions, such as the oxidative coupling of methane or hydrogenation of mesitylene [51], demonstrated a significant increase in both process yield and selectivity. The design of both moving-bed and simulated movingbed chromatographic reactors – and, indeed, of all separating reactors – is based on a general materials balance that requires the formulation of multi-component adsorption equilibrium models, kinetic models, and considers mass transfer, and backmixing. An elegant simplified theory [50] was devised which neglected all of the dissipation phenomena and assumed that product separation controls the process, so that sizing of the SMBR is very similar to the corresponding chromatographic separator. The technique was subsequently developed for the design of separators [52] and is also applied for the design of SMBRs. The critical factors in the design are the conversion and separation in Section 3 (see Fig. 4c), between the feed point and the withdrawal of extract, as well as in Section 2, between the feed point and the withdrawal of raffinate. Dimensionless ratios between the fluid and the solid velocities in these two central sections (designated as m2 and m3 ), along with the dimensionless residence times or Damk¨ohler numbers (Da2 and Da3 ), govern the performance of the unit. The hypothetical solid velocity for the SMBR [usolid ; (m s−1 )] is inversely proportional to the port switching time tswitch , usolid = (Lbed /tswitch ), where Lbed is the length of the bed (m). A triangle diagram [50, 52] constructed on the plane defined by velocity ratios in two central sections (m2 , m3 ) for a given residence time (at constant Da2 and Da3 ) establishes an area corresponding to complete conversion of reactants and separation of the products.

2171

Flow-Reversal Adsorptive Reactors If the bed contains an adsorbent, or if the catalyst itself has adsorptive properties, the reverse-flow reactor (see Section 10.4.8.1) allows for the selective trapping of adsorbed compounds in the same way as the hot zone is trapped for exothermic reactions. The high ammonia adsorption capacity of a vanadium oxide catalyst has been employed for selective NOx reduction in a reverse-flow reactor (see Ref. [53]). This process also utilizes periodic changing of the inlet gas composition. In the first part of a cycle, NH3 injected in stoichiometric excess fills the bed, thus creating an adsorption wavefront that gradually moves toward the outlet. In the second part of the cycle, no ammonia is supplied, although the waste gas containing nitrogen oxide continues to enter the bed. NOx reacts with the excess of ammonia previously adsorbed, after which the valves are switched and the first and second parts of the cycle are repeated, but with a different gas-flow direction. Due to the ammonia stored on the catalyst, the process is relatively insensitive to fluctuations of gas flow and NOx concentration and can provide for very low ammonia slip. Other examples for a potential beneficial combination of adsorption and reaction include sulfur recovery by the Claus process combined with water adsorption [54], and the two-stage Deacon process for hydrogen chloride recycling over a copper oxide catalyst [55]. 10.4.10.2

Pressure-Swing Reactors All chromatographic reactors employ flushing with an eluent or carrier gas to desorb the products, and thus regenerate the adsorbent. However, desorption can also be performed either by reducing the pressure or raising the temperature. Separating reactors which employ pressure reduction are referred to as ‘‘pressure swing reactors’’, while if a temperature increase is used they are known as ‘‘temperature-swing reactors’’. Both types of swing reactor employ a fixed-bed of catalyst mixed with adsorbent, and both resemble closely (in terms of operation) their respective separating systems. The pressure swing separating systems are used widely in industry for large throughputs. For example, the fixed-bed reactor is operated under periodically changing pressure (see Fig. 5), the simplest cycle including three steps: 10.4.10.3

• feeding of the reactor at high pressure, with simultaneous low-pressure delivery of a fraction of the gaseous mixture from the opposite end of the reactor • a delay, when only delivery is allowed • exhaust, when the reaction products are emitted at low pressure from both ends of the bed. Theoretically based studies [56] have shown that conversion and selectivity improvements can be expected References see page 2172

2172

10.4 Unsteady-State Reactor Operation

in the case of single or multiple reversible reactions. The high-pressure pumping of reactants enhances separation and increases throughput, although pressure swing reactors, in general, can only separate cleanly the weakly adsorbed reaction product. The strongly adsorbed product is contaminated by unconverted feed and some of the weakly adsorbed product. When the adsorption affinities of the products are quite similar, the separation and purity of the extract can be substantially improved by using a four- or five-step cycle. These extra steps include back-flushing of the bed with the delivery stream product and, if needed, repressurization of the mixed catalyst and adsorbent bed by this product. However, throughput is reduced when this approach is taken. Pressure swing reactors have been used experimentally to produce hydrogen by steam reforming of natural gas [47]. A synthetic dolomite, ‘‘hydrotalcite’’, adsorbs the CO2 formed in reforming so that the delivery stream contains only hydrogen and steam. A pilot pressure swing reactor used a five-step cycle which employed steam purging and repressurization. Unfortunately, the adsorption capacity of hydrotalcite for CO2 was lowered such that the reactor performance deteriorated markedly in successive cycles.

low inlet temperature, several steady states may exist in a wall-cooled, tubular reactor. The conventional steadystate operation is possible only for two outer stable states, which are often unsuitable because of very high temperatures or very low conversion of the reactant. However, appropriate periodic variations of the inlet concentration, flow rate, or temperature of the cooling medium allow an unstable middle temperature to be attained. In particular, this approach may be useful for partial oxidation, where the activation energy for complete oxidation is higher than that for the desired reaction. An example, namely ethylene oxidation over a silver catalyst, was discussed in Ref. [59]. The control system necessary to stabilize the unstable state should be of a positive feedback type [60]. In another example [61], a circulating loop reactor was developed and applied to the exothermic oxidation of VOCs. The reactor consisted of a cocurrent heat exchanger filled with a catalyst, and the inlet fraction of the catalyst bed was heated up by the outlet gas through the exchanger wall. The reactor was unstable within a broad range of operating parameters, so that a migrating reaction zone periodically initiated at the catalyst bed inlet.

NOx Storage Reduction Systems Recently developed for automotive emission control, the NOx storage and reduction (NSR) converters [57, 58] represent a unique example of a successfully commercialized unsteady-state reaction–adsorption process. The automotive engine is tuned to operate in a lean mode with a large air excess that provides for substantial fuel economy compared to the stoichiometric combustion usually applied in gasoline-fueled cars. For NOx removal the engine is periodically switched to stoichiometric or fuel-rich combustion that generates pulses of oxygen-depleted exhaust. The bifunctional NSR catalyst combines a noble metal (Pt) catalyst and an NOx adsorbent based on an alkali or alkali-earth metal. During the lean combustion cycle the adsorbent traps NOx and, simultaneously, the catalytic oxidation of NO to NO2 increases the trapping efficiency. During the rich fraction of the cycle, the adsorbent is regenerated by oxygen-depleted exhaust gas and, simultaneously, the released NOx is reduced to nitrogen over the catalyst. The catalyst and process details are reviewed in Ref. [58].

References

10.4.10.4

10.4.11

Stabilization of Unstable Steady States

Periodic changes of inlet parameters can be applied to maintain an intrinsically unstable state of a chemical reactor. Under conditions of high heat of reaction and

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10.5

Short Contact-Time Reactors .. Gotz Veser∗

10.5.1

Introduction

Short contact-time reactors (SCTRs) and reactions are catalytic processes in which the superficial contact time between the reactants and the catalyst is in the range of about 100 ms (10−1 s) to well below 1 ms. This approach to conducting a reaction has undergone rapid development during the past decade, based largely on the seminal studies of Schmidt and coworkers since the early 1990s [1–6]. With very few exceptions, short contact-time reactions typically involve the conversion of saturated, unsaturated, or oxygenated hydrocarbons with oxygen in the presence of a noble metal-based catalyst at very high reaction temperatures (T > 1000 K). Through the use of short catalyst beds with lengths lcat in the order of 10−4 to 10−2 m, and very high linear gas flow velocities of u ≈ 0.1 to 10 m s−1 , catalyst contact times of τ = lcat /u = 10−5 to 10−1 s are attained in these systems. It should be noted that, as the calculation of exact contact times would require a detailed knowledge of temperature and species profiles throughout the catalyst zone, investigations of SCTRs often state reactor space-time yields or so-called ‘‘gas-hourly space velocities’’ (GHSV); that is a reciprocal nominal residence time based on feed gases at standard conditions and total catalyst reactor volume. These space-time yields are in the order of several 10 000 h−1 to well above 107 h−1 . These extreme reaction conditions, however, render detailed investigations of the fast kinetics coupled with heat and mass transport based on steep temperature and concentration gradients extremely difficult, and a true understanding for short contact-time processes has only begun to emerge in recent years. The aim of this chapter is to illustrate some of the fundamental characteristics of SCTRs and ∗

Corresponding author.

reactions, using select examples from the literature. A comprehensive review of the field is neither within the scope nor the aim of this chapter. 10.5.2

Brief History

In spite of all the attention that SCTRs have received in recent years, they are not new. In fact, some of the oldest large-scale processes in the chemical industry are typical SCTRs: for example, NO production by catalytic ammonia oxidation (the ‘‘Ostwald process’’) was developed at the beginning of the 20th century and is thus among the oldest large-scale industrial processes. In this reaction, ammonia is oxidized in excess air at temperatures between 600 ◦ C and 700 ◦ C according to: 4NH3 + 5O2 −−−→ 4NO + 6H2 O (HRo = +906 kJ mol−1 ) Since the reaction product NO is not thermodynamically stable at reaction temperatures, the reaction is conducted over thin layers of Pt−Rh-gauzes, which results in contact times in the order of a few milliseconds and leads to a fast quenching of the metastable NO. The production of hydrogen cyanide via ammoxidation of methane (the ‘‘Andrussow process’’) was developed during the 1940s and proceeds at nearly identical process conditions, albeit at significantly higher temperatures (∼1000–1200 ◦ C) [7]: CH4 + NH3 + 1.5O2 −−−→ HCN + 3H2 O (HRo = −473 kJ mol−1 ) Again, the incentive for using a SCTR is the need to isolate an unstable reaction product (HCN) by rapid quenching. Based on this process, Degussa later developed the ‘‘BMA process’’ (from the German ‘‘Blaus¨aureMethan-Ammoniak’’, i.e., ‘‘hydrogen cyanide-methaneammonia’’), which is the oxygen-free conversion of methane with ammonia at slightly higher temperatures (∼1200–1300 ◦ C) and otherwise similar process conditions [8]: CH4 + NH3 −−−→ HCN + 3H2 (HRo = +251 kJ mol−1 ) While the BMA process can be regarded as the Andrussow process without oxygen addition, the complementary process – that is, the conversion of methane with oxygen but without ammonia addition – was first investigated during the 1940s, concurrently with the development of the Andrussow process [9]. The intent of this reaction

10.5.3 Short Contact-Time Reactions

route is the production of synthesis gas, a mixture of CO and H2 – that is, the catalytic partial oxidation of methane, rather than combustion: CH4 + 0.5O2 −−−→ CO + 2H2 (HR = −37 kJ mol−1 ) However, the contact times in those early studies were well outside the short contact range (≈1–50 s), and problems with the Ni-based catalysts due to strong temperature excursions at the entrance to the catalyst bed were most likely responsible for a lack of further investigation. While a homogeneous (non-catalytic) process was subsequently developed and commercialized by Texaco and Shell [8, 10], the catalytic process route lay essentially dormant until the early 1990s. During the early 1990s, a number of virtually simultaneous publications from several research groups around the world demonstrated the feasibility of the catalytic partial oxidation of methane (CPOM) over a range of catalysts and process conditions [1, 11–17], and this led to the subsequent explosive development of SCTRs, in particular for the production of hydrogen and synthesis gas from alkane feedstocks. Some of the main results of these studies will be discussed in the following section to illustrate the current state of the art for SCTRs. 10.5.3

Short Contact-Time Reactions

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on the integration of large amounts of process off-heat into the network of a large chemical production site. This limits the applicability (or the efficiency) to remote resources and locations where such networks may not be available. The CPOM overcomes this problem as the reaction route is mildly exothermic and the process can hence be operated autothermally without the need for additional heat supply. It has furthermore the advantage that it yields a stoichiometric ratio of H2 : CO = 2, which is the desired ratio for major downstream consumers such as methanol and Fischer–Tropsch synthesis. For thermodynamic reasons, CPOM must be operated at very high reaction temperatures: at temperatures above about 1200 K, virtually no thermodynamic limits exist on syngas yields for a stoichiometric feed gas mixture of CH4 : O2 = 2.0 (Fig. 1). However, due to competing total oxidation reactions, which stoichiometrically consume more oxygen, the best syngas yields are in practice achieved at a slightly lower ratio of CH4 : O2 ≈ 1.6–1.8. Most importantly, these high reaction temperatures, in combination with the intrinsically high reaction rates of catalytic oxidation reactions, allow the process to be operated at very short contact times in the order of 10 to 50 ms, in contrast to the typical contact times in the order of 1 s for a steam reformer. Despite the fact that syngas is an equilibrium product, however, the amount of total oxidation occurring in the system must be minimized in order to achieve (near) thermodynamic equilibrium

Hydrogen and Syngas Formation from Methane Since the early 1990s, the increased interest in new routes for the conversion of methane to synthesis gas (a mixture of CO and hydrogen) has been based on the need for more efficient utilization of the increasingly expensive natural gas resources, and the desire to access natural gas resources which are either too remote or too small to be efficiently exploited with existing technologies. As all attempts to convert methane directly into methanol or other usable chemicals have so far not resulted in a commercially viable process, much effort has focused on the further development of indirect methane utilization via the syngas route [18, 19]. The main disadvantage of the dominant industrial syngas production via steam reforming of methane is the very large amount of heat exchange required in tubular reformers to fire the highly endothermal steam reforming route, and the associated high capital costs. 10.5.3.1

CH4 + H2 O −−−→CO + 3H2 (HRo = +206 kJ mol−1 ) Furthermore, while commercial steam reformer units have a very high thermal efficiency, this efficiency is based

1.0

0.8

Yield

0.6

0.4 YH2 0.2 YCO 0.0 500

1000

1500

2000

T/K Yields of synthesis gas (CO and H2 ) from thermodynamic equilibrium calculations for a methane/oxygen mixture with CH4 : O2 = 2.0. Calculations were conducted using STANJAN.

Fig. 1

References see page 2186

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10.5 Short Contact-Time Reactors

yields within the millisecond contact times in a SCTR, as secondary reactions such as steam reforming or water-gas shift reactions are significantly slower than the primary oxidation reactions. Although many academic and industrial groups are currently actively investigating CPOM, the groundbreaking studies were conducted during the early 1990s by Schmidt and coworkers [1, 5, 16, 17, 20, 21]. Their studies demonstrated the feasibility of the CPOM route in an exceedingly simple tubular reactor using a Rh-coated monolithic catalyst support. Hickman and Schmidt demonstrated syngas selectivities in excess of 90% at complete oxygen conversion, and >90% methane conversion for a process operated with pure oxygen and mildly preheated feed gases [1, 5, 16, 20]. Subsequent studies improved these yields up to methane conversions of ≈96% with syngas selectivities of ≈98% via careful insulation of the compact reactors, sufficient pre-heat of the feed, and the use of sufficiently active catalysts [22]. Even for the air-blown CPOM process, methane conversions of up to 90% at syngas selectivities of ≈95% could be realized to date, through carefully designed catalysts and optimized reactor concepts (see also Section 10.5.5) [23]. Despite all these advances, however, the prospects for an industrial implementation of CPOM are still unclear, which demonstrates how difficult market penetration is for a new technology in a field with well-established technologies. Hydrogen and Syngas from Higher Alkanes The concept of short contact-time catalytic oxidation was subsequently extended by Schmidt and others to the partial oxidation of higher hydrocarbons, ranging from ethane up to diesel components, as well as the partial oxidation and autothermal reforming of alcohols. The conversion of short-chain hydrocarbons, in particular of ethane, to synthesis gas is of interest as these hydrocarbons form minor components of natural gas. Hence, for a commercially viable process of converting natural gas to syngas, these components must be converted to synthesis gas with similar efficiency at the same conditions as CPOM. Huff and Schmidt demonstrated that this is indeed possible in a SCTR: at essentially identical conditions to CPOM, these authors found that the partial oxidation of ethane over Rh-coated alumina monoliths gave ethane conversions of up to 95% with syngas selectivities of about 70%. Similar results were found for propane and butanes [24, 25]. Subsequently, syngas formation via partial oxidation of long-chain hydrocarbons such as octane, decane, and hexa-decane was demonstrated at near-identical SCT conditions [26, 27]. These components are of interest as model transportation fuels, where – beyond hydrogen production for fuel cells from these fuels – the 10.5.3.2

preconversion of some fuels is of interest for cleaner combustion processes as the addition of hydrogen has been shown to result in increased efficiency and cleaner combustion [28]. Hydrocarbon conversions of about 90% with syngas selectivities of about 80–85% were demonstrated for all of these hydrocarbon fuels. Interestingly, Schmidt and others found that for hydrocarbons beyond methane, the selectivity of short contact-time partial oxidation can be switched from mainly syngas-forming reactions to olefin production through changes in catalyst selection [24, 29–34]. Production of Olefins The application of SCTR to the production of olefins from ethane and higher hydrocarbons seems, at first glance, to be a simple and straightforward extension of CPOM. If short contact-times favor the production of the partial oxidation product synthesis gas in the catalytic oxidation of methane, then it might be expected that sufficiently short contact times and otherwise similar conditions should yield some amount of olefins from higher alkanes. In contrast to synthesis gas, however, olefins are not present to any significant degree at equilibrium for any temperature range in an oxygencontaining atmosphere. Hence, for SCTR to produce olefins efficiently, the catalytic (or non-catalytic) reaction must yield olefins as highly reactive primary oxidation products, and the SCT conditions must result in rapid quenching of this non-equilibrium composition of the product stream. Indeed, Schmidt and others showed that the catalytic conversion of alkane–oxygen mixtures under SCT conditions can produce very high yields of olefins [32, 35–38]. Here, we will focus on the most simple of these reaction systems – the production of ethylene (ethene) from ethane – to illustrate the main features of this class of SCT reactions, and particularly the differences to the above-discussed CPOM reaction. Results for the conversion of ‘‘higher’’ alkanes (>C3 ) will be summarized only briefly. Ethylene (like other small olefins) is a key intermediate in the petrochemical industry, and is currently produced industrially by steam-cracking of ethane. In this reaction, a pre-mixed and pre-heated ethane–steam mixture is fed to an externally fired multitubular reactor where it is thermally cracked at temperatures around 750–850 ◦ C to ethylene and H2 : 10.5.3.3

C2 H6 −−−→ C2 H4 + H2 ,

HRo = +136 kJ mol−1

The strong endothermicity of this reaction requires very high heat-transfer duties and hence very high wall temperatures in the fired tubular reactor, requiring expensive specialty steels. Coke deposition inside the tubes (which is somewhat reduced, but not avoided,

10.5.3 Short Contact-Time Reactions

by cofeeding large amounts of steam) further reduces the heat transfer. Finally, NOx formation in the furnace causes environmental concerns. In all of these regards, the oxidative dehydrogenation (ODH) of ethane is an attractive alternative. The reaction is exothermal due to the oxidation of the hydrogen, thus allowing for autothermal process conditions: C2 H6 + 0.5O2 −−−→ C2 H4 + H2 O, HRo = −106 kJ mol−1 Hence, no external heat transfer is necessary, which drastically simplifies the reactor, reduces the maximum temperatures, eases materials problems, and completely avoids NOx issues. However, debate persists as to whether the use of pure oxygen for this reaction may reduce or even completely obliterate the economic incentive that results from these advantages [39]. Shortly after their initial report on CPOM at SCT conditions, Huff and Schmidt reported the extension of their SCT approach to oxidative dehydrogenation of ethane to ethylene [35]. Using the same reactor configuration (a noble metal-coated alumina monolith reactor) as for their CPOM studies, these authors found that while Rh produced essentially only syngas from ethane, Pt gave very high olefin yields at slightly higher C : O feed ratios and otherwise identical SCT conditions. They took this as a strong indication that, like CPOM, the reaction must be truly catalytic. Carbon-based ethylene selectivities of ≈65% at 70% ethane conversion were reported at contact times in the order of 1 to 10 ms and slightly oxygen-rich conditions (C2 H6 /O2 = 1.7) [35]. Ethylene selectivity as well as ethane conversion could subsequently be improved by about 5% through adding Sn as additional component to the Pt/alumina catalyst [40], although this catalyst showed a slow loss of activity due to Sn volatilization [41]. A remarkable improvement in ethylene yield was found when adding H2 to the reactor feed [2]. At unchanged ethane conversion of 70%, the ethylene selectivity could be improved to about 85%, making ethylene yields competitive with the industrial steam cracking route. As almost all the H2 that is fed to the reactor is recovered in the product stream – that is, the reaction produces as much hydrogen as it consumes – H2 can be regarded as another (homogeneous) ‘‘catalyst’’ for this reaction system and, more importantly, the cost of cofeeding hydrogen should not pose an economic concern. Subsequently, the same group showed that a number of higher hydrocarbons could also be converted to olefins in the same SCTR, with similar catalysts and similar reaction conditions. The oxidative dehydrogenation of propane [37], butanes [37, 38], pentanes and hexanes [32], and even decanes [42] produced a range of olefins, although the product spectrum broadened with increasing

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hydrocarbon chain length. Generally, higher oxygen contents in the feed gas stream favored short olefins (mainly ethylene), while lowering the oxygen content increased the olefin chain length up to the parent olefin, albeit at the expense of alkane conversion. Production of Oxygenates The concept of SCTRs was most recently extended to the formation of oxygen-containing hydrocarbon derivatives (so-called oxygenates), such as alcohols, aldehydes and ketones, although this group of reactions has to date been studied to a much lesser degree than the abovediscussed production of synthesis gas or olefins. Current interest in oxygenate production is largely driven by the demand for clean combustion engines. The phasing out of lead additives since the 1970s first increased the demand of methyl tert-butyl ether (MTBE). However, since the discovery during the 1990s that MTBE accumulates in drinking water, interest in other oxygenate additives to gasoline and diesel fuel has increased. In 1996, Schmidt and coworkers demonstrated that, by reducing the contact times in the catalytic partial oxidation to the microsecond range, propane, butane, and pentane could be converted into a range of oxygenated products, mainly aldehydes [3]. In contrast to previous studies, a single Pt−10%Rh gauze woven from 90 µm-thick wires was used as catalyst, and this resulted in contact times of approximately 200 to 500 µs. Under these conditions, hydrocarbon conversions were less than 40% – that is, significantly lower than for the above-discussed syngas and olefin processes, although oxygen conversion was still about 80% at all conditions. Due to the low conversions, temperatures were also significantly reduced to about 650–900 K. Under these conditions, up to 60% selectivity to oxygenates could be attained, largely a mixture of C1 –C3 aldehydes [43, 44]. The highly complex nature of these heterogeneous–homogeneous reactions (see Section 10.5.4.4) is most likely the reason for a lack of detailed follow-up studies by the same or other research groups to date. Clearly, there is significant undeveloped potential in this process route. 10.5.3.4

Ketenes from Carboxylic Acids All of the above-discussed reaction systems constitute mildly exothermic partial oxidation reactions which can be conducted autothermally in a SCTR. Indeed, the vast majority of studies with SCTRs to date has been conducted on these partial oxidation reactions. However, Barteau and coworkers have recently shown that the concept of SCT conditions can also be applied to endothermic reaction 10.5.3.5

References see page 2186

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10.5 Short Contact-Time Reactors

systems [45, 46]. These authors used functionalized silica monoliths to dehydrate carboxylic acids to ketenes according to: R−CH2 −COOH −−−→ R−C=C=O + H2 O As these dehydration reactions are endothermic, the tubular reactor was externally heated to temperatures between 650 K and 850 K – that is, a similar temperature range to the oxygenate process discussed above. However, contact times were comparatively long at 10 to 100 ms. The authors demonstrated the dehydration of acetic acid, isobutyric acid, hexanoic acid and octanoic acid to the corresponding ketenes, with selectivities around 90% and conversions of up to 90%. Comparison with empty tube experiments showed that homogeneous reactions play only a minor role, and that the availability of surface hydroxyl groups (i.e., a functionalized silica catalyst with large surface area) is crucial for the attainment of good conversions. Increasing the contact time led to significantly decreasing ketene selectivities, presumably due to further decomposition of the unstable ketenes. Thus, the main purpose of the SCT conditions in these systems is rapidly to isolate (quench) an unstable reaction product before further reactions can occur. 10.5.4

Reaction Mechanisms and Catalysts for SCTR

In the following sections, some of the catalysts used in the above-presented SCTR will be described, along with the current status of understanding for the reaction mechanisms being utilized. Catalysts for Syngas Formation The requirements for a suitable catalyst for CPOM are dictated by the harsh conditions at which these processes are operating. Typically, the catalyst must be stable at very high reaction temperatures, it must be highly active in order to allow catalytic reactions to precede homogeneous gas-phase reactions (which are always possible at these temperatures), and it needs to ascertain a low pressure drop at very high flow rates. The combination of these requirements is best met by noble metals as the active component, as these combine robustness with high activity for hydrocarbon oxidations, and ceramic monoliths as supports, which combine high-temperature stability with a low pressure drop. Correspondingly, noble metals supported on alumina monoliths (extrudates or foams) are the preferred and most studied catalyst configuration for SCTR in general, and for CPOM in particular [1, 5, 16, 47–51]. Examples of these materials are illustrated in Fig. 2. 10.5.4.1

(a)

(b)

Typical catalysts for short contact-time reactions. Left: extruded straight-channel alumina monoliths before and after coating with noble metal. Right: a Pt-coated alumina foam monolith.

Fig. 2

While many catalysts have been screened, Pt and Rh are the most studied systems, with Rh being generally reported to be the most selective catalysts for syngas formation [1, 12, 16]. Based on elementary-step kinetic modeling, Hickman and Schmidt explained the difference in performance between Pt and Rh by the difference in the activation energy for OH formation on the two catalyst surface which makes hydrogen oxidation less favorable on Rh catalysts [1]. Among the many other metals studied, only Ru and Ni are generally reported to show similarly high activity with good selectivity and stability at SCT conditions [48, 52–56]. However, some disagreement exists with regard to the stability of the different catalysts [25], with volatilization of the metal component and coking being the main reasons for rapid deactivation seen in many cases. Different supports and support structures were also tested for this reaction, although the influence on syngas yields were generally minor [57] as long as the support is an irreducible oxide [58]. Hohn and Schmidt reported that Rh-coated alumina spheres outperformed the Rh-coated alumina monoliths, despite comparable surface areas and preparation procedures [59]. They explained this observation by the higher convective heat transport in flow direction and lower conductive heat transport against the flow direction in the monolith reactor; this results in earlier blow-out of the reaction front in the monolith reactor and hence lower syngas yields. While the pressure drop is higher in the fixed-bed reactor compared to the monolith reactor, this may not raise concerns due to the shallowness of the required bed (≈5–10 mm). The importance of heat conductivity for this reaction system may also explain recent results with noble metals supported on Si3 N4 [60] and SiC [61]. Beyond these comparatively simple catalysts, some more-complex catalyst formulations have been tested in recent years. Promising results were found with

10.5.4 Reaction Mechanisms and Catalysts for SCTR

hydrotalcite-supported [62] and with hexaaluminatesupported noble metals [63–65]. In both cases, a high dispersion of the active metal led to improved catalyst performance. Schicks et al. found that nanocomposite noble metal–hexaaluminate catalysts with very high surface areas and good high-temperature stability gave syngas yields far above those of alumina-supported metals, allowing for a reduction in the noble metal content of the catalyst bed by up to two orders of magnitude [64, 65]. Most recently, Mitri et al. found that the differences in activity and selectivity between different aluminasupported metal catalysts virtually disappeared upon preheating the reactant gases (see Fig. 3) [66]. These findings question the relevance of the catalyst’s nature for conversion and selectivity, and also questions the reaction mechanism occurring under these conditions. Syngas Formation: The Reaction Mechanism The reaction mechanism for the catalytic partial oxidation of methane to synthesis gas at SCT conditions has been the subject of much debate that even today has not been completely settled. Due to the extreme reaction conditions, SCTRs are not amenable to typical spectroscopy-based mechanistic in-situ investigations at 10.5.4.2

100

90

SH2 / %

80

70

60

RFR:

50

SS:

40

1.0

Pt Rh Ir Ni Pt Rh Ir Ni

1.5

2.0

2.5

CH4 /O2 Hydrogen selectivities in catalytic partial oxidation of methane in a conventional steady-state tubular reactor (‘‘SS’’, open symbols and dotted lines) and with preheating in a heat-integrated reverse-flow reactor (‘‘RFR’’, full symbols and solid lines) for Pt, Rh, Ir, and Ni catalysts. The difference in catalytic performance is strongly reduced through preheating. (Adapted from Ref. [66].)

Fig. 3

2179

process conditions. Hence, studies must be conducted at ‘‘unrealistic’’ conditions (lower flow rates, pressure and/or temperatures), or conclusions about the mechanism must be inferred from ‘‘global’’ measurements (inlet versus exit concentrations, temperatures). Furthermore, as the reaction systems are highly sensitive even to small changes in the experimental set-up, great care must be taken to achieve well-defined experimental conditions in order to avoid apparent contradictions with results from previous studies. A focus of the existing disagreement has surrounded the issue of direct versus indirect reaction pathways in the formation of synthesis gas via CPOM in a SCTR. In direct oxidation, synthesis gas is formed in a single step via CH4 + 0.5O2 → CO + 2H2 . In the indirect oxidation mechanism, part of the methane feed is initially combusted (CH4 + 2O2 → CO2 + 2H2 O), and synthesis gas is subsequently formed via steam reforming of the remaining fuel (CH4 + H2 O → CO + 3H2 ), as well as some water-gas shift reaction (CO + H2 O → CO2 + H2 ). Hickman and Schmidt suggested that syngas is formed at SCT conditions exclusively via direct oxidation, based largely on the striking difference between their results and previous studies. Previously, the syngas yields showed a strong decrease with decreasing contact time, while Hickman and Schmidt observed an initial increase followed by a broad plateau in their yields [5, 67]. They explained this observation based on a direct oxidation mechanism, in which methane is pyrolyzed on the catalyst surface to C and H atoms, H radicals recombine to molecular H2 which desorbs from the catalyst, and the remaining carbon reacts with (dissociatively adsorbed) oxygen to CO. In contrast, several studies have claimed that indirect oxidation must be responsible for the formation of synthesis gas. Typically, evidence from these studies is based on the observed temperature profiles in the catalyst bed, or on the observation of reduced syngas yields and increasing total oxidation products with decreasing contact times [11, 12, 47, 68–71]. The results of a recent study in our own laboratory led to the conclusion that the distinction between indirect and direct oxidation pathways might be an artificial distinction, and that much of the previous disagreement may be due to different reaction conditions. This was proposed because results from carefully controlled experiments varying residence time over Pt/alumina monoliths at autothermal conditions indicated that CO is in fact predominantly a direct oxidation product, while H2 is formed predominantly via indirect oxidation (i.e., steam reforming of methane; see Fig. 4) [72]. References see page 2186

2180

10.5 Short Contact-Time Reactors

CH4 H2

Relative concentration /a.u.

0.5

0.4

CO

0.3

O2 0.2

H2O 0.1

CO2 0.0 0

1

2

3

4

5

t /10−3 s Relative concentrations of reactants in catalytic partial oxidation of methane with CH4 /O2 = 1.8 as a function of catalyst contact time. The S-shape of the hydrogen curve is typical for a sequential (secondary) reaction product, while all other products (CO, CO2 , and H2 O) show typical behavior for primary reaction products. This suggests that hydrogen is a product of indirect oxidation, while CO is predominantly a product of direct oxidation. (From Ref. [72].) Fig. 4

The relevance of these reaction mechanisms is put into question, however, by claims that CPOM at SCT conditions is dominated by mass transport limitations, heat transport limitations, or a mixture of both [73–75]. Correspondingly, many numerical simulation studies showed that syngas production over different catalysts at SCT conditions could be successfully modeled with reactor models which assume no transport limitations [76–78], a completely mass transport-limited system [79], or even a strongly simplified two-zone equilibrium model [80]. No resolution of these contradictory results, other than the differences in model assumptions and reaction conditions, is currently available. In this context, much effort has also been devoted to the development of elementary-step kinetics and detailed reactor models [20, 49, 50, 76–78, 81–85], although at this point much uncertainty remains with regard to the accuracy of the kinetic parameters used in these models [86]. Another important question that arises at the very hightemperature conditions of CPOM concerns the possible occurrence of homogeneous gas-phase reactions. In a typical catalytic process, the catalyst is used to allow a chemical reaction at conditions at which homogeneous (uncatalyzed) reactions are too slow to be of practical relevance. The extremely high temperatures of CPOM, however, open virtually any reaction pathway – both

thermodynamically, due to the large entropic contribution to the Gibbs energy of reaction, and kinetically, by overcoming even the large activation energy barriers of some of the homogeneous reactions. Schmidt and coworkers argued early on that the experimental results in CPOM studies are in agreement with an explanation purely based on catalytic chemistry [1, 5, 24], and a number of later studies – mostly based on detailed numerical simulations – confirmed this hypothesis [76, 77]. The absence of homogeneous contributions can be attributed to the relatively long ignition delay time characterizing methane ignition, which is in the same order as the residence time for SCTR [87]. Only at contact times below 10−3 s did the occurrence of C2 components in the product stream indicate the onset of post-catalytic homogeneous reactions. As contact times become insufficient for complete catalytic conversion in the catalyst zone, CH3 · radicals (possibly generated on the catalyst surface) undergo coupling reactions resulting in C2 components [88–90]. However, elevated pressures change the relative rates of catalytic and homogeneous reactions, as the first typically show a linear or sublinear dependency on pressure, while the latter scale typically roughly with the square of the reaction pressure due to the predominance of bimolecular reaction steps. Hence, it might be expected that the onset of homogeneous gas-phase reactions ultimately limits the applicability of CPOM towards higher pressures. Very few experimental studies have investigated this issue, but the few which did seem to indicate that, at least up to pressures around 2 × 106 Pa, catalytic reactions remain dominant [91], in fair agreement with predictions from numerical simulation studies [76–78, 85]. Simulations by Goralski et al. indicate that the process could be operated with air up to significantly higher pressures, without the onset of homogeneous reactions, though this would be achieved at significant expense in terms of space-time yield and, depending on the use of the syngas, necessitate an additional separation step downstream [76]. Catalysts and Reaction Mechanism for Olefin Production The topics of catalyst selection and reaction mechanism are intimately intertwined for this class of SCT reactions, and are hence discussed jointly. Initial studies of the ODH of ethane focused on noble metal catalysts, based on the results of CPOM at SCT conditions. In these studies, Pt was identified as the most selective catalyst, with Rh yielding mostly syngas, and most other noble metals showing either low activity or stability problems [35–37, 92]. Ethylene yields were shown to improve upon alloying the Pt catalyst with Sn or Co [40, 41]. These results – and in particular the 10.5.4.3

10.5.4 Reaction Mechanisms and Catalysts for SCTR

strong dependence of the ethylene selectivities on the choice of catalyst – were initially interpreted as indications that the process is a true catalytic process; that is, the ethylene is formed on the catalyst surface rather than in homogeneous gas-phase reactions. Since then, several studies have shown that similar ethylene yields could be obtained over chromia [93], doped rare earth oxides [94, 95], and even fairly conventional combustion catalysts [96], thereby casting serious doubt on the importance of the catalytic reaction for olefin formation. In agreement with this, Huff et al. showed, in simulation studies, that the reaction yields are in agreement with the assumption that the catalyst serves predominantly, if not exclusively, as a heat source for the homogeneous reaction [97] (see Fig. 5). This was further confirmed by a series of experimental studies by Beretta et al., which indicated that this heterogeneous–homogeneous mechanism, in which the catalyst mainly serves the purpose to ignite the homogeneous reaction, is responsible for the formation of olefins not only in the ODH of ethane but similarly also for ODH of propane [33, 98–101]. In further studies, it was found that the addition of a wash-coat onto the monolith walls [57], as well as the use of high-surface-area alumina pellets [102], strongly reduced ethylene selectivity. These observations agreed well with a heterogeneous–homogeneous mechanism, as increased catalyst surface areas increase the relative contributions of catalytic versus gas-phase reactions, and the results hence again indicate that catalytic reaction are not significantly contributing to the formation of olefins. Similar results were found for olefin production from long-chain hydrocarbons [42, 103]. In contrast to shortchain alkanes, Pt is a poor catalyst for syngas formation C2H4 + H2

C2H6 O2

Heat CO, CO2, H2O

C H O

CO CO2 H2O Catalyst (Pt)

Schematic of olefin production at short contact-time conditions, using ethylene production from ethane as example. All oxygen is consumed in some initial catalytic combustion of ethane, resulting in the formation CO, CO2 , and H2 O. The heat from this catalytic reaction then ignites the homogeneous dehydrogenation of the remaining ethane to ethylene and hydrogen. The catalyst thus serves as a catalytic ignitor for the homogeneous process.

Fig. 5

2181

as well as for olefin production from longer chain alkanes, due to a relatively low activity and hence low conversion. However, selectivity in the partial oxidation of long-chain hydrocarbons can nevertheless be tuned from olefin production to syngas formation by decreasing the porosity of the monolith, wash-coating the monolith, and increasing the catalyst downstream temperature. All of these effects again indicate that the olefin-producing reactions are largely homogeneous gas-phase reactions which, nevertheless, require the presence of an efficient catalytic ignitor. Overall, it appears therefore that the oxidative dehydrogenation of alkanes to olefins at SCT conditions is not a truly catalytic process, despite the crucial role that the catalyst plays as an ultra-fast ignitor of homogeneous reactions. Mechanistically, the process is hence not even a true oxidative dehydrogenation, but rather a catalytically supported pyrolysis. For this reason, Lange et al. suggested recently the term ‘‘oxycracking’’ to describe the oxidative production of olefins at high-temperature, SCT conditions, reserving ‘‘oxidative dehydrogenation’’ for the truly catalytic, low-temperature reaction over oxide catalysts [39]. Catalysts and Reaction Mechanism for Oxygenate Production Oxygenate production from alkanes proceeds exclusively over single gauze catalysts. However, due to the limited availability of such gauzes, studies have been restricted to the use of Pt–10%Rh gauzes which, due to their use as commercial catalysts in the Ostwald and Andrussow processes (see Section 10.5.2), are widely and cheaply available. At an early stage, the results of oxygenate production from alkanes were interpreted as being dominated by homogeneous reactions. Iordanoglou et al. found that more-open gauzes give better oxygenate yields, and even slight increases in contact time (e.g., through using a fivegauze pack) strongly reduce oxygenate yields [43, 44]. Both observations were taken as strong indications that catalytic reactions play only a minor role, if any at all. Iordanoglou and coworkers suggested that alkanes are catalytically combusted to produce heat, which then initiates the homogeneous gas-phase production of oxygenates. The mixing of the hot product gases with cold, unconverted feed gas which passes through the large void of the gauze, leads to rapid quenching and hence suppresses further oxidation of the products, despite the availability of unconverted gas-phase oxygen. More recently, Marengo et al. further elucidated the role of gas-phase reactions for oxygenate formation in SCTR [104]. By carefully controlling and measuring 10.5.4.4

References see page 2186

2182

10.5 Short Contact-Time Reactors

temperature profiles inside the tubular reactor and on the gauze, it was shown that the gauze does indeed play the role of a thermal ignitor of a homogeneous reaction which then produces oxygenates. Surprisingly, however, these authors showed that essentially unchanged oxygenate yields can also be obtained in an empty reactor, albeit ignition temperatures are increased by about 150 K in comparison to the catalytic process. Most interestingly, it was found that even if the gauze was present in the reactor, the reaction front propagated to the reactor entrance after ignition if the oven temperature was kept above the ignition temperature of the homogeneous reaction. At this point, contributions from catalytic reactions could no longer be identified. Marengo and coworkers did not report whether the homogeneous reaction would sustain itself once ignited – that is, whether a purely homogeneous process could hence be designed to operate autothermally. 10.5.5

Reactor Concepts

One of most often-stated advantages of SCTR in comparison to existing industrial processes is the simplicity of the reactor. This becomes apparent in a comparison of CPOM at SCT conditions with the industrially predominant steam reforming process. As CPOM proceeds adiabatically, no external firing of the reactor is necessary, which not only simplifies the design but also reduces the demands on the materials used. The main issue for SCTR is the avoidance of excessive pressure drop at the very high gas flow rates in the catalyst bed. Therefore, the vast majority of SCTR studies have used monolith reactors – that is, structured reactors similar to the catalysts used in automotive catalytic converters (Fig. 6a). These reactors combine a reasonably large, easily accessible surface area with a highly open structure and hence minimal pressure drop. Both extruded, straight-channel monoliths as well as foam monoliths have been investigated (see also Fig. 2), with foam monoliths generally being preferable as the irregular tortuous nature of the channels in foam monoliths suppresses the formation of a laminar boundary layer above the catalyst surface and hence reduces mass transport limitations. Additionally, the increased axial mixing in foam monoliths improves axial heat transport and makes these reactors less prone to flow maldistribution problems which often mar extruded monoliths. Typically, the monolith is made from a lowsurface-area ceramics such as α-alumina (typical surface area ≈1 m2 g−1 ). High-surface-area ceramics are not used as they usually do not show the required high-temperature stability. Moreover, the porous nature of these materials

would result in an undesired broadening of the residence time distribution, with potential detrimental effects on process yields. Beyond monolith reactors, a few studies have reported the use of conventional shallow-bed reactors – that is, fixed-bed reactors with pelletized catalysts and bed heights of the order of a few millimeters to avoid excessive pressure drop (Fig. 6b) [59, 102]. Increased heat conduction against the flow direction, and hence a reduced danger of blow-out of the reaction front, have been stated as advantages of this configuration for CPOM. It seems questionable, however, whether these advantages will outweigh the lower pressure drop and ease of handling with monolithic catalysts. Bharadwaj and Schmidt reported the use of a fluidizedbed reactor for CPOM, with generally comparable results to monolith reactors, although the contact times were increased by about an order of magnitude (≈50–200 ms) [105, 106] (Fig. 6c). As fluidization leads to a significant degree of backmixing in the reactor, these results are interesting with regard to the reaction mechanism as they indicate that, for syngas formation, secondary reactions and rapid quenching of reaction products is not a prime concern as syngas is an equilibrium product. Beyond these simple and comparatively conventional reactor concepts, a number of studies investigated the use of integrated reactor concepts in which either heatexchange or product separation were integrated into the catalytic reactor. Heat-integrated reactors (Fig. 6d) were investigated by Verykios and coworkers [107, 108], as well as by Veser and colleagues [66, 85, 109–112], although for different reasons. Based on the assumption of an indirect oxidation mechanism in which part of the methane feed is combusted in the first step and the remaining methane is then undergoing steam reforming in a second step, Verykios et al. demonstrated the use of a wall reactor in which the inside of the reactor tube is coated with a combustion catalyst for the first step and the outside with a reforming catalyst for the subsequent step. The feed gases are (partially) combusted on the inside of the reactor tube and then flown back through an outer tube where steam reforming occurs on the outside of the inner reactor tube. The very efficient heat transfer between the exothermic combustion and the endothermic reforming through the thin reactor wall is shown to reduce the temperature maximum in the combustion zone significantly. A similar reactor configuration was later used by Venkataraman et al. to couple methane combustion with homogeneous dehydrogenation of ethane to ethylene [113]; that is, to couple two separate endothermal and exothermal reactions rather than have subsequent steps in a single reaction. The principle of coupling endothermal

10.5.5 Reactor Concepts

(a)

Monolith reactor (extruded or foam)

(c)

(b)

Shallow-bed reactor

(d)

Heat-exchange reactor

(f)

Single-gauze reactor

2183

Fluidized-bed reactor t= [n ·T…(n + 0.5)T ] t= [(n + 0.5) T…n ·T ]

(e)

Reverse-flow reactor Fuel

Heater

Air

Premix

(g)

Different short contact-time reactor configurations used to date. (a) Monolith reactor; (b) shallow-bed reactor; (c) fluidized-bed reactor; (d) countercurrent heat-exchange reactor; (e) reverse-flow reactor; (f) single-gauze reactor; (g) tubular reactor with liquid fuel injector.

Fig. 6

and exothermal reactions in autothermal reactors is skillfully explained and discussed in great detail in a series of publications by Kolios, Eigenberger and coworkers, and will not be further elaborated here [114–117]. In contrast to these studies, the use of a heat-integrated reactor by Veser et al. was motivated by the desire

to suppress the initial combustion zone. This was based on an analysis of the CPOM system, which showed that the adiabatic temperature rise of the partial oxidation reaction alone is not sufficient to yield the References see page 2186

2184

10.5 Short Contact-Time Reactors

observed (and thermodynamically necessary) very high reaction temperatures. Instead, Veser and coworkers concluded that the low entrance temperatures into the catalyst zone are responsible for the occurrence of total oxidation which hence could be suppressed if the feed gases were to be sufficiently preheated, resulting in improved syngas yields [85, 111]. The functionality of this operating principle was demonstrated first in a simple countercurrent reactor with integrated recuperative heatexchange (Fig. 6d) [85, 118], and then in a reverseflow reactor (RFR) with regenerative heat-integration (Fig. 6e) [23, 66, 110, 111]. The RFR concept relies on the principle of a moving reaction front. The front part of the catalyst bed acts as a regenerative heat-exchanger and increases the cold feed gas temperature while the reaction front creeps towards the exit of the reactor. To prevent the reaction from extinguishing once the reaction front reaches the outlet of the reactor, the flow through the reaction tube is reversed and the reaction front travels in the opposite direction [119, 120]. If this flow-reversal is conducted at an appropriate frequency, very efficient heatintegration can be achieved, surpassing the efficiency of more conventional recuperative heat-integration [121]. In both reactor configurations strong improvements in syngas yields were obtained and, by combining the RFR operation with appropriate catalysts, methane conversions in excess of 90% at syngas selectivities greater than 90% could be attained in a simple air-blown process – that is, without the need to operate with pure oxygen [23]. The RFR furthermore allowed an at least threefold reduction in contact times, emphasizing that this reactor concept is particularly suitable for SCT conditions and hence small-scale processes with very high throughputs. The RFR concept was recently also applied to the ODH of ethane to ethylene, as well as autothermal reforming of methane at SCT conditions, and similar improvements were reported [122–124]. An important aspect in the design of SCTRs with preheat concerns the mixing of the reactor feed [22], as the precatalytic ignition of flammable hydrocarbon–oxygen mixtures must be avoided. Suggested approaches include mixing the two feed gases right in front of the catalyst zone [109], or carefully design and operation of the reactor in a regime where this preignition is avoided [110]. Another integrated reactor concept that has found application to SCTR is the membrane reactor (MR). In these reactors, permselective membranes are typically used to separate a reaction product from the reactive mixture in order to shift equilibrium limitations or to remove the product from further reactions. Huff and coworkers demonstrated mild improvements in olefin yield by using a MR to remove hydrogen from the product stream during oxidative dehydrogenation of butane in a Pd-membrane reactor [125]. Generally, however, the very

short contact times in SCTRs do not lend themselves easily to MR configurations, as the comparatively long characteristic time scales for hydrogen transport through permselective membranes compare unfavorably with the extremely short time scales of reaction kinetics and reactor residence times [126, 127]. Furthermore, the removal of hydrogen from dehydrogenation reactions is known often to result in coke formation, although this was not reported for the studied system. The most ‘‘radical’’ SCTR concept is the single-gauze reactor as discussed above (see Section 10.5.4.4; see also Fig. 6f). This reactor can essentially be seen as an ultra-short version of a monolith reactor. While this configuration yields lower conversions due to the extremely short contact times, this low conversion can be desirable as it allows the postcatalytic quenching of unstable reaction products via fast mixing with cold, unconverted feed gases. However, it has been shown that similar quenching can be achieved in a simple tubular reactor by operating sufficiently close to the extinction temperature of the reaction, and a purely homogeneous SCT process could hence potentially pose an alternative for a gauze reactor [104]. Finally, an elegant and daring configuration for SCTR with liquid feeds was used by Schmidt and coworkers [42]. Instead of vaporizing the liquid fuels, followed by premixing with air or oxygen, these authors used a standard automotive fuel injector to spray the liquid directly onto the preheated walls of a premixing zone (see Fig. 6g). The liquid fuel vaporizes upon hitting the hot reactor walls, while being mixed with an incoming oxygen stream. Although one might expect this to result in highly flammable (if not explosive) fuel–oxygen mixtures somewhere in the premixing zone, Schmidt and coworkers reported on the safe operation of this reactor over a rather wide range of operating conditions. While this configuration requires additional external heating for the vaporization–premixing zone (and thus no longer strictly constitutes an autothermal reactor configuration), it is a perfect match for the simple and compact nature of SCTR. 10.5.6

Summary and Outlook

Short contact-time reactors and reactions still represent a relatively young field of reactor and reaction engineering, and it can be expected that further uses for SCTRs will be developed over the next decade. Likewise, the use of existing SCTRs will become more prevalent as novel reactor concepts will need to be implemented, particularly in the rapidly developing fields of energy and environmental engineering.

10.5.6 Summary and Outlook

In this chapter, an attempt was made to demonstrate the flexibility of SCTRs, based on a brief (and incomplete) overview of some of the main activities in this field over the past decade. Perhaps the best illustration of the flexibility of the concept, however, is provided by a final example from the studies of Schmidt and colleagues on the oxidative conversion of cyclohexane at SCT conditions [128]. By changing the catalyst (and adjusting the carbon:oxygen ratio), they could ‘‘tune’’ the product spectrum of partial oxidation of this single fuel with oxygen in an autothermal SCTR from essentially pure syngas over Rh-coated monoliths, to olefins (mostly cyclohexene and benzene) over Pt/Sn-coated monoliths, to the formation of different oxygenated products over different monolithic or gauze catalysts (see Fig. 7). Although cyclohexane conversions vary widely for the

2185

different reaction pathways, selectivities to the main products are respectable in all cases, and the spectrum of products certainly demonstrates at the same time the versatility and the complexity of SCTRs. Figure 8 summarizes, in a somewhat simplified schematic, the main reactions for which SCTRs have been successfully demonstrated to date, along with the corresponding reaction conditions: contact times decrease from about 10−2 s (10 ms) for syngas formation to about 10−3 s (1 ms) for olefin production, and finally into the submillisecond range for oxygenate formation. In parallel, reaction temperatures decrease from well above 1200 K to below 1000 K. The corresponding C : O molar ratio in the reactor feed (not shown in the schematic) varies References see page 2186

+ O2

Rh (monolith)

Pt / 10%Rh (gauze)

Pt-Sn + H2 (monolith)

Ag, Mo, or Co (monolith)

Syngas

O

(H2 + CO)

O

O

OH Xc 6 = 97%

Xc 6 = 25%

Xc 6 = 11%

Xc 6 = 3%

Xo 2 > 99%

Xo 2 > 99%

Xo 2 > 99%

Xo 2 = 3%

S > 95%

S = 60%

S = 94%

S = 80%

Example of the high flexibility of short contact-time reactions. Major cyclohexane partial oxidation pathways are shown schematically as a function of catalyst used along with cyclohexane and oxygen conversions and major product selectivities. (Adapted from Ref. [128].)

Fig. 7

Catalytic reaction

Fuel CH4

Rh

Syngas

Homogenous reaction

Contact time t

T

(none)

> 10−2s

> 1200 K

Olefins

~10−3−10−2s

~ 1000 − 1200 K

< 10−3s

< 1000 K

Alcohols (C1−C3) pt-sn

CO, CO2,H2O, heat CO, CO 2,H2O, C4 − C10 Alkanes heat Pt/10%Rh gauze C2H6, C3H8

or:LaMnO3

Oxygenates

Schematic summarizing the main groups of (oxidative) short contact-time reactions studied to date, along with the contributions of catalytic versus homogeneous reaction pathways and main reactor operating conditions (temperature and contact times).

Fig. 8

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10.5 Short Contact-Time Reactors

significantly for different hydrocarbon feedstocks, but generally decreases from C : O ≈ 1 for syngas formation, to higher C : O (typically ≈ 1.5–3) for olefin production, and even richer fuel–oxygen mixtures with higher C : O ratios for oxygenate formation. Overall, perhaps the most interesting and surprising insight that emerges from the sum of SCTR investigations to date is the fact that, despite often being referred to as ‘‘short contact-time catalysis’’, the only true catalytic reaction product in all of the studies on oxidative conversions at SCT conditions is synthesis gas. In all other cases (olefins, oxygenates), the catalytic reaction appears to play little (if any) role beyond being an ultra-fast ignitor of the homogeneous reaction. Nevertheless, as the hexane example illustrates, this role is crucial to attaining the demonstrated variety of products at autothermal reactor operation. Beyond the inherent product quenching due to the ultra-short catalyst contact times in SCTRs, the steep temperature gradients that develop based on this initial catalytic oxidation step are hence key to a true understanding of SCTRs. Due to the fast flow rate and high reaction rates, SCTRs proceed near-adiabatically at autothermal conditions. While the adiabatic temperature rise for the desired partial oxidation reactions typically amounts to only a few hundred Kelvin, combustion reactions – which consume most (if not all) of the oxygen within the first few millimeters of the catalyst bed – produce an adiabatic temperature rise of 2000 K and above. Hence, SCTRs are characterized by temperature gradients at the front end of the catalyst bed of up to 106 K m−1 . The combination of these extreme temperature gradients with fast product quenching due to ultra-short contact times is what truly sets SCTR apart from more conventional chemical processes. The coupling of heat- and mass-transport processes which are governed by these extreme gradients with the extremely fast reaction kinetics of catalytic oxidation reactions, as well as the fine interplay between the catalytic ignition and the subsequent homogeneous production of chemicals, make SCTRs a formidable challenge for chemical engineers. Detailed numerical simulations will become an increasingly important tool in the further development of SCTRs, and these simulations will depend strongly on the availability of reliable reaction kinetics for the (product-forming) homogeneous reactions. Detailed reaction kinetics for the catalytic reaction pathways, which have been the focus of intense research during the past decade, are probably less important for a deeper understanding of these reactions, as surface reactions appear to contribute not much more than heat (as well as some CO, CO2 , and H2 O) to the overall process.

Finally, while more advanced reactor concepts will help the industrial implementation of SCTRs, these concepts must aim to maintain the simplicity of SCTRs while addressing crucial issues such as optimizing the heat load in the reactors and minimizing safety concerns while handling premixed hydrocarbon–air mixtures. Overall, the industrial implementation of SCTRs for hydrocarbon conversion seems only a matter of time, although it remains to be seen how far the SCT concept can be extended to further, novel areas of application. References 1. D. A. Hickman, L. D. Schmidt, Science 1993, 259, 343. 2. A. S. Bodke, D. A. Olschki, L. D. Schmidt, E. Ranzi, Science 1999, 285, 712. 3. D. A. Goetsch, L. D. Schmidt, Science 1996, 271, 1560. 4. G. A. Deluga, J. R. Salge, L. D. Schmidt, X. E. Verykios, Science 2004, 303, 993. 5. L. D. Schmidt, M. Huff, S. S. Bharadwaj, Chem. Eng. Sci 1994, 49, 3981. 6. L. D. Schmidt, J. Siddall, M. Bearden, AIChE J. 2000, 46, 1492. 7. A. Andrussow, Angew. Chem. 1935, 48, 593. 8. K. Weissermel, H.-J. Arpe, Industrial Organic Chemistry, 4th Ed., VCH-Wiley, Weinheim, 2003, 511 pp. 9. M. Prettre, C. Eichner, M. Perrin, Trans. Faraday Soc. 1946, 43, 335. 10. J. S. Falsetti, Hydrogen Technol. Intl. 1993, 57. 11. A. T. Ashcroft, A. K. Cheetham, J. S. Foord, M. L. H. Green, C. P. Grey, A. J. Murrell, P. D. F. Vernon, Nature 1990, 344, 319. 12. P. D. F. Vernon, M. L. H. Green, A. K. Cheetham, A. T. Ashcroft, Catal. Lett. 1990, 6, 181. 13. D. Dissanayake, M. P. Rosynek, K. C. C. Kharas, J. H. Lunsford, J. Catal. 1991, 132, 117. 14. V. R. Choudhary, A. S. Mamman, S. D. Sansar, Angew. Chem. 1992, 104, 1273. 15. V. R. Choudhary, A. M. Rajput, B. Prabhaka, Angew. Chem. 1994, 106, 2179. 16. D. A. Hickman, E. A. Haupfear, L. D. Schmidt, Catal. Lett. 1993, 17, 223. 17. D. A. Hickman, L. D. Schmidt, J. Catal. 1992, 138, 267. 18. I. Wender, Fuel Proc. Technol. 1996, 48, 189. 19. M. A. Pena, J. P. Gomez, J. L. G. Fierro, Appl. Catal. 1996, A 144, 7. 20. D. A. Hickman, L. D. Schmidt, AIChE J. 1993, 39, 1164. 21. P. M. Torniainen, X. Chu, L. D. Schmidt, J. Catal. 1994, 146, 1. 22. S. C. Reyes, J. H. Sinfelt, J. S. Freeley, Ind. & Eng. Chem. Res. 2003, 42, 1588. 23. D. Neumann, M. Kirchhoff, G. Veser, Catal. Today 2004, 98, 565. 24. M. Huff, P. M. Torniainen, L. D. Schmidt, Catal. Today 1994, 21, 113. 25. M. Huff, P. M. Torniainen, L. D. Schmidt, Catal. Today 1994, 21, 113. 26. R. Subramanian, G. J. Panuccio, J. J. Krummenacher, I. C. Lee, L. D. Schmidt, Chem. Eng. Sci. 2004, 59, 5501.

References 27. L. D. Schmidt, E. J. Klein, C. A. Leclerc, J. J. Krummenacher, K. N. West, Chem. Eng. Sci. 2003, 58, 1037. 28. Y. Jamal, M. I. Wyszynski, Intl J. Hydrogen Energy 1994, 19, 557. 29. L. D. Schmidt, M. Huff, Catal. Today 1994, 21, 443. 30. S. S. Bharadwaj, C. Yokoyama, L. D. Schmidt, Applied Catal. A 1996, 140, 73. 31. T. Ioannides, X. E. Verykios, Catal. Lett. 1996, 36, 165. 32. A. G. Dietz, A. F. Carlsson, L. D. Schmidt, J. Catal. 1998, 176, 459. 33. A. Beretta, P. Forzatti, J. Catal. 2001, 200, 45. 34. B. Silberova, H. J. Venvik, A. Holmen, Catal. Today 2005, 99, 69. 35. M. Huff, L. D. Schmidt, J. Phys. Chem. 1993, 97, 11815. 36. M. Huff, P. M. Torniainen, D. A. Hickman, L. D. Schmidt, in Natural Gas Conversion II, H. E. Curry-Hyde, R. F. Howe (Eds.), Elsevier, Amsterdam, 1994, pp. 315. 37. M. Huff, L. D. Schmidt, J. Catal. 1994, 149, 127. 38. M. Huff, L. D. Schmidt, J. Catal. 1995, 155, 82. 39. J.-P. Lange, R. J. Schoonebeek, P. D. L. Mercera, F. W. v. Breukelen, Appl. Catal. A 2005, 283, 243. 40. C. Yokoyama, S. S. Bharadwaj, L. D. Schmidt, Catal. Lett. 1996, 38, 181. 41. A. S. Bodke, L. D. Schmidt, Catal. Lett. 1999, 63, 113. 42. J. J. Krummenacher, K. N. West, L. D. Schmidt, J. Catal. 2003, 215, 332. 43. D. I. Iordanoglou, A. S. Bodke, L. D. Schmidt, J. Catal. 1999, 187, 400. 44. D. I. Iordanoglou, L. D. Schmidt, J. Catal. 1998, 176, 503. 45. R. Martinez, M. C. Huff, M. A. Barteau, Appl. Catal. A 2000, 200, 79. 46. R. Martinez, M. C. Huff, M. A. Barteau, J. Catal. 2004, 222, 404. 47. K. Heitnes, S. Lindeberg, O. A. Rokstad, A. Holmen, Catal. Today 1995, 24, 211. 48. Y. Boucouvalas, Z. L. Zhang, X. E. Verykios, Catal. Lett. 1996, 40, 189. 49. G. Veser, L. D. Schmidt, AIChE J. 1996, 42, 1077. 50. G. Veser, J. Frauhammer, L. D. Schmidt, G. Eigenberger, in Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis, G. F. Froment, K. C. Waugh (Eds.), Studies in Surface Science and Catalysis, Vol. 109, Elsevier, Amsterdam, 1997, pp. 273. 51. K. H. Hofstad, J. Hoebink, A. Holmen, G. B. Marin, Catal. Today 1998, 40, 157. 52. V. A. Tsipouriari, Z. Zhang, X. E. Verykios, J. Catal. 1998, 179, 283. 53. V. A. Tsipouriari, X. E. Verykios, J. Catal. 1998, 179, 292. 54. C. Elmasides, T. Ioannides, X. E. Verykios, AIChE J. 2000, 46, 1260. 55. C. Elmasides, X. E. Verykios, J. Catal. 2001, 203, 477. 56. L. Basini, A. Aragno, G. Vlaic, Catal. Lett. 1996, 39, 49. 57. A. S. Bodke, S. S. Bharadwaj, L. D. Schmidt, J. Catal. 1998, 179, 138. 58. E. Ruckenstein, H. Y. Wang, J. Catal. 1999, 187, 151. 59. K. L. Hohn, L. D. Schmidt, Appl. Catal. A 2001, 211, 53. 60. F. Monnet, Y. Schuurman, F. C. S. Aires, J. C. Bertolini, C. Mirodatos, Catal. Today 2001, 64, 51. 61. S. Ramani, D. M. Minahan, Y. Jiang, WO Patent 2004/ 072209, assigned to ConocoPhillips Company, 2004. 62. F. Basile, L. Basini, G. Fornasari, M. Gazzano, F. Trifiro, A. Vaccari, Chem. Commun. 1996, 2435. 63. L. Majocchi, G. Groppi, C. Cristiani, P. Forzatti, L. Basini, A. Guarinoni, Catal. Lett. 2000, 65, 49.

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64. J. Schicks, D. Neumann, U. Specht, G. Veser, Catal. Today 2003, 83, 287. 65. M. Kirchhoff, U. Specht, G. Veser, in Innovation in the Manufacture and Use of Hydrogen, G. Emig (Ed.), DGMK Publishing, Hamburg, 2003, pp. 33. 66. A. Mitri, D. Neumann, T. F. Liu, G. Veser, Chem. Eng. Sci. 2004, 59, 5527. 67. S. S. Bharadwaj, L. D. Schmidt, Fuel Proc. Technol. 1995, 42, 109. 68. S. C. Tsang, J. B. Claridge, M. L. H. Green, Catal. Today 1995, 23, 3. 69. A. P. E. York, T. Xiao, M. L. H. Green, Topics Catal. 2003, 22, 345. 70. W. J. M. Vermieren, E. Blomsma, P. A. Jacobs, Catal. Today 1992, 13, 427. 71. K. Heitnes, S. Lindeberg, O. A. Rokstad, A. Holmen, Catal. Today 1994, 21, 471. 72. T. Liu, C. Snyder, G. Veser, J. Catal. 2006, submitted. 73. C. R. H. de Smet, M. de Croon, R. J. Berger, G. B. Marin, J. C. Schouten, Appl. Catal. A 1999, 187, 33. 74. C. R. H. de Smet, M. de Croon, R. J. Berger, G. B. Marin, J. C. Schouten, AIChE J. 2000, 46, 1837. 75. A. Bizzi, L. Basini, G. Saracco, V. Specchia, Chem. Eng. J. 2002, 90, 97. 76. C. T. Goralski, R. P. O’Connor, L. D. Schmidt, Chem. Eng. Sci. 2000, 55, 1357. 77. G. Veser, J. Frauhammer, Chem. Eng. Sci. 2000, 55, 2271. 78. O. Deutschmann, L. D. Schmidt, AIChE J. 1998, 44, 2465. 79. M. Bizzi, L. Basini, G. Saracco, V. Specchia, Ind. & Eng. Chem. Res. 2003, 42, 62. 80. P. M. Biesheuvel, G. J. Kramer, AIChE J. 2003, 49, 1827. 81. L. L. Raja, R. J. Kee, O. Deutschmann, J. Warnatz, L. D. Schmidt, Catal. Today 2000, 59, 47. 82. M. Bizzi, G. Saracco, R. Schwiedernoch, O. Deutschmann, AIChE J. 2004, 50, 1289. 83. P. A. Bui, D. G. Vlachos, P. R. Westmoreland, Surf. Sci. 1997, 385, L1029. 84. P. Aghalayam, Y. K. Park, D. G. Vlachos, AIChE J. 2000, 46, 2017. 85. G. Veser, J. Frauhammer, U. Friedle, Catal. Today 2000, 61, 55. 86. J. D. Taylor, M. D. Allendorf, A. H. McDaniel, S. F. Rice, Ind. & Eng. Chem. Res. 2003, 42, 6559. 87. R. J. Berger, G. B. Marin, Ind. & Eng. Chem. Res. 1999, 38, 2582. 88. K. H. Hofstad, T. Sperle, O. A. Rokstad, A. Holmen, Catal. Lett. 1997, 45, 97. 89. P. M. Witt, L. D. Schmidt, J. Catal. 1996, 163, 465. 90. K. L. Hohn, P. M. Witt, M. B. Davis, L. D. Schmidt, Catal. Lett. 1998, 54, 113. 91. L. Basini, K. Aasberg-Petersen, A. Guarinoni, M. Ostberg, Catal. Today 2001, 64, 9. 92. M. C. Huff, L. D. Schmidt, AIChE J. 1996, 42, 3484. 93. D. W. Flick, M. C. Huff, Appl. Catal. A 1999, 187, 13. 94. O. V. Buyevskaya, D. Wolf, M. Baerns, Catal. Today 2000, 62, 91. 95. S. A. R. Mulla, O. V. Buyevskaya, M. Baerns, J. Catal. 2001, 197, 43. 96. F. Donsi, R. Pirone, G. Russo, J. Catal. 2002, 209, 51. 97. M. C. Huff, I. P. Androulakis, J. H. Sinfelt, S. C. Reyes, J. Catal. 2000, 191, 46. 98. A. Beretta, P. Forzatti, E. Ranzi, J. Catal. 1999, 184, 469. 99. A. Beretta, L. Piovesan, P. Forzatti, J. Catal. 1999, 184, 455.

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100. A. Beretta, E. Ranzi, P. Forzatti, Chem. Eng. Sci. 2001, 56, 779. 101. A. Beretta, E. Ranzi, P. Forzatti, Catal. Today 2001, 64, 103. 102. D. W. Flick, M. C. Huff, J. Catal. 1998, 178, 315. 103. J. J. Krummenacher, L. D. Schmidt, J. Catal. 2004, 222, 429. 104. S. Marengo, P. Comotti, G. Galli, Catal. Today 2003, 81, 205. 105. S. S. Bharadwaj, L. D. Schmidt, J. Catal. 1994, 146, 11. 106. S. S. Bharadwaj, L. D. Schmidt, J. Catal. 1995, 155, 403. 107. T. Ioannides, X. E. Verykios, Catal. Lett. 1997, 47, 183. 108. A. Piga, X. E. Verykios, Catal. Today 2000, 60, 63. 109. U. Friedle, G. Veser, Chem. Eng. Sci. 1999, 54, 1325. 110. D. Neumann, V. Gepert, G. Veser, Ind. & Eng. Chem. Res. 2004, 43, 4657. 111. D. Neumann, G. Veser, AIChE J. 2005, 51, 210. 112. G. Veser, in Process Intensification and Microreaction Technology, Y. Wang, J. Holladay (Eds.), ACS Symposium Series, Vol. 914, ACS, Washington, 2005, p. 145. 113. K. Venkataraman, J. M. Redenius, L. D. Schmidt, Chem. Eng. Sci. 2002, 57, 2335. 114. G. Kolios, J. Frauhammer, G. Eigenberger, Chem. Eng. Sci. 2000, 55, 5945. 115. G. Kolios, J. Frauhammer, G. Eigenberger, Chem. Eng. Sci. 2001, 56, 351. 116. G. Kolios, J. Frauhammer, G. Eigenberger, Chem. Eng. Sci. 2002, 57, 1505. 117. B. Glockler, A. Gritsch, A. Morillo, G. Kolios, G. Eigenberger, Chem. Eng. Res. & Design 2004, 82, 148. 118. U. Friedle, J. Schicks, G. Veser, in Synthesis Gas Chemistry, D. Hoenicke (Ed.), DGMK, Hamburg, 2000, pp. 53. 119. G. K. Boreskov, Y. S. Matros, Appl. Catal. 1983, 5, 337. 120. G. Boreskov, Y. Matros, Catal. Rev. Sci. Eng. 1983, 25, 551. 121. Y. S. Matros, G. K. Bunimovic, Catal. Rev. Sci. Eng. 1996, 38, 1. 122. T. Liu, V. Gepert, G. Veser, Chem. Eng. Res. Design 2005, 83, 611. 123. T. Liu, H. Temur, G. Veser, in preparation 2006. 124. V. Gepert, G. Veser, in preparation 2006. 125. T. M. Raybold, M. C. Huff, Catal. Today 2000, 56, 35. 126. G. Saracco, V. Specchia, Catal. Rev. Sci. Eng. 1994, 36, 304. 127. G. Saracco, H. W. J. P. Neomagus, G. F. Versteeg, W. P. M.v. Swaaij, Chem. Eng. Sci. 1999, 54, 1997 128. R. P. O’Connor, E. J. Klein, D. Henning, L. D. Schmidt, Appl. Catal. A 2003, 238, 29.

to perform a single specific operation. The possibility of coupling reaction and separation in a single vessel, simultaneously performing two operations, offers potential for reductions in investment costs (fewer vessels needed) and operating costs (e.g., the direct exploitation of heat generated by an exothermic reaction). As fractional distillation is the most commonly applied separation technique, R&D engineers have been working since the early 1920s to develop a novel unit operation, to combine chemical reaction and fractional distillation in a single vessel; this process is referred to as ‘‘reactive distillation’’. When a catalyst, either molecular or solid, is used to enhance the reaction rate, this unit operation is also known as ‘‘catalytic distillation’’. A schematic representation of a catalytic distillation column for a reaction of the type: ← −C+D A+B− −− −− → is shown in Fig. 1, assuming that at least one of the products has a non-negligible difference in volatility with respect to the other compounds. The figure shows one of the possible configurations, including the following sections: • a rectification section in the upper zone • a reactive distillation section in the middle, containing the catalyst • a stripping section in the lower zone. The history of catalytic distillation begins with the pioneering studies of Backhaus [1], who filed 11 patents

C

10.6

Catalytic Distillation A+B

Ivano Miracca∗ , Arvids Judzis, Jr., Nishit Sahay, and Domenico Sanfilippo

10.6.1

Introduction

A chemical process transforms reactants into products with the required degree of purity. Chemical reaction and separation/purification usually take place in different equipment units, the design of which is optimized ∗

Corresponding author.

D

Fig. 1

A schematic representation of a catalytic distillation column.

10.6.2 Potential versus Constraints

covering its use in reactions of esterification. A detailed review on the subject was published as early as 1932 [2], and application was soon extended to other types of reaction, such as nitrations, sulfonations, and saponifications [3]. All of these reactions were homogeneously catalyzed, with the catalyst in the liquid phase flowing down the column to be recovered from the residue. In spite of this early start, progress was slow and applications to heterogeneously catalyzed reactions were devised only during the early 1970s [4]. This gap in time serves as a measure of the difficulty of placing what is effectively a solid catalyst inside a distillation column. The past 30 years have witnessed a vast R&D effort in the field, leading to the production of hundreds of patents and publications, as well as a few successful commercial applications, such as the production of methyl-tert-butyl ether (MTBE), which is widely used as an octane booster and oxygen carrier in reformulated gasoline. This process was pioneered by Catalytic Distillation Technologies (CDTECH), and about 70 units have been built in the past 20 years. More recently, CDTECH have added other successful histories of development, with their desulfurization technologies (as of 2005, about 20 units licensed) and selective hydrogenation (35 units licensed). Today, following the optimism of the 1980s, during which time widespread application seemed to be at hand, catalytic distillation is regarded – sometimes skeptically [5] – as a niche technology, the high potential of which is subjected to several constraints that often prevent its commercial implementation. The development of more efficient hardware and software, including systematic procedures based on heuristic rules to assess the feasibility and attractiveness of the application of catalytic distillation to particular reactions, has in the past few years, led to renewed interest and development efforts. 10.6.2

Potential versus Constraints

Catalytic distillation is not applicable to every type of chemical reaction, the main constraint being that the reaction must proceed at a reasonable rate in conditions of vapor–liquid equilibrium (i.e., at relatively low temperature compared to the majority of chemical processes). On the other hand, catalytic distillation has potential for a breakthrough compared to the conventional approach, when withdrawal of a product as soon as it is formed may enhance the yield of the reaction. In this regard, it is possible to distinguish two main cases:

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• Reactions with severe limitations imposed by chemical equilibrium; the separation of one of the products by evaporation (or condensation) as it is formed can drive the reaction to completion. This is usually achieved by feeding one of the reactants in large excess, while catalytic distillation in principle allows the reaction to be carried out by feeding reactants in stoichiometric ratio. • Reactions in which a high concentration of one of the products or reactants may push undesired side reactions. For instance, in consecutive reactions of the type A + B −−−→ C (desired product) C + B −−−→ D (undesired by-product) if B and C have different volatilities, the concentration of one of them in the reacting phase may be kept at a very low level, thus hindering the undesired reaction. Application to exothermic reactions may result in an optimized treatment of reaction heat through its easy removal by vaporization, and direct exploitation decreasing the external duty supplied by the reboiler. Moreover, the reaction can be run almost isothermally with the temperature strictly controlled by the vaporization rate. Evaluation of the opportunity to apply catalytic distillation to a certain reactive system may now be performed by specialized process synthesis software tools which couple numeric calculations to the application of heuristic rules, thereby allowing a quick identification of attractive options as well as the rejection of non-feasible ones [6]. Heuristic rules based on expertise have been developed by several companies. For example, BASF [7] has recently proposed a set of principles based on the evaluation of reaction rate versus relative volatility. Shell also claims to have developed a set of 15 guidelines to assess attractiveness and feasibility of catalytic distillation [8]. On the other hand, the design of column internals must deal with different requirements both for reaction and for distillation. Even in the case of homogeneous catalytic distillation, the high liquid hold-up on the tray required to maximize conversion is in conflict with the high interfacial area required by distillation. Moving on to heterogeneous catalytic distillation poses considerably more challenges, as the reaction requires a small particle size of the catalyst in order to avoid diffusion limitation, and a high catalyst loading to maximize efficiency (in a fixed-bed reactor, about two-thirds of the volume is occupied by the catalyst), while distillation needs high open areas for gas and liquid flow to minimize pressure drop and avoid a too-close approach to ‘‘flooding’’. For References see page 2197

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10.6 Catalytic Distillation

instance, state-of-the-art packings for catalytic distillation only contain about 20–25 vol.% of catalyst. The design of efficient hardware for catalytic distillation has been the greatest challenge for developers and, in spite of the outstanding achievements in the mitigation of design conflicts, the past few years have seen a trend towards a new concept of catalytic distillation which is characterized by a lower degree of integration, in which the reaction and distillation are no longer run in a single vessel. Rather, the reaction zones are confined in side reactors that are fed by side liquid drawn from the distillation column [9], thereby offering greater flexibility and robustness in the process scheme. 10.6.3

Practical Realization of Catalytic Distillation

The location of the reaction zone inside an heterogeneously catalyzed column is one of the main concerns for the optimization of catalytic distillation operations. Depending on the type of reaction and the relative volatilities of the components, it may be advantageous to carry out the reaction in one or more regions, and allow the remainder of the column to operate as a conventional separation unit. Practical rules can be derived either from experience or from the simulation of different configurations. A comprehensive theoretical investigation was performed by Balashov and coworkers [10–13], who formulated a set of criteria relevant to determine the extension of the reaction zone to reach a desired conversion. The mechanical arrangement of the catalyst inside the column, and the design of a correct shape, are of primary importance, as an optimal performance for both reaction and distillation should be achieved. On the one hand,

(a)

(b)

a catalytic distillation zone should offer enough open space for vapor and liquid flow, to perform an efficient distillation, because a large pressure drop in the vapor phase can have a negative effect on physical equilibria. On the other hand, reactions take place efficiently when a good contact is achieved between reactants and catalyst. Moreover, a liquid hold-up higher than is strictly necessary for distillation should be used if the reaction is slow. The optimal solution must be a compromise between these requirements. Several arrangements have been proposed either for tray or packed columns. The first attempts led to quite simplified solutions, for example placing the catalyst on shelves with layers of distillative packing above it [4], or filling the down-comers of a tray tower with catalyst [14]. In time, the packing solution seems to be favorite, as most – if not all – existing commercial applications are based on packings, in spite of several patents covering trays that have been developed specifically for catalytic distillation. Tray Towers The first arrangement proposed for tray towers consisted of closed containers [15, 16] of porous catalyst, and a clip which held and supported them on the trays. The trays had weirs to provide a liquid level to cover the containers (Fig. 2a). In this way, it is only the liquid phase that makes full contact with the catalyst while moving across the tray, and the gas flows axially through the tray holes. Confining a particulate catalyst above the tray by using appropriate containing screens represents another option [17]. The catalyst is fluidized by the vapor flowing axially, whereas the liquid flows transversally, following the surface of the tray. 10.6.3.1

(c)

Some examples of arrangements proposed for tray catalytic distillation columns. (a) A perforated sieve tray with a contact catalyst structure positioned thereon. (b) A reactive tray with radial flow of liquid phase. (c) A fluid bed tray.

Fig. 2

10.6.3 Practical Realization of Catalytic Distillation

Subsequent developments in the field of reactive trays led to arrangements in which the vapor phase substantially bypasses, through one or more chimneys, the fixed catalytic bed supported above a tray. This is a particularly desirable solution for liquid-phase reactions, allowing the use of a simple fixed catalytic bed with low pressure drop and easy catalyst replacement. Such a bypass can be achieved with countercurrent axial flow of liquid and vapor phases [18, 19], with radial flow of the liquid phase through the bed (Fig. 2b) [20], or with an upward liquid flow in the catalytic bed [21]. All of these arrangements are claimed to improve the efficiency of liquid-phase reactions because the liquid phase becomes continuous with the vapor bubbling through it. Another proposal coupled a reactive tray supporting a fixed-bed of catalyst to a distillative tray, with the vapor phase bypassing the reactive tray and the liquid phase recirculating between them and completely covering the catalyst [22]. Another development reported in the field of reactive trays is a fluid-bed tray [23], upon which the catalyst bed is partly fluidized by the upward gas flow (Fig. 2c). A hood is included at the upper end of the chimney to direct the flow back onto the tray. The liquid from the tray above is distributed by pipes which extend below the static catalyst level on the tray. This solution is claimed to provide better contact compared to other reactive trays, but does not appear to be suitable for equilibrium-limited reactions, favored by plug-flow. Work on the development of reactive trays continued up to the end of the 1990s, always with the perspective of achieving an efficient separation of the vapor and liquid phases, allowing only the liquid phase to flow through the catalyst. For example, Koch-Glitsch tried to improve the concept of ‘‘catalyst bales’’ (see Section 10.6.3.2), by proposing containers for catalyst particles that were placed conveniently in the column in order to avoid the vapor phase that may flow through them [24]. By contrast, the research group at the University of Alberta developed a reactive tray based on the concept of perforated chevron grids (Fig. 3), on which preferential distinct vertical paths are provided for the two phases, allowing contact only in the non-catalytic zone on the tray surface [25]. Since the year 2000, tray towers in catalytic distillation seem to be present only in cases where reaction and distillation are carried out in different vessels, and the scheme is optimized by side draws and feeds, both in the reactor and in the distillation column. A first example of this approach was proposed as early as 1993 [26], and later developed and commercially applied to etherifications, claiming higher flexibility and lower maintenance costs compared to catalytic distillation columns [27].

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Liquid

Liquid overflow

Vapor

A tray for catalytic distillation, based on the concept of perforated chevron grids.

Fig. 3

Packed Towers The first packed-bed reactor which was claimed to be particularly appropriate for reactive distillation consisted of catalyst pellets or beads placed in pockets in a cloth belt supported by an open-mesh, knitted stainless steel wire [28]. The cloth belt was then coiled into a spiral, and the wire mesh placed between the coils. These spirals assumed the form of large cylinders that could be stacked inside the column (Fig. 4, left); they have subsequently been commercialized as ‘‘catalyst bales’’ [29], and provide a preferential route for vapor flow. In the process, the liquid wets the catalyst, and the gas represents the true continuous phase; this may be a disadvantage in cases where chemical equilibrium is disfavored in the gas phase, but this arrangement achieved great commercial success, and most commercial plants currently in operation still employ this type of internal device. The direct coating of normal distillation packings, as Raschig rings or Berl saddles, with a catalytic layer was also claimed [30], but such a solution is probably feasible only for fast reactions, because of the small surface area available. Starting in the early 1990s, companies with a background in the field of distillative packings began the 10.6.3.2

References see page 2197

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Fig. 4

10.6 Catalytic Distillation

Left: A plan view of a catalyst bale. Right: the catalytic packing KATAPAK , as developed and commercialized by Koch-Glitsch.

development of structured packings containing a catalyst, specifically optimized for catalytic distillation [31–33]. The basic elements of these structured packings have double walls that are permeable to the reactants and can be filled with a catalyst (Fig. 4, right). Several packings for reactive distillation have been proposed in recent years, the aim always being to increase voidage with respect to catalytic packed beds [34–38]. Structured packings have been available commercially since the mid-1990s, with trade-mark names such as Katapak (Sulzer) or Katamax (Koch-Glitsch), and have been successfully applied in several commercial units. CDTECH has also developed their own commercial structured packing (CDModule ) (Fig. 5) that has been optimized for application in liquid-phase catalytic distillation

Reflux

Vapor Catalyst

Wire mesh

Vapor

Fig. 5

processes. Compared to the catalyst bales, the new generation of structured packings is significantly more efficient for distillation than earlier versions, and offers the advantage of reusing the structures by replacing the spent inactive catalyst with fresh catalyst. Catalyst replacement seems to be the only remaining drawback of this type of packing, as removal of the whole packing is needed to replace spent catalyst. The development of these packings has also powered a new modeling effort by the academic world, by combining the rate-based approach with computational fluid dynamics to provide reliable simulation tools to process engineers, as described in the following section. Finally, an interesting approach has been patented by BASF [39], in which a technique is claimed enabling a

Liquid

Simplified schematics of the CDModule developed and commercialized by CDTECH.

10.6.4 Modeling Catalytic Distillation

layer of catalyst to be deposited directly onto a structured packing for distillation. 10.6.4

Modeling Catalytic Distillation

The development of suitable models for design and simulation of catalytic distillation goes in parallel with the development of suitable models for multi-component distillation, with the additional problem of dealing with reactive terms in the equations. Equilibrium-Stage Models The traditional equilibrium-stage approach is based on the following assumptions: 10.6.4.1

• The system consists of a known number of stages. • Each stage may have feed or withdrawal of vapor, liquid and heat. • The liquid and vapor phases are in equilibrium, or stage efficiency is considered. • Reaction may be present in some or all of the stages and may be considered at equilibrium or kinetically controlled. The resulting mathematical model for steady-state operation consists of a system of non-linear equations, the solution of which allows the determination of flow rates, compositions, temperatures, and conversions for each stage. A variety of algorithms has been proposed, and the computations required to model catalytic distillation columns can be considered either as an extension of the procedures used for pure distillation columns, or as a particular case of mass and enthalpy balance where both reaction and phase separation are present. The introduction of reaction kinetics and equilibria into the mass and energy balances of a column leads to difficulties in solving the long sequence of interconnected equations. These difficulties arise from the strong non-linearities introduced by chemical kinetics or equilibria. As in the case of reactive distillation, algorithms can be classified into three main categories: • Relaxation techniques that express the balance equations in dynamic form and use time as a convergence parameter towards the steady state [40, 41]. • Tearing and partitioning of equations aiming at a size reduction of the large original system [42–44]. • Global approaches to the large sparse system of equations using Newtonian algorithms coupled to convergence accelerators and optimization techniques [45, 46]. An important item to help discriminate between different approaches is the ratio between reaction

2193

and mass transfer rates, as represented by the Hatta number [47]. If the reaction is fast compared to the masstransfer rate, then the system can be described assuming reaction at equilibrium. If the reaction is slow compared to mass-transfer rate, a kinetic expression must be integrated into the mass and energy balances, with considerable convergence problems. Several studies on the subject have been published [48, 49]. The introduction of parameters such as tray or packing efficiencies simplifies the treatment of binary distillation systems, but often fails in the case of multicomponent (or catalytic) distillation as diffusion interactions cause a strong dependency of these parameters on the local concentrations of each component [50]. Another feature of catalytic distillation is the acceleration of interphase mass transfer due to chemical reaction. Empirical enhancement factors [47] have been introduced in equilibrium-stage models to account for this effect, but their validity is questionable. Rate-Based Models Rate-based models represent a more rigorous approach to catalytic distillation modeling, as actual rates of multicomponent mass and heat transfer and chemical reactions are taken directly into account [51, 52]. Different theoretical correlations are available to describe mass transfer at the gas–liquid interface, the best known being the two-film model [53] and the penetration/surface removal model [54]. Model parameters should, however, be estimated by experimental correlations. Application of the two-film model allows the exploitation of a number of correlations available in the literature for several types of internals, while the choice is more limited for the penetration/surface removal model. The theory of the two-film model is based on the assumption that resistance to mass transfer is concentrated in thin films close to the phase interface, and transfer takes place through these films by molecular diffusion. Equations deduced by Stefan–Maxwell, and derived from the kinetic theory of gases, may be used to describe this type of multicomponent diffusion [55]. The system of equations for the multicomponent rate-based models has been widely described in the literature [47, 56]. A typical simplification compared to the more generalized form is that chemical reactions usually take place in one phase only (often the liquid phase). A comprehensive review on application of rate-based modeling to catalytic distillation, including treatment of flow non-idealities in catalytic packings and several case studies, was published in 2003 [57]. An interesting approach in the application of ratebased models to the structured packings, and recently 10.6.4.2

References see page 2197

2194

10.6 Catalytic Distillation

developed for heterogeneous catalytic distillation, is the use of computational fluid dynamics (CFD) techniques. The prediction of key parameters for the rate-based models, as for example mass-transfer coefficients, may be achieved using CFD simulations rather than expensive tracer experiments. It has been shown convincingly [58] that CFD may simulate well the flow patterns achieved in structured catalytic packings, and application has been proposed [59, 60] to the design of ‘‘intelligent’’ internals, optimized for each specific process. This type of development could lead to a wider adoption of heterogeneous catalytic distillation in commercial processes. Rate-based models may also be extended to the dynamic simulation of catalytic distillation columns. Dynamic modeling is useful when optimizing start-up and shutdown procedures, as well as solving process control issues. Dynamic modeling has also proved to be a powerful tool in the study of transitions between multiple steady states [61], the existence of which has been recognized in the synthesis of ethers through catalytic distillation [45]. Of course, the dynamic formulation of model equations requires a careful analysis of the whole system to minimize convergence problems during the solution [62]. 10.6.5

Selected Processes

The reactions between alkenes and alcohols to produce ethers were considered to be very promising from the very outset of heterogeneous reactive distillation. Etherification reactions are equilibrium-limited, and take place in the liquid phase on strong acid resins at low temperatures (50–100 ◦ C). Under these conditions, liquid–vapor equilibrium is established at moderate pressures (0.5–1 MPa). These reactions can be schematized as: C4 −C5 isoalkene + C1 −C2 alcohol ← − alkyl-tert-alkyl ether − −− −− → MTBE and ethyl-tert-butyl ether (ETBE) are prepared from isobutene and methanol and ethanol, respectively (see Chapter 13.10). The ethers have been extensively used as gasoline blending components due to their oxygenated nature and superior octane characteristics. Usually, isobutene is not pure, but is mixed with other C4 hydrocarbons of very close boiling points that are practically inert under the process conditions. In fact, even before ethers could assume their current key role in the gasoline pool, these etherification reactions were proposed to separate mixtures of close-boiling hydrocarbons, taking advantage of the inert behavior of alkanes and n-alkenes. The first patent to claim the

use of reactive distillation in etherification was filed by Chevron in 1971 [63], and considered to be a process to separate iso- from n-alkenes. As a consequence of such a difficult separation, in etherification processes unreacted alkenes cannot be recycled. Moreover, in order to obtain a complete conversion, a series of conventional reactors with intermediate separations is needed. This target can be achieved, in principle, using a single catalytic distillation column exploiting the high volatility of the reactants compared to the product. CDTECH, which until 1988 was known as the Chemical Research and Licensing Co. (CR&L), has developed since the early 1980s several processes employing heterogeneous reactive distillation ranging from the separation of close-boiling alkenes [64] to the production of MTBE [65–67] and, more recently, ETBE [68], tertamyl-methyl ether (TAME) [69], and tert-amyl-ethyl ether (TAEE). The CDTECH etherification processes achieved wide commercial implementation; a scheme of the reaction section is shown in Fig. 6. A reactive distillation column, in which catalytic distillation structures (CDModules ) are packed between two tray distillation zones, is used as a finisher, placed downstream of a conventional fixedbed reactor, which converts most of the feed. This is a common feature of every reactive distillation process of etherification that has been commercialized: the potential presence of poisons and the large quantity of heat generated by the exothermic reactions still poses problems regarding the thermal control of the distillation zones, in case the entire conversion takes place in the column. This means that the potential advantages of reactive distillation have not been fully realized in etherification. In spite of this limitation, etherification technologies based on catalytic distillation have rapidly achieved a dominant position in the market. Other licensors and users of conventional etherifications have patented, and in some cases commercialized, reactive distillation processes, including IFP [70], Snamprogetti [20], and UOP [71]. Remarkable activity has taken place in the field of alkylation of aromatics with light alkenes, exploiting their high volatility. For example, CDTECH has developed processes to produce cumene (from propene and benzene) [72] and ethylbenzene. Patents in this field have also been filed by Chevron and the China Petrochemical Corporation. The alkylation of alkenes (in particular the reaction of isobutane with butenes to form isooctanes) is a key process that converts highly volatile C3 –C4 hydrocarbons to valuable products with high octane numbers. This liquidphase reaction is catalyzed by strong acids (H2 SO4 , HF), but this poses both safety and environmental problems. Thus, considerable effort is being exerted to identify

10.6.5 Selected Processes

Feed wash

Boiling point reactor

Catalytic distillation

Methanol extraction

2195

Methanol Recovery

Recycle methanol

Fresh methanol

C4 Raffinate Methanol and C4s Water

Water

Mixed C4s Water and contaminants

Fig. 6

MTBE

Simplified process scheme for the CDMtbe technology developed and commercialized by CDTECH.

suitable solid catalysts for this reaction (see Chapter 13.8), and ExxonMobil has indeed proposed a heterogeneous catalytic distillation [73]. Other reactions involving light hydrocarbons proposed by CDTECH include the desulfurization of gasoline, the selective hydrogenation of dienes and alkynes in streams of hydrocarbons [74], and the isomerization of alkenes [75]. The CDTECH process for gasoline desulfurization includes the CDHydro and CDHDS technologies that provide refiners with the ability to remove sulfur from FCC gasoline while maximizing octane retention and minimizing hydrogen consumption. Both processes are based on reactive distillation technology, and as such they offer unique advantages over fixed-bed technologies. Along with maximum octane retention, these technologies offer an unmatched ability to increase the blending flexibility and control the boiling range of the gasoline. The CDHydro process typically splits the FCC gasoline feed into two fractions; this process also simultaneously removes mercaptan sulfur from light FCC gasoline (LCCS) by converting the mercaptans into heavier sulfides. Once formed in the catalyst zone of the distillation column, the sulfides distill into the bottom product, leaving a low-sulfur light gasoline. The bottom product from the CDHydro unit is fed into the CDHDS unit where the sulfur compounds are hydrogenated to hydrocarbons and H2 S. An unavoidable side reaction is the hydrogenation of part of the olefins, but by using catalytic distillation the saturation of olefins can be minimized. Distilling the gasoline creates a temperature profile in the column; this naturally occurring profile sets up a different reaction severity for various boiling ranges of the gasoline. The heavy portion of the gasoline

tends to distill downward. For the heavy cuts, where the highest percentage desulfurization is required to meet the specification, the reaction temperature is the highest. For lighter cuts, the percentage hydrodesulfurization (HDS) required is lower, and the reaction temperature is also lower. As one reduces the severity and the percentage HDS, the conversion of olefins is also reduced [76]. A simplified scheme of the assemblage of catalytic distillation columns peculiar to this technology by CDTECH is shown in Fig. 7. Other companies have studied the application of reactive distillation to the metathesis of light alkenes [77], and to the HDS of heavy oil residues [78].

LCN CDHydro Hydrogen MCN FCC C5+ Gasoline MCN/HCN

CDHDS

Hydrogen HCN

Simplified process scheme for the CDHydro /CDHDS catalytic distillation columns developed and commercialized by CDTECH. LCN, MCN, HCN indicate light catalytic naphtha, medium catalytic naphtha, and heavy catalytic naphtha, respectively. Fig. 7

References see page 2197

2196

10.6 Catalytic Distillation

Currently, there are several industrial applications of heterogeneous reactive distillation in such fields as the production of dimethyl terephthalate and the transesterification of dimethyl carbonate to diphenyl carbonate, all of which are intermediates for polymers. Mitsubishi holds many patents covering the esterification of aromatics and transesterifications, and a large number of applications have been proposed in the field of fine chemicals. The most recent years have witnessed a continuing wider patenting effort in heterogeneous catalytic distillation, in spite of the still relatively limited field of commercial applications. It is, however, remarkable to note how several large oil, chemical and utility companies have shown interest in this unit operation by filing patents proposing new reactions of interest. Among some of the most recent applications, the hydroxylation of benzene proposed by Shell [79] is worthy of mention, as are the production of lube base stocks from lower molecular-weight feedstocks by Chevron [80] (who were among the pioneering companies in catalytic distillation), and the hydrolysis of ethers by ExxonMobil. Also very interesting as a future perspective might be the application to the synthesis of methanol proposed by ConocoPhillips [81], due to the importance of methanol as intermediate in the production of chemicals. Another interesting patent has been filed by General Electric [82], concerning the production of diaryl carbonates with

a process scheme including three catalytic distillation columns in series. As mentioned above, the technology of reactive distillation was first applied to homogeneously catalyzed esterification and transesterification reactions. Many esters serve as important intermediates in the production of various polymers, and are produced by the liquidphase reaction of the corresponding acid and alcohol, catalyzed by a strong acid (H2 SO4 , H3 PO4 ). Today, highlevel research activities are seeking to substitute these strong liquid acids by solid catalysts, the aim being to avoid problems related to the recovery of catalyst in the reactor effluent and thereby reduce the environmental impact of these processes which, in many cases, have already reached industrial application. The process used to produce methyl acetate from acetic acid and methanol, as developed by Eastman Kodak (now Eastman Chemical Company), was the first breakthrough in catalytic distillation. However, the involved reaction: − MeOAc + H2 O HOAc + MeOH ← −−− −− → suffers in several ways, including chemical equilibrium limitations, difficult separation between acetone and water, and the presence of azeotropes. Conventional processes used one or more liquid-phase reactors with a large excess of one reactant in order to push

MeOAc Separation acetic acid/Methyl acetate

AcOH

Separation water/Methyl acetate H2SO4

Reactive separation

MeOH Separation methanol/water

H2 O

Schematic representation of the catalytic distillation column in the methyl acetate synthesis process developed and commercialized by Eastman Chemical Company.

Fig. 8

References

the conversion of the other reactant. The reaction section was followed by a series of eight distillation columns and one liquid–liquid extraction. In the catalytic distillation technology developed by Eastman [83], the entire process is carried out in a single catalytic distillation column (Fig. 8) which has stoichiometrically balanced feeds. The lighter reactant (methanol) is fed at the bottom, and the heavier acetic acid at the top. The reaction takes place in the middle section, below the injection point of the catalyst (sulfuric acid). In the bottom section methanol and water are stripped from the product, while in the top section the azeotrope between methyl acetate and methanol is broken by the addition of acetic acid. This single column represents an entire chemical plant, but with capital costs and energy consumption only about 20% of that for conventional units. The latest developments in the field of homogeneous catalytic distillation have been focused on the synthesis of alkanolamines [84, 85]. 10.6.6

Future Perspectives

In spite of the potential advantages and several proposed applications, catalytic distillation has been commercialized in only a few, nonetheless important, chemical processes. The reason for this is that several conditions must be met simultaneously in order to make its implementation attractive. The introduction of catalytic distillation into heterogeneous catalysis involves certain requirements, however, including: • the availability of catalysts active at the operating conditions of the distillation • in the case of fixed beds, the availability of catalysts with a sufficient longevity (in the order of one year, at least) and resistance to poisons, so as to avoid excessive shut-downs, catalyst replacements, and start-ups that may be more complex than in conventional operation • a strong difference in volatility between at least one of the compounds involved and the others • a relatively low reaction heat, in order to avoid adverse effects when controlling the distillation process. Nonetheless, the past decade has witnessed major development in three very important fields: • The development and application of catalytic structured packings that should become standard for all applications of heterogeneous catalytic distillation. • The development of a more rigorous and efficient approach to modeling, based on mass-transfer rates and computational fluid dynamics.

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• The development of heuristic rules, which may be applied to specific softwares, to determine rapidly the attraction of catalytic distillation for any specific system. Today, these developments are helping to lower the barriers to the commercial implementation of catalytic distillation. Indeed, the longer-term forecast for its application in the petrochemical industry is that this technique will be used industrially for a much greater number of reactions and processes. For the interested reader, a recently published book by Sundmacher and Kienle [86] provides valuable reference material for a comprehensive treatment of the subject. References 1. A. Backhaus, US Patents 1,400,849; 1,400,850; 1,400,851, 1921. 2. D. B. Keyes, Ind. Eng. Chem. 1932, 24, 1096. 3. D. F. Othmer, Ind. Eng. Chem. 1941, 33, 1106. 4. O. Kageyama, A. Kawano, US Patent 3,506,408, assigned to Daicel Ltd., 1970. 5. E. H. Stitt, Chem. Eng. Sci. 2002, 57, 1537. 6. G. Schembecker, S. Tlatlik, Chem. Eng. Proc. 2003, 42, 179. 7. H. G. Schoenmakers, in European Symposium on Computer Aided Process Engineering, CAPE-12, J. Grievink, J. van Schijndel (Eds.), Elsevier, Amsterdam, 2002, p. 9. 8. G. J. Harmsen, in Proceedings of the 18th World Petroleum Congress, The Energy Institute, London, 2005, ref. RFP-5 (electronic format). 9. K. Jakobsson, A. Pyhalahti, S. Pakkanen, K. Keskinen, J. Aittamaa, Chem. Eng. Sci. 2002, 57, 1521. 10. M. I. Balashov, Theor. Found. Chem. Eng. 1980, 14, 119. 11. M. I. Balashov, L. A. Serafimov, Theor. Found. Chem. Eng. 1980, 14, 332. 12. M. I. Balashov, V. P. Patlasov, L. A. Serafimov, Theor. Found. Chem. Eng. 1980, 14, 406. 13. M. I. Balashov, V. P. Patlasov, L. A. Serafimov, Theor. Found. Chem. Eng. 1980, 14, 495. 14. W. Haunschild, US Patent 3,634,535, assigned to Chevron, 1971. 15. E. M. Jones, US Patent 4,439,350, assigned to CR&L, 1984. 16. E. M. Jones, US Patent 4,536,373, assigned to CR&L, 1985. 17. F. G. Franklin, US Patent 4,471,154, assigned to Chevron, 1984. 18. D. V. Quang, P. Amigues, J. Gaillard, J. Leonard, J. Nocca, US Patent 5,026,459, assigned to IFP, 1991. 19. X. Hao, J. Wang, Z. Yang, Y. Xing, X. Xu, T. Wang, Chinese Patent 1,042,664, assigned to Qilu Petrochemical Corporation, 1990. 20. D. Sanfilippo, M. Lupieri, F. Ancillotti, US Patent 5,493,059, assigned to Snamprogetti, 1996. 21. L. Asselineau, P. Mikitenko, J. Viltard, M. Zuliani, US Patent 5,368,691, assigned to IFP, 1994. 22. E. M. Jones, EU Patent 0,461,855, assigned to CR&L, 1991. 23. W. T. Evans, K. Stork, EU Patent 0,571,163, assigned to CR&L, 1993. 24. N. Yeoman, R. Pinaire, M. A. Ulowetz, O. J. Berven, T. P. Nace, D. A. Furse, US Patent 5,914,011, assigned to KochGlitsch, 1999.

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25. K. T. Chuang, Z. Xu, US Patent 6,045,762, assigned to Governors of the University of Alberta, 2000. 26. L. A. Smith, US Patent 5,345,006, 1994. 27. J. Aittamaa, J. Jakkula, P. Lindqvist, M. Koskinen, J. Linnekoski, O. Krause, M. Sourander, J. Ignatius, A. Pyh¨alahti, US Patent 6,613,108, assigned to Fortum Oil & Gas, 2003. 28. L. A. Smith, US Patent 4,215,011, assigned to CR&L, 1980. 29. R. Arai, in Proceedings of the AIChE Spring National Meeting 1986, American Institute of Chemical Engineers, New York, 1986. 30. L. A. Smith, US Patent 4,250,052, assigned to CR&L, 1981. 31. A. B. Gelbein, M. Buchholz, US Patent 5,073,236, 1991. 32. R. Shelden, J. P. Stringaro, US Patent 5,417,938, assigned to Sulzer, 1995. 33. J. L. De Garmo, V. N. Parulekar, V. Pinjala, Chem. Eng. Progr. 1992, 88(3), 43. 34. J. R. Adams, US Patent 5,057,468, assigned to CR&L, 1992. 35. L. A. Smith, US Patent 5,262,012, assigned to CR&L, 1993. 36. G. P. Hagen, D. A. Palmer, US Patent 5,325,102, assigned to Amoco, 1993. 37. D. Hearn, US Patent 5,266,546, assigned to CR&L, 1993. 38. B. Paikert, J. Lloyd, T. A. Griffin, US Patent 6,277,340, assigned to ABB Lummus Global, 2001. 39. F. J. Brocker, K. Flick, C. F. Erdbrugger, G. Kaibel, G. Meyer, H. J. M¨uller, P. Polanek, E. Schwab, US Patent 6,297,415, assigned to BASF, 2001. 40. H. Komatsu, J. Chem. Eng. Japan 1977, 10, 200. 41. J. H. Grosser, M. F. Doherty, M. F. Malone, Ind. Eng. Chem. Res. 1987, 26, 983. 42. M. F. Doherty, G. Buzad, Comp. Chem. Eng. 1994, 18, S1. 43. A. Izarraz, G. W. Bentzen, R. G. Anthony, C. D. Holland, Hydrocarbon Process. 1980, 59(4), 195. 44. M. Kinoshita, I. Hashimoto, T. Takamatsu, J. Chem. Eng. Japan 1983, 16, 370. 45. S. A. Nijhuis, F. P. J. M. Kerkhof, A. N. S. Mak, Ind. Eng. Chem. Res. 1993, 32, 2767. 46. J. Law, Hydroc. Techn. Int. 1992, 201. 47. L. L. Doraiswamy, M. M. Sharma, Heterogeneous Reactions: Analysis, Examples and Reactor Design, Wiley, New York, 1984, p. 57. 48. I. Suzuki, H. Yagi, H. Komatsu, M. Hirata, J. Chem. Eng. Japan 1971, 4, 26. 49. Y. A. Chang, J. D. Seader, Comp. Chem. Eng. 1988, 12, 1243. 50. H. L. Toor, Am. Inst. Chem. Eng. J. 1964, 10, 448. 51. J. D. Seader, Chem. Eng. Prog. 1989, 85(10), 41. 52. S. S. Katti, Trans. Inst. Chem. Eng. 1995, 73, 595. 53. W. K. Lewis, W. G. Whitman, Ind. Eng. Chem. 1924, 16, 1215. 54. R. Higbie, Trans. Am. Inst. Chem. Eng. 1935, 31, 365. 55. J. O. Hirschfelder, C. F. Curtiss, R. B. Bird, Molecular Theory of Gases and Liquids, Wiley, New York, 1964, p. 539. 56. R. Taylor, R. Krishna, Multicomponent Mass Transfer, Wiley, New York, 1993, 616 pp. 57. C. Noeres, E. Y. Kenig, A. Gorak, Chem. Eng. Proc. 2003, 42, 157. 58. J. M. van Baten, R. Krishna, Catal. Today 2001, 69, 371. 59. M. Kl¨oker, E. Y. Kenig, A. Gor´ak, Catal. Today 2003, 79–80, 479. 60. A. Gor´ak, E. Y. Kenig, P. Moritz, Chem. Eng. Process. 2005, 44, 607 and entire issue No. 6, p. 607–700. 61. K. D. Mohl, A. Kienle, E. D. Gilles, P. Rapmund, K. Sundmacher, U. Hoffmann, Chem. Eng. Sci. 1999, 54, 1029. 62. C. C. Pantelides, SIAMJ Sci. Stat. Comp. 1988, 9, 213. 63. W. Haunschild, M. Willard, US Patent 3,629,478, assigned to Chevron, 1971.

64. 65. 66. 67. 68. 69. 70. 71. 72. 73. 74. 75. 76. 77. 78. 79. 80. 81. 82. 83. 84. 85.

86.

L. A. Smith, US Patent 4,302,356, assigned to CR&L, 1981. L. A. Smith, US Patent 4,336,407, assigned to CR&L, 1982. E. M. Jones, US Patent 4,504,687, assigned to CR&L, 1985. L. A. Smith, US Patent 5,118,873, assigned to CR&L, 1992. A. Bakshi, E. M. Jones, B. A. Strain, US Patent 5,248,836, assigned to CR&L, 1993. E. M. Jones, H. J. Semerak, L. A. Smith, US Patent 5,248,837, assigned to CR&L, 1993. J. Nocca, J. Leonard, J. Gaillard, P. Amigues, US Patent 4,847,431, assigned to IFP, 1989. C. P. Luebke, T. L. Marker, US Patent 5,243,102, assigned to UOP, 1993. J. D. Shoemaker, E. M. Jones, Hydrocarbon Process 1987, 66(6), 57. A. Hass, C. R. Kennedy, US Patent 4,935,557, assigned to Mobil Oil, 1990. D. Hearn, R. P. Arganbright, US Patent 6,169,218, assigned to CDTECH, 2001. D. Hearn, R. P. Arganbright, E. M. Jones, L. A. Smith, US Patent 6,495,732, assigned to CDTECH, 2002. G. G. Podrebarac, Preprints, Div. Fuel Chem., Am. Chem. Soc. 2003, 48(2), 251. C. W. Jung, P. E. Garrou, G. R. Strickler, US Patent 4,709,115, assigned to Dow Chemical, 1987. A. J. Dennis, G. E. Harrison, D. H. McKinley, US Patent 5,252,198, assigned to Davy McKee, 1993. D. M. Hamilton, US Patent 6,900,358, assigned to Shell Oil, 2005. R. R. Krug, D. J. O’Rear, US Patent 6,841,711, assigned to Chevron, 2005. J. S. Allison, H. A. Wright, T. H. Harkins, D. S. Jack, US Patent 6,723,886, assigned to ConocoPhillips, 2004. P. R. De Bruin, J. S. Law, V. A. Vriens, US Patent 6,767,517, assigned to General Electric, 2004. V. H. Agreda, L. R. Partin, W. H. Heise, Chem. Eng. Progr. 1990, 86(2), 40. W. Peschel, A. Hildebrandt, B. Bessling, US Patent 6,696,610, assigned to BASF, 2004. D. Garg, S. N. Shah, M. J. Okasinski, A. S. Drayton-Elder, US Patent 6,846,959, assigned to Air Products and Chemicals, 2005. K. Sundmacher, A. Kienle, Reactive Distillation. Status and Future Directions, Wiley-VCH, Weinheim, 2003, 288 pp.

10.7

Catalytic Membrane Reactors .. Roland Dittmeyer and Jurgen Caro∗

10.7.1

Introduction

Catalytic membrane reactors have received increased attention over the past two decades, initially due to progress made in the field of inorganic membranes. This is because conventional polymeric membranes, despite having ∗

Corresponding author.

10.7.2 Types of Membrane Reactor

achieved a degree of commercial success in membrane reactors (mainly in biotechnology and in other lowtemperature applications [1–7]), have a limited thermal, chemical and mechanical stability. Inorganic membranes, on the other hand, can withstand high temperatures, they are chemically much more stable under the often harsh environments of industrial reactors, and they are more robust. Advanced inorganic membranes which hold promise for selective separations under reaction conditions, for example based on the molecular dimensions of the reactants, or on component-specific chemical interactions with the membrane material, have become increasingly available. Examples include crystalline zeolite films, amorphous microporous oxide membranes, and mixed ion- and electron-conducting solid oxides. These materials have achieved impressive separation factors for different systems – not always as high as that of palladium alloy membranes for hydrogen, as proposed during the 1960s by Gryaznov and others [8–10], but high enough to initiate dreams of advanced engineered catalysts with built-in shape-selectivity effects. In addition, from an engineering standpoint, the vision of process intensification through multifunctional reactors first conceived during the early 1990s [11, 12], has breathed new life into research on catalytic membrane reactors, such that they may permit the elimination of process steps and hence lead to more compact and cost-efficient plants than a conventional design with separate units. The field of catalytic membrane reactors is an interdisciplinary research area that mainly connects the membrane material sciences with catalysis research and chemical engineering. The field is discussed on a regular basis at the International Conferences on Catalysis in Membrane Reactors (ICCMR), a series which commenced in 1994 in Lyon [13], and with follow-up events in Moscow (1996) [14], Copenhagen (1998) [15], Zaragoza (2000) [16], Dalian (2002) [17], Lahnstein (2004) [18], and Cetraro (2005) [19]. In addition, topical sessions on catalytic membrane reactors are frequently held at many other international conferences. Catalytic membrane reactors are addressed in numerous publications (Fig. 1), and in many review articles and book chapters [20–47], including that by Dalmon in the First Edition of this Handbook. In addition, comprehensive treatments have been provided in the books by Hsieh [48] and by Sanchez-Marcano and Tsotsis [49]. In this chapter it is impossible to cover all of the developments and applications described; rather, the aim is to: (i) illustrate the generic concepts of how to combine catalytic reactions and membrane separations in one unit; and (ii) show, by example of selected applications, how far the field has evolved and which problems are still to be solved. In order to achieve this, Section 10.7.2 introduces

2199

a classification of catalytic membrane reactors according to functional aspects, while Section 10.7.3 focuses on achievements in the field of membrane reactors equipped with hydrogen-selective membranes (i.e., mainly dense metal membranes based on palladium and its alloys). Section 10.7.4 is devoted to the use of oxygen-selective membranes, above all mixed ion- and electron-conducting ceramics, for selective oxidation of hydrocarbons. Section 10.7.5 analyzes applications not relying on selective membrane transport; this concerns membranes for distributed feeding of reactants and catalytic membrane contactors. Section 10.7.6 addresses some issues related to the design of catalytic membrane modules, especially for high-temperature applications, and in Section 10.7.7 the development of the field is assessed in total. 10.7.2

Types of Membrane Reactor

According to the IUPAC definition, a membrane reactor is a device combining a membrane-based separation and a chemical reaction in one unit [50]. Separation thereby implies that the membrane shows a preferred permeation for one or several of the constituents of the reactant mixture. For this functional integration various possibilities exist (Fig. 2). The most common concept is the selective removal of products from the reaction zone (Fig. 2a), known as the extractor membrane reactor. This is mainly applied to equilibrium-restricted reactions to increase the yield of desired products beyond the equilibrium value. Dehydrogenation reactions have been among the first applications. In Section 10.7.3.2.1, the current status of this field is discussed. On the basis of the intention to overcome the equilibrium restriction it may be obvious to note that the conversion must not be limited by the reaction kinetics; that is, the reaction must be sufficiently fast compared to the mass transport through the membrane, a feature summarized by the term ‘‘kinetic compatibility’’. The ultimate target would be to drive the reaction to completion, but for this those products not removed through the membrane (e.g., the olefin in alkane dehydrogenation) must not inhibit the reaction kinetically, for example by competing with the educts for adsorption sites. An especially favorable situation is found if the desired product is removed through the membrane selectively enough for the target application. Then, the membrane reactor not only benefits from overriding the equilibrium constraint but also provides an integrated product purification which may allow the number of process units to be decreased (i.e., reduced investment costs). This is particularly attractive if it eliminates interstage temperature and/or the pressure changes required References see page 2241

2200

10.7 Catalytic Membrane Reactors

300

300

250

250

Patents

200

Total

Total

200 150

150

100

100

50

50

0

1970

1980

1990

0

2000

Patents + CATALY? in any of the text fields

1970

1980

1990

2000

Left: Number of publications found by a search in the CAS database for the keywords Membran? and React? in the title. Right: Number of publications and patents with the additional keyword Cataly? in any of the text fields.

Fig. 1

IMCR or CMR S (opt.)

S/P2

Permeate

CNMR (opt.)

F2

(opt.)

F2 Reaction zone F

P1

Reaction zone (Retentate)

(a) Selective removal of products via the membrane

F1

P

Reaction zone

(b) Selective supply of reactants via the membrane

Permselective membrane

Removal or supply

(c) The membrane defines the reaction zone

Catalytic membrane P2

P2

F2 or S P1

or

F2 or S

P2

Catalyst F1

P

Non-permselective membrane

Inert membrane F2 or S

F1

Packed-bed catalyst

P1

(d) Packed-bed catalytic reactor - Packed-bed catalyst - Removal of products or supply of reactants

F1

Removal or supply

P2

Catalyst P1

(e) Catalytic membrane reactor - Membrane acts as sole catalyst - Selective permeation of P1 or P2

F1

Packed-bed catalyst

P1

(f) Packed-bed catalytic membrane reactor

- Catalytic membrane and packed-bed catalyst - Removal of products or supply of reactants

IMCR = Inert membrane catalytic reactor; CMR = catalytic membrane reactor; CNMR = catalytic non-permselective membrane reactor; opt. = optional; F1, F2 = reactant feed; S = sweep; P1, P2 = products Fig. 2

Generic concepts of catalytic membrane reactors. (Reproduced from Ref. [41]; reprinted with permission from Elsevier.)

10.7.2 Types of Membrane Reactor

in a conventional process scheme. A higher yield, selectivity, and/or productivity of the membrane reactor due to the beneficial effects of product removal on the reaction kinetics may lead to additional savings on energy and raw materials. The most prominent examples are membrane reformers for the production of fuel cell-grade hydrogen from hydrocarbons or alcohols (cf. Section 10.7.3.2.3). In a different approach, one or several of the reactants are fed through a membrane to the reaction zone (Fig. 2b). This so-called distributor membrane reactor does not necessarily require a selective membrane, provided that the trans-membrane flux can be properly controlled by the differential pressure. There may be different aims: one would be to establish a more uniform concentration of the dosed species along the reactor, if this provided a higher selectivity or yield, or an improved safety. The total amount of reactants that can be supplied without exceeding the concentration limit in the feed for safe operation can also be increased. If suitable membranes were available, or a design in sections were feasible, a specific shape of the concentration profile could be attempted, if this paid back in terms of improved reactor performance due to kinetic reasons. One additional effect observed when applying a distributed feed along the reactor is an increasing flow rate downstream of the reactor. Together with the fact that fluid elements entering the reactor further downstream have a reduced residence time compared to reactants supplied at the entrance, this alters the residence time distribution and may improve the selectivity or yield of intermediate products in multiple reactions. An additional benefit on the process level results also in the distributor membrane reactor from a membrane being able to extract the desired reactants selectively from a mixture with undesired components; then again, one process unit may be eliminated. Oxygen ion-selective ceramic membranes for catalytic partial oxidation fall into this category (cf. Section 10.7.4.2). Finally, a third concept refers to the case where the membrane is used to set the reaction zone. For reactions relying on a catalyst, an active material could be incorporated into the membrane uniformly or only in a certain region, for example a surface layer. Two reactant streams could then be passed along the different sides and would mix in the catalytic zone by diffusion (Fig. 2c). One consideration for this type of configuration might be a porous membrane and a solid active material coated to the pore walls, due to the higher trans-membrane fluxes that such systems allow compared to dense membranes. Yet, other types are also conceivable, for example a molecular catalyst dissolved in a liquid filling the pores. Even if no catalyst is required, the same principle can be applied if the reaction is fast enough to reach complete conversion within the membrane. An alternative to relying on diffusion as the transport mechanism is to push a premixed reactant stream through

2201

a catalytic membrane by a pressure difference. In the more likely case of a porous membrane, convection will occur inside the pores and this will result in a very efficient contact between the fluid and the active phase (which may be present in the form of attached nanosized particles) on the pore walls. The pore size of such membranes would be chosen in the micrometer or submicrometer range, and the contact time per pass would be short. Thus, the arrangement can be viewed as a short contacttime catalytic microreactor or nanoreactor. The aim of both principles – that is, the catalytic diffuser and forced through-flow catalytic membrane – is to optimize contact between the reactants and the active phase to exploit its intrinsic catalytic properties. In the latter variant, the membrane has no dosing function and can be viewed as a special type of structured catalyst. As mentioned above, the mass transport across a membrane can be either permselective, if only some components of a mixed feed permeate through it (Fig. 2a,b), or non-permselective, if all species pass through at comparable rates (Fig. 2c). Permselective transport, which is found in dense membranes, is governed by a solution–diffusion mechanism. Nonpermselective transport normally occurs in macroporous and mesoporous membranes; in the latter, Knudsen diffusion is often the dominating transport mechanism. Microporous membranes show both activities, with both permselective and non-permselective transport being possible depending on the size of the permeating molecules compared to the pore size, and on their interaction with the membrane material. When the membrane reactor is used for carrying out a catalyzed reaction, the question arises as to whether the membrane itself has a catalytic function. If it does act as a catalyst, this is referred to as a catalytic membrane reactor (CMR; Fig. 2e,f), but if not it is known as an inert membrane catalytic reactor (IMCR; Fig. 2d). The CMR case is further subdivided into two categories: (i) when the membrane acts as the only catalyst (Fig. 2e); or (ii) when a conventional catalyst is present in addition (Fig. 2f). The interested reader should consult Refs. [25, 36] for a systematic classification of membrane reactor concepts. Besides the elimination of process units (process level) and the utilization of synergy effects from the integration of reaction and mass transport into one unit (unit level), a third level exists where membranes may offer advantages for catalyzed reactions. This is the reaction level; it requires that the catalyst is an integral part of the membrane. With dense membranes employed in a distributor membrane reactor it is possible that the membrane, due to its chemical nature, supplies one of the reactants in a special form, perhaps more References see page 2241

2202

10.7 Catalytic Membrane Reactors

active and/or selective in the reaction that one wishes to catalyze than in its usual form. An example is a ceramic oxygen ion-conducting membrane, which can pass oxygen ions to a solid catalyst attached to it; thus, the use of molecular oxygen from the reactant gas phase is avoided. Other examples include silver membranes that selectively permeate atomic oxygen, and membranes constructed from palladium or its alloys which show exclusive permeability to atomic hydrogen. Polymeric protonexchange membranes and ceramic high-temperature proton-conducting membranes represent another class of membranes that could be used for such purpose. 10.7.3

Catalytic Reactors with Hydrogen-Selective Membranes Membrane Types Selective permeation of hydrogen through membranes requires dense or microporous membranes, of which different types may be distinguished under material and functional aspects. 10.7.3.1

Organic Polymers As dense organic polymer membranes function according to a solution–diffusion mechanism, the permeability of a certain compound is determined by the product of its solubility and its diffusivity in the membrane material. The separation factor therefore depends both on the ratio of the diffusivities and on the ratio of the solubilities of the species to be separated; often, these two function in opposite directions. Consequently, organic polymer membranes offer only modest selectivity for the separation of hydrogen from gas mixtures. They are, however, used in industry to recover hydrogen from purge gas streams in refineries (e.g., hydrocracking, hydrotreating, hydrogenation units) and petrochemical plants (e.g., ammonia and methanol synthesis), and also for adjusting the H2 : CO ratio in syngas production. Although they compete with cryogenic separation and pressure swing adsorption, they have established a significant market share, with many hundred installations. PRISM membranes (distributed by Air Products) may serve as a typical example. These are phase-inversion polysulfone-based hollow fibers, coated with a thin layer of silicon rubber which serves to seal the defects that are not completely avoidable in the manufacture of very thin separation layers. The permselectivity (cf. Ref. [50] for a definition) for hydrogen from carbon monoxide is around 35 with this type of membrane [51]; the maximum temperature is limited to about 100 ◦ C. 10.7.3.1.1

10.7.3.1.2 Solid Polymer Electrolytes The second group of materials are the solid organic polymer electrolyte

membranes; these are dense membranes that are permeable only to protons. Nafion , which is prepared by sulfonation of polytetrafluorethylene (PTFE), is the best known example. Depending on the copolymer side chains holding the HSO3 end groups, different types of Nafion may be distinguished. The proton in the HSO3 group is bonded ionically, and so the end of the side chain is actually an SO3 − ion; for this reason, the resulting structure is called an ‘‘ionomer’’. Transport of protons through the ionomer requires cleavage of the SO3 − H+ ionic bond, and this is facilitated by the presence of water collecting around clusters formed by aggregation of the hydrophilic sulfonated side chains within the hydrophobic PTFE structure. Nafion requires close to 100% relative humidity to reach high proton conductivity. Therefore, its use as a membrane, mainly in low-temperature fuel cells, is limited to temperatures below 100 ◦ C (at normal pressure). Some more recently developed solid polymer electrolyte membranes tolerate higher operating temperatures of up to 150–200 ◦ C; one material, which may function without humidification is polybenzimidazole, developed by PEMEAS (former Celanese). However, as these materials cannot be applied at temperatures above 200 ◦ C, their use in catalytic membrane reactors is restricted. Moreover, when using ionomer membranes in a membrane reactor it is necessary to use electrodes and current collectors, and this complicates the reactor design. Metals Dense membranes made from metals that selectively absorb atomic hydrogen are well-suited for high-temperature hydrogen separation as they offer infinite selectivity for hydrogen. Palladium and its alloys (e.g., Pd77 Ag23 , Pd58 Cu42 , Pd94 Ru6 (wt.%)) must be mentioned first in this context. Alloying of the palladium reduces the lattice distortion upon hydriding and dehydriding, and this improves the resistance against hydrogen embrittlement (palladium fails below 300 ◦ C). The alloy with 23 wt.% silver is the most common material for hydrogen-selective membranes. This has higher hydrogen permeability than pure palladium, and it is not damaged by hydrogen embrittlement down to room temperature. However, grain coarsening takes place at high temperature (>420–450 ◦ C), and this may lead to pinhole formation, especially in thin films. The main benefit of the alloy with 42 wt.% copper is a better resistance against hydrogen sulfide, while the alloy with 6 wt.% ruthenium has excellent high-temperature stability and strength. Metal membranes are produced in the form of selfsupporting foils or thin tubes. Owing to inclusions which lead to pinholes, their thickness normally is at least 10–50 µm in order to reach sufficiently low leak rates 10.7.3.1.3

2203

10.7.3 Catalytic Reactors with Hydrogen-Selective Membranes

(typical thicknesses of foils are 25–100 µm) that limit the hydrogen flux and lead to high material costs. It has been known since the late 1960s that Group Vb metals (V, Nb, Ta) exhibit an even higher hydrogen permeability than palladium [52], but to facilitate the entry and exit of hydrogen their surface must be protected against oxide formation [53]. Consequently, composite membranes have been developed which consist of a supporting foil from a Group Vb metal and a thin (0.2–0.3 mm) still are limited to low hydrogen fluxes (i.e., below ca. 0.3 mN 3 m−2 h−1 ) at 950 ◦ C and 105 Pa (1 bar) hydrogen feed pressure [114]. Composites consisting of separate ceramic proton- and electron-conducting phases did not show much better performance. In contrast to this, cermets with BCGO or La2 Zr1.8 Y0.2 O6.9 as the proton-conducting phase and nickel alloys as the metallic phase have reached promising hydrogen fluxes in the range of 0.6 mN 3 m−2 h−1 at temperatures around 700–950 ◦ C [119]. The metal phase thereby also improves the structural stability and the surface catalysis for hydrogen dissociation. If highly hydrogen-permeable metals are used in the cermet instead of a nickel alloy, the hydrogen flux is further increased and the application range extended towards lower temperatures (i.e., down to 550 ◦ C) [120]. Very recently, supported thin-film cermet membranes from Pd−YSZ with thicknesses down to ca. 22 µm were prepared which reached a hydrogen flux of 12 mN 3 m−2 h−1 at 900 ◦ C and 90 kPa hydrogen feed pressure [121]. The highest hydrogen fluxes of all types of dual-phase membrane were obtained with socalled intermediate-temperature cermets. These consist of a Group Vb (or another low-cost hydrogen-permeable metal) as the hydrogen-conducting phase and a metal oxide to provide the required structural stability [122]. The formation of a cermet is difficult due to the higher reactivity of these metals compared to palladium, and the temperature window is narrow (340–440 ◦ C). However, supported thin-film membranes based on such materials showed hydrogen fluxes up to 254 mN 3 m−2 h−1 at 400 ◦ C and at a hydrogen feed pressure of 33 × 105 Pa [120]. 10.7.3.1.6 Final Remarks The transition between these membrane categories is fluent. Composites have been developed to improve the limitations of certain materials, for example the so-called ‘‘mixed matrix membranes’’. In these systems, inorganic materials such as zeolites

or other adsorbents are incorporated into polymeric membranes to improve the thermal stability or the separation factor [123, 124]. The typical performances of different hydrogenselective membranes are listed in Table 1. Selective Removal of Hydrogen The ability to drive a hydrogen-producing equilibrium reaction to complete conversion by removing the hydrogen from the reactor through a membrane is an aged concept that dates back to the pioneering studies of Gryaznov [8, 9] and Pfefferle [10] on palladium membranes. The concept is outlined in Fig. 2a. The hydrogen-producing reaction occurs in the gas phase over a conventional catalyst surrounded by the membrane. Transport through the membrane is driven by the hydrogen partial pressure difference between the catalyst side – that is, the retentate, and the permeate side. Three different options exist to create the driving force: 10.7.3.2

• An inert sweep gas in the permeate compartment (e.g., nitrogen, helium, steam). • A pressure difference between the retentate and permeate compartments, created either by an overpressure in the retentate or by evacuation of the permeate, or by both. • A reactive sweep gas consuming the permeated hydrogen (e.g., oxygen/air, CO, unsaturated hydrocarbons). Inert sweep gas is often the worst choice, because it has to be provided and compressed, and diluted hydrogen is produced to be used only as a fuel gas with low calorific value. Steam is an exception because pure hydrogen is easily obtained by condensation if the membrane is sufficiently permselective. If the kinetics permits the hydrogen-producing reaction to be run at elevated pressure (i.e., in the range of 1–3 MPa), this is the preferred option, because the pressure difference is then high enough to drive off the permeated hydrogen. A general drawback of membranes for hydrogen separation is the pressure loss over the membrane, which could make costly recompression necessary. If pure hydrogen, for example for fuel cells, is the target product, then an increased retentate pressure is the only viable option. If not, reacting the permeated hydrogen with the sweep gas is promising. Generally, this is the most effective way to create a high driving force for the permeation. Alkane dehydrogenations to olefins represent one example (cf. Section 10.7.3.2.1A). These are endothermic reactions where it would be favorable to generate the required heat directly in the permeate compartment. An exothermic hydrogen-consuming reaction – for example, oxidation of the permeated hydrogen by air or hydrogenation of an organic compound – could be carried out for

10.7.3 Catalytic Reactors with Hydrogen-Selective Membranes

Tab. 1

2207

Typical performance of different hydrogen-selective membranes

Membrane type

T/ ◦ C

Thickness of the separation layer/µm

H2 -flux at
pH2 = 2–1 bar/ m3N m−2 h−1

Separation factor

References and comments

Organic polymer

0.5 h−1 . Even if all the hydrogen were to be removed instantaneously (irreversible reaction; see Fig. 10c), the predicted styrene yield would not increase by more than 7% (WHSV = 1 h−1 ) or 13% (WHSV = 0.5 h−1 ) over that of the conventional packed bed under these conditions. The highest predicted styrene yield for instantaneous hydrogen removal is 93%. However, for the permeance measured with H2 /EB/ST mixtures (cf. Fig. 7), the simulation predicts a maximum styrene yield of 66.7% for a space velocity close to 0.3 h−1 (Fig. 10b). The hydrogen removal under these conditions is only 18%, which shows that the performance was limited by the hydrogen permeance of the membrane. An attempt was made to remove more hydrogen by reducing the S/O-ratio from 2 to 1, as this would increase the hydrogen partial pressure in the packed bed and thus the driving force for the permeation. Although the concept worked, the catalyst deactivated quickly and the styrene yield fell by 15% within 20 h. To increase the membrane area per gram catalyst, the amount of catalyst was reduced to one-fifth by diluting the packed bed with inert alumina spheres. While this allowed 50% of the hydrogen to be removed, the catalyst deteriorated slowly due to loss of the promoter potassium to the alumina spheres. B Alcohols Among alcohol dehydrogenation in extractor membrane reactors, the conversion of ethanol to acetaldehyde has received the greatest interest. The reaction was proposed as an alternative to the Wacker process for the manufacture of acetaldehyde by oxidation of References see page 2241

100 90 80

XEB YST SST XEB YST SST

Equilibrium conversion (calculated by aspen plus, release 10.0)

70 60

Conventional packed-bed Membrane reactor

50 40 0

0.25

0.5

0.75

1

1.25

WHSV/h−1 Performance of a palladium composite membrane reactor for ethylbenzene (EB) dehydrogenation to styrene (ST) versus a conventional packed-bed. EB conversion, ST yield and ST selectivity as a function of the space velocity. Conditions: T = 580 ◦ C; retentate pressure 110 kPa; permeate pressure 15 kPa; nitrogen sweep gas flow 900 mLN min−1 ; steam to EB ratio S/O = 2 kg kg−1 ; membrane .. thickness 8.5 µm; membrane area per gram catalyst 1.2 cm2 g−1 ; conventional K-promoted iron oxide catalyst (Sud-Chemie).

Fig. 9

2214

10.7 Catalytic Membrane Reactors

Simulation of ethylbenzene (EB) dehydrogenation to styrene (ST) in a laboratory membrane reactor. Conversion, yield, selectivity and removed fraction of hydrogen as a function of the space velocity. (a) Conventional packed-bed. (b) Membrane reactor with hydrogen permeance as measured in H2 /EB/ST mixtures (5 × 10−5 mol m−2 s−1 Pa−0.65 ; cf. Fig. 7). (c) Ideal case of 100% hydrogen removal (→ irreversible reaction). Same conditions as in Fig. 9.

Fig. 10

ethene [171]. Amandusson et al. [172] used a 25 µm-thick palladium membrane at 350 ◦ C, whereupon dehydrogenation occurred at the palladium surface and led to the formation of a carbonaceous layer. However, it was found that by adding oxygen to the ethanol feed the deactivation could be avoided and continuous operation becomes possible. Later, the same group extended their studies to palladium–silver foils, which proved to be much more efficient for ethanol dehydrogenation, and also to methanol as the reactant [175]. Thin-film palladium–silver membranes were applied to ethanol dehydrogenation by Keuler and Lorenzen [176]. The membranes were prepared by electroless plating on the inner side of porous α-alumina tubes, such that the palladium thickness was about 2 µm. Instead of relying on the catalytic activity of the palladium membrane surface, a Cu/SiO2 -catalyst was packed into the membrane tube. The best results were obtained at

275 ◦ C, with the membrane reactor giving a significant increase in conversion versus the conventional packed bed – from 45% to 60% at lower flow rates, and from 36% to 46% at higher flow rates. The selectivity to acetaldehyde was in excess of 90% at 275 ◦ C [176]. The membranes were also applied to the dehydrogenation of 2-butanol to methyl-ethyl-ketone (MEK) [169] in the temperature range 190 to 240 ◦ C. Again, the membrane reactor performed significantly better than the conventional packed bed; that is, the conversion increased from 80% to 93% at 240 ◦ C, and the selectivity to MEK was above 96%. In addition, membranes from nickel-based amorphous alloys were proposed for ethanol dehydrogenation. Examples include NiB [174] as well as NiP and NiP−Cu [177], which were prepared by a modified electroless plating method. These membranes also led to an improved reactor performance due to a combination of the catalytic effect of the membrane material with its hydrogen separation

10.7.3 Catalytic Reactors with Hydrogen-Selective Membranes

capability. Another example is methanol dehydrogenation to methyl formate; this reaction was performed at a temperature of ca. 250 ◦ C on mesoporous silica membranes which were modified by the incorporation of palladium [179] or copper [180] as the catalytically active component. 10.7.3.2.2 Water Gas Shift Reaction The water gas shift (WGS) reaction [Eq. (2)] is an important step in the manufacture of hydrogen by steam reforming, partial oxidation, or autothermal reforming of hydrocarbons or alcohols, or by the gasification of solid fuels such as coal, biomass, or organic waste. It eliminates the majority of the unwanted CO produced in the primary reaction and generates additional hydrogen. This facilitates the following gas cleaning steps if pure hydrogen is the target.

← − CO2 + H2 CO + H2 O − −− −− →

(2)

The reaction is reversible and mildly exothermic. In order to achieve a low CO content, it is normally performed in two steps – that is, a high-temperature shift (320–470 ◦ C, Fe−Cr-oxide catalyst) followed by a lowtemperature shift (180–270 ◦ C, Cu−ZnO catalyst). The reason for this is a conflict between the thermodynamics and the kinetics typical of exothermic equilibriumrestricted reactions; an increased temperature gives a higher space time yield, but reduces the equilibrium conversion. With combined high-temperature and lowtemperature shift, CO contents of 0.3 to 0.8% are reached in the product gas; higher conversions are prevented by a marginal reaction rate at further reduced temperature. The moderate temperature and the absence of vigorous coking make a membrane reactor an attractive option for the WGS reaction, because through hydrogen removal the equilibrium constraint at high operating temperature can be overcome and so an increased space time yield is expected. The concept was successfully demonstrated in many studies targeted at hydrogen production from hydrocarbons or coal [181, 182, 184, 193, 194, 196], the recovery of tritium from tritiated water in the fusion reactor fuel cycle [183, 185, 189], the removal of CO2 for integrated gasification combined-cycle power stations [187], and at the conversion of residual CO and hydrogen in the solid oxide fuel cell anode off-gas [249]. Both dense and porous membranes were used. Among the latter, mesoporous membranes (i.e., γ -Al2 O3 , Vycor glass and similar) were unsuccessful due to their poor separation factors for hydrogen versus carbon oxides [186, 188]. Microporous amorphous silica membranes, as yet, have also failed to demonstrate sufficiently high separation factors over more than a few days at 350 ◦ C in a WGS reaction environment, due to

2215

their poor hydrothermal stability [83, 196]. More recently developed microporous silicon carbide membranes show more promise, however, with hydrogen permeances of 0.5 to 3 mN 3 m−2 h−1 bar−1 and a permselectivity for hydrogen versus nitrogen of 50 to 200 being reported for 200–250 ◦ C, without aging over a period of six months under hydrothermal conditions [197, 198]. Based on these data, it was predicted that a silicon carbide membrane reactor for low-temperature WGS could operate with a much lower steam to CO ratio than a conventional packed bed, and this would be associated with substantial energy savings. Thus, with a small amount of steam as sweep gas (ca. one-fourth of the feed), which is easily removed by condensation, the majority of the hydrogen could be recovered, and a CO content of ca. 200 ppm expected. This should be compared to 0.3 vol.% for a conventional WGS reactor [197]. Carbon molecular sieve membranes also showed promising performance in low-temperature WGS membrane reactors at 250 ◦ C [197]. Unfortunately, porous membranes in general do not allow the necessary stringent requirements to be met for supplying proton-exchange membrane (PEM) fuel cells with hydrogen (99.99%

PSA (Pressure swing adsorption)

CO2 + H2

180 to 220 °C

Membrane reformer system

Steam (H2O)

Membrane reformer (Steam reforming) (Separation) 550 °C

City gas (CH4)

Hydrogen Desulfurizer

H2: >99.99%

Schematic process design of a conventional natural gas reformer system with hydrogen purification by water gas shift followed by pressure swing adsorption, and comparison with a membrane reformer system. (Source: Tokyo Gas [225].)

Fig. 12

2218

10.7 Catalytic Membrane Reactors

occurs in the same unit; thus, it runs at a much higher temperature than in a conventional WGS reactor, while the equilibrium constraint is eliminated in the same way by hydrogen removal through the membrane. In addition to the studies conducted by Oertel et al. [199], Uemiya et al. [200, 201], Kikuchi et al. [202] and others during the late 1980s and early 1990s (cf. Table 2), which demonstrate conversions above the equilibrium value in membrane reformers equipped with palladium-based membranes for methane steam reforming, many reports have also been published on hydrocarbon reforming in membrane reactors. A Packed-Bed Membrane Reformers The majority of experimental studies on membrane reformers has focused on steam reforming of methane, followed by dry reforming of methane and steam reforming of methanol; higher hydrocarbons such as n-butane, n-hexane or octane were used in some cases. Most experiments were performed in small-scale, packed-bed membrane reactors. An example is the double-jacketed design used by Lin and Rei [240, 242] for methanol reforming (Fig. 13), which enables coupling of the heat released by the oxidation of the residual hydrogen as well as the unconverted methanol and CO in the retentate with the endothermic reforming. A composite membrane, namely a 20 µm-thick palladium film on the outside of a porous stainless steel tube, was used for hydrogen removal. The system was operated at 350 ◦ C and 1.5 MPa, with a WHSV of 1 h−1 and a permeate pressure of 0.1 MPa for 900 h, without degradation. Air was not supplied in this experiment. The unit delivered a stable hydrogen flux of 3.5 mN 3 m−2 h−1 with hydrogen purity above 99.9%;

the methanol conversion reached 95%. With air supplied to the outer compartment, methanol and CO conversion was complete. The retentate pressure and the methanol feed flow rate are the main operating parameters that determine hydrogen recovery and productivity. For example, at 1.5 MPa and 100 mol m−2 h−1 methanol feed flow rate, the hydrogen recovery was 73% for a WHSV of 5 h−1 , and the unit delivered a hydrogen flux of 5.4 mN 3 m−2 h−1 [240]. Reducing the load increased the hydrogen recovery, for example up to 97% for a methanol feed flow rate of 50 mol m−2 h−1 . It was deduced from the energy balance that the system becomes autothermal for 74% hydrogen recovery if heat losses to the surroundings are minimized. This technology has been developed further by the company Green Hydrotec, of Kweishang, Taiwan, which offers compact so-called ‘‘3-in-1’’ hydrogen generators with hydrogen capacities up to 10 mN 3 h−1 operating on aqueous methanol solutions [254]. The size of the largest unit is 0.6 × 0.6 × 0.6 m, giving a volume of 0.22 m3 and thus a system space time yield of 46.3 mN 3 m−3 h−1 . In methanol reforming there is no equilibrium constraint above 250 ◦ C. Commercial Cu/ZnO/Al2 O3 catalysts operated at 260 ◦ C, 0.1 MPa and a WHSV of 5 h−1 reach close to 100% conversion, which translates into a hydrogen productivity of ≈3200 mN 3 m−3 h−1 (undiluted catalyst bed). Moreover, CO contents below 0.5 vol.% are reached [255]. As there is no large synergy between methanol reforming and hydrogen removal at the reaction level, what remains as an advantage of the membrane reformer is the integrated purification of the hydrogen. However, this might also be achieved in a separate permeation module following a conventional packed-bed reformer.

Permeated H2 MeOH/H2O feed

Exhaust gas 2nd compartment (oxidation, Pd/Al2O3) 1st compartment (reforming, Cu/ZnO/Al2O3) Pd membrane tube 1st jacket SS tube 2nd jacket SS tube

Electric furnace Retentate

Air

Fig. 13 Schematic design of the double-jacketed membrane reformer used by Lin and Rei [242] for methanol reforming. (Reproduced from Ref. [242]; reprinted with permission from Elsevier.)

10.7.3 Catalytic Reactors with Hydrogen-Selective Membranes

Many experimental studies on methane steam reforming employing palladium-based membranes, either as self-supporting foils or tubes or as thin films on porous substrates, showed that the conversion can be increased by hydrogen removal, and that this synergy effect is most pronounced at moderate temperatures of 500–600 ◦ C. Above 700 ◦ C, the equilibrium conversion is already greater than 80%, so there is little potential for further improvement through hydrogen removal. Below 400 ◦ C, the reaction would not reach a satisfactory conversion until a very large portion (>90%) of the produced hydrogen is removed [203]. Within this temperature window, self-supporting metal membranes have no commercial perspective in membrane reformers due to the high costs of these metals and the low hydrogen permeance. Therefore, thin-film composite membranes have been increasingly evaluated for methane steam reforming over the past decade (for details, see Ref. [219]). Dense palladium–alloy films on porous supports were reported with thicknesses down to less than 1 µm on high-quality asymmetric porous ceramic supports which have a smooth surface with uniform pore size. However, the membranes must offer stable long-term performance at high temperature under a significant pressure gradient, which brings about the need for robustness. Moreover, they must be integrated into modules without too-elaborated sealing. Consequently, metallic supports are still preferred, along with film thicknesses above 6–7 µm, and in most cases even between 10 and 20 µm [219]. Membranes with metallic supports for use above 400 ◦ C, on the other hand, require a porous inorganic intermediate layer to prevent intermetallic diffusion [256]. Different strategies are being pursued to form a well-adherent, highly permeable intermediate layer for thin and durable metallic membranes. In this respect, surface oxidation of the support [257], deposition of a porous ceramic layer by thermal spraying, sputtering or wet-powder spraying [67], slip-coating and other methods have each been applied. Tong et al. recently reported a series of interesting results on methane steam reforming [222–224, 226]. The group developed an improved technique, called ‘‘multidimensional plating’’ [258], to prepare thin palladium membranes on rough surfaces as are typical of stainless-steel supports. The membranes (for descriptions, see Ref. [222]) had a 6 µm-thick palladium layer directly on the steel tube. Excellent hydrogen fluxes of 21–24 mN 3 m−2 h−1 at 500 ◦ C and 105 Pa hydrogen pressure difference were reached, and infinite permselectivity for hydrogen. Membrane reactor experiments were performed with 15 g of a conventional Ni/Al2 O3 catalyst and an installed membrane area of 20 cm2 . The influence of the main operating variables on conversion and hydrogen recovery was studied in the range of 450 to 550 ◦ C, 0.1 to 0.5 MPa, WHSV 0.29 to 1.43 h−1 , S/C of

2219

2–4, and a sweep gas flow of 100 to 500 mLN min−1 . The performance of the membrane reactor was clearly superior to that without hydrogen removal. The conversion reached 98% at 500 ◦ C, 0.3 MPa, and S/C 3 at the lowest space velocity and the highest sweep gas flow rate; hydrogen recovery under these conditions was ca. 97%, but without hydrogen removal the conversion was P2 Feed

Permeate

Air

Air O2−

O2 (g) + 4

e−

→2

O2−

2 O2− (s) → O2 (g) + 4 e−

(s) e−

O2-depleted air

O2-enriched air

P1

P2 Membrane

Principle of the production of O2 -enriched air by mixed ionic and electronic conductor (MIEC) membranes. (Reproduced from Ref. [332]; reprinted with permission from Wiley-VCH.)

Fig. 19

10.7.4.2.2 Partial Oxidation of Methane to Synthesis Synthesis gas (H2 and CO) is usually preGas pared by the strongly endothermic methane steam ◦ reforming (MSR), CH4 + H2 O   3H2 + CO, R H = +206 kJ mol−1 , at pressures between 15 and 30 × 105 Pa, and temperatures between 850 and 900 ◦ C. An alternative is the slightly exothermic partial oxidation of methane (POM), CH4 + 1/2O2 → 2H2 + CO, R H ◦ = −36 kJ mol−1 , giving a lower H2 /CO-ratio of 2, as required for methanol synthesis or the Fischer–Tropsch process. The POM reaction can take place in a MIEC membrane reactor with oxygen from the surrounding air (Fig. 20). Due to the potential economic and operational benefits of the production of synthesis gas in a MIEC

2229

membrane reactor, worldwide alliances coordinated by Air Products and Praxair have launched major R&D programs. The key to success in the industrial implementation of this process will be the development of a ceramic membrane with high long-term operational stability. In numerous previous POM studies using MIEC membranes, the membrane stability transpired to be a critical problem. In a tubular La0.6 Sr0.4 Co0.2 Fe0.8 O3−δ membrane reactor with a packed-bed steam-reforming catalyst, excellent conversions (>96%) and CO-selectivities (>97%) were reported for diluted methane (6% methane in He) as the feed [334, 335]. However, after a few hours of operation mechanical failure of the membrane occurred. A tubular La0.2 Sr0.8 Co0.8 Fe0.2 O3−δ membrane with a Rhbased catalyst was tested in the POM, and the brittle membrane broke a few minutes after introduction of the methane [336]. Both the oxygen partial pressure and grain size were found to influence the mechanical failure behavior of La0.2 Sr0.8 Cr0.2 Fe0.8 O3−δ perovskites [337, 338]. In another study with La0.2 Sr0.8 Cr0.2 Fe0.8 O3−δ , a loss of strength under mild conditions was reported [339, 340], while the operation of a tubular SrCoFeO3−δ perovskite membrane lasted for only a few minutes due to lattice expansions which resulted in cracking [341]. On SrCoFeO3−δ perovskites two types of fractures were observed: one type resulted from a strong oxygen gradient across the membrane, and the other was the result of chemical decomposition References see page 2241

MIEC membrane

CH4 e− O2 + 2Vo + 4e− → 2O2−

2CO + 4H2 + 4e− → 2CH4 + O2 O2−

Air CO + H2

Partial oxidation catalyst Fig. 20 Mixed ionic and electronic conductor (MIEC) membrane in the partial oxidation of methane to synthesis gas. (Reproduced from Ref. [346]; reprinted with permission from Elsevier.)

2230

10.7 Catalytic Membrane Reactors

of the perovskite in a reducing atmosphere [342]. A La0.2 Sr0.8 Co0.1 Cr0.1 Fe0.8 O3−δ membrane cracked in the POM after 350 h at 900 ◦ C [329]. Different concepts were developed to prolong the membrane lifetime at low O2 partial pressure. A bilayered membrane consisting of a Sm-doped CeO2 protective layer on the reducing side of the tubular La0.6 Sr0.4 Co0.8 Fe0.2 O3−δ membrane led to an improved stability in the POM [343]. Recently, bilayered membranes with a protective layer of Ce0.8 Gd0.2 O2 on La1−x Srx CoO3 and La0.9 Sr0.1 FeO3 were proposed [344]. Using a tubular YSZ-promoted SrCo0.4 Fe0.6 O3−δ perovskite membrane, failure was observed after 4 h, but by adding a small amount of oxygen to the methane feed the lifetime of the membrane was extended to 70 h [345]. Currently, several concepts are being followed to improve membrane stability, such as reducing the relative amount of cobalt in the perovskite and co-doping the material with less-reducible metals such as Zr4+ or Ga3+ [346]. Moreover, new cobalt-free perovskite membrane materials such as BaCe0.15 Fe0.85 O3−δ [347] and Ba0.5 Sr0.5 Zn0.2 Fe0.8 O3−δ [348] have been developed. During the past few years, remarkable progress has been made in the development of stable MIEC materials (cf. Table 3). Using La0.2 Ba0.8 Fe0.8 Co0.2 O3−δ in a disk-shaped reactor for the POM reaction, oxygen permeation fluxes of 4.4 mLN min−1 cm−2 over 850 h were obtained [349]. Impressive results in the POM were obtained recently by the Dalian group when studying Ba0.5 Sr0.5 Co0.8 Fe0.2 O3−δ tubular [350] and disk [351, 352] membranes. When using a CH4 feed diluted by 20% He, the CO selectivity was 95% at 94% methane conversion, and the stability of the membranes was of the order of 500 h. For better stability under syngas conditions, the material BaCo0.4 Fe0.4 Zr0.2 O3−δ was developed [353], and subsequently stable syngas production for more than 2200 h was achieved with a membrane reactor of this type [307]. Eltron Research have developed a material with

a Brownmillerite structure of the general composition A2 B2 O5 . The corresponding membrane reactor was continuously operated over one year in syngas production at 900 ◦ C, and a very high oxygen flux of 10–12 mLN min−1 cm−2 , linked to a syngas production rate of about 60 mLN cm−2 min−1 was reported [354]. The reaction of methane with oxygen in a membrane reactor is called ‘‘partial oxidation’’. However, there is experimental evidence that methane is first oxidized to CO2 and H2 O, after which these products of total oxidation are reduced by unreacted methane on conventional Ni-based catalysts to CO and hydrogen according to dry reforming (CO2 ) and steam reforming (H2 O), respectively. Therefore, synthesis gas formation from methane in a MIEC perovskite membrane reactor is referred to as an ‘‘oxidation-reforming process’’ [355]. 10.7.4.2.3 Oxidative Dehydrogenation of Light Alkanes to Olefins Ethene and propene are produced by steam cracking of ethane and propane or naphtha in a highly endothermic and energy-consuming process in which the dilution of hydrocarbons with steam reduces the formation of coke. Oxidative dehydrogenation of ethane (ODE) and of propane (ODP) to the respective olefins is considered a promising alternative, due mainly to the reduced demand for energy (lower temperature, no need for steam) and to the fact that there is no equilibrium constraint on the conversion; moreover, less coke formation is expected. However, the yields attained in the oxy-dehydrogenation in conventional co-feed reactors are, as yet, too low for an industrial application. During the past few years remarkable progress has been observed in the classical co-feed ODE, especially by reaction engineering measures, leading to ethene yields up to 56% at 71% selectivity in autothermal oxidative dehydrogenation at short contact time of ca. 45 ms using catalysts as igniters [360]. Schmidt and coworkers [361, 362] also used the concept of very short contact times (in

Examples for newly developed perovskite materials, their oxygen flux and stability in the partial oxidation of methane to synthesis gas (POM)

Tab. 3

Material

Initial La0.2 Sr0.8 Co0.4 Fe0.6 O3−δ Improved La0.2 Sr0.8 Co0.4 Fe0.6 O3−δ La0.8 Sr0.2 Co0.1 Fe0.6 Cr0.1 O3−δ La0.3 Sr1.7 Ga0.6 Fe1.4 O6−δ SrCo0.5 Fe0.5 O3−δ SrCo0.8 Fe0.2 O3−δ Ba0.5 Sr0.5 Co0.8 Fe0.2 O3−δ BaZr0.2 Co0.4 Fe0.4 O3−δ Ax A# 2−x By B# 2−y O5−δ

Temperature/ ◦ C

Oxygen permeation flux/mLN min−1 cm−2

850 850 900 900 850 87 875 850 900

– 3–4 14.5 2–4 2–4 – 11.5 5.6 10–12

Stability in the POM/h 0.04 ; χ= + ; for circular tubes Peax Re · Sc χ dt Dm 192

(10)

Empty tube, turbulent flow: 1 3 · 107 1.35 u · dt = + 1/8 ; Peax = Peax Re2.1 Dax Re

(11)

Packed bed, gas flow: 0.008 < Rep < 400 dp 1 0.3 0.5 = > 15; + ; 3.8 Peax Rep · Sc dt 0.28 < Sc < 2.2 1+ Rep · Sc

(12)

Packed bed, liquid flow: ε · Peax = 0.2 + 0.011 · Re0.48 p ;

dp > 15; 10−3 < Rep < 103 dt

Flow in microchannels with diameters between 10 and 500 µm is mostly laminar, and has a parabolic velocity profile. Therefore, the molecular diffusion in axial and radial directions plays an important role in RTD. The diffusion in the radial direction tends to diminish the spreading effect of the parabolic velocity profile, while in the axial direction the molecular diffusion increases the dispersion [58, 59]. The RTD in MSR will be compared with the RTD in conventional fixed-bed reactors, assuming the abovedefined equivalent design criteria. The RTD in packed beds and tubular reactors with laminar flow can

8 7 6

be estimated with Eq. (12) and Eq. (10), respectively. The axial P´eclet number for a packed bed and a structured microchannel reactor as function of Re · Sc is shown in Fig. 12. In packed-bed reactors, Peax as function of Re · Sc passes through a small maximum, and becomes constant for Re · Sc > 20. The maximum is much more pronounced for laminar flow in microchannels. The maximum Pe number corresponds to a minimal √ axial dispersion is obtained for Re · Sc = χ. The Bodenstein number in microchannels can be determined with Eq. (9) and Eq. (10): Dm L 1 dt2 u Dm L 1 4· u 1 = 2 + = 2 + Bo L u 192 Dm L L u 192 Dm L τ 1 1 tD,rad = ; + Bo tD,ax 48 τ with: tD,ax =

Structured microreactor

Peax

5 4 3 2

Packed bed reactor

1 0 0

10

20

30

40

50

60

(13)

70

80

90

100

Re Sc Fig. 12 P´eclet number as function of Re Sc in microchannel reactor and in a packed-bed reactor for ε = 0.45.

L2 R2 ; tD,rad = t Dm Dm

(14)

The first term in Eq. (14) corresponds to the ratio between space time and the characteristic axial diffusion time. The diffusion coefficient lies in the order of 10−5 m−2 s−1 and 10−9 m−2 s−1 for gases and liquids, respectively. Typical lengths of MSR are several centimeters, and the space time is in the range of seconds. Therefore, the axial dispersion in microchannels is mainly determined by the second term in Eq. (14), and the Bodenstein number can be estimated with Eq. (15) Bo ∼ = 48 ·

Dm τ ∼ = 50 · τ · 2 tD,rad Rt

(15)

10.8.3 Main Characteristics of Microstructured Reactors

Tab. 2

Characteristic radial diffusion time as function of channel diameter

Channel diameter dt /µm

2000 1000 500 200 100 50 20

Characteristic radial diffusion time tD,rad /s gas: Dm = 10−5 m2 s−1

128 · η · V˙tot · L π · N · d¯t4 · (1 + 6σˆ d2 )

Characteristic radial diffusion time tD,rad /s liquid: Dm = 10−9 m2 s−1

100 × 10−3 25 × 10−3 6.3 × 10−3 1 × 10−3 0.25 × 10−3 0.06 × 10−3 0.01 × 10−3

It follows that axial dispersion can be neglected (Bo ≥ 100), if the space time is at least twice the radial diffusion time. The characteristic radial diffusion times for different channel diameters are summarized in Table 2. Accordingly, axial dispersion of gases in microchannels can be neglected, if their diameters are less than 1000 µm and the space time is longer than 0.1 s. This could also be proven experimentally [27, 28]. This approach can also be used for MSR. Due to the small volume of a single channel, many channels must be used in parallel in order to obtain sufficient reactor performance. A uniform distribution of the reaction mixture over thousands of microchannels is necessary to obtain an adequate performance of the MSR [60]. Flow maldistribution will enlarge the RTD in the multitubular reactor, and lead to a reduced reactor performance along with reduced product yield and selectivity [28, 61]. Therefore, several authors have described design studies of flow distribution manifolds [55, 62, 63]. Besides maldistribution, small deviations in the channel diameter introduced during the manufacturing process may cause an enlargement of the RTD. These small deviations might be due to a non-uniform coating of the channel walls with catalytic layers. If the number of parallel channels is large (N > 30), a normal distribution of the channel diameters with a standard deviation σ can be assumed. The relative standard deviation, σˆ d = σd /d¯t influences the pressure drop over the microreactor [64]: p =

2257

(16)

The relationship in Eq. (16) shows that a variation of the channel diameter leads to a decrease of the pressure drop at a constant overall volumetric flow. As the pressure drop for each channel is identical, the variation of the diameter results in a variation of the individual flow rates, V˙i and the space time, τi = Vi V˙i . Supposing plug-flow in each channel (Boi → ∞), the overall dispersion is inversely proportional to the relative

1000 250 63 10 2.5 0.6 0.1

standard deviation, and can be estimated by Eq. (17) [64]: d2 Boreactor ∼ = t2 2σd

(17)

In consequence, the plug-flow behavior in a multichannel microreactor (Boreactor ≥ 100) can be assumed for the relative standard deviation of σd /dt ≤ 0.07. Mass Transfer Performance As mentioned above, a multichannel MSR allows a narrow RDT and near-isothermal operation. For the reactions involving a solid catalyst, which is introduced as a porous layer immobilized on the channel walls, the mass transfer must be considered. Prior to reaction, the reacting molecules must reach the catalytic surface, and therefore the influence of mass transfer on the overall kinetics must also be taken into account. 10.8.3.3

10.8.3.3.1 Internal Mass Transfer The potential influence of internal mass transfer on the global reaction rate depends on the intrinsic kinetics and the thickness of the porous catalytic layer, δcat , in microchannel wall reactors. Internal mass transfer limitation should be avoided, as it reduces the reactor performance and influences strongly the obtainable product selectivity [65, 66]. In general, the importance of mass transfer is expressed in form of an effectiveness factor, ηr , defined as the ratio of the observed reaction rate to that which would occur if the temperature and concentrations were constant throughout the catalytic layer. Different criteria for the absence of significant diffusion effects are proposed in the literature and summarized by Maers [67]. To ensure an effectiveness factor of ηr ≥ 0.95 in an isothermal catalyst, the following criterion, modified for a catalytic layer in microchannels, results:  Deff · cs δcat,max ≤ b (18) reff References see page 2263

2258

10.8 Microstructured Reactors

where Deff and reff are the effective diffusion coefficient and the observed reaction rate, respectively, and cs is the reactant concentration on the outer catalyst surface. The parameter b is dependent on the formal reaction order, m : b = 0.8 (m = 0); b = 0.3 (m = 1); b = 0.18 (m = 2). Strong exothermal or endothermal reactions may provoke a temperature profile within the catalytic layer that influences the reaction rate and selectivity. The importance of the temperature profile depends, besides the reaction rate and the layer thickness, on the reaction enthalpy, HR , the activation energy, E, and the thermal conductivity of the porous catalyst, λeff . For quasi-isothermal behavior, the observed rate, reff , must not differ from the rate that would prevail at constant temperature by more than about 5%. The resulting criterion, estimated for a catalytic layer in microchannels, is given in Eq. (19).  R λeff · Ts2 δcat,max ≤ 0.3 (19) E |HR |reff where Ts corresponds to the temperature on the catalyst surface and R is the gas constant. In order to ensure a high reactor performance of the microstructured wall reactor, the maximum acceptable thickness of the catalytic layer, δcat,max , is requested. 10.8.3.3.2 External Mass Transfer In general, the thickness of the catalytic layer will be kept sufficiently small to avoid the influence of internal mass transfer on the kinetics. In this case, only the transfer of the reactants from the bulk to the catalytic wall must be considered. The radial velocity profile in a single channel develops from the entrance to the position where a complete Poiseuille profile is established. The length of the entrance zone depends on the Re number, and can be estimated from the following empirical relationship [68, 69]:

Le ≤ 0.06 · Re · dt

(20)

Within the entrance zone, the mass-transfer coefficient diminishes, reaching a constant value. The dependency can be described with Eq. (21) in terms of Sherwood numbers, Sh = kD dt Dm [70, 71]:
0.45 dt (21) Sh = B 1 + 0.095 Re · Sc L The constant B in Eq. (21) corresponds to the asymptotic Sh number for constant concentration at the wall, which is identical to the asymptotic Nu-number characterizing the heat transfer in laminar flow at constant wall temperature. The constant B depends on the geometry of the channel, as summarized in Table 3.

Mass and heat transfer characteristics for different channel geometries [70]

Tab. 3

Geometry

Constant B in Eq. (21)

Circular Ellipse; width: height = 2 Parallel plates Rectangle; width: height = b/e = 4 Rectangle; width: height = b/e = 2 Square Equilateral triangle Sinusoidal Hexagonal

3.66 3.74 7.54 4.44 3.39 2.98 2.47 2.47 3.66

If the entrance zone in the tube can be neglected, the mass transfer is constant and given by B. As for the heat transfer, for a circular tube-shaped reactor: Sh∞ = 3.66; for L ≥ 0.05Re · Sc · dt (constant wall concentration)

(22)

If the mass transfer is accompanied by a chemical reaction at the catalyst surface taking place on the reactor wall, the mass transfer depends on the reaction kinetics [72]. For a zero-order reaction, the rate is independent of the concentration and the mass flux from the bulk to the wall is constant, whereas the reactant concentration at the catalytic wall varies along the reactor length. For this situation the asymptotic Sh number in circular tube reactors becomes Sh∞ = 4.36 [72]. The same value is obtained for low reaction rates as compared to the rate of mass transfer. If the reaction rate is high (very fast reactions), the concentration at the reactor wall becomes zero within the whole reactor and the final value for Sh is Sh∞ = 3.66. As a consequence, the Sh number in a reacting system depends on the ratio of the reaction rate to the rate of mass transfer characterized by the 2nd Damk¨ohler number (DaII): DaII =

ks · dt ks = ; kg Dm · 2

for a first order reaction and laminar flow

(23)

Villermaux [72] proposed a simple relationship to estimate the asymptotic Sh number as function of DaII (Fig. 13): 
1 DaII 1 1 1 = + − (24) Sh∞ Sh∞ DaII + 1.979 Sh∞ Sh∞ For the following discussion, we will assume that the mass transfer is the rate-limiting step (DaII ≥ 100). Under this condition, the wall concentration is

10.8.4 Design Criteria for Microstructured Multichannel Reactors

with two different particle Re-numbers. The Re-number is calculated for a particle diameter of dp = 1 mm, and is related to the channel dt based on the equivalent design criteria indicated by Eq. (1). As the mean residence time and specific surface areas are identical in the packed-bed reactor and in the MSR, the reactor performances are proportional to the mass-transfer coefficient. The MSR attains higher performances for the reactions limited by mass transfer.

4.4 4.3 4.2 4.1

Sh∞″

2259

4.0 3.9 3.8

Heat Transfer Performance Knowledge of the temperature profiles along nonisothermal tubular reactors is necessary for the design and operation of such reactors. If the locally exchanged heat does not correspond to the heat produced or consumed by the chemical reaction, then hot spots, and respectively cold spots, are formed. The great advantage of microstructured catalytic wall reactors stems from the fact that the reaction takes place in a thin layer near the wall. Therefore, the reaction heat can be transferred directly to the channel wall and the cooling medium. In general, the resistance to heat transfer of the channel wall and of the cooling medium is small compared to resistance within the porous catalytic layer. Therefore, only an excess temperature in the porous layer can be expected. If the condition of Eq. (19) is respected, the efficient reaction rate deviates at a maximum 5% compared to isothermal conditions. 10.8.3.4

3.7 3.6 0.01

0.1

1

10

100

Dall Fig. 13 Variation of the asymptotic Sherwood number with the 2nd .. Damkohler number (DaII). First-order reaction, circular channel.

ci,s ∼ = 0 = const and the asymptotic Sh number becomes Sh∞ ∼ = Sh∞ = 3.66. To compare the performance of a catalytic wall reactor with a catalytic packed bed, the mass transfer between the bulk and the particle surface as function of the particle Reynolds number, Rep , can be estimated from [73]: 1/2

Shp = 2.0 + 1.8Rep Sc13

(25)

In Fig. 14, the mass-transfer coefficient in a circular tube as function of the tube diameter is compared to the mass-transfer coefficient estimated for a packed bed

e = 0.45 Sc = 1 Dm = 10−5 m2 s−1

kg m/s

1

Microchannel

0.1

Packed bed

10 Re1mm = 1

0.01 0

200

400

600

800

1000

d t /µm Fig. 14 Mass-transfer coefficient in microchannel and packed-bed reactors (DaII > 100).

10.8.4

Design Criteria for Microstructured Multichannel Reactors

The main characteristic of multichannel MSR is the small diameter of the channels ensuring short radial diffusion time. This leads to a narrow RTD, and high heat and mass transfers. In addition, MSRs possess a high surface-to-volume ratio of around 10 000 to 50 000 m2 m−3 , allowing efficient heat removal and high molar fluxes to avoid internal mass-transfer limitations for heterogeneous catalytic reactions. Hence, MSRs are suitable for portable devices. The small reactor dimensions facilitate the use of production units at the place of consumption, thus avoiding the transport and storage of dangerous materials. In the following we will discuss the design criteria for achieving high specific MSR performance considering simple irreversible reactions. The specific reactor performance, LP ,V , is defined as the production of a desired product per time and unit volume: LP ,V =

n˙ product VR

References see page 2263

=

V˙0 c1,0 X c1,0 X = VR τ

(26)

2260

10.8 Microstructured Reactors

N1

N2

}e

b

N4 N5

N3 N6

N15

N7

Reaction channels (microstructured) Mixing (microstructured)

N9

N11 N10

N2

H2O2 evaporation (microstructured)

N12 N19

N14 N16 N17

DN51

N25

N21

Fig. 15

Schematic outline (left) and external view of the DEMIS reactor (right). (Illustration courtesy Uhde GmbH, Degussa AG [10].)

and the Damk¨ohler number becomes:

where X is the conversion of the key reactant: n˙ 1,0 − n˙ 1,s X= n˙ 1,0

(27)

The conversion obtained in the reactor depends on the ratio of the characteristic time of the operation, top to the space time, τ = VR /V˙0 , known as the first Damk¨ohler number, DaI. The characteristic time of operation depends on the limiting step of the process: this may be the chemical reaction, the mass transfer between fluid and a catalyst surface, or the heat transfer for highly exothermic or endothermic reactions. Heterogeneous Catalytic Reactions Limited by Chemical Transformation If the intrinsic kinetics of the reaction determines the reactor performance, the characteristic time of operation is given by the characteristic reaction time: top = tr . For an irreversible reaction of mth order, the characteristic reaction time is defined by 10.8.4.1

tr =

m k · c1,0

c1,0

m−1 = k · c1,0

(28)

m−1 DaI = k · c1,0 ·τ

(29)

For a given inlet concentration and rate constant, the necessary space time for a required conversion in a tubular reactor depends on the residence time distribution. The highest reactor performance is obtained for plug-flow behavior: Bo → ∞. In practice, a Bodenstein number of Bo ≥ 100 is sufficient to assume a plug-flow pattern. The Bo number is related to the P´eclet number by Eq. (30). With the channel diameter dt and its length L, we obtain: Bo = P eax

L dt

(30)

Under laminar flow conditions, the P´eclet number depends on Re · Sc as shown in Fig. 12. If it is assumed that L/dt = 20, a Pe number of P eax ≥ 5 is required to assure plug-flow behavior. Therefore, Ret Sc = u · dt /Dm must be between 6 and 30. For heterogeneous reactions the volume-based rate constant is proportional to the area of the porous catalytic layer per reactor volume: k = ks

A = ks · a VR

(31)

10.8.4 Design Criteria for Microstructured Multichannel Reactors

For a catalytic wall reactor, the specific catalytic surface area is related to the channel diameter: a = 4/dt and the Damk¨ohler number becomes DaI = ks · τ · 4/dt . For a required conversion and constant Damk¨ohler number, the space time is proportional to the channel diameter: τ=

DaI dt 4 · ks

(32)

As a consequence, the specific reactor performance increases with decreasing channel diameter [Eq. (26)]. However, changing the channel diameter will influence the Re-number and, as a consequence, the RTD; this effect must be taken into account for the reactor design. Heterogeneous Catalytic Reactions Limited by Mass Transfer If the external mass transfer becomes rate limiting (DaII > 100), the characteristic time of operation is given by the characteristic time of mass transfer in the channel: top = tD = 1/(kg · a). For fully developed velocity and concentration profiles, the Sherwood number becomes constant and the mass-transfer coefficient is inversely proportional to the channel diameter: kg = Sh∞ · Dm /dt . As the specific surface area increases also with decreasing diameter, the characteristic time of mass transfer and the Damk¨ohler number become: 10.8.4.2

dt dt τ · 4 · Sh∞ · Dm tD = ; DaI = Sh∞ · Dm 4 dt2

Bo ∼ = 48 ·

τ tD,rad

and endothermic reactions by providing process intensification and safety. Based on the fundamentals of chemical reaction engineering, the design and range of operation conditions which benefit by reactor miniaturization have been determined. The flexibility in constructing microstructured devices allows the process structure and equipment to be adapted for efficient and safe chemical transformation. Moreover, production units can be created simply by the integration of small-scale MSRs into a large-scale plant. A first example of the successful implementation of microstructures into a pilot unit of industrial scale has been demonstrated by the DEMIS project [10]. Compared to conventional systems, the MSR significantly improved the control of reactant concentrations and the temperatures allowed for the safe production of methyloxirane in the gas phase with gaseous hydrogen peroxide. In DEMIS, the rectangular channels have a height (e) of between 100 and 1000 µm, and are up to 0.1 m wide (b) and 5 m long. The design allows the catalyst to be developed in the laboratory and to be transferred to the industrial scale. The slit-like design, with only one dimension in the microscale, leads to a reduced number of parallel channels, but ensures a hydraulic diameter in the microrange, corresponding to the channel height: dt = 4

(33)

The space time for a required conversion decreases with dt2 , and in consequence the specific reactor performance increases with 1/dt2 . High conversions (X ≥ 0.9) can be obtained at space times in the order of 10−2 s and, as might be expected, the RTD can no longer be neglected under those conditions. Equation (14) can be used for an estimation of the Bo-number in microchannels with L/dt ≥ 20.

e·b ∼ = 2·e 2(e + b)

Symbols and Abbreviations Symbols

m2 m−3 m2

ci

– kmol m−3

ci,o

kmol m−3

c¯0

kmol m−3

Conclusions and Outlook

cp

J kg−1 K−1

Microreactors with continuous flow have been shown to improve the performance of fast and highly exothermic

References see page 2263

Bo ∼ = 48 ·

DaI ∼ = 13 · DaI Sh∞

(34)

For conversions in the order of 0.9 < X < 0.99, Bonumbers of 30 > Bo > 60 are obtained. This leads to a ≈10% reduced specific performance compared to ideal plug-flow reactors.

(35)

The structured multiscale design is a promising conceptual approach in chemical processes development, with significant potential to improve the competitiveness of the chemical industry.

a A Ai , Ak

Dm DaI · dt2 = 48 · τ · 2 ; with τ = 4 · Dm Rt

2261

10.8.5

specific surface area surface area reactant molar concentration of compound i initial, feed concentration of compound i mean initial concentration of a tracer heat capacity at constant pressure

2262

10.8 Microstructured Reactors

C, C(θ)

–

referred concentration of a tracer diameter (particle, tube) hydraulic diameter: section dh = 4 surconference effective diffusion coefficient molecular diffusion coefficient axial, radial dispersion coefficient exit age distribution activation energy

d, dp, dt dh

m m

Deff

m2 s−1

Dm

m2 s−1

Dax , Drad

m−2 s−1

E(t) E, Ea fi = ni /ni,0 = (1 − Xi )

s−1 J mol−1 −

F(t ) =  t 0 E(t)dt

remaining fraction of reactant Ai

−

HR h hov

J mol−1 W m−2 K−1 W m−2 K−1

k=

kL ,kg

k0 e(−E/RT ) reaction rate constant (different) m s−1

cumulative residence time distribution reaction enthalpy heat-transfer coefficient overall heat-transfer coefficient moll−n m3(n−1) s−1 k = k∞ e(−E/RT )

L Lp

m mol s−1

mcat Mi m ni n˙ i p

kg kg mol−1 − mol mol s−1 Pa

p pi

Pa Pa

rj R

mol m−3 s−1 J mol−1 K−1

ko

Rt S t ¯t = 

∞ 0 t

· E(t)dt

m m2 s

pre-exponential factor mass-transfer coefficient, liquid, gas length production performance mass of catalyst molar mass reaction order amount of compound i molar flow total pressure, 1 Pa = 1 N m−2 pressure drop partial pressure (1 bar = 105 Pa) reaction rate (reaction j) gas constant: R = 8.314 J mol−1 K−1 tube radius cross-section time

s

mean residence time

tD,ax ; tD,rad

s

T u V

K m s−1 m3

VR V˙ X= nio − nis nio z Z

m3 m3 s−1

axial, radial diffusion time temperature superficial velocity volume (reaction mixture) reactor volume volumetric flow

−

conversion

m −

δ ε

m −

η η λ

Pa s − W m−1 K−1

µ ρ θ = t/τ ρ σ σ τ = V /V˙0 φ

axial coordinate referred axial coordinate layer thickness fractional volume of fluid phase in fluid/solid systems dynamic viscosity Effectiveness factor heat conductivity (effective)

m2 s−1

kinematic viscosity

− kg m−3 N m−1 − s

referred time density surface tension variance space time geometric factor

ν=

Abbreviations

CSTR MSR PFR RTD

continuous stirred-tank reactor microstructured reactor plug-flow reactor residence time distribution

Dimensionless groups

uL Dax m−1 DaI = k · c1,0 ·τ ks DaII = kg Bo =

u·d ε · Dax u·d Re = νν Sc = Dm kg dt Sh = Dm P eax =

Bodenstein number .. first Damkohler number .. second Damkohler number (1st order reaction) axial P´eclet number Reynolds number Schmidt number Sherwood number

References

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47. S. Thybo, S. Jensen, J. Johansen, T. Johannessen, O. Hansen, U. J. Quaade, J. Catal. 2004, 223, 271–277. 48. S. Schimpf, M. Bron, P. Claus, Chem. Eng. J. 2004, 101, 11–16. 49. P. Tribolet, L. Kiwi-Minsker, Catal. Today 2005, 105, 337–343. 50. P. Tribolet, L. Kiwi-Minsker, Catal. Today 2005, 102–103, 15–22. 51. P. Tribolet, Nanofibres de carbone sur filtre m´etallique comme support catalytique structur´e, PhD Thesis Ecole polytechnique f´ed´erale de Lausanne, 2006. 52. R. Knitter, M. A. Liauw, Lab on a Chip 2004, 4, 378–383. 53. J. M. Commenge, Mod´elisation de micror´eacteurs en g´enie des proc´ed´es, PhD Thesis, ENSIC, INPL Nancy, Nancy, 2001. 54. J.-M. Commenge, L. Falk, J.-P. Corriou, M. Matlosz, Chem. Eng. Technol. 2005, 28, 446–458. 55. J. M. Commenge, L. Falk, J. P. Corriou, M. Matlosz, AICHE J. 2002, 48, 345–358. 56. VDI-W¨armeatlas. 9th Ed. Springer, Berlin, Heidelberg, New York, Verein Deutscher Ingenieure, 2002. 57. M. Baerns, H. Hoffmann, A. Renken, Chemische Reaktionstechnik, Wiley-VCH, Weinheim, 1999, p. 428. 58. G. Taylor, Proc. Roy. Soc. London 1953, A219, 186–203. 59. R. Aris, Proc. Roy. Soc. London 1956, A235, 67–77. 60. O. W¨orz, K.-P. J¨ackel, T. Richter, A. Wolf, Chem. Eng. Technol. 2001, 24, 138–142. 61. A. Rouge, A. Renken, in Reaction kinetics and the development and operation of catalytic processes, G. F. Froment. K. C. Waugh (Eds.), Studies in Surface Science and Catalysis, Vol. 133, Elsevier, Amsterdam, 2001, pp. 239–246. 62. J. M. Commenge, A. Rouge, A. Renken, J. P. Corriou, M. Matlosz, R´ec. Prog. G´enie des Proc´ed´es 2001, 15, 329–336.

63. E. R. Delsman, M. H. J. M. De Croon, A. Pierik, G. J. Kramer, P. D. Cobden, C. Hofmann, V. Cominos, J. C. Schouten, Chem. Eng. Sci. 2004, 59, 4795–4802. 64. E. R. Delsman, M. H. J. M. De Croon, G. D. Elzinga, P. D. Cobden, G. J. Kramer, J. C. Schouten, Chem. Eng. Technol. 2005, 28, 367–375. 65. M. Baerns, A. Renken, Chemische Reaktionstechnik, in Winnacker-K¨uchler Chemische Technik: Prozesse und Produkte. Wiley-VCH, Weinheim, 2004, pp. 453–643. 66. C. N. Satterfield, Mass Transfer in Heterogeneous Catalysis. M.I.T. Press, Cambridge, 1970. 67. D. E. Mears, Indust. Eng. Chem. Process Design Development 1971, 10, 541–547. 68. D. F. Sherony, C. W. Solbrig, Int. J. Heat Mass Transfer 1970, 13, 145–146. 69. J. H. B. J. Hoebink, J. M. A. Harmsen, C. M. L. Scholz, G. B. Marin, J. C. Schouten, Modeling of Automotive Exhaust Gas Converters, in Structured Catalysts and Reactors, A. Cybulski, J. A. Moulijn (Eds.), Marcel Dekker, Inc., New York, 1998, pp. 311–354. 70. A. Cybulski, J. A. Moulijn, Catal. Rev. -Sci. Eng. 1994, 36, 179–270. 71. R. E. Hayes, S. T. Kolaczkowski, Chem. Eng. Sci. 1994, 49, 3587–3599. 72. J. Villermaux, Int. J. Heat Mass Transfer 1971, 14, 1963–1981. 73. J. Villermaux, G´enie de la r´eaction chimique. TEC & DOC, Paris, 1993. 74. D. Schweich, G´enie de la r´eaction chimique, Edition TEC & DOC, London, Paris, New York, 2001.

HANDBOOK OF HETEROGENEOUS CATALYSIS Second, Completely Revised and Enlarged Edition Volume 5 Gerhard Ertl, Helmuth Knözinger, Ferdi Schüth, Jens Weitkamp (Editors) Wiley-VCH Verlag GmbH& Co. KGaA, Weinheim, Germany, 2265 ISBN: 978-3-527-31241-2, 2008

11

Environmental Catalysis 11.1

Catalysis in Environmental Protection Wilhelm Keim∗

11.1.1

Introduction

Within the discussion of this chapter, environmental catalysis has been defined as the combined fields of heterogeneous, homogeneous, and biocatalysis. However, due to the purpose of this Handbook of Heterogeneous Catalysis, the main emphasis will be focused on heterogeneous catalysis. It should be borne in mind, however, that catalytic solutions concerning environmental protection can embrace all three fields, and that often the interplay provides technical solutions. Historically, catalysis has always been a domain of the chemical and petroleum industries, with catalysis having been applied to introduce new or improved processes and products. The major success of the chemical and petroleum industries has resulted in an immense growth, and the sheer size of the industry has impacted upon the environment not only by the emissions of the manufacturing plants but also by some of the products. A growing public awareness, which began during the early 1970s, has demanded better protection, and the search for solutions has become an important issue. Catalysis offered a variety of solutions to these problems, and the introduction of exhaust catalysts for cars has heralded this development. Today, we hold special conferences that are dedicated to environmental catalysis. For example, the first International Conference on Environmental Catalysis (ICEC) was held in 1995 in Pisa, followed by Miami Beach in 1998, Tokyo in 2001, and Heidelberg in 2005. At these conferences, scientists and engineers from both industry and academia interchange their experiences and discuss ∗

Corresponding author.

Handbook of Heterogeneous Catalysis, 2nd Ed. .. .. Edited by G. Ertl, H. Knozinger, F. Schuth, and J. Weitkamp Copyright  2008 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim ISBN: 978-3-527-31241-2

recent developments and technological opportunities. A number of specialized scientific journals have also become devoted to environmental issues, for example Green Chemistry or Applied Catalysis B: Environmental. The aim and emphases can be visualized from the table of contents of the latter journal: • Catalytic elimination of environmentally hazardous effluents, nitrogen oxides, sulfur compounds, carbon monoxide, chlorinated and other organic compounds, used catalysts, etc., from stationary or mobile sources, in air, water, soil, etc. • Basic understanding, modeling and characterization of catalysts and processes used in environmental applications. • Catalytic processes occurring in the environment itself. • Catalytic reactions in which wastes are converted to useful products. • Clean manufacturing with catalysts replacing toxic chemicals with environmentally friendly catalysts, solvents, process conditions, etc. • Production of clean energy by mechanisms other than combustion, for instance, via fuel cells and fuel processing of hydrocarbons to hydrogen. • Photocatalysis as applied to environmental problems. • New catalytic combustion technologies and catalysts. In addition, almost every journal dealing with catalysis contains reports into environmental catalysis. Indeed, on occasion complete issues are devoted to this topic, for example November 2003 issue of Catalysis Today, ‘‘Catalysis for environmental friendly technologies’’. In addition, government agencies and scientific organizations have issued special reports [1–4], a number of books covering environmental catalysis have been produced [5–11], and various universities now offer special courses and programs on environmental catalysis [12–15]. Most of the emissions that pollute the environment are assumed to come from chemical plants, vehicular References see page 2273

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11.1 Catalysis in Environmental Protection

exhausts, and from large power plants. However, many other sources for pollution exist, including volatile organic compounds (e.g., solvents released into the atmosphere), wood- and coal-burning stoves, nitrates and residues in ground water, and aluminum and steel production. The concept of the environmental burden, EB, which may be expressed by Eq. (1), refers to the underlying reasons EB = P × Wp × Bw

(1)

where P is the population of the earth, Wp is the wealth per person, and Bw is the burden per unit of wealth. Wealth per person (Wp) contains the aspirations of the developing world, which will demand its share of the wealth of First-World countries, whilst behind burden per unit of wealth (Bw) also stands the impact of technology. Environmental catalysis is intimately intertwined with: • • • • •

automotive exhaust catalysis catalysis for clean air, water, and soil catalysis for sustainable energy conversion environmentally benign routes to chemicals the use of renewables.

Environmental catalysis became an integrated part of declarations (catch words) such as: Green Chemistry, Sustainability, Environmentally Benign Chemistry, Clean Chemistry, Atom Economy, Design Chemistry, Green Catalysis, Green Engineering, and many more. Catalysts work functions are: selectivity, activity and catalyst life (productivity), all three of which coincide with the goals of environmental protection. Activity Activity reduces catalytic reactor size and capital expenditure, which is often directly related to the tonnage of materials used for the construction of reactors, and also to the reactor dimensions. An increase in feedstock conversion per pass is a consequence, thus minimizing the recycling of unconverted feed. In addition, plant throughput can be increased, thereby decreasing energy requirements. More active systems also allow the temperature of operation to be lowered, thereby enhancing selectivity whilst maintaining conversion. Tab. 1

E-factors in the chemical and oil industries

Industry segment

Product tonnage/t a−1

kg byproduct/kg product

Oil refining Bulk chemicals Fine chemicals Pharmaceuticals

106 to 108 104 to 106 102 to 104 10 to 103

90% in 1990 [32]. On occasion, the chemical industry uses harmful chemicals in their production processes, but this can be circumvented by the substitution of catalytic pathways. For example, the use of phosgene in the manufacture of polyurethanes is under consideration for a catalytic route which avoids phosgene use. The traditional route to dimethylcarbonate (DMC) has been to react methanol with phosgene, according to Eq. (10). DMC can serve as an eco-friendly substitute for toxic reactants, such as phosgene and chloroformates in carbonylations: COCl2 + 2CH3 OH → CH3 OCOOCH3 + 2HCl

(10)

In order to avoid the use of the extremely toxic phosgene [Eq. (10)], the use of which also burdens the process with high costs to provide safety, Eni Chem has developed a route to DMC based on catalytic oxidative carbonylation [Eq. (11)] [25]: 2CH3 OH + CO + 12 O2 Cu-catalyst

−−−−−−→ CH3 OCOOCH3 + H2 O

(11)

Here, water and CO2 are the only byproducts, and the high selectivity obtained translates into a high product purity. In addition, the use of phosgene leads to chlorinated byproducts, which are not present in the oxidative carbonylation. Often, environmental problems arise from the use of chlorine. Historically, propylene oxide was prepared via the chlorine route, as in Eq. (9) (the synthesis of

11.1.5 Catalysis for Environmental Protection: Future Prospects

ethylene oxide). Today, an alternative catalytic route based on ethylbenzene oxidation has been developed, but unfortunately this process creates styrene as a byproduct, which in turn leads to problems of economy with regards to a two products process. However, a new process based on propylene epoxidation with H2 O2 (Uhde-Degussa) is currently on the developmental horizon [25]. The fine chemical industry, with markets in both pharmaceuticals and agricultural agents, offers great potential for the use of environmentally benign catalytic pathways (see Table 1). Traditionally, in this field of interest, stoichiometric, multistep reactions have been used, but these yield substantial amounts of byproducts. In addition, more severe medical challenges have led to significant increases in the molecular complexity of modern drugs, which in turn has caused the number of steps in a chemical synthesis to increase from an average of four in the period 1960–1979, to an average of nine today. The synthesis of ibuprofen (Fig. 3), using a new process developed by Hoechst-Celanese, demonstrates an approach to eliminating processing steps. For ibuprofen synthesis, the ‘‘old’’ process used six stoichiometric steps, with low atom efficiency, such that for each 1000 tons of product, a total of 760 tons of wastewater was produced. In contrast, the new process uses three catalytic steps, all of which are environmentally more friendly. It can be anticipated that the fine chemical industry will enrich its technological toolbox by using catalysis in various roles, thereby reducing the number of steps required and limiting the number and amounts of byproducts. Today, the traditional sequential approach – catalyst design, catalyst optimization, reactor concept and process design – has been replaced by an integrated approach, in which environmental compatibility will play a key role at all stages of development (Fig. 4). In future, the catalytic reaction, the preparation of the catalyst, catalyst disposal, and indeed the whole catalytic process will be analyzed in light of environmental compatibility. In recent years, environmental laws have become increasingly stringent with regards to the use of metal catalysts, mainly because their correct disposal is expensive. Thus, catalyst recycling, extending the lifetime of the catalyst and the recovery of spent catalysts represent crucial areas of process development, process design, and process operation. In the past, the majority of chemical processes were carried out with pollution prevention being introduced at ‘‘the end of the pipe’’. Indeed, sewage treatment plants, deposits and incinerations were all characteristic of the chemical industry of the past [33]. Today, however, the search for pollution prevention begins at the point, where waste, sewage and waste air originate; hence,

2271

we speak of ‘‘integrated pollution prevention, processintegrated environmental protection’’ which is aimed at: (i) the optimal use of resources; (ii) an avoidance and minimization of waste (solid, liquid, gaseous); and (iii) making use of waste products. In this respect, catalysis plays a key role in the integrated pollution prevention process. The manufacture of bisphenol A (BPA) represents one such example of integrated pollution prevention [see Eq. (12) and Fig. 5], where the disposal of an acidic mixture of HCl/H2 SO4 is replaced by using an ion-exchange resin. CH3 2

OH + Acetone

Acid

HO

C

OH

CH3

(12) 11.1.5

Catalysis for Environmental Protection: Future Prospects

Today, catalysts play an important and innovative role in environmental pollution prevention, with great success having been achieved especially in the control of air pollution (e.g., cleaner vehicular exhausts). Unfortunately, water and soil are inherently more difficult to deal with, although catalytic routes to provide such control are in the pipeline. The importance of catalysis in environmental improvements is also manifested in the production of chemicals. Currently, about 90% of all new chemical processes are catalytic and, as highlighted above, improvements in catalysts with regards to activity, selectivity, and catalyst life are environmentally beneficial goals. New, environmentally beneficial conversions include: • the direct conversion of natural gas to alcohols, epoxides or nitriles • the direct conversion of nitrogen to nitric acid • fine chemical synthesis, with enantiomeric control (e.g., drug, agricultural agents, odor and fragrance chemicals, aroma chemicals) • the conversion of alkanes to chemicals (e.g., methane to methanol) • dehalogenation processes • polymer recycling (back to the chemical feedstock). With regards to future technical developments, the conversion of renewable sources (biomass) into chemicals, and also the renaissance of ‘‘coal to chemicals technology’’ will introduce many more opportunities for the application of environmental catalysis. References see page 2273

2272

11.1 Catalysis in Environmental Protection

Old process

New Process

Ac2O/AlCl3

Ac2O/HF

O C

O

Base CICH2COO Et

H2 Pd/C

O COO Et

OH

H2O/H+

CHO

CO/Pd cat.

NH2OH

NOH

−H2O

CN H2O/H+ COOH 6-steps Fig. 3

3-steps

The synthesis of ibuprofen.

The switch from oil and gas to lignocellulose feedstocks, which in a broader sense also represents catalysis in terms of environmental protection, can follow a variety of approaches, including: (i) the use of chemical structures as produced by Nature; (ii) the modification of these structures; and (iii) the total reduction to

C1 fragments and controlled build-up of the desired molecules. The obvious advantages of this approach will be: (i) economic benefits from reduced emissions into water, air, and soil, with specific emphasis on climate change; (ii) economic benefits from increases in energy

References

Catalytic reaction

Reaction engineering

Mathematic modeling

Fig. 4

Catalyst preparation

Environmental compatibility

Catalyst analysis

An integrated approach to process development. Old process

New Process

Phenol, acetone HCl, H2SO4 (100%), NaOH (100%) H2 O

Phenol, acetone

Reaction HCl/H2SO4-solution

Reaction catalyst: ion-exchange resin

1.0 t BPA 3.3 t waste water 0.3 t salts emission 6.0 kg COD

1.0 t – 90% – 99% – 94%

BPA waste water salts emission COD

A comparison of the ‘‘old’’ and ‘‘new’’ manufacture of bisphenol A (BPA) [34]. COD = Chemical oxygen demand.

Fig. 5

and material efficiency; and (iii) cost and safety advantages resulting from the smarter design of processes, manufacturing facilities, and by process intensification. An example of this is the synthesis of 1,3-propanediol, for which three processes are currently under consideration: • The hydroformylation of ethylene oxide [35] • The oxidation of propene to acrolein, followed by hydration (Degussa) [36] • Sugar fermentation (Du Pont) [37]. While the route via sugar fermentation appears to be the most feasible, this example emphasizes the

2273

delicate interplay that exists between homogeneous, heterogeneous, and biocatalysis in the process/product development of environmental catalysis. Clearly, the chemical and related industries realize that their sustainable development will depend critically upon the introduction of new and improved catalytic processes that should be not only more efficient but also environmentally benign with regards to raw materials and energy utilization. Indeed, a reduced or zero production of unwanted byproducts will be the aim of these investigations, while the creation of new catalytic technologies as the key to technical solutions will help to build a sustainable economy and society. Environmental legislation will also play an important role in this development, with perhaps the most wellknown example being the fuel formulations required by law. It is clear that integrated environmentally friendly product and process design will help to determine our future based on catalysis for environmental protection. This is more than just fashion! References 1. Netherland Ministry of Economic Affairs ‘‘Catalysis, key to sustainability’’, Netherlands 2001, 53 pp. (www. technologyroadmapping.com). 2. ConNeCat Competence Network Catalysis Roadmap of German Research Catalysis, February 2003, 44 pp. (www.connecat.de). 3. American Chemical Society, Catalyst Technology Roadmap Report, Washington DC 1997, 50 pp. 4. M. S. Reisch, US Technology Vision 2020, Chem. Eng. News 1999, August 9, 10. 5. P. T. Anastas, T. C. Williamson, Green Chemistry, ACS Symposium Series 626, American Chemical Society, Washington DC, 1996, 251 pp. 6. F. J. J. G. Janssen, R. A. van Santen, Environmental Catalysis, Imperial College Press, London, 1999, 369 pp. 7. G. Ertl, H. Kn¨ozinger, J. Weitkamp, Environmental Catalysis, Wiley-VCH, Weinheim, 1999, 217 pp. 8. R. A. Sheldon, H. van Bekkum, Fine Chemicals through Heterogeneous Catalysis, Wiley-VCH, Weinheim, 2000, 610 pp. 9. The Handbook of Environmental Chemistry, Vols. 1–5, SpringerVerlag, Heidelberg, 2005. 10. V. Ahluwalia, M. Kidwai, New Trends in Green Chemistry, Springer-Verlag, Heidelberg, 2004, 250 pp. 11. J. N. Armor, ACS Symposium Series 552, American Chemical Society, Washington DC, 1994, 552. 12. Special course, University of Twente, Catalytic Materials and Processes, http://cpm.tnw.utwente.nl/teaching/Course. http://cpm.tnw.utwente.nl. 13. Special course, ETH Z¨urich, Environmental Catalysis, http://baiker.ethz.ch/research/Env.Cat. 14. Special course, Kungl. Tekniska H¨ogs Kolan, Environmental Catalysis, http://www.kth.se/student/studiehandbok/03/Kurs. 15. Special course, Northwestern Univ., Institute of Environmental Catalysis, http://www.iec.northwestern.edu/ 16. R. A. Sheldon, Chemistry & Industry, 1992, 903. 17. G. Petrini, G. Leofanti, M. A. Mantegazza, F. Pignataro, EP Patent 208 311, assigned to Eni Chem., 1988.

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11.2 Automotive Exhaust Treatment

18. P. T. Anastas, T. C. Williamson, Green Chemistry, ACS Symposium Series 626, American Chemical Society, Washington DC, 1996, p. 33. 19. F. M. Dautzenberg, Chem. Tech. 1995, 3, 54. 20. B. M. Trost, Y. Shi, J. Am. Chem. Soc. 1991, 113, 701. 21. F. M. Dautzenberg, P. J. Angevine, Catal. Today 2004, 93–95, 3. 22. H. U. Blaser, Chem. Commun. 2003, 293. 23. H. U. Blaser, M. Studer, Appl. Catal. A: General 1999, 189, 191. 24. D. E. DeVos, I. F. J. Vankelecom, P. A. Jacobs, Chiral Catalyst Immobilization and Recycling, Wiley-VCH, Weinheim, 2000, 280 pp. 25. J. F. Jenck, F. Agterberg, M. J. Droescher, Green Chemistry 2004, 6, 544. 26. J. P. Pradier, C. M. Pradier, Carbon Dioxide Chemistry Environmental Issues, The Royal Society of Chemistry, Cambridge, 1994, 403 pp. 27. A. Behr, Carbon Dioxide Activation by Metal Complexes, VCHVerlagesellschaft, Weinheim, 1988, 154 pp. 28. N. Eisberg, Chemistry & Industry 2005, 16 May, 12. 29. W. Keim, Chimia 1981, 344. 30. W. Keim, Angew. Chem. Int. Ed. Eng. 1990, 29, 235. 31. F. Cavana, G. Centi, C. Perego, A. Vaccani, Catal. Today 2005, 99, 1. 32. S. Rebsdat, D. Mayer, in Ullmann’s Encyclopedia of Industrial Chemistry, 5th Ed., Vol. A10. VCH Verlagsgesellschaft mbH, Weinheim, 1987, p. 117. 33. European IPPC Bureau, Integrated Pollution Prevention and Control, Sevilla, February 2002, 48 pp. 34. U. Bornewasser, U. M¨uller-Eisen, J. Wiesner, Integrierter Umweltschutz, in Winnacker-K¨uchler, Chemische Technik, 5th Ed., Vol. 2. Wiley-VCH, Weinheim, 2004, pp. 1–39. 35. B. Cornils, in Hydroformylation, Oxo Synthesis, Roelen Reaction, J. Falbe (Ed.), New Syntheses with Carbon Monoxide, Reactivity and Structure Concepts in Organic Chemistry, Vol. 11. SpringerVerlag, Berlin, 1980, p. 131. 36. D. Arntze, T. Haas, A. Sch¨afer-Sindlinger, US Patent 5,364,984, assigned to Degussa AG, 1994. 37. D. C. Cameron, N. E. Altaras, M. L. Hoffmann, A. J. Shawm, Biotechnol. Prog. 1998, 14, 116.

11.2

Automotive Exhaust Treatment1 Egbert S. J. Lox∗

11.2.1

Introduction Origin of Emissions The industrialization of the Western world was accompanied by a drastic increase in the consumption of fossil fuels. The energy stored in fossil fuels was freed mostly by flame combustion, which is the reaction between the 11.2.1.1

1 A list of abbreviations/acronyms used in the text is provided at the end of the chapter. ∗ Corresponding author.

carbon-containing constituents of the fossil fuel and the oxygen of the air, according to the reaction ← − mCO2 + (n/2)H2 O Cm Hn + (m + n/4)O2 − −− −− →

(1)

Carbon dioxide and water are the main products of this reaction. However, incomplete combustion causes some emissions of unburned hydrocarbons, as well as intermediate oxidation products such as alcohols, aldehydes, and carbon monoxide. As a result of thermal cracking reactions that take place in the flame, especially with incomplete combustion, hydrogen is formed and emitted, as well as hydrocarbons that are different from those present in the fuel. Most fossil fuels also have some amount of sulfurcontaining and nitrogen-containing constituents, that will yield some emissions of sulfur oxides (mainly SO2 ) and nitrogen oxides, commonly denoted as NOx but consisting mainly of NO, and a small amount of N2 O. During flame combustion, temperatures in excess of 1700 K occur. At those temperatures, the reaction between the air constituent’s nitrogen and oxygen is thermodynamically favored, resulting in the formation of nitrogen oxides, according to the overall reaction equation − 2NO N2 + O2 ← −−− −− →

(2)

With fuels used in internal combustion engines, the reaction in Eq. (2) is the major cause of nitrogen oxides emissions. Of course, the amount of CO, hydrocarbons, and nitrogen oxides that are emitted is dependent on the detailed composition of the fuel, as well as on the way the combustion is performed. However, as an order of magnitude, the exhaust gas of a gasoline-powered spark-ignited internal combustion engine will have the composition shown in Fig. 1. Importance of Road Transport To date, vehicles equipped with internal combustion engines that use fuels derived from fossil fuels perform the majority of on-road transport. The vehicles range from passenger cars to heavy-duty trucks, and the engines used are either spark-ignited or compression-ignited reciprocating internal combustion devices. An order of magnitude of the relative importance of transport in the total emissions of CO, hydrocarbons, nitrogen oxides, sulfur oxides and dust is shown in Fig. 2. From these data, it is clear that road traffic is one of the major sources of CO, hydrocarbons, and nitrogen oxides emissions [1]. In 2005, the annual worldwide production of vehicles amounted to about 60 million, which yielded about 500 million vehicles on the road worldwide. Figure 3 shows the historical evolution and the projections of the 11.2.1.2

11.2.2 Legislation

2275

HC 0.05 Vol.% NOx 0.08 Vol.%

N2 71.0 Vol.%

Particle 0.005 Vol.% Pollutants 1.0 Vol.%

CO 0.85 Vol.%

H2O 9.2 Vol.% CO2 18.0 Vol.%

O2 and noble gases 0.7 Vol.%

Fig. 1

Typical composition of the exhaust gas of a gasoline-powered spark-ignition internal combustion engine.

10

Total emission (106 t/year)

9

8.2

8 7 6 5 4 3

2.7

2.4

2

1

1

0.6

0 HC

CO Other sources

NOx

SO2

Dust

Caused by road traffic

Total emissions of hydrocarbons (HC), carbon monoxide, nitrogen oxides (NOx ), sulfur dioxide and dust in the Federal Republic of Germany in 1989, and relative importance of traffic.

Fig. 2

total vehicle population [2]; their increasing use, expressed as average distance driven per vehicle and per year, should also be noted. 11.2.2

Legislation History From the data presented above, it is obvious that traffic has an important impact on the air quality, which affects both the environment and human health. Therefore, legislation was introduced to limit the emission of CO, hydrocarbons, nitrogen oxides and particulates caused by traffic emissions. The initiative to introduce legislation to limit exhaust gas emissions from vehicles was taken in the USA. In 1966, California introduced limits for the exhaust gas emission of CO and hydrocarbons from passenger cars equipped with spark-ignition engines; similar regulations were then extended to other states. 11.2.2.1

The US Federal Clean Air Act, introduced in 1970, prescribed a further drastic reduction in the allowable exhaust gas emissions from passenger cars. An historical overview is given in Table 1. Similar measures were then taken in Japan, Australia, and Switzerland. In 1985, the European Community passed respective strict legislation for passenger cars with spark-ignition engines, to be followed by South Korea in 1987 and Brazil in 1988. Present and Future The definition of emission legislation is a continuous process, and at present a variety of new steps are being taken worldwide to further improve the air quality. The first step is that an increasing number of countries beyond the USA, Japan and the European Union (EU) are introducing stringent emission standards for passenger 11.2.2.2

References see page 2343

2276

11.2 Automotive Exhaust Treatment

Tab. 1

Historical overview of the limits for CO, HC and NOx emissions from passenger cars in the USA

Year

1966–67 1968–69 1970 1972

1981 1993

California Federal & California Federal & California California Federal California Federal California Federal Federal & California California

∗ Non-methane

hydrocarbons.

1975 1980

Tab. 2

Pass levels (g mile−1 except where stated)

Area

CO

HC

NOx

1.5% 1.5% 23 39 39 9 16 8 7 3.4 3.4

275 ppm 275 ppm 2.2 3.2 3.4 0.9 1.5 0.41 0.41 0.41 0.25∗

– – – 3.2 – 2 2 1 2 1 0.4

Test method

7 Mode

FTP-72

FTP-75

Legislation on exhaust emission limits (g mile−1 ) for passenger cars and light-duty trucks in California (test weights 2000 h), and at catalyst temperatures from about 1000 K up to about 1300 K. Their pore structure is also designed to balance the requirement for a high internal surface area with the requirement for a limited intraparticle diffusion resistance. The aluminum oxide is typically formed in the desired modification and with the desired chemical composition before it is added to the washcoat. Today, some of the washcoat internal surface area is also provided by cerium and zirconium oxides, both of which are available in a modification that exhibits a moderately high internal surface area (typically 20 to 100 m2 g−1 ), together with an appreciable stability of this internal surface area at typical catalyst operating temperatures. The main task of the cerium oxide washcoat component is oxygen storage, because the cerium ion is easily reduced and oxidized under typical catalyst operating conditions, according to the reaction ← − Ce2 O3 + 0.5O2 2CeO2 − −− −− →

Fig. 29

Scanning electron microscopy image of a washcoat layer.

Fig. 30 Transition electron microscopy image of precious metals on a washcoat particle.

2295

(41)

The oxygen-storage ability widens the range of exhaust gas compositions at transient operation under which the closed-loop-controlled three-way catalyst has optimal conversion for CO, HCs and NOx simultaneously. Indeed, considering a catalyst that was first operated under a slightly lean exhaust gas composition, then the cerium ion will be present in the highest valence state (+4), thus having oxygen ‘‘stored’’. If the exhaust gas composition becomes slightly rich, the ‘‘stored’’ oxygen is released and becomes available for a few seconds to support the conversion of CO and HCs under this oxygen-deficient exhaust gas composition. The cerium ion is now in the lower (+3) valence state, and it will store oxygen again for a few seconds, as soon as the exhaust gas composition changes back to slightly oxidizing. In doing so, the conversion of NOx is enhanced at the slightly oxidizing exhaust gas composition during this transition period. The cerium washcoat component also influences the stability of the dispersion of some of the precious metals, and thus the catalytic function of the precious metal component. Zirconium oxides are the preferred supports for the precious metal component rhodium. The cerium oxide and/or the zirconium oxide are added to the washcoat either as preformed oxides or as oxide precursors, such as their respective carbonates or nitrates; the oxides are then formed in situ during washcoat drying and calcination. References see page 2343

2296

11.2 Automotive Exhaust Treatment

Various methods exist for applying the washcoat to the support. With ceramic monolithic supports, the washcoat is typically applied as an aqueous slurry by a slip coating process, followed by drying and reactive calcination [29]. The same procedure can be used for preshaped metallic supports, the only difference being that some metal foil surface treatments may be used prior to the washcoat application. For some metallic supported catalyst designs, however, the washcoat is applied to the metal foil by a continuous coil slip coating or spraying process, prior to the formation of the monolithic structure. This has the advantage that there is no handling of the many individual support units throughout the washcoating step and, eventually, the precious metal impregnation step. Also, a coil coating process usually leads to a homogeneous thickness of the washcoat layer over the metal foil. The major disadvantage is that it is much more difficult to form a very rigid monolith structure from an already coated metal foil, as brazing or welding processes can no longer easily be applied. Precious Metals The early days of automotive catalytic converter research and development, targeted the use of base-metal catalysts. Numerous publications have described the results obtained with catalysts that contain, for example, the oxides of Cu, Cr, Fe, Co and Ni [7, 8]. Despite this major effort, no real breakthrough was achieved. One reason for this is that the precious metals have a much higher intrinsic activity (i.e., activity per gram component) for the simultaneous conversion of CO, HCs and NOx than base metal catalysts. Furthermore, it is easier not only to obtain but also to keep precious metals in a very finely divided state. Precious metal-based catalysts are also much more resistant to sulfur poisoning at temperatures below 750 K than are base-metal catalysts. The precious metals currently used in three-way catalysts are platinum, palladium, and rhodium. In the past, also ruthenium and iridium have been tested, but because of the volatility and/or the toxicity of these metals or their oxides, neither has yet found practical application in three-way catalysts. Until recently, most of the monolithic three-way catalysts contained platinum and rhodium, in mass ratios of about 5 to 20 : 1 Pt : Rh. The total precious metal loading is typically 0.9 to 2.2 g L−1 catalyst volume. These are only typical values, as the amount of precious metals and the mass ratio of platinum to rhodium depend on the specific application of the catalyst, and is also governed by factors such as the composition of engine-out emissions, the emissions targets to be reached, the catalyst operating conditions, and the properties of the fuels used. In the past, palladium was a common element in oxidation catalysts where it was used together with 11.2.4.5.6

platinum, in a mass ratio of about 5 : 2 Pt : Pd at a total precious metal loading of about 1.5 g L−1 catalyst volume. Some early designs of three-way catalysts used palladium together with platinum and rhodium, in which the palladium was a partial replacement of platinum. The loading used was 0.9 to 3.1 g Pt, zero to 3.1 g Pd, and 0.15 to 0.5 g Rh per converter [9]. Later on, three-way catalyst formulations were used with the majority or even all of the platinum replaced by palladium. These catalysts had a higher total precious metal loading, typically in the range 2 to 5.5 g L−1 catalyst volume, and a mass ratio of Pt : Pd : Rh of about 0–1 : 8–16 : 1 [30]. Today, with the advanced catalyst technology and the much-improved boundary conditions for the catalyst application – such as clean fuel with ultra-low sulfur content and performance engine management systems – rhodium can be combined with either platinum or palladium to achieve a highperformance catalyst. In addition, catalysts have been developed that use only palladium as the precious metal component, with Pdloadings of between 1.8 and 10.6 g L−1 catalyst volume. These catalysts are mainly used as light-off catalysts; these typically are small catalysts units located close to the engine outlet, and are used in combination with a main catalyst that contains rhodium together with platinum or palladium and which is located in the vehicle underbody downstream of the light-off catalyst. These palladium-only catalysts are sometimes also used as the main catalyst, in conjunction with particular engine management systems. The precious metal composition is typically uniform in the macroscopic radial and axial directions of the monolith structure, although different designs have been described in the patent literature, and have even been used in some selected applications. On a microscopic scale, a non-uniform distribution of the precious metals within the washcoat layer is a commonly used architecture. One example of such a non-uniform distribution is that the amount of one precious metal component decreases from the part of the washcoat that is in contact with the gas phase towards the part of the washcoat that is in contact with the monolith wall, and eventually vice-versa for the second precious metal component. Another example of a non-uniform distribution within the washcoat is that each precious metal component is selectively deposited on a different washcoat component. These non-uniformities are intentional and are desirable for kinetic reasons or because of specific beneficial interactions between the precious metals and the washcoat oxides. The type of nonuniformity that can be achieved depends also substantially on the production procedure of the catalyst. Within a single secondary washcoat particle, the distribution of the precious metals can be assumed to be relatively homogeneous. The precious metals are typically present in a highly dispersed state. Dispersions measured

11.2.4 Catalytic Systems for Gasoline-Fueled, Indirect Injection Spark-Ignition Engines

by CO chemisorption methods are typically in the range 50 to 100% for fresh catalysts. This means that the precious metals are present as single atoms or as small clusters of a few atoms. For a catalyst with about 1.8 g precious metal per liter of catalyst volume, this corresponds to a precious metal surface area in the range of about 3 to 30 m2 L−1 catalyst volume. Compared to a washcoat surface area of about 20 000 m2 L−1 catalyst volume, it is apparent that the precious metal surface area is several orders of magnitude lower than monolayer coverage of the washcoat surface. These values are only approximations of the order of magnitude; they are valid only for fresh catalysts and are different for each of the individual precious metals components. The precious metals are generally introduced in the catalyst by wet chemical methods such as incipient wetness impregnation, typically using aqueous solutions of the precious metal salts, followed by a drying step to remove the water, and then by a reactive calcination step to decompose the precious metal salts. Sometimes, a reduction step is applied to convert the precious metal oxides into the metallic state. Other

Tab. 14 Relative importance of some countries in the supply and reserves of Pt, Pd, and Rh (in 2005)

Metal

Country

Demand/tons

Share of use/%

247 253 31

51 49 87

1990

1997

2004

Metal

Pt Pd Rh Pt Pd Rh Pt Pd Rh Pt Pd Rh

Total supply/ tons 77 78 6a 130 135 9.8 163 239 20 214 283 23

Platinum/ Palladium/ Rhodium/ % % %

South Africa

Tab. 13 Historical overview of the price and the market supply of Pt, Pd, and Rh

1983

preparation procedures are described in the patent literature. Because of the large number of automotive exhaust catalysts produced annually, they constitute an important fraction of the worldwide precious metal consumption (Table 12). Changes in emission legislation and in automotive emission control catalyst technology therefore might affect both the supply and the price of platinum and palladium, and especially of rhodium (Table 13). Platinum, palladium and rhodium are mined in a limited number of countries, the most important being South Africa, the former USSR, the USA, and Canada. These countries report huge reserves, albeit with different mine ratios of Pt to Pd to Rh (Table 14) [31]. The concentration of Pt, Pd and Rh in the raw ore is reported to be in the order of magnitude of a few parts per million, whereas their concentration in automotive catalysts is in the order of magnitude of about 3000 wt. ppm for Pt and about 600 wt. ppm for Rh. These aspects contribute to the increasing interest in the recycling of the precious metals from used automotive emission control catalysts. The evolution of the contribution of recycling to the demand of precious metals for the automotive emission control application is indicated in Table 15 [32, 33].

Tab. 12 Annual demand of Pt, Pd, and Rh in the Western world, and share of use in automotive emission control catalysts (2005)

Platinum Palladium Rhodium

Year

2297

Automotive catalyst use/%

Average price price/US$ per troy ounce

31 11 11a 39 9 75 33 41 80 47 54 88

424 135 300 472 115 3726 390 150 350 850 200 1000

Supply Reserves Russia Supply Reserves USA & Canada Supply Reserves Other Supply Reserves

77 74 13 6 5 4 5 16

30 50 55 19 11 10 4 21

Tab. 15 Relative importance of recycling in the supply of Pt, Pd, and Rh for automotive emission control catalysts (in 2005)

Precious metal

Platinum Palladium Rhodium

References see page 2343

Share of recycling in supply/% 20 16 17

83 – 12 – 3 – 2 –

2298 11.2.4.6

11.2 Automotive Exhaust Treatment

Aspects of Three-Way Catalyst Performance

11.2.4.6.1 Introduction The performance of three-way catalysts depends upon numerous factors, as summarized in Fig. 31. These can be grouped into factors related to the chemistry of the catalyst (e.g., the washcoat, precious metals, age, and preparation), the physics of the catalyst (e.g., support and converter design), and the chemical engineering aspects of the catalyst (e.g., reaction temperature, residence time, gas composition and dynamic conditions). These factors are not independent and vary for each particular application of the three-way catalyst. It is therefore not possible at this point to provide a scientifically sound and complete description of all the factors that influence the catalyst performance, nor to derive general rules on the extent of their influence. Hence, only selected topics are described in the following sections, all highlighting ceramic monolithic three-way catalysts used in closed-loop-controlled gasoline spark-ignition engines. 11.2.4.6.2 Measurement of Catalyst Performance The ultimate test of catalyst performance is the vehicle test described in Section 11.2.2.3. In such a vehicle test, all of the reaction conditions that influence the conversion reached over the catalyst vary simultaneously in an interdependent fashion, as they are fixed by the speed and load of the vehicle at each moment of the test. For research and development purposes, however, it is useful to evaluate the catalyst performance using fewer parameters, or parameters that can be varied independently. Several simplified test procedures have therefore been developed. Engine tests are performed with an engine mounted on a computer-controlled brake, with an exhaust gas cooler or heater installed in between the engine outlet and the catalytic converter inlet. Several catalytic activity tests can be performed using this set-up.

Support design

Converter design

Reaction temperature

Space velocity

Washcoat Three – way catalyst performance

Gas composition

Precious metals

Preparation

Age

Dynamic conditions

Fig. 31 Factors that affect the performance of three-way catalysts. (Reproduced with permission from Ref. [42];  1991 Society of Automotive Engineers, Inc.)

Light-off tests aim to measure catalyst performance at various settings of the exhaust gas temperature. One way of doing this is to fix the engine speed, load and A/F value, and then to adjust the exhaust gas temperature in front of the catalyst by the exhaust gas cooler or heater. The catalyst performance is determined by measuring the conversion of CO, HCs and NOx at each exhaust gas temperature setting. The test can be repeated at various engine A/F settings, and allows measurement of the catalyst activity at fixed exhaust gas composition and fixed overall space velocity used. There are of course other ways to evaluate the light-off behavior, but usually these feature a more complex portfolio of contributing phenomena and simultaneously varying parameters, thus complicating the cause–effect analysis. For example, an engine can be operated at a given speed and load setting, with the exhaust gas bypassing the catalyst. A valve is then switched to guide the exhaust gas through the catalyst, and the time needed to reach 50% conversion for a given exhaust gas constituent is recorded. With such a procedure, both the intrinsic chemical activity of the catalyst function, as well as the thermal properties of the catalyst substrate, such as thermal conductivity and mass, are addressed. In the A/F scan the engine is again operated at a fixed speed and load, and the composition of the exhaust gas at the catalyst inlet is then varied. This test can be performed by at first successively increasing and then at successively decreasing set points of the A/F ratio to detect possible hysteresis phenomena. Sometimes, the traverse of the A/F ratio range of interest is performed in a few seconds, in which case the dynamics of the transition of the conversion between the various A/F settings is accounted for. Another way is to wait at each A/F set point until a stable conversion value is obtained. The A/F scan allows comparison of the catalyst performance at fixed space velocity and exhaust gas temperature. The main differences between these two types of engine test and the vehicle test are summarized in Table 16. Engine tests offer the additional advantage over vehicle tests, that several catalysts can be evaluated at about the same time under really identical engine operation conditions. This is achieved with a multichamber converter, in which small cores drilled out of a full-size converter are used (Fig. 32). The size of the cores is such that each is operated at the same space velocity, as the full-size catalyst would experience. A further simplification in the procedure to measure the catalytic activity is the use of a model gas reactor (Fig. 33). A small piece of catalyst is mounted in a reactor tube, the temperature of which can be externally adjusted. Most commonly, a monolithic core of diameter 2.54 cm and length 7.50 cm is used. Mixing either pure gases or mixtures of the desired exhaust gas component with nitrogen simulates the desired exhaust gas composition.

11.2.4 Catalytic Systems for Gasoline-Fueled, Indirect Injection Spark-Ignition Engines

2299

Tab. 16 Major differences between vehicle and the engine test procedures used to measure the activity of three-way emission control catalysts

Operating conditions

Vehicle test

Engine tests

Gas temperature Space velocity Feedstock partial Pressure Redox ratio A/F modulation frequency A/F modulation amplitude

Variable Variable

Variable Fixed

Fixed Fixed

Variable Variable

Fixed Fixed

Fixed/variable Variable

Variable

Fixed

Fixed

Variable

Fixed

Fixed

Light-off

A/ F scan

Such a model gas reactor test gives the highest possible flexibility, as each of the characteristic parameters of the exhaust gas, such as composition, temperature and space velocity, can be varied in a truly independent fashion. 11.2.4.6.3 Influence of Catalyst Operating Conditions In the chemical and petrochemical industries, solid catalysts are typically operated in a narrow range of reaction conditions that are chosen so as to achieve an optimal feedstock conversion at minimal catalyst deactivation. The conditions are typically either constant over time or are slightly modified to compensate for feedstock conversion loss due to deactivation of the catalyst.

FI

FIRC

Fig. 32 Engine dynamometer equipped with a multichamber catalytic converter, for the simultaneous aging of eight catalyst samples.

Quite the opposite is true for automotive catalytic converters. None of the operating conditions can be chosen to guarantee optimal conversion at minimal deactivation, as engine speed and load fix the operating conditions. Further, the engine operation conditions depend upon the driving conditions. A typical example of the influence of engine operation upon the reaction conditions of a three-way catalyst on a ceramic monolith with 62 cells cm−2 is shown in Table 17. Figure 34 illustrates, in the reaction temperature versus reactant space-time diagram, the range of operation conditions of a three-way catalyst between an engine at idle and an References see page 2343

PIC

PI

PIC

PI

FIRC

HC

PIC

PI

FIRC

SO2

PIC

PI

FIRC

PIC

PI

FIRC

PIC

PI

FIRC

PIC

PI

FIRC

Wet chemistry

Mode gas test unit

N2 CO2 HC/N2

NH3

NO/N2

H2 TIRC

SO2 /N2

O2

CO / H2 /N2

CO

CO2

O2 /N2 PIC

PI

FI

Fig. 33

FIRC NO / NOX

N2 HCllquld

TIRC

TIRC HPLC – pump

Model gas reactor to study the activity of automotive emission control catalysts.

N2 O

2300

11.2 Automotive Exhaust Treatment Tab. 17 Influence of engine operation conditions upon the reaction conditions over a ceramic monolithic three-way catalyst (gasoline-fueled spark-ignition engine with four cylinders and a total displacement of 1.8 L; ceramic monolith catalyst with 62 cells cm−2 ; total volume 1.24 L)

Operation conditions Torque Rotational speed Catalyst inlet temperature Space velocity Linear velocity Residence time Mass flow Gas composition CO HC NOx O2 λ-Value Dimensionless quantities Reynolds number Nusselt number

Units

Idling

Nm rpm K N L−1 h−1 m s−1 s−1 kg h−1

0 900 553 5700 0.54 0.3 9.14

18 2000 698 16 100 1.92 0.075 25.81

26 3000 778 27 100 3.62 0.045 43.45

41 4000 918 47 100 7.54 0.021 75.52

114 5000 1183 123 900 25.2 0.006 198.65

0.76 689 162 1.17 1.0147

0.55 527 980 0.69 1.0047

0.68 514 1820 0.70 1.0035

0.78 521 2820 0.67 1.0015

1.05 380 2670 0.43 0.9855

13 0.11

36 0.21

61 0.31

108 0.46

282 0.896

vol.% vol.ppm vol.ppm vol.%

Partial load

Engine idling

10 000

Propylene ammonoxidation Phthalic anhydride from o – xylene Engine partial load

1000

Catalytic cracking

W / F° kg ⋅ s mol

Full load

Engine full load HC – Automotive

100 Hydrodesulfurization NaphthaReforming

10

NOx – Automotive

Fischer – Tropsch

CO – Automotive

Steam reforming 0 0

100 200 300 400 500 600 700 800 900 1000 1200 1300

Temperature/°C Comparison between the operating range of catalyst temperature and reactant space time W/F ◦ for a three-way catalyst at two engine operation conditions (idling and full load operation), and for various heterogeneous catalytic processes used in the chemical and petrochemical industries. (Reproduced with permission from Ref. [34],  1991 Society of Automotive Engineers, Inc.)

Fig. 34

engine at full load operation. For comparative purposes, the catalyst operation range for some major chemical and petrochemical processes is also provided [34]. The data in Table 17 also show that the exhaust gas flow is laminar inside the channels of a ceramic monolith with 62 cells cm−2 , at all engine operation conditions. With a Reynolds number in the range of about 10 up

to about 300, it is not surprising that both the limiting Sherwood and Nusselt numbers also assume low values, which means that there is only a limited contribution of convection to the transfer of heat and mass from the gas phase to the catalyst surface. However, as the exhaust gas flow is turbulent in front of the catalyst, there is a region of flow pattern transition

11.2.4 Catalytic Systems for Gasoline-Fueled, Indirect Injection Spark-Ignition Engines

at the inlet of the monolith. The length of this transition zone can be estimated from le = (0.05) · dk · Re

770 K, the rate of the conversion is controlled by the rate of the intraparticle diffusion, which is the diffusion within the pores of the washcoat. The corresponding apparent activation energy is about 25 kJ mol−1 . Between 770 K and 1200 K, the catalyst is operated under interphase diffusion control, which is the rate of mass transfer between the gas phase and the washcoat boundary surface. The apparent activation energy is now about 6 kJ mol−1 . Finally, above 1200 K, for some reactions a non-catalytic, homogeneous gas-phase reaction occurs [36]. As described above, the catalyst light-off, which is typically defined as the temperature at which under the chosen reaction conditions the reactant conversion reaches the value of 50%, occurs between about 470 K and 570 K, for a catalyst operated at a space velocity of 60 000 N L−1 h−1 . Although the phenomenon of catalyst light-off is complex and still not well understood, numerous simulations have shown that, depending on the difference between the temperature of the gas phase and the temperature of the catalyst, light-off can occur either at the inlet of the monolith and then progress to the downstream part of the catalyst. Alternatively, the light-off can occur in the middle of the catalyst, and then progress both upstream and downstream through the catalyst [37, 38]. With the present generation of three-way catalysts, the light-off phenomenon is in many applications the factor that governs the overall conversion in the vehicle test procedures. This is demonstrated in Table 18. The raw emissions of CO, HCs and NOx are of the same order of magnitude in each of the three phases, but their conversion is below about 90% only in the first phase. The two main reasons for this are: (i) that the catalyst needs to be heated up from ambient temperature until the light-off temperature during this phase; and (ii) that the exhaust gas composition is rich during the first minutes following the cranking of the engine. The oxygen sensor must also exceed a minimum temperature before it starts to function. As shown in Fig. 36, once the catalyst light-off has occurred, the exhaust gas temperature downstream of the catalyst will exceed that in front of the catalyst, because of heat released by the exothermic combustion

(42)

where le is the length of the flow transition zone (m), dk is the width of the monolith channel (m), and Re is the Reynolds number in the monolith channel. It can be calculated with this equation that the length of the transition zone is typically less than 20% of the monolith length [35]. Because of the wide range of catalyst operation conditions, it is to be expected that a variety of kinetic regimes will occur during the use of the catalyst. This is exemplified in the Arrhenius diagram for the Eq. (13) measured in a Berty reactor with a monolithic catalyst (Fig. 35) [36]. Below a gas temperature of about 470 K, the reaction rate is so small that almost no conversion is reached over the catalyst. Above this temperature, and up to about 570 K, the extent of conversion is governed by the rate of the chemical reaction, with an apparent activation energy of about 100 kJ mol−1 . The catalyst light-off occurs in this temperature range. In the temperature range 570 K to

pCO = 0.005 bar pNO = 0.005 bar

3

−4

ln rs

2

−6 Reaction:

−8

CO2 + ½N2

CO + NO

1

−10 −12

900°C 600 °C 500 °C 400 °C

0.8

1.0

1.2

1.4

300 °C

1.6

1.8

2301

200 °C

2.0 ⋅ 10−3

1 −1 [K ] T Fig. 35 Arrhenius diagram for Eq. (13), recorded in a Berty reactor experiment with a fresh three-way catalyst (monolith catalyst with 62 cells cm−2 ; partial pressure CO 500 Pa, partial pressure NO 500 Pa, balance N2 ; Pt: 1.1 g L−1 , Rh: 0.2 g L−1 ). (Reproduced from Ref. [36] with kind permission of Elsevier Science.)

References see page 2343

Engine-out and tailpipe emissions of a passenger car equipped with a closed-loop three-way emission control catalyst in the three phases of the US-FTP 75 vehicle test cycle (engine-aged catalyst)

Tab. 18

Emission

Engine-out Tailpipe Conversion

Unit

g mile−1 g mile−1 %

Phase 1

Phase 2

Phase 3

Total

CO

HC

NOx

CO

HC

NOx

CO

HC

NOx

CO

HC

NOx

18.6 9.6 48

2.0 0.6 69

3.8 0.5 85

11.6 0.14 99

1.9 0.02 99

2.8 0.10 96

10.0 0.72 93

1.5 0.07 95

4.2 0.14 96

12.6 2.3 82

1.8 0.2 91

3.4 0.2 94

11.2 Automotive Exhaust Treatment

Temperature / °C

2302

550 500 450 400 350 300 250 200 150 100 50 0

‘‘Cold start’’

‘‘Stab. portion’’

‘‘Hot start’’

-Sample no.-

−1−

−2−

−3−

−4−

−5− −6− 1972

505

−7−

−8−

Test length / s Behind catalyst

In front of catalyst

500

Temperature before cat., °C

Temperature before cat., °C

Fig. 36 Evolution of the exhaust gas temperature in front of the catalytic converter, and behind the catalytic converter, as a function of the time in the US-FTP 75 vehicle test cycle for a US Tier 1-calibrated vehicle.

400 300 200 100 0

Lambda

1 0.9

300 200 100 0

1.2

0.8

1

6000

6000

Concentration HC before cat., ppm

Lambda

400

1.4

1.1

Concentration HC before cat., ppm

500

4000 2000 0 0

1

2

3

4 min

Vehicle I

4000 2000 0 0

1

2

3

4 min

Vehicle II

Evolution of the exhaust gas temperature in front of the catalytic converter, the exhaust gas lambda value, and the hydrocarbon concentration in the exhaust gas, for two different US Tier 1-calibrated engine designs, as a function of the time during the first 4 min of the US-FTP 75 test cycle. (Reproduced with permission from Ref. [39],  1993 Society of Automotive Engineers, Inc.)

Fig. 37

of CO and of the HCs. This exotherm is between about 50 K and 100 K. The catalyst temperature is rather constant during the second phase of the test cycle, and is maintained to some extent after the engine has stopped.

The time taken for a catalyst to reach 50% conversion of the reactants depends heavily on the design of both the engine and the exhaust aftertreatment system, in addition to the catalyst formulation. The effect of the engine design is exemplified in Fig. 37. At a similar distance between

11.2.4 Catalytic Systems for Gasoline-Fueled, Indirect Injection Spark-Ignition Engines

Converson of CO, %

100 80 60 40

Static, l = 0.995 Dynamic 1 Hz +/− 0.068 l Dynamic 1 Hz +/− 0.034 l

20 0 200

250

300

350

400

450

Temperature before cat., °c Fig. 38 Influence of the dynamics in the exhaust gas composition on the conversion of carbon monoxide over a three-way catalyst at various settings of the exhaust gas temperature (monolith catalyst 62 cells cm−2 ; engine bench light-off test, space velocity 60 000 N L−1 h−1 ; Pt: 1.42 g L−1 , Rh: 0.28 g L−1 ).

100

CO conversion / %

90 80 70 60

0.4% O2 0.6% O2 0.8% O2 1.0% O2 1.1% O2

50 40 30 20 10

75 10 0 12 5 15 170 5 20 0 22 5 25 0 27 5 30 0 32 5 35 0 37 5 40 0 42 5 45 0 47 5 50 0

0 25 50

0

Mean catalyst temperature / °C

(a) 100

HC conversion / %

90 80 70 60

0.4% O2 0.6% O2 0.8% O2 1.0% O2 1.1% O2

50 40 30 20 10 75 10 0 12 5 15 0 17 5 20 0 22 5 25 0 27 5 30 0 32 5 35 0 37 5 40 0 42 5 45 0 47 5 50 0

0 25 50

0

Mean catalyst temperature / °C

(b) 100 90

NOx conversion / %

80 70 60

0.4% O2 0.6% O2 0.8% O2 1.0% O2 1.1% O2

50 40 30 20 10 75 10 0 12 5 15 0 17 5 20 0 22 5 25 0 27 5 30 0 32 5 35 0 37 5 40 0 42 5 45 0 47 5 50 0

0 0 25 50

the engine outlet collector and the catalyst inlet, the ramp of the exhaust gas temperature in front of the catalyst is about 180 K min−1 for vehicle I, and about 300 K min−1 for vehicle II, in which the engine-management was calibrated for quick heating of the exhaust gas following engine cranking. Furthermore, vehicle II was equipped with a so-called secondary air-pump, which adds air to the exhaust gas in front of the catalyst for some time after the cranking of the engine. In doing so, the composition of the exhaust gas in front of the catalyst is adjusted to lean, which favors the rate of the oxidation reactions, as will be shown below [39]. Another major difference between the operation of solid catalysts in chemical and petrochemical conversion processes on the one hand, and in automobile catalytic converters on the other hand, is that the exhaust gas composition in front of the three-way catalysts varies periodically. This is because of the feed-back control of the engine. This dynamic nature of the exhaust gas composition also affects the conversion reached over the catalyst, as is exemplified in Fig. 38. These phenomena are related to the dynamics of the coverage of the catalytic sites by the reactants and/or reaction products. The overall composition of the exhaust gas also has a drastic effect upon the conversion, which can be reached over the catalyst. As stated above, the optimal simultaneous conversion of CO, HCs and NOx along with minimal secondary emissions is reached with a stoichiometric exhaust gas composition. The operation of the engine outside of the stoichiometric point affects the amount of CO, HCs, NOx and O2 emitted simultaneously, as well as the composition of the HC fraction. The oxygen content of the exhaust gas is of particular interest, as this is measured by the lambda sensor and is used to control the engine. Figure 39 shows as an example the influence of the exhaust gas oxygen content on the conversion of CO, HCs and NOx , at fixed CO,

2303

(c)

Mean catalyst temperature / °C

Fig. 39 Influence of the exhaust gas oxygen content on the conversion of: (a) CO; (b) HC; and (c) NOx reached over a three-way catalyst at various settings of the catalyst temperature (monolith catalyst with 62 cells cm−2 , three-way formulation with Pt: 0.38 g L−1 , Rh: 0.16 g L−1 ), fresh; model gas reactor; space velocity 60 000 N L−1 h−1 ; static conditions.

HCs and NOx content. This experiment was performed in a model gas reactor in which a stoichiometric composition is reached at about 1.0 vol.% O2 . With an oxygen content above the stoichiometric value, the conversion of CO and HCs at a given temperature is increased, whereas the conversion of NOx is decreased with respect to the values obtained at stoichiometry. With an oxygen content below the stoichiometric value, the conversion References see page 2343

11.2 Automotive Exhaust Treatment

HC conversion / %

90

0.6% O2

40

CH4 C2H6 C2H4 C3H8 C3H6 C4H10

30

C6H14

80 70 60 50

20 10 0 50

75 10 0 12 5 15 0 17 5 20 0 22 5 25 0 27 5 30 0 32 5 35 0 37 5 40 0 42 5 45 0 47 5

0 25 50

0

Mean catalyst temperature /°C

(a) 100

HC conversion / %

90

0.8% O2

40

CH4 C2H6 C2H4 C3H8 C3H6 C4H10

30

C6H14

80 70 60 50

20 10

0 25 50

75 10 0 12 5 15 0 17 5 20 0 22 5 25 0 27 5 30 0 32 5 35 0 37 5 40 0 42 5 45 0 47 5 50 0

0

Mean catalyst temperature /°C

(b) 100 90

1.0% O2

40

CH4 C2H6 C2H4 C3H8 C3H6 C4H10

30

C6H14

80 70 60 50

20 10 75 10 0 12 5 15 0 17 5 20 0 22 5 25 0 27 5 30 0 32 5 35 0 37 5 40 0 42 5 45 0 47 5 50 0

0 25 50

0

(c)

Mean catalyst temperature /°C 100 90

1.1% O2

40

CH4 C2H6 C2H4 C3H8 C3H6 C4H10

30

C6H14

80 70 60 50

20 10

(d)

75 10 0 12 5 15 0 17 5 20 0 22 5 25 0 27 5 30 0 32 5 35 0 37 5 40 0 42 5 45 0 47 5 50 0

0 0 25 50

Influence of Substrate and Converter Design In the application of catalytic converters to vehicles, an extremely broad range of different converter designs is used. The reasons for this are that each vehicle has different raw emissions, different catalyst operation conditions, and different – in most cases limited – space available in the vehicle underbody to accommodate the catalyst. As a rule of thumb, the volume of the converter is such that the geometric volume of the ceramic monolith corresponds approximately to the engine displacement. At 11.2.4.6.4

100

HC conversion / %

of HCs shows a maximum as a function of the exhaust gas temperature. This maximum can be caused by the occurrence of HC-reforming reactions by which HCs are formed that have a drastically different response factor in the flame ionization detector (FID). The extent to which this maximum occurs depends on the type of HC and on the oxygen content (Fig. 40). It should be emphasized that the HC fraction in the exhaust gas is composed of typically more than 100 different individual HCs. The composition of the HC fraction depends on numerous parameters, such as the composition of the fuel, the engine operating conditions, and the design of the engine (Fig. 41). It was shown in Fig. 40 that the HC conversion reached over the catalyst at fixed reaction conditions depends on the HC type. Typically, alkenic and aromatic HCs are more reactive than alkane HCs; the reactivity of alkane HCs also increases with the number of carbon atoms in the molecule. The nature of the HC also affects the conversion of CO and NOx (Fig. 42). With a lean exhaust gas composition, the level of 50% conversion of an alkane HC is reached at a higher temperature than the corresponding temperature for CO and NOx . Alkene HCs reach the 50% conversion level at the same temperature as CO. These phenomena depend heavily on the detailed kinetic mechanism of the conversion reactions, and are therefore influenced by the composition and the design of the catalyst. From the description given above, it is apparent that the tailpipe emissions of CO, HCs and NOx , as recorded in a vehicle test, depend upon numerous factors, the variation of which is caused by the changing operation conditions of the vehicle. From this description, it would also be expected that the performance of one and the same catalyst will depend on the particular application – which explains why each vehicle might need a different catalyst to meet the emission legislation in an optimal way. Nevertheless, with the present generation of catalytic emission control systems, very low tailpipe emission values for CO, HCs and NOx are reached, at all the typical vehicle operation conditions (Fig. 43) [40].

HC conversion / %

2304

Mean catalyst temperature / °C

Fig. 40 Influence of the exhaust gas oxygen content on the conversion of hydrocarbons, as a function of hydrocarbon type, at various settings of catalyst temperature: (a) 0.6 vol.% O2 ; (b) 0.8 vol.% O2 ; (c) 1.0 vol.% O2 ; (d) 1.1 vol.% O2 . (Experimental details as for Fig. 39.)

11.2.4 Catalytic Systems for Gasoline-Fueled, Indirect Injection Spark-Ignition Engines

2305

30 27 24

[vpm]

21 18 15

Vehicle A Vehicle B Vehicle C Vehicle D Vehicle E Vehicle F Vehlcle G

12 9 6 3 0

Methane

Alkanes

Alkenes

Alkynes

Aromatics

Fig. 41 Concentration of various types of hydrocarbon in the engine-out exhaust gas of different gasoline-fueled spark-ignition engines (vehicles A–F) and one diesel vehicle (vehicle G) during the first bag of the US-FTP 75 vehicle test cycle. All vehicles are EU1 design. (Reproduced with permission from Ref. [39],  1993 Society of Automotive Engineers, Inc.)

average engine speed and load, this corresponds to a space velocity of about 60 000 N gas L−1 catalyst h−1 . The space velocity is generally defined as the ratio between the total exhaust gas volume stream at normal conditions (105 Pa, 273 K) and the geometric volume of the monolithic support. It is therefore only a rough basis for comparison, as it does not account for differences in the substrate characteristics and in the catalyst composition or for differences in the exhaust gas composition. As expected, the space velocity – and thus the catalyst volume – has a major impact on the conversion reached over the catalyst. Especially with aged catalysts this is apparent both from the light-off temperature and the dynamic conversion at a fixed temperature (Figs. 44 and 45). Even at a fixed space velocity, a different conversion is obtained depending on how the total catalyst volume is assembled. As shown in Fig. 46, a lower tailpipe emission of CO is recorded when the total catalyst volume is reached by placing two smaller pieces at a distance of typically around 25 mm from each other, as compared to a single, longer piece of the same volume. The reason for this is that the enhanced heat and mass transfer during the transition from turbulent to laminar exhaust gas flow is used twice with the dual-catalyst converter, which occurs when the distance between the two catalysts is sufficiently long to assure again turbulent flow conditions at the inlet of the second catalyst. The same space velocity can be reached by simultaneously varying the diameter and the length of the monolith, keeping the total catalyst volume constant. As shown in Figs. 47 and 48, in accordance with numerical simulations, an increase in the catalyst diameter from 3.66 inch (9.3 cm) to 5.66 inch (14.4 cm) brings about a slight improvement in the dynamic conversion, especially for NOx with aged catalysts, but has a slightly negative influence on the catalyst light-off temperature. The diameter/length

relationship also affects deactivation of the catalyst at very high values of exhaust gas temperature. Indeed, the catalyst solids temperature – which governs the thermal deactivation phenomenon – is the net result of two other phenomena. On the one hand, there is heating through the energy input from the exhaust gas, including heat released during the conversion of some of the exhaust gas constituents along exothermic reaction pathways. On the other hand, there is energy loss through conduction, convection and radiation cooling through the frontal sides of the monolith body. The contribution of the latter becomes significant at an exhaust gas temperature above about 1070 K. The quantity of energy removed via the radiation cooling pathway increases with the increasing diameter of the monolith body. In extreme comparisons (see Fig. 49), the body temperature may be up to 50 K colder, which has a significant impact on the catalyst deactivation at these high baseline values. At identical catalyst volume, the pressure drop over the monolithic substrate is lower for the catalyst with the biggest diameter, and therefore with the shortest length. Indeed, under laminar flow conditions, the pressure drop over the monolith follows from
p = 32 · µ · Vk · G · (dk + wk )2 /(dk4. · Sk2 )

(43)

where
p is the pressure drop over the monolith (Pa), µ is the dynamic viscosity (kg m−1 ), Vk is the monolith geometric volume (m3 ), G is the gas flow (m3 h−1 ), Sk is the monolith frontal area (m2 ), dk is the diameter of the monolith channel (m), and wk is the thickness of the monolith wall (m). A given space velocity can also be reached with monoliths that differ in cell density, wall thickness, References see page 2343

2306

11.2 Automotive Exhaust Treatment

100 90

C3 H 8

HC −conversion /%

80 +

70

CO NOx O2

60

HC +

50

+ +

40

+

30

+

+

+

20

+

10 +

0

0

25

50

+

75 100 125 150 175 200 225 250 275 300 325 350 375 400 425 450 475 500

Mean catalyst temperature/°C

(a) 100 90

C 3H 6

80 +

Conversion / %

70

CO NOx O2

60

HC +

50 40

+ 30

+ +

20

+

10 +

0

0 (b)

25

50

+

75 100 125 150 175 200 225 250 275 300 325 350 375 400 425 450 475 500

Mean catalyst temperature/°C

Conversion of CO, O2 , NOx and the hydrocarbon component reached over a three-way catalyst at various settings of catalyst temperature, as a function of hydrocarbon type: (a) C3 H8 ; (b) C3 H6 (experimental details as for Fig. 39). (Reproduced with permission from Ref. [34],  1991 Society of Automotive Engineers, Inc.)

Fig. 42

and wall porosity. The cell density and wall thickness of the monolith mainly affect the pressure drop over the monolith [Eq. (43)]. These parameters, together with the porosity of the monolith walls, also affect the bulk density of the monolith, and therefore both the weight and the thermal mass of the converter. Although these substrate-related parameters of course do affect the performance of the catalyst, their influence proved to be less important than the influence of the parameters of a chemical nature, such as precious metals formulation and washcoat composition. This is particularly apparent for the aged state of the catalytic converter. Real-world experience

showed that the optimal cost–benefit relationship is obtained with cordierite monoliths with a cell density in the range of 62 to 93 cells cm−2 and a wall thickness of between 0.075 and 0.16 mm. Finally, the role of the design of the converter cones should be mentioned. Because of limitations in the space available at the vehicle underbody, catalytic converters are often fitted with short inlet and outlet cones. Short inlet cones tend to cause a non-uniform distribution of the exhaust gas flow over the inlet surface of the monolith (Fig. 50). In accordance with numerical simulations, a nonuniform distribution of the exhaust gas flow might

11.2.4 Catalytic Systems for Gasoline-Fueled, Indirect Injection Spark-Ignition Engines

2307

1.5

Emission (g/km)

CO + NOx HC 1.0

0.5 +

+

+ +

+ 0.0 30

40

50

60

70

80

90

100

110

120

130

Average sped (km/h)

Tailpipe emission of CO, HC and NOx for a EU1-calibrated vehicle with a gasoline-fueled spark-ignition engine, equipped with a three-way catalyst, as a function of the vehicle speed. (Adapted from Ref. [40].)

Influence of Catalyst Formulation The catalyst formulation is fixed by the microscopic composition and the properties of its washcoat components, by the nanoscopic composition and properties of the various precious metals used, and by the details of the catalyst preparation procedure. The composition and the properties of the washcoat are one of the key factors that govern the performance and durability of the catalyst. The washcoat is the part of the catalyst, which is in direct contact with the gas phase. Its microscopic geometric – and thermodynamic – properties therefore control, for example, the rate at which the heat of the exhaust gas is transferred to the active sites of the catalyst. An example of this is shown in Fig. 51. As the washcoat actively takes part in the catalytic function, it is to be expected that the amount of washcoat used also has a significant influence on the performance of the catalyst. Figure 52 shows how the NOx conversion 11.2.4.6.5

450

50%

400

90%

350 300

278

291 274

292

289 264

277 254

308 288 282

268

250 200

SV 85.000 72.000 H-1 H-1

53.000 H-1

35.000 H-1

72.000 H-1

35.000 H-1

450 50%

HC/°C

400

70%

350 300

282 289

277 285

296 303 267 274

275

258 265

283

250 200

450 NOx 50%/90% /°C

have a positive influence on the catalyst light-off, as an increased fraction of the energy contained in the exhaust gas is concentrated on a smaller portion of the catalyst. However, flow maldistributions have a negative impact on the conversion at a temperature above the catalyst light-off, as well as on the deactivation of the catalyst by deposition of poisoning elements. Finally, flow maldistributions increase the pressure drop over the converter. For example, it has been shown that the design of the inlet and outlet cones may account for up to 50% of the total pressure drop over the converter [41, 42]. Overall, the design target for optimal and sustainable performance proved to be the homogeneous distribution of the exhaust gas flow over the diameter of the catalyst body structure.

CO 50%/90%/°C

Fig. 43

50% 90%

400 352

346

350 300

341 322

296

300

283

279

298 281

294

258

250 200 Fresh

200 H Engine −

Influence of space velocity on gas temperature needed to reach 50% resp. 90% conversion of CO and NOx , and 50% resp. 70% conversion of HC over a fresh and an engine-aged three-way catalyst (monolith catalyst with 62 cells cm−2 , three-way formulation with Pt: 1.42 g L−1 , Rh: 0.28 g L−1 , engine bench light-off test at lambda 1.02 for CO and HC, and at lambda 0.986 for NOx ; engine bench aging during 200 h).

Fig. 44

References see page 2343

2308

11.2 Automotive Exhaust Treatment

90 80 70

Fresh 200 h engine

60 50 25 000

50 000

75 000

CO / g mile−1

CO conversion/%

100

1 00 000

2.0 1.8 1.6 1.4 1.2 1.0 0.8 0.6 0.4 0.2 0.0

90

0.4

HC / g mile−1

0.5

HC conversion / %

100

80 70 60 50 25 000

50 000

75 000

0.3

0.1

0.8

NOx / g mile−1

90

NO conversion / %

1.0

60 50 25 000

50 000

75 000

1 00 000

Influence of space velocity on the conversion of CO, HC and NOx , reached over a three-way catalyst in the fresh state and after engine-aging, at fixed exhaust gas temperature and exhaust gas composition (monolith catalyst with 62 cells cm−2 , three-way formulation with Pt: 1.42 g L−1 , Rh: 0.28 g L−1 ; engine bench test at 723 K exhaust gas temperature; exhaust gas composition lambda 0.995; dynamic frequency 1 Hz; amplitude 1 A/F; engine bench aging during 200 h). (Reproduced with permission from Ref. [76],  1990 Society of Automotive Engineers, Inc.) Fig. 45

is affected by the amount of washcoat present on the monolith. The influence is particularly noticeable for aged catalysts. It is difficult to separate the detailed influence of the precious metals both from each other and from the influence of the washcoat composition on the catalyst performance. Each of the precious metals platinum, palladium and rhodium, which preferentially catalyze a different reaction pathway, interact with each other during the operation and aging of the catalyst. This causes their overall influence on the catalyst performance in most cases not to be additive. This finding follows from the data in Fig. 53, which compares the performance of a fully formulated Pt–Rh catalyst to the performance of catalysts having either Pt or Rh alone in the same loadings and on the same washcoat as the fully formulated catalyst. By comparing the performance of each of the precious metals, Pt, Pd, and Rh, at an equimolar loading on

0.61

Length 6.3"

Length 3.12" + 3.12"

Fresh 20 h 1000 °C engine 0.15

0.2

100

70

1.11 0.81

0.15 0.08

0.06

0.0

1 00 000

80

1.68

0.6 0.4 0.2

0.07

0.12

0.14

0.16

0.0

Tailpipe emission of CO, HC and NOx from a US Tier 1-calibrated gasoline-fueled passenger car equipped with a three-way catalytic converter, in the US-FTP 75 vehicle test, as a function of the number of catalysts in the converter at fixed total catalyst volume (monolith catalyst with 62 cells cm−2 ; three-way formulation with Pt: 0.83 g L−1 , Rh: 0.16 g L−1 ; fresh condition and after high-temperature aging for 20 h on an engine bench).

Fig. 46

the same washcoat, Fig. 54 shows that the hydrocarbon conversion reached changes with the reaction conditions in a different way. Increasing the total loading of Pt–Rh catalysts, at a fixed platinum to rhodium ratio, directly affects the lightoff temperature (Fig. 55). This is consistent with the fact that the catalytic activity is under kinetic control in the temperature range where the light-off occurs. The conversion recorded on a fresh catalyst at temperatures above light-off is practically independent of the total amount of precious metals within the indicated range. This is consistent with the fact that the catalytic activity is under interface mass transfer control under these conditions. However, as shown in Fig. 56, after aging of the catalyst, the reaction kinetics again become ratecontrolling at the test conditions (especially in relation to exhaust gas temperature), and therefore the total amount of precious metals strongly affects the conversion. Finally, the influence of the total amount of precious metals on the conversion depends also on the exhaust gas

11.2.4 Catalytic Systems for Gasoline-Fueled, Indirect Injection Spark-Ignition Engines

2309

450

CO 50% / °C

400 350 320 300

305

288

292

3.66'' × 6''

4'' × 5.03''

315 288 294

292

250 200

4.66'' × 3.7''

5.66'' × 2.5''

450 Fresh Aged

HC 50% / °C

400 350 300

325 297

293

314 291

307

321 297

250 200 450

NOx 50% / °C

400 350 300

328 301

315 309

305 301

309 310

250 200

Influence of the monolith radius at fixed catalyst volume on the temperature needed to reach 50% conversion of CO, HC and NOx , reached over a three-way catalyst in the fresh state and after engine-aging for 200 h (monolith catalyst with 62 cells cm−2 , three-way formulation with Pt: 1.42 g L−1 , Rh: 0.28 g L−1 ; engine bench light-off test with a space velocity 60 000 N L−1 h−1 at lambda 1.02 for CO and HC, and at lambda 0.986 for NOx ).

Fig. 47

composition. Figure 57 shows that, after engine aging, the HC light-off under lean conditions is improved by about 40 K upon increasing the Pd loading from 3.5 g L−1 to 40 g L−1 , whereas the improvement is only about 20 K under stoichiometric conditions. Each of the precious metals will interact in a different way with the major washcoat constituents, Al2 O3 and CeO2 . The extent of these interactions is controlled by the catalyst operation conditions, especially the temperature and the net oxidizing power of the exhaust gas. These interactions affect the valence state of both the precious metals and of some of the washcoat constituents. Furthermore, the dispersion and the stability of the dispersion of the precious metals is influenced, as well as the solid-state reactions between the various washcoat constituents. Figure 58 shows the differences in the valence state of Pt and Rh as a function

of the type of washcoat oxide on which they are deposited [43, 44]. Beyond this, the precious metals also catalyze the solid-state interactions in-between some of the washcoat constituents. As an example of this, Fig. 59 shows the effect of both the amount and type of precious metal on the extent of the solid-state reaction between aluminum oxide and cerium oxide, leading to the formation of ceriumaluminate. The overall effect of these interactions is that the catalytic activity and the durability depends strongly on the microscopic and nanoscopic details of the catalyst formulation [45–47]. When multibrick converters with different catalysts are used, the sequence in which the catalysts are mounted in the converter will also affect the overall conversion. References see page 2343

2310

11.2 Automotive Exhaust Treatment

100

98.4

97.3 89.5

92.1

3.66'' × 6''

4'' × 5.03''

90.3

92.3

97.4 97.1

98.0 98.4

4.66'' × 3.7''

5.66'' × 2.5''

92.2 89.2

92.3 90.6

CO / %

80 60 40 20 0 100

85.1

85.2

HC / %

80 Fresh Aged

60 40 20 0 100

NOx / %

80

99.5

99.3

99.4

99.3 97.5

92.1 76.7

84.3

60 40 20 0

Influence of the monolith radius at fixed catalyst volume on the conversion of CO, HC and NOx , reached over a three-way catalyst in the fresh state and after engine-aging for 200 h (monolith catalyst with 62 cells cm−2 , three-way formulation with Pt: 1.42 g L−1 , Rh: 0.28 g L−1 ; engine bench test with a space velocity 60 000 N L−1 h−1 ; exhaust gas temperature 723 K; exhaust gas composition lambda 0.995; dynamic frequency 1 Hz; amplitude 1 A/F).

11.2.4.7

Deactivation of Three-Way Catalysts

Frontal face temperature / K

Fig. 48

1150 1140

Introduction In the application of solid catalysts to the chemical and petrochemical industries, either all precautions are taken to minimize deactivation of the catalyst, or the process is designed for regular regeneration of the catalyst. In contrast, automotive emission control catalysts are a mass application in which the operating conditions cannot be controlled and in which pretreatment of the ‘‘feedstock’’ is rarely possible. Despite this, legislation calls for catalyst durability to be the same order of magnitude as the vehicle’s lifetime. Because of these particularities in the use of automotive emission control catalysts, these materials experience a great number of deactivation phenomena (see Fig. 60).

Fig. 49 Calculated body temperature of the frontal surface area of the monolith during high-temperature aging as a function of the monolith radius at fixed total monolith volume. (Monolith catalyst with 62 cells cm−2 .)

11.2.4.7.2 Assessment of Durability The catalytic converter is only one part in the complex emission control

system. Many parameters outside of the catalytic converter have a consequence on its durability, and therefore

11.2.4.7.1

1130 1120 1110 1100 1090 50

70

85

125

Monolith radius / mm

11.2.4 Catalytic Systems for Gasoline-Fueled, Indirect Injection Spark-Ignition Engines

2311

the most convincing and complete way to prove the durability of a catalyst is to perform a durability test with a fleet of vehicles of identical design on the road under uncontrolled driving conditions. As expected for multivariate systems, the fleet emission test result obeys the Gaussian distribution, but with different characteristics for CO, HCs and NOx , as shown in Fig. 61 [48]. As time and cost constraints make it impossible to use a fleet test as the test procedure in the development of emission control catalysts, several simplified durability tests have been developed. In order of decreasing complexity, these include: • a vehicle driven on the road according to a well-defined driving schedule • a vehicle mounted on a vehicle dynamometer driven according to a well-defined procedure by robots • tests with an engine mounted on an engine dynamometer • tests with a fuel burner bench • tests in a laboratory furnace.

Fig. 50 Influence of design of converter inlet cones on the distribution of exhaust gas flow velocity over the radius of a monolith. Upper: converter with long inlet cones. Lower: converter with short inlet cones.

References see page 2343

Gas inlet temperature

220 Temperature / °C

Which of the deactivation phenomena are addressed by each of these test procedures are summarized in Table 19. It is apparent that the level of abstraction from reality increases substantially with decreasing complexity of the test procedure and corresponding test equipment. The durability tests performed in laboratory furnaces mainly aim at studying thermal deactivation effects, and are usually carried out at temperatures between about 1170 K and 1320 K, under an oxidizing (air) or an inert (nitrogen with or without water and carbon dioxide) atmosphere.

Washcoat C1 Washcoat C2

190 160 130 100 70 40 100

75

50

25

0

25

50

75

100

Radial position / % Fig. 51 Influence of the washcoat heat-transfer coefficient on the distribution of the solid-phase temperature over the radius of a ceramic monolith at the outlet frontal area, for a washcoat with a high heat-transfer coefficient (C2) and, for a washcoat with a low heat-transfer coefficient (C1). (Reproduced with permission from Ref. [34],  1991 Society of Automotive Engineers, Inc.)

2312

11.2 Automotive Exhaust Treatment

100 +

90

+

Conversion %

80 70 60 50 40

100% WC-loading + 85% WC-loading 70% WC-loading 55% WC-loading

30 20 10 0 0

20

40

60

80

100

120

140

160

Simulated aging (1000 km)

Influence of washcoat loading of a ceramic monolith on conversion of NOx as a function of the aging state of a three-way catalyst (monolith catalyst with 62 cells cm−2 , three-way formulation with Pt: 1.42 g L−1 , Rh: 0.28 g L−1 ; engine bench test with a space velocity 60 000 N L−1 h−1 ; exhaust gas temperature 723 K; exhaust gas composition lambda 0.999; dynamic frequency 1 Hz; amplitude 1 A/F; high-temperature engine bench aging). (Reproduced with permission from Ref. [34],  1991 Society of Automotive Engineers, Inc.) Fig. 52

Conversion / %

100

99 96 81

80

CO HC NOx

93 90 80

75

59

60

40

40 20 0

Pt: 41.7 g/ft3

Pt: 40 g/ft3

g/ft3

g/ft3

Rh: 8.3

Rh: 0

Pt: 0 g/ft3 Rh: 8.3 g/ft3

The function of platinum and rhodium on a fully formulated three-way catalyst washcoat in the conversion of CO, HC and NOx (monolith catalyst with 62 cells cm−2 , three-way formulation; engine bench test with a space velocity 60 000 N L−1 h−1 ; exhaust gas temperature 673 K; exhaust gas composition lambda 0.999; dynamic frequency 1 Hz; amplitude 1 A/F; high-temperature engine bench aging during 20 h). (Reproduced with permission from Ref. [34]  1991 Society of Automotive Engineers, Inc.) Fig. 53

The durability test performed on engine dynamometers may differ from one application to the other. As shown in Fig. 62, some tests target high temperature stability, whereas others target poisoning effects and are therefore performed at lower temperatures. The common feature of the present-day engine durability tests is that a multitude of engine operation conditions are used, in which the composition of the exhaust gas and its temperature are varied according to a well-defined protocol. Today, this can to some extent also be simulated on fuel burner test benches, which are substantially more cost-effective than engine test benches. Usually, correlations can be established between the engine-, burner-, and furnace tests on the one hand, and the vehicle durability tests on the other hand. In many cases however, these correlations differ

between vehicle platforms, so that the assessment of the aging phenomenon remains a lengthy and costly process. Selected Deactivation Phenomena A drastic and common deactivation phenomenon with automotive emission control catalysts is the irreversible mechanical destruction of the support during road use. In the case of ceramic monoliths, this occasionally occurs through breakage following direct mechanical stress, or also indirectly as a result of thermal stresses following sudden temperature changes. In the case of some metallic substrates this would result from telescoping of the matrix or breakage of the foil. Finally, in the case of bead catalysts, this resulted from extended attrition of the beads. 11.2.4.7.3

11.2.4 Catalytic Systems for Gasoline-Fueled, Indirect Injection Spark-Ignition Engines

500 0.6 v.% O2 1.0 v.% O2 1.5 v.% O2

L.o. temperature / °C

Butene 400 366

3.0 v.% O2 300

291

300

271 263

260 260 240

250

240 225 200

200 Pt/Al2O3

Rh/Al2O3

Pd/Al2O3

500 L.o. temperature / °C

450

Butane

400

370 374 350 300

300

300

300 293 300 300

280 252

200 Pt/Al2O3

Rh/Al2O3

Pd/Al2O3

Fig. 54 Effect of platinum, palladium and rhodium at an equimolar loading on the temperature needed to reach 50% conversion of butene and butane, as a function of the exhaust gas oxygen content (monolith catalyst with 62 cells cm−2 ; γ -Al2 O3 washcoat; fresh state; precious metal loading 8.8 mmol L−1 ; model gas light-off test at a space velocity of 60 000 N L−1 h−1 ; model gas composition is stoichiometric at 1.0 vol.% O2 ). (Reproduced with permission from Ref. [30],  1994 Society of Automotive Engineers, Inc.)

A broad range of deactivation phenomena are related to the operation temperature of the catalyst. Operation at both low and very high temperatures can induce specific catalyst deactivations (Fig. 63). Usually, deactivation phenomena occurring at low temperatures are reversible, which means that they are removed by operating the catalyst at a higher temperature, eventually in conjunction with a different net oxidizing or reducing power of the exhaust gas. Examples of lowtemperature deactivation phenomena are: • physisorption of both, reactants and reaction products, such as CO2 • chemisorption of both, reactants and reaction products, such as the reactions between the various sulfur oxides SO2 and SO3 and the washcoat oxides • oxidation of the precious metals. Deactivation phenomena occurring at higher temperatures are typically irreversible. These include solid-state reactions (as explained above): (i) between the individual

2313

washcoat components; (ii) between the precious metals resulting in the formation of less-active alloys and components not targeted at in the concept of the catalyst; and (iii) between the different precious metals and the washcoat oxides. A common example of the latter group is the migration of Rh3+ ions in the framework of the transitional aluminum oxide, as the crystal structure Rh2 O3 is isomorphic with the crystal structure of γ -Al2 O3 . The most important deactivation phenomena occurring at a high catalyst operational temperature are the loss of the internal surface area of the washcoat components and the loss of the precious metals dispersion. The sintering of the washcoat oxides decreases the internal surface area of the washcoat; this dramatic process also leads to the inclusion of precious metals, as they are supported on the washcoat oxides (Fig. 64). The most important parameter that governs the loss of the washcoat internal surface area is the temperature, whereby it should be noted that each of the washcoat constituents has a different temperature stability. Sintering of the precious metals leads to a loss in precious metals surface area, and also to the establishment of a broader distribution of the precious metals particle diameter. As shown in Fig. 65, a catalyst in the fresh state features a rather homogenous size of the precious metal particles, whereas after aging, very large precious metal crystallites coexist with small ones. The extent to which precious metals sinter depends upon: • their initial dispersion • the nature of the interaction between the precious metal and the washcoat • the type of the precious metal • the net oxidizing power of the exhaust gas. For example, platinum and rhodium sinter much faster under an oxidizing exhaust gas, whereas palladium sinters faster under a reducing exhaust gas. This behavior is consistent with the oxidation state of these elements under the respective atmospheres, and is explained by the difference in the vapor pressure of the metals and their respective oxides (Table 20) [50, 51]. With spark-ignition engines, the solids temperature range during use can be substantially broader than the gas temperature range. This is caused by the exothermic nature of the conversion reactions, and also by the fact that under ignition failures and vehicle deceleration conditions in which the fuel supply to the engine is suddenly cut off, some amount of unburned fuel can still reach the catalyst. The unburned fuel will be oxidized by the catalyst, thereby dramatically raising its temperature. Figure 66 shows as an example the evolution of the temperature of the gas References see page 2343

2314

11.2 Automotive Exhaust Treatment

50% CO temperature / °C

400

Fresh After 200h

After 300h After 400h

350 300 250 200

40 g ft−3

30 g ft−3

20 g ft−3

10 g ft−3

Precious metal content

50% HC temperature / °C

400 350 300 250 200 50% NOx temperature / °C

400 350 300 250 200 Fig. 55 Effect of total precious metal loading at a fixed 5 : 1 Pt : Rh mass ratio on the temperature needed to reach 50% conversion of CO, HC and NOx , as a function of the duration of the engine aging (monolith catalyst with 62 cells cm−2 , three-way catalyst formulation; engine bench light-off test with a space velocity 60 000 N L−1 h−1 , at lambda 1.02 for CO and HC, and at lambda 0.984 for NOx ). (Reproduced with permission from Ref. [76],  1990 Society of Automotive Engineers, Inc.)

phase and the solid phase at deceleration conditions. Despite the fact that the temperature of the exhaust gas in front of the catalyst decreases rapidly, the temperature of the catalyst itself increases by the exothermic combustion of the unburned fuel. The solids temperature may even exceed 1600 K, causing the melting of the substrates [52]. Finally, deactivation of the catalyst by poisoning elements should be mentioned. Precious metal-based catalysts are poisoned by sulfur oxides (which mainly originate from the combustion of sulfur-containing fuel constituents), by phosphorus and zinc (which mainly originate from additives in the engine lubricating oil), and by silicium (which sometimes is present in engine seals). In the past, traces of lead which are present

even in unleaded gasoline fuel due to contamination of the fuel supply chain, have made an important contribution to the deactivation of the catalyst. The order of magnitude of the amounts of the most important poisoning elements to which catalysts were exposed under the boundary conditions of the EU1 legislation step are listed in Table 21. The fraction of the poisoning elements retained by the catalyst is typically distributed non-uniformly over the length of the catalyst and the depth of the washcoat layer, although the highest concentrations occur at the entrance to the monolith (Fig. 67). Each of the poisoning elements interact in a different way with the washcoat constituents and the precious

11.2.4 Catalytic Systems for Gasoline-Fueled, Indirect Injection Spark-Ignition Engines

Fresh After 200 h

2315

After 300 h After 400 h

CO conversion / %

100 90 80 70 60 50

40 g ft−3

30 g ft−3

20 g ft−3

10 g ft−3

Precious metal content

HC conversion / %

100 90 80 70 60 50

NOx conversion / %

100 90 80 70 60 50

Effect of total precious metal loading at a fixed 5 : 1 Pt : Rh mass ratio on the conversion of CO, HC and NOx , reached over a three-way catalyst at a stoichiometric exhaust gas composition, as a function of the duration of the engine aging (monolith catalyst with 62 cells cm−2 , three-way catalyst formulation; engine bench activity test with a space velocity 60 000 N L−1 h−1 , exhaust gas temperature 723 K; exhaust gas composition lambda 0.995; dynamic frequency 1 Hz; amplitude 0.5 A/F). (Reproduced with permission from Ref. [76],  1990 Society of Automotive Engineers, Inc.) Fig. 56

metals. The interaction phenomena are rather complex, and sometimes the simultaneous presence of several poisoning elements causes specific interactions, that would not occur if each of these elements were present alone. Examples of this are the effects caused by the simultaneous presence of Pb and S, or of P and Zn [53]. The type and extent of interaction between the poisoning elements and the catalyst constituents depend strongly upon the operation temperature of the catalyst both during and after the deposition of the poisoning elements. As a rule of thumb, Pb and Si will mainly affect the function of the precious metals, by the

formation of alloys or chemical compounds between the precious metals and these poisoning elements, whereas P, Zn and S will mainly affect the function of the washcoat components [54, 55]. For example, phosphorus will interact with the aluminum oxide of the washcoat and cause considerable loss in the stability of its internal surface area at a high operating temperature. Another example is the joint interaction of phosphorus and zinc with the aluminum oxide of the washcoat, leading to a mechanical obstruction of the washcoat pores. References see page 2343

L.o. temperature / °C (Stoichiometric)

2316

11.2 Automotive Exhaust Treatment

Fresh

400

350

330 340 300

300

Aged 340 310

320

320

320

297 280

330 310

280

200 3.5

5.3

7

10

20

40

Pt/Rh

L.o. temperature / °C (Lean)

Pd-loading g L−1

300

285 226

240

250

226 230 220 230 220 220 190

200 3.5

5.3

7

10

NOx during those vehicle driving conditions where the exhaust gas temperature is below about 770 K. Figure 68 shows as an example the conversion of CO, HCs and NOx for a fresh and a vehicle aged catalyst in different assigned parts of the EU driving cycle. The difference in the conversion reached over the catalyst in the fresh and in the aged state is most obvious during the first phase. The relatively low average catalyst operation temperature causes this during this phase. In the third phase of the driving cycle, only small differences occur between the conversion over the fresh and the engine aged catalyst, because of the relatively high average catalyst operation temperature during this phase.

210

20

190

200

40

Future Developments and Concepts As discussed above, legislation keeps moving, aiming at further reductions in the tailpipe emissions of CO, HCs and NOx during all operation conditions of the vehicle over its entire lifetime. In order to meet these requirements, changes in the design of the engine and catalytic exhaust gas aftertreatment system are also continuously ongoing, usually along with reformulations of the fuel. One important development direction is the ongoing improvement of the performance of the three-way catalytic converter applied to closed-loop-controlled spark-ignition engines during the first few minutes following cranking of the engine. Numerous technologies have been – and are still being – developed to enhance this so-called ‘‘cold-start performance’’. One approach is to increase the temperature of the exhaust gas in front of the catalyst during the cold-start phase. This can be achieved by mounting the catalyst closer to the engine outlet, or by the incorporation of a so-called ‘‘light-off catalyst’’ in the exhaust gas system, upstream of the main catalyst and close to the engine outlet. These measures call for the ongoing development of catalysts with excellent stability at high operation temperatures. Another approach is to minimize losses of exhaust gas heat by insulating the exhaust gas pipe. A further approach is quickly to increase the catalyst temperature by the supply of additional heat. This is achieved either by resistively heating the catalyst support or by the installation of a small fuel burner in front of the catalyst. The technologies that lead to a quicker increase of the catalyst temperature will improve the catalytic performance for the conversion of all three exhaust gas components under consideration. However, some of the new legislation focuses on the reduction of the tailpipe HCs emission, so that catalytic aftertreatment systems have been developed that mainly improve HC conversion. An example is the incorporation of HC adsorbers into the aftertreatment system, in addition to the three-way 11.2.4.8

Pt/Rh

Pd-loading g L−1 Fig. 57 Effect of the palladium loading in a ‘‘palladium-only’’ three-way catalyst formulation on the temperature needed to reach 50% conversion of hydrocarbons with a stoichiometric and with a lean exhaust gas, for fresh and for engine-aged catalysts (monolith catalyst with 62 cells cm−2 , dedicated washcoat for ‘‘palladium-only’’ catalyst; reference three-way catalyst formulation with Pt: 1.42 g L−1 , Rh: 0.28 g L−1 ; engine bench light-off test with a space velocity 60 000 N L−1 h−1 ; exhaust gas composition stoichiometric and lean at lambda 1.15; high-temperature engine bench aging). (Reproduced with permission from Ref. [30],  1994 Society of Automotive Engineers, Inc.)

The net result of the chemical and thermal deactivation phenomena is a decrease in catalyst activity for the main reactions that govern the conversion of CO, HCs and NOx , as well as for the side reactions that cause the formation of both desired (e.g., H2 ) and undesired (e.g., H2 S) secondary emissions under operation of the catalyst with non-stoichiometric exhaust gas compositions. The latter effect can be – and has been – used to cure some unwanted side effects that occurred with the very active fresh state of the catalysts. As some of the main reactions consist of several sequential steps, deactivation of the catalyst can cause an enhanced concentration of intermediate reaction products in the tailpipe exhaust gas. These intermediates include aldehydes, which may be formed as intermediate oxidation products of alcohols, as well as N2 O, which can be the intermediate reduction product of the nitrogen oxides. Again, the extent to which these intermediate reaction products will occur in the tailpipe exhaust gas depends upon the formulation and the operating conditions of the deactivated catalyst. The decrease in catalytic activity by deactivation has the strongest influence on the conversion of CO, HCs and

2317

11.2.4 Catalytic Systems for Gasoline-Fueled, Indirect Injection Spark-Ignition Engines

Rh 3d,g

550 500 450 400 350 300 250 200 150 100 50

Intensity (cps)

Intensity (cps)

Pt 4d,a

500 450 400 350 300 250 200 150 100 50 340

335 330 325 Binding, energy [eV]

(a)

320

Rh 3d,e Rh 3d,a

322

320

(a)

316 314 318 Binding, energy [eV]

400 350 300 250 200 150 100 50

Rh 3d,b

2400 Intensity (cps)

Intensity (cps)

Pt 4d,a

312

Pt 4d,b

2000 1600

Rh 3d,d

1200 800 400

350

345

(b)

340 335 330 Binding, energy [eV]

326

325

324

(b)

322 320 318 Binding, energy [eV]

316

X-ray photoelectron spectroscopy (XPS) spectra showing the valency states of platinum and rhodium as a function of the washcoat composition (model catalysts on monoliths with 62 cells cm−2 ; washcoat (a) γ -Al2 O3 ; washcoat (b) CeO2 ; loading Pt or Rh at 1.76 g L−1 , fresh catalysts). (Adapted from Ref. [43].)

Fig. 58

CeAlO+

200

100 +

CeAlO 2

CeAlO+ relative SIMS - intensity

SIMS intensity / cps

300

100 %

1.76

(a)

Reference

Pt

Pd

Rh

(b)

6.36 7.91

10

g (Pt, Rh) L−1

Secondary ion mass spectra (SIMS) showing the extent of solid-state reactions between alumina and ceria in a three-way catalyst washcoat: (a) as a function of the type of precious metal; and (b) as a function of the total precious metal loading (model catalysts on monoliths with 62 cells cm−2 ; Pt, Pd or Rh loading at 1.76 g L−1 ; various total loadings at fixed 5 : 1 Pt : Rh mass ratio; model washcoat with 70 wt.% alumina and 30 wt.% ceria, after air-aging at 973 K).

Fig. 59

2318

11.2 Automotive Exhaust Treatment

Tab. 19

Main differences between the procedures used to study the deactivation of three-way catalysts

Deactivation mechanism

Vehicle test

Engine test

Burner test

Oven test

On road

On Dynamo meter

Yes Yes

No Yes

No Yes/No

No No

No No

Yes Yes Yes Yes Yes

Yes/No Yes/No Yes/No Yes Yes/No

Yes Yes Yes Yes/No Yes

Yes Yes Yes No Yes

Yes Yes Yes/No No Yes/No

Yes Yes Yes/No

Yes/No Yes/No Yes/No

Yes/No Yes/No Yes

Yes No Yes

No No Yes/No

Mechanical Disintegration Fouling Thermal Sintering Phase change Compounds formation Hot spots Component migration Chemical Poisoning Blocking by products Changes in oxidation state

catalyst. The goal of these adsorbers is to remove the HCs from the exhaust gas by adsorption onto suitable materials, as long as the catalyst temperature is lower than its light-off temperature, and then to supply these adsorbed HCs back to the exhaust gas by desorption as soon as the catalyst temperature is high enough to ensure their conversion. Examples of suitable adsorption materials are activated carbon and zeolites. If adsorption materials are used that withstand the high temperatures of exhaust gases containing oxygen and water vapor, the corresponding HC adsorbers can be either incorporated into the three-way catalyst or mounted in front of the three-way catalyst. Figure 69 compares the HC emission performance of such a combined adsorber-catalyst system to the performance of the catalyst alone. If the adsorption materials cannot withstand the harsh operating conditions, they can be

Substrate disintegration

Fuel poison adsorption

Product adsorption

Washcoat phase changes

Catalyst loss

Precious metal sintering

Deactivation

Oxidation state changes

Active component migration

Fouling

Coking

I R R E V E R S I B L E R E V E R S I B L E

Fig. 60 Overview of the reversible and the irreversible deactivation phenomena of three-way catalysts.

mounted in a bypass to the exhaust gas system, and the exhaust gas passed through the adsorber only during the cold-start phase. Initial designs of HC adsorber systems found an application on a limited number of upperclass passenger car platforms during the EU2 legislative timeframe. In these designs, the HC trap was incorporated in the exhaust system upfront of the three-way catalyst, and a small electrically heated catalyst was placed in between them. This was done to ensure a high degree of HC conversion following their desorption over the lifetime of the vehicle, even with aged, three-way catalysts. Surprisingly, these systems vanished from the EU market as the continuously increasing sophistication of the fuel injection devices (and especially the engine management systems) enabled dedicated engine operation strategies following engine cranking to enable a fast catalyst lightoff. However, they were reintroduced in about 2003 on some vehicles sold in the USA market, the aim being to reach the most demanding SULEV emission legislation level [56–59]. 11.2.5

Catalytic Systems for Gasoline-Fueled Direct Injection Spark-Ignition Engines Introduction The operation of a spark-ignition gasoline engine above the stoichiometric point reduces the engine’s fuel consumption and therefore its CO2 emission. At the same time, the engine-out emissions of CO, NOx and, up to some defined A/F ratio also of HCs, are decreased. These so-called ‘‘lean-burn’’ engines still require exhaust gas aftertreatment, but the conversion level needed to meet the legislative requirements is somewhat lower 11.2.5.1

11.2.5 Catalytic Systems for Gasoline-Fueled Direct Injection Spark-Ignition Engines

as compared to engines operated at the stoichiometric point. As of today, three types of lean-burn engine have been proposed (Fig. 70) [60, 61]. The first type always operates in the lean-burn range, with a lambda value between 1.5 and 1.6. The exhaust gas composition will be net oxidizing at all vehicle-driving conditions. The second type operates under some driving conditions in the lean-burn mode, and for other driving conditions in the stoichiometric mode; the latter conditions occur, for example, during acceleration. The third type of engine also operates under both lean-burn and stoichiometric conditions, the difference from the second type being that the operation mode does not depend upon vehicle driving conditions, but is enforced by the engine management system. These engines will operate for a few minutes under lean-burn conditions,

2319

Tab. 20 Sintering behavior of Pt, Pd, and Rh as a function of the aging atmosphere (washcoat: La2 O3 -doped Al2 O3 , precious metal content 0.14 wt.%)a

Parameter

Pt

Particle size/nm 21 N2 , 1373 K Exhaust gas, 1373 K 78 Air, 1373 K 97 Vapor pressure at 1073 K/torr Metal 9.1 × 10−17 Oxide 1.2 × 10−5

Pd

Rh

97 68 n.a.

14 88 n.a.

1.2 × 10−9 ∼0

2.9 × 10−17 5.8 × 10−6

a Reproduced

from Refs. [50, 51] with kind permission of Elsevier Science. n.a., not available.

References see page 2343

Conversion/%

40 30 HC

20 10 0

0.5

1

1.5

2

2.5

FTP emissions - g/mile

Conversion/ %

40 30

CO 20 10 0

5

10

15

20

25

Conversion/ %

40 30

NOx

20 10 0

0.7

1.4

2.1

2.8

3.5

Emission of CO, HC and NOx from a sample of 532 identical vehicles equipped with three-way catalysts, used on the road since 1985, as measured in the US-FTP75 vehicle test cycle. (Reproduced with permission from Ref. [48],  1988 Society of Automotive Engineers, Inc.)

Fig. 61

2320

11.2 Automotive Exhaust Treatment

100 Aging cycle A Aging cycle B Aging cycle C

90 80 Frequency / %

70 60 50 40 30 20 10 0 300

Fig. 62

400

500

600 700 800 Catalyst inlet temperature / °C

900

1000

1100

Temperature histogram of three engine bench aging cycles using gasoline-fuelled spark-ignition passenger car engines.

Tab. 21 Order of magnitude of the poisoning effects during the aging of three-way catalysts

Compound Catalyst formulation Al2 O3 CeO2 Pt Rh Poisons offered during 80 000 km use Pb S P Zn

Quantity/g L−1 catalyst

by the third type of lean-burn engine, for which special catalysts – the so-called NOx -adsorber systems – were developed; these are described in more detail below. Aspects of Operation of NOx -Adsorber Systems NOx -adsorber systems are catalytic exhaust aftertreatment devices that have the ability to store and release nitrogen oxides during dedicated engine operation conditions. To achieve this, they include three distinct functions. The first function is to oxidize NO with O2 to NO2 on a first platinum group metals (PGM) component, which is optimally platinum at a loading of up to 3 g L−1 catalyst volume, according to the reaction 11.2.5.2

100–200 40–80 1–2 0.1–0.4 10–130 1600–5000 20–50 80–120

after which the engine management system will adjust to stoichiometric or even rich operation for a few seconds. Each of these engine types requires a different catalytic aftertreatment system. With conventional three-way catalysts, sufficient conversion can be reached for CO and for the non-methane HCs under the lean exhaust gas compositions of lean-burn engines. However, three-way catalysts are unable to convert NOx under these net oxidizing exhaust gas compositions. Intensive research programs have shown that elements such as Co and Cu, supported on specific zeolites, are able to convert the NOx to some extent under these conditions [71, 72, 85]. However, the durability of corresponding catalysts, especially their high-temperature stability and their resistance against sulfur poisoning, proved not to be sufficient under the current boundary conditions. This resulted ultimately in a discontinuation of the production of the first and also the second types of lean-burn engine from the EU4 emission legislation level onwards. They were succeeded

← − 2NO2 2NO + O2 − −− −− →

(44)

This reaction is thermodynamically favored at a temperature up to about 773 K, which thus represents the upper end of the temperature window in which such a device can be operated. The second function is to store the NO2 formed. This is usually achieved by the incorporation of up to about 100 g L−1 catalyst volume of one or more alkali and/or alkaline earth components in the catalyst. During operation, these alkali and alkaline earth components are usually present in the form of their corresponding carbonates, which will react with the NO2 supported by O2 according to the reaction − 2‘AE (NO3 )2 + 2CO2 4NO2 + O2 + 2‘AE CO3 ← −−− −− → (45) where ‘AE’ represents a bivalent alkaline earth element. This reaction works well at a temperature above typically 573 K, which thus represents the lower end of the

11.2.5 Catalytic Systems for Gasoline-Fueled Direct Injection Spark-Ignition Engines

2321

Exhaust gas temperature range Catalyst temperature range 1500 Support melting 1400 1300

Support phase change Washcoat deterioration

1200

Formation of a-Al2O3

Thermal deactivation range

1100

Temperature /°C

1000 900 800

Diffusion of Rh2O3 in Al2O3

700

Sintering of Pt

600

Reaction of Zn and P with washcoat Desorption of SO3, SO2

500 400

Chemisorption of SO3

300

Desorption of CO2

200

Chemisorption of SO2 Desorption of H2O

100 0

Fig. 63

Alumina phase change

Operation range

No conversion

Adsorption of H2O; CO2

Schematic overview of some deactivation phenomena with three-way catalysts, as a function of catalyst temperature.

(a)

(b)

Fig. 64 Transmission electron microscopy image of a washcoat alumina in the fresh state (a) and after air-aging for 24 h at 1373 K (b). (Original magnification, ×200 000.)

2322

11.2 Automotive Exhaust Treatment

1200 °C 1100

T2 T3

T 1000

T4

T5 T6

900 800

700

T7

Transmission electron microscopy image of platinum crystallites on alumina. Left: for a fresh catalyst. Right: for a catalyst aged at 1323 K in air. Fig. 65

600

T1

500

temperature window in which the NOx -adsorber systems can be operated. The reactions in both Eqs. (44) and (45) require a net oxidizing exhaust gas composition, and will be applicable until the saturation of the alkali/alkaline earth function is achieved. Under the typical practical application conditions, for a system targeting the EU4 emission legislation level, this takes up to about 2 min. Once this situation is reached, however, the engine operation conditions must be changed to generate a net reducing exhaust gas composition. At the typical prevailing temperature, this will provoke decomposition of the alkali/alkaline earth nitrates, which subsequently can be reduced using the third catalytic function according to the reactions − 4NO2 + O2 + 2‘AE CO3 2‘AE (NO3 )2 + 2CO2 ← −−− −− → (46)

400

0

1

2

3

4

5

6

7

8

9 10

Fig. 66 Evolution of exhaust gas temperature in front of the catalyst (T1) and behind the catalyst (T7), as well as solids temperature at different positions inside a monolithic three-way catalyst (T2–T6) as a function of the time at deceleration of an EU1-calibrated vehicle equipped with a gasoline-fuelled spark-ignition engine without a fuel-cut device. Deceleration starts at time zero. (Adapted from Ref. [52].)

← − 4CO2 + N2 2NO2 + 4CO − −− −− →

(47)

The preferred catalytic function for this is now again based upon PGMs, typically a combination of Pt and Rh in compositions and loadings similar to those used for three-way catalysts. With a properly designed system, it

0.8 0.7

Catalyst entrance Catalyst middle Catalyst exit

0.73

Concentration / %

0.6 0.5 0.4 0.29

0.3 0.2

0.14

0.11

0.1

0.09 0.05

0.02 0.02

0.11 0.05

0.02

0.02

0.0 Zn

Pb

S

P

Fig. 67 Concentrations of Zn, Pb, S and P as a function of the distance from the catalyst inlet for a vehicle-aged monolithic three-way catalyst used in the Federal Republic of Germany in 1985–1989.

11.2.5 Catalytic Systems for Gasoline-Fueled Direct Injection Spark-Ignition Engines

98

100

88

Fresh Aged

90

98

99 94

95

85 78

80

Conversion/%

99 96 98

97

2323

70 60 50

50 40

50

40

30

25

24

20

20

10 0

CO

HC

NOx

CO

Part 1

HC

NOx

CO

Part 2

HC

NOx

Part 3

Conversion of CO, HC and NOx in three different parts of the European vehicle test cycle, reached over a three-way catalyst in the fresh state and after high-temperature engine bench aging. Test performed with a vehicle equipped with an EU1-calibrated gasoline-fueled spark-ignition engine. Part 1 = first city driving cycle; part 2 = second, third and fourth city driving cycle; part 3 = extra-urban driving cycle).

Fig. 68

typically takes only a few seconds to complete these reactions, after which the whole cyclus can be restarted. Upon repeating this cyclus, an overall NOx conversion of typically more than 95% can be achieved, of which the time-resolved pattern is shown in Fig. 71.

Total HC emissions / %

100

80

60

40

20

0 0

20

40

60

80

Time / s TWC(50) + TWC(50), fresh TWC(50)LP + TWC(50), fresh TWC(50)LP + TWC(50), aged

Fig. 69 Integrated emissions of hydrocarbons as a function of time during the first 80 s of the US-FTP 75 vehicle test cycle, for an US Tier 1-calibrated vehicle equipped with a spark-ignition gasoline engine, using a conventional three-way catalyst converter system (TWC (50) + TWC (50)) and a converter system in which a hydrocarbon adsorber is combined with a three-way catalyst (TWC (50)LP + TWC (50)). (Monolith catalyst with 62 cells cm−2 , three-way catalyst formulation with Pt: 1.42 g L−1 , Rh: 0.28 g L−1 ; fresh and aged on an engine bench; engine without a secondary air supply.) (Reproduced with permission from Ref. [39],  1993 Society of Automotive Engineers, Inc.)

11.2.5.3 Aspects of Deactivation of NOx -Adsorber Systems NOx -adsorber systems also lose their activity through both thermal and chemical phenomena. One such thermal deactivation phenomenon is the loss of dispersion of the PGM – the first catalytic function – thus causing a slowing down of the oxidation of NO. This affects the lower end of the temperature window of operation of the NOx -adsorber. Another thermal deactivation phenomenon is the loss of internal surface of the alkali/alkaline earth compounds – the second catalytic function – thus causing a decrease of the capacity to store NO2 . This results in the need for a more frequent regeneration, and thus leads to a decrease in fuel consumption benefit of the lean-burn engine. Both phenomena typically start to show effect from a temperature of 773 K onwards under net oxidizing exhaust gas conditions, and lead to a substantial loss of catalytic function at a temperature of about 1173 K (see Fig. 72). The most important chemical deactivation phenomenon is caused by the similarity in behavior of sulfur oxides and nitrogen oxides. Indeed, using the same catalytic function, the SO2 present in the engine-out exhaust gas will be oxidized to SO3 , which subsequently will References see page 2343

2324

11.2 Automotive Exhaust Treatment

Constant lean burn

Mixed lean burn

NOx

NOx

CO

NOx

CO

CO

HC

0.8

Adsorption type

HC

1.0

1.2

Rich Lean

1.4

1.6

1.8

0.8

HC 1.0

1.2

1.4

1.6

Rich Lean

1.8

0.8

1.0

1.2

1.4

1.6

1.8

Rich Lean

Emission of CO, HC and NOx , at the outlet of a gasoline-fuelled spark-ignition engine, as a function of the engine lambda value. The range of lambda values in which three different catalytic exhaust aftertreatment concepts for lean-burn vehicles operate is also shown. (Adapted from Ref. [60].)

Fig. 70

100

NOx Conversion / %

also be stored on the alkali/alkaline earth oxide compounds along reaction pathways similar to the reactions in Eqs. (44) and (45). This competition for the same adsorption sites reduces the capacity to store NO2 . Unfortunately, the sulfates formed accordingly are thermodynamically more stable than the corresponding nitrates, so that they will not be decomposed during the typical regeneration phases. If no specific measures were to be taken, ultimately a complete blockage of the NO2 adsorption sites and thus a full loss of activity of the NOx -adsorber would occur. To prevent this, specific additional desulfurization events are implemented into the engine management system. These events need a net reducing exhaust gas composition at a temperature typically above 873 K, and should occur with the current system design and boundary operation conditions typically once per tank filling. The optimal frequency of these desulfurization events is guided, amongst others, by the boundary conditions of NOx engine-out emission and fuel sulfur content. Every desulfurization event causes a significant decreases in fuel consumption benefits of this type of engine. One particular challenge that needed to be solved to enable the practical application of this technology already during the EU4 legislation timeframe was to avoid the potential tailpipe emission of H2 S during the regeneration phases of the NOx adsorber. This is realized through an ingenious engine calibration strategy, which imposes an A/F modulation between lean and rich exhaust gas compositions even during the overall net reducing regeneration phases, so that any H2 S formed will be reoxidized to SO2 before it leaves the NOx adsorber [77, 78].

75 50 25 0 0

200

400

600

800

Time / s Time-resolved pattern for the conversion of NOx on a NOx -adsorber system.

Fig. 71

11.2.6

Catalytic Systems for Diesel-Fueled Compression-Ignition Engines Introduction Compression-ignition engines differ from spark-ignition engines in terms of their air and fuel mixing, and ignition of the air/fuel mixture. Furthermore, they use different operating pressures, operating temperatures and A/F ratios (Table 22). A variety of compression engine designs exist, with distinctions being made on the basis of both the fuel and the air admission principles. With regards to fuel admission, the two most important designs are: (i) indirect injection (IDI) engines, where the fuel is injected into a precombustion or a swirl chamber connected to the main cylinder; and (ii) direct injection (DI) engines, where the fuel is injected directly into the 11.2.6.1

11.2.6 Catalytic Systems for Diesel-Fueled Compression-Ignition Engines

cylinder. With regards to air admission, a distinction can be made between naturally aspirated (NA) engines, in which the air is sucked into the engine, and turbocharged (TC) engines, where the air is compressed before it enters the cylinder. The latter is in some cases combined with a cooling of the compressed air (TCI). Most passenger car diesel engines designed for the EU4 emission legislation level belong to the direct-injected, turbo-charged type (DI-TC). These modern diesel engines are in most cases also equipped with highly sophisticated exhaust gas recirculation (EGR) devices, which cause part of the exhaust gas to be returned to the combustion chamber. Clearly, the time-resolved detailed exhaust emission characteristics depend strongly upon the design of the engine. The exhaust gas of diesel engines has a complex composition, as gaseous components are present together with liquid and even with solid components (Table 23). The solid exhaust gas components are denoted particulate matter (PM), and defined as any matter that can be collected on a Teflon-coated filter paper from diluted Tab. 22 Main design and operation differences between compression-ignition engines and spark-ignition engines

Feature

Compression ignition

Spark ignition

Process type

Internal combustion

Combustion type Air-fuel mixing Ignition type Operating pressure Temperature at compression Lambda during operation Exhaust gas oxygen content

Cyclic Heterogeneous Auto 30–55 h Pa 973–1173 K

Internal combustion Cyclic Homogeneous External 15–25 h Pa 673–873 K

099% for a loaded

Pressure-drop / mbar

11.2.6 Catalytic Systems for Diesel-Fueled Compression-Ignition Engines

2337

120 110 100 90 80 70 60 50 40 30 20 10 0 0

1

2

3

4

Soot loading / g L

5

6

7

−1

Fig. 92 Graphical representation of the evolution of the pressure drop of the exhaust gas passing through a wall-flow diesel particulate filter as a function of the mass of trapped diesel particulate matter.

Engine load

• Late post injection • Retarded main injection

• Post injection • Retarded main injection • Decreased boost pressure

• Late post injection • Retarded main injection

No additional measures needed

• Post injection • Retarded main injection • Decreased boost pressure • Decreased air mass

No regeneration possible

• Combustion stabilizing

Engine speed

Schematic representation of the various measures needed to initiate the regeneration of a diesel particulate filter loaded with diesel particulate matter, as a function of the engine operation point in the engine speed- and load map. (Adapted from Ref. [87].)

Fig. 93

filter. However, as the filter cake builds up the backpressure increases, as shown for a typical case in Fig. 92. The maximum allowable back-pressure, combined with a maximally tolerated exothermic during the regeneration phase to avoiding destruction of the filter structure, fixes the duration of this first phase. With the current boundary conditions such as engine-out emission of PM, exhaust gas temperature and filter system design, this phase will last for about 1000 km of real road driving. The second operation phase is regeneration of the loaded filter by combustion of the diesel particulates. This combustion process needs to be initiated, as it occurs only at a gas temperature of at least about 673 K when using oxygen as the oxidative agent. As this temperature is higher than the exhaust gas temperature during typical vehicle operation, measures must to be taken to raise the exhaust temperature on a temporary basis.

Typically, measures on the level of engine operation are combined with catalytic solutions. In terms of engine operation, changes in the EGR rate and fuel injection timing usually are combined with some degree of exhaust gas throttling. In addition, an extra amount of fuel can be injected either into the combustion chamber at the end of the combustion cycle, or directly into the exhaust gas pipe, thus substantially increasing the concentration of HCs (and to some extent also of CO). These reactants can be combusted using a dedicated oxidation catalyst placed upstream of (and ideally close to) the filter. The oxidation catalyst function incorporated inside the filter structure can also be used, either in addition to this upstream catalyst or even stand-alone, replacing the upstream catalyst. Figure 93 shows an example of how References see page 2343

11.2 Automotive Exhaust Treatment

100 NOx conversion/ %

these various measures can be combined in a typical application of catalytic filter systems to passenger cars [87]. In addition to these measures, a dedicated fuel additive based on cerium and/or iron can be used, to serve as a particulate combustion catalyst that is built inside the particulates during the diesel combustion process within the engine’s combustion chamber. This additional catalytic function has the advantage of speeding up the regeneration process substantially, but the disadvantage of leaving some inorganic ash on the filter, and this adds to the pressure drop over the filter. In several real-world applications, provisions were taken to wash out this ash a few times during the vehicle’s lifetime, typically after at least 80 000 km for the early designs, and 250 000 km for more recent designs of such an application. With all of the above-described measures, the duration of the second operation phase is limited typically to a few minutes, after which the filtration cycle can be restarted [81, 82].

75 50 25 0 100

HC conversion/ %

2338

75 50 25 0 125

175

225

275

325

375

425

475

Exhaust gas temperature / °C

11.2.6.4

Diesel NOx -Reduction Systems

Zeolite-based catalysts (16 h, 750 °C, air)

11.2.6.4.1 Introduction Whereas the diesel oxidation catalyst combined with the diesel particulate filter became the accepted technology to limit the emission of CO, gaseous HCs and PM, the major remaining challenge is to limit the NOx emission of diesel engines. One effective approach which found widespread use was to lower the engine-out emission of NOx using EGR systems. This system had limited use, however, as it was found to cause an increase in the engine-out emission of PM, such that it required complementary measures to achieve the very low emission levels featured in advanced gasoline-powered vehicles. Thermodynamically, nitrogen oxide is known to be unstable with respect to N2 and O2 in the exhaust gas composition, and at the exhaust gas temperatures that occur at typical diesel engine operating conditions. Unfortunately, as of today no solid catalyst has been found that can catalyze the direct decomposition of NO in N2 and O2 within the boundary conditions (space velocity, durability) that apply to on-road vehicles. Therefore, much research effort is required to identify alternative methods, of which the HC-DeNOx -, Diesel NOx -adsorber- and SCRsystems have proved to be the most promising [83]. 11.2.6.4.2 HC-DeNOx -Systems HC-DeNOx systems are based upon dedicated catalysts that enable the conversion of NOx with, for example HCs and alcohols according to the reactions in Eqs. (48) and (49):

← − (3m + 1)N2 (6m + 2)NO + 2Cm H2m+2 − −− −− → + 2mCO2 + (2m + 2)H2 O

(48)

Cu Au

Pd Ag

Pt Ir

Rh

Conversion of nitrogen oxides and gaseous hydrocarbons reached over different NOx -reduction catalyst formulations, as a function of exhaust gas temperature (monolith catalyst with 62 cells cm−2 , dedicated NOx -reduction catalyst formulations with zeolites and with different active components; after laboratory oven-aging in air at a temperature of 1023 K for 16 h; light-off test in a model gas reactor at a space velocity of 50 000 N L−1 h−1 ; model gas simulates the exhaust gas composition of an IDI/NA passenger car diesel engine at medium load and speed, except for the hydrocarbon concentration, which was increased to reach yHC /yNOx 3/1 (mol/mol)).

Fig. 94

− 3mN2 6mNO + 2Cm H2m+1 OH ← −−− −− → + 2mCO2 + (2m + 2)H2 O

(49)

As the concentration of HCs and alcohols in the engineout exhaust gas of diesel engines is too low to fulfill the stoichiometry of Eqs. (48) and (49), the systems need provisions to enable adding appropriate quantities of these components to the exhaust gas. Promising results have been obtained with catalysts based upon zeolites, together with elements from Groups Ib and VIII as well as PGMs. Figure 94 shows the HC and NOx conversion obtained in model gas reactor experiments with a series of air-aged catalysts. About 50% NOx conversion is obtained both with a platinum-based catalyst at temperatures below about 570 K, and with a copper-based catalyst at temperatures above about 570 K. These temperatures coincide with the temperature range in which these two catalysts convert

11.2.6 Catalytic Systems for Diesel-Fueled Compression-Ignition Engines

11.2.6.4.3 Diesel NOx -adsorber systems The principle of NOx -adsorber systems (as described for lean-burn gasoline engines in Section 11.2.5) can also be applied to advanced diesel engines. A major challenge to be solved was the periodic generation of net-reducing exhaust gas conditions for short periods of time to enable regeneration of the NOx -adsorber. Modern engine technology enables to do so at engine operation conditions below full load and full speed, without causing excess emission of smoke and PM. In addition, systems have been developed that generate the required exhaust gas composition for the regeneration event by adding fuel or even other reductants such as CO and hydrogen to the exhaust gas after it has left the engine. For the latter, small fuel reformers can be used that generate these reactive reductants on-board and on-demand. The detailed chemistry of the NOx -adsorbers for diesel engines is of course substantially different from that

NOx conversion / %

60 Vehicle A Vehicle B

40 20

Exhaust gas temperature / °C

0 500 Vehicle A Vehicle B

400 300 200 100 0 120 100

Speed/km h−1

HCs. However, whereas the platinum-containing catalyst is able to convert CO efficiently, the copper-containing catalyst generates CO, probably by the incomplete oxidation of the added HCs [71]. Taking all of these characteristics together, an optimal system combines two catalysts in series, with the first catalyst, for example, based on copper, and the second based on PGMs [84, 85]. Platinum-based catalysts to reduce NOx in the exhaust gas of diesel engines with HC addition have also shown promise under the dynamic conditions that occur during vehicle driving. As shown in Fig. 95, the degree of NOx conversion depends primarily on the exhaust gas temperature during the different phases of the vehicle test. In addition to the exhaust gas temperature, the space velocity, the NOx concentration, and the HC concentration and its type have a major impact on the overall NOx conversion This is apparent from the data shown in Fig. 96. [71, 72]. The detailed conversion mechanism is still not very well understood. Initial selectivity studies point to nitrogen as the primary product of the reaction, and both N2 O and NO2 as secondary products (Fig. 97). Together with information gathered from temporal analysis of products (TAP) experiments, and from in-situ catalyst characterization, a dual-site reaction mechanism is proposed, a simplification of which is shown in Fig. 98 [72–74]. With regards to the durability of the HC-DeNOx catalysts, it was found that aging decreases the maximum NOx conversion and shifts the onset of the conversion to higher temperatures. The underlying aging phenomena were found to be similar to those identified for the deactivation of diesel oxidation catalysts.

2339

80 60 40 20 0 40 195

195

195

195

400

1220

Time/s

Conversion of nitrogen oxides reached over a NOx -reduction catalyst mounted on two different passenger cars equipped with an IDI/NA and DI/TC diesel engine in three different phases of the European vehicle test procedure. The exhaust gas temperature in front of the catalyst is also shown as a function of time in the vehicle test procedure (monolith catalyst with 62 cells cm−2 , dedicated NOx -reduction catalyst formulations containing Pt; fresh state; vehicle dynamometer test according to the European vehicle test procedure with passenger cars equipped either with a DI/TC diesel engine with four cylinders and an engine displacement of 2.5 L (vehicle A) or with an IDI/NA diesel engine with four cylinders and an engine displacement of 1.9 L (vehicle B). In both cases, a 2/1 molar mixture of 1-butene/n-butane was added to raise the concentration of gaseous hydrocarbons in the exhaust gas to 800 vol.ppm). (Reproduced with permission from Ref. [71],  1993 Society of Automotive Engineers, Inc.)

Fig. 95

of gasoline engines, due mainly to the generally lower operation temperature of the diesel engine. Finally, the application of NOx -adsorbers to diesel engines will be facilitated to a large degree by the widespread availability of fuels with an ultra-low sulfur content. References see page 2343

2340

11.2 Automotive Exhaust Treatment

70 60 NOx conversion / %

SNO2, SN2O, SN2

225 °C

100

50 40 n-alkanes

30

80

Alkanes Alkenes Alcohols Aromatics

N2O NO2 N2

60 40

20

iso -octane

10

20

0 0

5 10 Number of C atoms per molecule

15

17

0 0

1

2

3

4

5

Inverse space velocity Conversion of nitrogen oxides reached over a NOx -reduction catalyst at 498 K for different classes of hydrocarbons and oxygenates added to the exhaust gas, as a function of the number of carbon atoms in the molecule of the organic component (monolith catalyst with 62 cells cm−2 , dedicated NOx -reduction catalyst formulation containing platinum; in the fresh state; test at a fixed temperature in a model gas reactor with a space velocity of 50 000 N L−1 h−1 ; model gas composition based upon the exhaust gas composition of an IDI/NA passenger car diesel engine at medium load and speed, having 270 vol.ppm NOx and modified by adding various hydrocarbons and oxygenates to a fixed concentration of 3200 vol.ppm C1 ). (Reproduced from Ref. [72] with kind permission of Elsevier Science.)

Fig. 96

11.2.6.4.4 SCR Systems SCR systems represent highperformance technology for eliminating nitrogen oxides from diesel engine exhaust gases, and have been used successfully for over 20 years to clean the exhaust gases of large industrial installations such as power plants (details of the process are provided elsewhere in this Handbook). A major advantage of SCR technology is that the combined high NOx conversion performance and use of external reducing agents places less strain on the compromise between a low-NOx tailpipe emission on the one hand, and maintaining the excellent fuel economy of modern diesel engines on the other hand. Translation of the basic technology as used on power plants towards application on mobile sources such as trucks and passenger cars induced several changes, both in the layout of the system and the detailed chemistry of the catalysts. A typical system (see Fig. 99) consists of up to four different catalytic functions, and a facility to add a reductant to the exhaust gas. The first catalyst to be contacted by the exhaust gas is a dedicated oxidation catalyst, usually based upon platinum, which promotes especially the oxidation of NO into NO2 . This step is taken because it was found that the overall rate of SCR reactions is increased if (typically) up to 50% of the nitrogen oxides are present as NO2 . This higher rate enables good performance despite the low exhaust gas temperature, especially on a passenger car diesel

Selectivity for nitrogen oxides SNO2 , dinitrogen oxide SN2O and nitrogen SN2 reached over a NOx -reduction catalyst as a function of the inverse of the space velocity (monolith catalyst with 62 cells cm−2 , dedicated NOx -reduction catalyst formulation containing platinum; in the fresh state; test at a fixed temperature of 498 K in a model gas reactor with various settings of the space velocity reached by varying the length of the catalyst sample; model gas composition based upon the exhaust gas composition of an IDI/NA passenger car diesel engine at medium load and speed, having 270 vol.ppm NOx and modified by adding n-hexadecane to a fixed concentration of 3200 vol.ppm C1 ; selectivity defined as moles of component formed per 100 mol NOx converted). (Reproduced from Ref. [72] with kind permission of Elsevier Science.) Fig. 97

CxHy + zO2

N2, H2O, CO2 NO Desorption − N2O − NO2

Adsorption CxHyO2z

Interaction

Adsorption + O2 NO Sl

Sll Support

Fig. 98 Proposed reaction mechanism for the reduction of nitrogen oxides in the exhaust gas of diesel engines over a NOx reduction catalyst. (Reproduced from Ref. [72] with kind permission of Elsevier Science.)

engine, and also helps to optimize the size of the catalytic system. Beyond this catalyst, the precursor of the reducing agent is added to the exhaust gas. Urea is a very commonly used precursor, and is carried typically as an aqueous solution in an extra tank in the vehicle. The second catalytic function is the so-called ‘‘hydrolysis catalyst’’, the main task of which is to speed up transformation of the precursor into the desired reducing agent, NH3 ,

11.2.6 Catalytic Systems for Diesel-Fueled Compression-Ignition Engines

according to: ← − 2 NH3 + CO2 (NH2 )2 CO + H2 O − −− −− →

(50)

This catalyst is typically free of PGMs. The exhaust gas is then mixed with the reducing agent and reaches the third catalytic function, which is responsible for the main SCR function according to: − 4N2 + 6H2 O 4NH3 + 4NO + O2 ← −−− −− →

(51)

− 2N2 + 3H2 O 2NH3 + NO + NO2 ← −−− −− →

(52)

− 7N2 + 12H2 O 8NH3 + 6NO2 ← −−− −− →

(53)

Typical catalysts for these reactions are free of PGMs and are based either on TiO2 –V2 O5 or on base metaldoped zeolites. Finally, the fourth catalytic function (which may be optionally incorporated) is another dedicated oxidation catalyst; typically, this is based upon PGMs and removes excess NH3 by oxidation to N2 according to: ← − 2N2 + 6H2 O 4NH3 + 3O2 − −− −− →

(54)

This type of NOx reduction technology has proven to yield typically 80% conversion of the engine-out nitrogen oxides during real-world operation conditions. An example of the conversion performance during testing on a heavy-duty diesel engine is shown in Fig. 100. Within the EU, the serial application of this technology to heavyduty vehicles commenced in 2005. SCR systems may be deactivated by the same mechanisms as diesel oxidation catalysts, and consequently particular attention must be paid to limit the operation

Urea (NH2)2CO

Exhaust gas

V

2341

temperature to about 1000 K in the case of TiO2 –V2 O5 based catalyst to avoid loss of the active component V2 O5 by volatilization. An ultra-low sulfur content of the diesel fuel is a strict necessity for this type of aftertreatment technology, in order to avoid SO3 being generated on the first oxidation catalyst and eventually even on the SCR catalyst. Indeed, SO3 could react with the reductant NH3 to form (NH4 )2 SO4 which would deposit on the catalysts and lead to their deactivation. 11.2.6.4.5 Future Developments and Concepts It is to be expected that continued research efforts will lead to the merging of at least some part of the NOx -elimination and PM filtration functions. Indeed, the first designs of a diesel particulate filter with a built in NOx -adsorber have been proposed as being commercially feasible. It may also be imagined that such a device would be complemented with a downstream SCR function in a subsequent generation of this technology. These technologies will emerge parallel to further developments in engine technology, such as the ‘‘homogeneous charge compression ignition’’ (HCCI) principle, for at least part of the engine operation map. As the latter will lead to even further reductions in the engine-out emission of both NOx and PM, it will facilitate the application of these combined aftertreatment concepts such that diesel engines will have the same low tailpipe emission levels as already achieved by gasoline-powered engines. References see page 2343

SCR catalyst (S) 4NH3 + 4NO + O2 → ← 4N2 + 6H2O 2NH3 + NO + NO2 → ← 2N2 + 3H2O → 8NH3 + 6NO2 ← 7N2 + 12H2O

H optional

S

O

Oxidation catalyst (V) 2NO + O2 → 2NO2

Hydrolysis catalyst (H)

2HC + 3O2 → 2CO2 + 2H2O

(NH2)2 CO + H2O → 2NH3 + CO2

2CO + O2 → 2CO2 Oxidation catalyst (O) 4NH3 + 3O2 → 2N2 + 6H2O

Layout of a modern SCR-catalyst system (VHSO system with V: oxidation catalyst dedicated for NO2 generation; H: hydrolysis catalyst; S: selective reduction catalyst and O: oxidation catalyst dedicated for NH3 removal).

Fig. 99

2342

11.2 Automotive Exhaust Treatment

1400 NOx engine-out emission NOx emission after VHSO system

NOI concentration / vol.-ppm

1200 1000 800 600 400 200 0 0

200

400

600

800 1000 Time / s

1200

1400

1600

1800

Concentration of NOx in the exhaust gas of a heavy-duty diesel engine as a function of time in the European transient engine test procedure (ETC) upstream and downstream of a SCR catalyst system (DI/TC engine with a displacement of 7 L).

Fig. 100

11.2.7

Conclusions and Outlook

The application of catalytic exhaust aftertreatment devices to passenger cars and trucks, both with spark ignition and with compression ignition engines, is an established method of achieving the increasingly demanding legislative emission limits for CO, HCs, NOx and PM. Dedicated research and development efforts made by both the automotive and catalyst industries has led, since the 1970s, to an ongoing introduction of a wide variety of new, better-performing and more cost-effective catalytic solutions to the worldwide market. As a result of this effort, automotive catalysis has today become one of the most innovative fields of catalytic science, and one of the most important applications of solid catalysts. Despite its complexity and harsh operation environment, automotive catalysis has been shown to be a robust technology, which is able to perform very well over a broad range of operation conditions and applications [40, 75]. Finally, a continuous increase in the atomistic understanding of the operation and deactivation phenomena is leading to new developments with the potential for successful application to advanced engine designs. In time, this will contribute substantially to the environmental compatibility and hence the sustainability of on-road transport. Acknowledgments

The author wishes to express his gratitude to all colleagues and coworkers at Umicore, who contributed to this vast amount of high-quality scientific investigations.

The author also thanks Elsevier Science Publishers B.V., Sara Burgerhart straat 25, NL-1055 KV Amsterdam, and the Society of Automotive Engineers, Inc., USA, for their kind permission to reproduce the figures. List of Abbreviations and Symbols Abbreviations

AMA BET

DI DIN DOC EDX EGR EU FID FTP HC HCCI IDI LEV MVEG NA NOx PAH PGM

American Motorists Association Internal surface area of a porous material as defined according to the theory of Brunauer, Emmett and Teller direct injection (for a diesel engine) Deutsche Ingenieurs Norm diesel oxidation catalyst energy dispersive X-ray exhaust gas recirculation European Union flame ionization detector Federal Test Procedure all type of hydrocarbon homogeneous charge compression ignition indirect diesel injection low-emission vehicle Motor Vehicle Emission Group naturally aspirated (of a diesel engine) all type of nitrogen oxides, mainly NO and NO2 polynuclear aromatic hydrocarbons platinum-group metals

References

PIXE PM SAE SCR SIMS SOF SOx SULEV TAP TC TCI TLEV TWC ULEV XPS XRF ZEV

proton-induced X-ray emission particulate matter Society of Automotive Engineers selective catalytic reduction secondary ion mass spectrometry soluble organic fraction (of diesel particulates) all type of sulfur oxides, mainly SO2 and SO3 super ultra-low-emission vehicle temporal analysis of products turbocharged turbocharged intercooled transitional low-emission vehicle three-way catalyst ultra-low-emission vehicle X-ray photoelectron spectrometry X-ray fluorescence zero-emission vehicle

m vol.%

A/F kg kg−1 b m D vol.%
p dk Dk F0

bar m m mol s−1

Gk G Hz le Lk m

Nm3 h−1 Nm3 h−1 s−1 m m

n

Nk Nu p r

Re

bar m m

m2

Sg

m2 m−3

Si Sk Sk0 T tk

mol mol−1 m2 m2 K h

Vk vk W W wk xS y

m3 m h−1 kg kg m mg L−1 vol.-ppm or vol.% m−2

z

geometric surface area per monolith geometric surface area per unit monolith volume selectivity for a component i total frontal area of a monolith open frontal area temperature residence time of a gas molecule in a monolith channel geometric monolith volume linear gas velocity per channel mass of a bead mass of catalyst as used in Fig. 34 monolith wall thickness mass of sulfur per volume of fuel volumetric concentration of a component in the gas mixture cell density of a monolithic substrate

References

Symbols

a A

Sg0

2343

major axis of a bead ellipsoid volumetric concentration as defined by Eq. (9) air-to-fuel ratio minor axis of a bead ellipsoid volumetric concentration as defined by Eq. (7) pressure drop monolith channel width diameter of a monolith molar flow of reactants as used in Fig. 34 gas flow per channel total exhaust gas flow frequency in Hertz length of the flow transition zone length of a monolith number of carbon atoms in an hydrocarbon molecule number of hydrogen atoms in an hydrocarbon molecule number of attrition runs according to Eq. (24) number of channels per monolith Nusselt number partial pressure; also total pressure radius of a bead catalyst impregnation depth in a bead catalyst Reynolds number

1. G. A. Ahrens, Verkehrsbedingte Luft- und Laermbelastungen, Umweltbundesamt, Berlin, 1991. 2. J. J. MacKenzie, M. P. Walsh, Driving Forces, World Resources Institute, Washington, 1990. 3. A. Lowi, W. P. L. Carter, Technical Paper Series 900710, Society of Automotive Engineers, Warrendale, 1990. 4. Association des Constructeurs Europeens d’Automobiles (ACEA), Future Exhaust Emission Standards of Passenger Cars, AE/71/91/VE/ACEA/2, Brussels, 1991. 5. D. M. Heaton, R. C. Rijkeboer, P. van Sloten, Proceedings, Symposium on Traffic-Induced Air Pollution, AVL List GmbH, Graz, 1992. 6. J. S. McArragher, Concawe Report 93/51, CONCAWE, Brussels, 1993. 7. J. T. Kummer, Prog. Energy Combust. Sci. 1979, 6, 177. 8. J. E. McEvoy, Catalysts for the Control of Automotive Pollutants, American Chemical Society, Washington, 1975. 9. K. C. Taylor, in Catalysis and Automotive Pollution Control I, A. Crucq, A. Frennet (Eds.), Studies in Surface Science and Catalysis, Elsevier, Amsterdam, 1987, p. 97 10. F. Schaefer, R. van Basshuysen, Die Verbrennungskraftmaschine, H. List, A. Pischinger (Eds.), Springer-Verlag, Wien, 1993, p. 7. 11. H. J. Foerster, Automobiltechnische Zeitschrift 1991, 93, 342. 12. D. Hundertmark, Automobil-Industrie, 1990. 13. E. S. Lox, B. H. Engler, E. Koberstein, in Catalysis and Automotive Pollution Control II, A. Crucq, A. Frennet (Eds.), Studies in Surface Science and Catalysis, Elsevier, Amsterdam, 1991, p. 291. 14. J. Brettschneider, Bosch Technische Berichte 1979, 6, 177 15. A. J. L. Nievergeld, Institute for Continuing Education, Eindhoven University of Technology, Eindhoven, 1994. 16. S. H. Oh, J. C. Cavendish, L. L. Hegedus, Chem. Eng. Prog. 1980, 26, 935. 17. S. H. Oh, J. C. Cavendish, AIChE J. 1985, 31, 935. 18. A. Renken, Int. Chem. Eng. 1994, 33, 61.

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11.2 Automotive Exhaust Treatment

19. C. N. Montreuil, S. C. Williams, A. A. Adamczyk, Technical Paper Series 920096, Society of Automotive Engineers, Warrendale, 1992. 20. L. L. Hegedus, J. J. Gumbleton, Chemtech 1980, 10, 630. 21. P. Nortier, M. Soustelle, in Catalysis and Automotive Pollution Control I, A. Crucq, A. Frennet (Eds.), Studies in Surface Science and Catalysis, Elsevier, Amsterdam, 1987, p. 275. 22. J. C. Summers, L. Hegedus, US Patent 4,152, 301, assigned to General Motors, 1979. 23. G. Kim, M. V. Ernest, S. R. Montgomery, Indust. Eng. Chem. Prod. Res. Dev. 1984, 23, 525. 24. J. S. Howitt, in Catalysis and Automotive Pollution Control I, A. Crucq, A. Frennet (Eds.), Studies in Surface Science and Catalysis, Elsevier, Amsterdam, 1987, p. 301. 25. J. P. Day, L. S. Socha, Technical Paper Series 881590, Society of Automotive Engineers, Warrendale, 1988. 26. J. Yamamoto, K. Kato, J. Kitagawa, M. Machida, Technical Paper Series 910611 Society of Automotive Engineers, Warrendale, 1991. 27. W. Maus, H. Bode, A. Reck, Technische Akademie Wuppertal, Wuppertal, 1989. 28. M. Nonnenmann, Automobiltechnische Zeitschrift 1989, 91, 4. 29. W. B. Kolb, A. A. Papadimitriou, R. L. Cerro, D. D. Leavitt, J. C. Summers, Chem. Eng. Prog. 1993, 2, 61. 30. B. H. Engler, E. S. Lox, K. Ostgathe, T. Ohata, K. Tsuchitani, S. Ichihara, H. Onoda, G. T. Garr, D. Psaras, Technical Paper Series 940928, Society of Automotive Engineers, Warrendale, 1994. 31. M. C. F. Steel, in Catalysis and Automotive Pollution Control II, A. Crucq, A. Frennet (Eds.), Studies in Surface Science and Catalysis, Elsevier, Amsterdam, 1991, p. 105. 32. C. Hagel¨ucken, Proceedings Second European Precious Metals Conference, IPMI, London, 1995. 33. W. Kuemmerle, R. Duesmann, Motortechnische Zeitschrift, 1994, 55, 708 34. E. S. Lox, B. H. Engler, E. Koberstein, Technical Paper Series 910841, Society of Automotive Engineers, Warrendale, 1991. 35. A. Cybulski, J. A. Moulijn, Chem. Eng. Sci. 1994, 49, 19. 36. E. Koberstein, G. Wannemacher, in Catalysis and Automotive Pollution Control II, A. Crucq, A. Frennet (Eds.), Studies in Surface Science and Catalysis, Elsevier, Amsterdam, 1991, p. 437. 37. D. Schweich, in Catalysis and Automotive Pollution Control II, A. Crucq, A. Frennet (Eds.), Studies in Surface Science and Catalysis, Elsevier, Amsterdam, 1991, p. 421. 38. J. P. Leclerc, D. Schweich, J. Villermaux, in Catalysis and Automotive Pollution Control II, A. Crucq, A. Frennet (Eds.), Studies in Surface Science and Catalysis, Elsevier, Amsterdam, 1991, p. 465. 39. B. H. Engler, D. Lindner, E. S. Lox, K. Ostgathe, A. Sch¨aferSindlinger, W. Muller, Technical Paper Series 930738, Society of Automotive Engineers, Warrendale, 1993. 40. E. Schwizer, D. Burch, P. Riedwyl, 10 Jahre Katalysator-Autos, Touring Club der Schweiz, Emmen, 1994. 41. D. W. Wendland, W. R. Matthes, Technical Paper Series 861554, Society of Automotive Engineers, Warrendale, 1986. 42. D. W. Wendland, P. L. Sovrell, J. E. Kreucher, Technical Paper Series 912372, Society of Automotive Engineers, Warrendale, 1991. 43. B. H. Engler, E. Koberstein, P. Schubert, Appl. Catal. 1989, 48, 71. 44. E. Koberstein, Dechema Monographien, 120, VCH, Weinheim, 1990, p. 67.

45. P. Albers, B. H. Engler, J. Leyrer, E. S. Lox, G. Prescher, K. Seibold, Chem. Eng. Tech. 1994, 17, 161. 46. B. H. Engler, E. Koberstein, D. Lindner, E. S. Lox, in Catalysis and Automotive Pollution Control II, A. Crucq, A. Frennet (Eds.), Studies in Surface Science and Catalysis, Elsevier, Amsterdam, 1991, p. 641. 47. B. H. Engler, D. Lindner, E. S. Lox, P. Albers, Proceedings IV th International Congress on Catalysis, L. Guczi, F. Solymosi, P. Tetenyi (Eds.), Akademiai Kiado, Budapest, 1993, p. 2701. 48. H. H. Haskew, J. J. Gumbleton, Technical Paper Series 881682, Society of Automotive Engineers, Warrendale, 1988. 49. E. S. Lox, B. H. Engler, E. Koberstein, Technical Paper Series 881682, Society of Automotive Engineers, Warrendale, 1988. 50. H. Skinjoh, H. Muraki, Y. Fujitani, in Catalysis and Automotive Pollution Control II, A. Crucq, A. Frennet (Eds.), Studies in Surface Science and Catalysis, Elsevier, Amsterdam, 1991, p. 617. 51. G. Mabilon, D. Durand, M. Prigent, in Catalysis and Automotive Pollution Control II, A. Crucq, A. Frennet (Eds.), Studies in Surface Science and Catalysis, Elsevier, Amsterdam, 1991, p. 569. 52. Forschungsvereinigung Verbrennungskraftmaschinen (FVV), Report 419, Frankfurt, 1988. 53. M. Shelef, K. Otto, N. C. Otto, Adv. Catal. 1978, 27, 311. 54. W. B. Williamson, H. K. Stephien, W. L. H. Watheins, H. S. Gandhi, Environ. Sci. Technol. 1979, 13, 1109. 55. K. C. Taylor, Automobile Catalytic Converters in Catalysis, Science and Technology, J. R. Anderson, M. Boudart (Eds.), Springer-Verlag, Berlin, 1984, p. 5. 56. K. Kollmann, J. Abthoff, W. Zahn, Automotive Eng. 1994, 10, 17. 57. M. Theissen, P. Langer, J. Mallog, R. Zielinski, Proceedings 4. Aachener Kolloqium Fahrzeug- und Motorentechnik, Forschungsgesellschaft Kraftfahrwesen Aachen mbH (FKA), Aachen, 1993. 58. E. Otto, W. Held, A. Donnerstag, P. Kuiper, B. Pfalzgraf, A. Wirth, Proceedings 16. Internationales Wiener Motorensymposium, H. P. Lenz (Ed.), VDI- Fortschrittberichte, Reihe 12,205, OEVK, Wien, 1995. 59. P. Hofmann, F. Indra, Proceedings 16. Internationales Wiener Motorensymposium, H. P. Lenz (Ed.), VDI- Fortschrittberichte, Reihe 12,205, OEVK, Wien, 1995. 60. T. Kreuzer, E. S. Lox, D. Lindner, J. Leyrer, in Proceedings Second Japan - EC Joint Workshop on the Frontiers of Catalytic Science and Technology for Energy, Environment and Risks Prevention, published in Catal. Today 1996, 29, 17. 61. M. Miyoshi, T. Tanizawa, S. Takeshima, N. Takahashi, K. Kasahara, Toyota Technical Review, Aichi 1995, 40, 21. 62. W. Held, A. K¨onig, T. Richter, L. Puppe, Technical Paper Series 900496, Society of Automotive Engineers, Warrendale, 1990. 63. J. R. Needham, D. M. Doyle, S. A. Faulkner, H. D. Freeman, Technical Paper Series 891949, Society of Automotive Engineers, Warrendale, 1989. 64. L. Burgler, P. L. Herzog, P. Zelenka, Proceedings Worldwide Engine Emission Standards and How to Meet Them, Institution of Mechanical Engineers, London, 1991, p. 57. 65. J. Leyrer, E. S. Lox, B. H. Engler, Proceedings Symposium Diesel engines, Technische Akademie Esslingen, U. Esser, K. H. Prescher (Eds.), Esslingen, 1993, p. 55. 66. M. Horiuchi, K. Saito, S. Ichihara, Technical Paper Series 900600, Society of Automotive Engineers, Warrendale, 1990.

11.3.1 Emissions from Stationary Sources 67. U. Graf, P. Zelenka, Proceedings 15. Internationales Wiener Motorensymposium, H. P. Lenz (Ed.), VDI- Fortschrittberichte, Reihe 12,239, OEVK, Wien, 1994. 68. B. H. Engler, E. S. Lox, K. Ostgathe, W. Cartellieri, P. Zelenka, Technical Paper Series 910607, Society of Automotive Engineers, Warrendale, 1991. 69. R. Beckmann, W. Engeler, E. Muller, B. H. Engler, J. Leyrer, E. S. Lox, K. Ostgathe, Technical Paper Series 922330, Society of Automotive Engineers, Warrendale, 1992. 70. P. Zelenka, K. Ostgathe, E. S. Lox, Technical Paper Series 902111, Society of Automotive Engineers, Warrendale, 1990. 71. B. H. Engler, J. Leyrer, E. S. Lox, K. Ostgathe, Technical Paper Series 930735, Society of Automotive Engineers, Warrendale, 1993. 72. B. H. Engler, J. Leyrer, E. S. Lox, K. Ostgathe, in Catalysis and Automotive Pollution Control III, A. Frennet, J. M. Bastin (Eds.), Studies in Surface Science and Catalysis, Elsevier, Amsterdam, 1995, p. 529. 73. J. Leyrer, E. S. Lox, Proceedings Dechema Jahrestagung, Dechema e.V., Frankfurt, 1995. 74. Ricardo Consulting Engineers, Automotive Diesel Engines and the Future, Ricardo, Shoreham-by-Sea, England, 1994. 75. R. C. Rijkeboer, M. F. van den Haagen, TNO-Report 730210064, TNO Industrial Research, Delft, Netherlands, 1992. 76. B. H. Engler, E. Koberstein, E. S. Lox, Technical Paper Series 900271, Society of Automotive Engineers, Warrendale, 1990. 77. K. H. Gl¨uck, U. G¨obel, H. Hahn, J. H¨ohne, R. Krebs, T. Kreuzer, E. Pott, Motortechnische Zeitschrift, 2000. 78. C. Enderle, C. Sch¨on, T. Ried, W. M¨uller, L. Ruwisch, M. K¨ogel, S. Franoschek, T. Kreuzer, E. S. Lox, Technical Paper Series 2003-01-11, Society of Automotive Engineers, Warrendale, 2003. 79. P. Tancell, G. Sivalingham, P. Retman, P. Spurk, F. W. Sch¨utze, G. Jeske, M. Moarder, S. Franoschek, Proceedings 14. Aachener Kolloquium Fahrzeug- und Motorentechnik, Forschungsgesellschaft Kraftfahrwesen Aachen mbH (FKA), Aachen, 2005. 80. A. Morlang, U. Neuhausen, K. V. Klementiev, F. W. Sch¨utze, G. Miehe, H. Fuess, E. S. Lox, Appl. Catal. B: Environmental 2004, 60, 191. 81. M. Pfeifer, P. C. Spurk, N. S¨oger, M. K¨ogel, E. S. Lox, Proceedings 2. FAD- Konferenz, Foerderkreis Abgasnachbehandlungstechnologien fuer Dieselmotoren e.V., Dresden, 2004. 82. M. Pfeifer, M. Votsmeier, M. K¨ogel, P. C. Spurk, E. S. Lox, J. F. Knoth, Technical Paper Series 2005-01-1756, Society of Automotive Engineers, Warrendale, 2005. 83. P. C. Spurk, E. S. Lox, 3. Int. Forum Abgas- und Partikelemissionen, AVL List GmbH, Graz, 2004. 84. S. Eckhoff, D. Hesse, J. van der Tillaart, J. Leyrer, E. S. Lox, in Catalysis and Automotive Pollution Control IV, N. Kruse, A. Frennet, J. M. Bastin (Eds.), Studies in Surface Science and Catalysis, Vol. 116, Elsevier, Amsterdam, 1997. 85. H. Klein, S. Lopp, E. S. Lox, M. Kawanami, M. Horiuchi, Technical Paper Series 992134, Society of Automotive Engineers, Warrendale, 1999. 86. J. Gieshoff, M. Pfeifer, A. Sch¨afer-Sindlinger, P. C. Spurk, Technical Paper Series 2001-01-0514, Society of Automotive Engineers, Warrendale, 2001. 87. B. Kampmann, III, Advanced Diesel Engine Technology Symposium, South Korea, 2003.

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11.3

Flue Gases from Stationary Sources .. Par Gabrielsson∗ and Henrik Guldberg Pedersen

11.3.1

Emissions from Stationary Sources

The increasing level of human industrial activity continuously requires improved environmental technologies to prevent uncontrolled pollution of our common environment, both local and global. Catalytic flue gas cleaning plays a major role worldwide for ensuring a sustainable environment. Pollutants are emitted continuously from both mobile and stationary sources worldwide. The stationary sources for gaseous emissions range from private homes over small industries and waste incineration facilities to enormous coal-fired power plants. All of these have no, little or extensive use of flue gas cleaning technologies, whether non-catalytic or catalytic [1–10]. Fossil fuels, biomass and waste are gasified or burned in power plants, process heaters or private homes for energy generation. The trends of making industry – and in particular power plants – more fuel flexible also pushes forward the development of new flue gas cleaning technologies. In addition, the increased operating efficiency will generally result in lower emission rates. Emissions of NOx , SOx , and NH3 cause acidification of the environment and smog formation in the atmosphere. A review of NOx emissions from coal and the effects is provided by Sloss [9]. Whilst greenhouse gases such as CO2 , N2 O and CH4 give rise to global warming, other gaseous emissions such as dioxins, volatile organic compounds (VOCs), H2 S, and heavy metals are directly harmful to humans. The types and typical concentrations of gaseous emissions from a number of stationary sources are listed in Table 1. Emissions may be expressed per amount of energy input (typically lower heating value of the fuel), per amount fuel consumed or as a concentration, normally under certain reference conditions which are not equal to the actual operating conditions. Correcting for oxygen and adding or subtracting water to give the desired reference concentration of oxygen and dry conditions must be carried out to compare emission concentrations. For example, the correction from a dry reference oxygen content to the actual conditions for NOx is: References see page 2382 ∗ Corresponding author.

2346

11.3 Flue Gases from Stationary Sources

Actual% O2 1 − XH2 O NOx (wet, actual% O2 ) = · 20.9 − Reference% O2 , dry (1 − XH2 O ) · NOx (dry, ref% O2 ), where XH2 O is the mole fraction of water. The NOx content is often given in equivalent mg NO2 per m3 . Conversion to ppm is:

250

20.9 −

Patents filed yearly

1 mol · NO (mg NO per m3 ) NOx (ppm) = x 2 g 46.01 mol 22.41

200

150

100

50

11.3.2 0 1950

Drivers for Flue Gas Cleaning Technology

Legislation is the one key driver for the development and implementation of flue gas cleaning technology. Yeh et al. [10] used patent data to show that, in the cases of Japan, Germany and the United States, innovations in NOx control technologies did not occur until stringent government regulations were in place (see Fig. 1). NOx emission regulations in Japan were not established until the mid-1970s, and selective catalytic reduction (SCR) installations started later during that decade. In Germany, legislation for large furnaces imposing control requirements on SO2 and NOx emissions was passed during the early 1980s, and this was followed by the installation of SCR on most boilers of more than 300 MW output. In the United States, SCR installations were introduced in the 1990s on new coal-fired boilers, and grew rapidly in numbers following the imposition of the NOx State Implementation Plans (SIPs) call in 1998. The USA The 1990 Clean Air Act was adopted in order to obtain the same basic air quality all over the USA. The Clean Air Act from 1990 was written on the basis of the Clean Air Act from 1977, but now with five new amendments. The law is federal, and the individual states carry out most of the enforcement. The individual states are also allowed to have stricter laws than the federal law, but not less stringent. The American states have developed SIPs which describe how each state plans to comply with the federal regulation. Each SIP is approved by the Environmental Protection Agency (EPA), which can overtake the enforcement from the state if the plan is insufficient. The EPA has developed criteria for two types of pollutant: (i) the primary standard for pollutants which are harmful for humans; and (ii) the secondary standard for pollutants that can harm the environment and properties. The standard is looked upon in geographical areas, where an area which meets the primary standards is referred to as an ‘‘attainment area’’, and the areas which do not fulfill the standards are called ‘‘non-attainment areas’’.

11.3.2.1

1960

1980

1970

1990

2000

Year

Numbers of NOx -related patents filed by year. (Adapted from Ref. [10].)

Fig. 1

The Clean Air Act contains a program where the power plants or other polluting sources require a permit to emit polluting emissions into the atmosphere. The permit includes what and how much the industries are allowed to emit. As an example, a power plant can be omitted by the acid rain reduction program, hazardous air pollution program and ozone reduction program, all within the Clean Air Act programs. The permits should cover all parts of the plant, and a New Source Review (NSR) is carried out every time a new plant is about to be erected. In 2003, an additional act was introduced called ‘‘Clear Skies’’; this was intended to be a cap-and-trade type of legislation, which focused on SO2 , NOx and mercury emissions. The intention was that these emissions should be reduced by 73, 67, and 67%, respectively, until 2018. Unfortunately, this act was not implemented in 2006, as it met some resistance both in political and legal instances, and hence its future legal status is currently unclear. Europe The European Commission has highlighted the following drivers for improving the air quality [11]: 11.3.2.2

• Human health: the air pollution affects the respiratory system • Acidification: causes damage to nature, ecosystems, buildings and historical sites • Eutrophication: higher levels of nitrogen nutrients • Material damage: buildings, including historical sites, are damaged by acidification and particulates. The European Commission furthermore initiated and developed the Clean Air for Europe (CAFE) program, the

11.3.2 Drivers for Flue Gas Cleaning Technology

Tab. 1

Typical ranges of gaseous emissions from stationary sources without flue gas cleaning

Industry

Source

Heat and power generation

Boiler

Fuel type Coal

Oil/Pet coke

Biofuel Gas turbine

Gas

Diesel engine

Oil

Gas emissions

Levels/ppmva

NOx SOx Particlesb NOx SOx Particlesb NOx SOx NOx CO NOx SO2 CO Hydrocarbons

150–700 300–2500 10 mg Nm−3 200–500 1000–5000 10 mg Nm−3 100–300 0–50 15–50 1–200 1000–1500 100–2000 100–1000 50–500 ppm C1

Incineration

Municipal waste, sewage sludge

NOx SO2 CO Dioxins Particlesb

150–300 10–100 5–20 1–10 ng Nm−3

Process Industry

Nitric acid plants

100–2000 100–1000 Trace Trace 100–3000 10–100 20 000–100 000 10–100 50 15–20

Gas 50–100 3–5 ≈0 12–15

as NO2 and SO2 .

11.3.3.7.1 Laboratory- and Bench-Scale Testing The widely accepted standard VGB scheme for SCR catalyst testing [81] defines a microreactor test as comprising at least four channels. For the purposes of catalyst development and intrinsic kinetic studies this is not always feasible; thus, in this context the term ‘‘microreactor’’ is used when testing catalyst samples weighing between a few milligrams and several grams. A ‘‘generic’’ layout for laboratory and bench SCR catalyst activity testing is shown in Fig. 25. Flows are controlled by mass flow controllers, water is added in a preheated gas stream, and the mixed gas is fed to the reactor placed inside a correctly designed furnace to ensure isothermal conditions. Preferably, a thermocouple is placed inside the

catalyst bed. The furnace should be able to heat to about 700 ◦ C to simulate the highest temperature applications for SCR. Typical gas compositions are shown in Table 6. Evaluation of the catalyst activity is normally carried out under standard test conditions, for example as described in the VGB guidelines [81]. The SCR activity is most commonly expressed in terms of a first-order rate constant (in Nm3 m−2 h−1 ): NH3 : NOx < 1 : kNH3 = −NHAV · ln(1 − XNH3 ) and NH3 : NOx > 1 : kNOx = −NHAV · ln(1 − XNOx ) where NHAV is the area space velocity given by the volumetric flow rate divided by the geometric area; that

2369

11.3.3 NOx Removal Technologies

Full-Scale Catalyst Testing For the owner of the plant in which the SCR catalyst is installed, the actual emission compared to the guaranteed emissions and the legally permitted emissions is of great importance. Therefore, testing of the catalyst performance at the site is required. Optimal performance of full-scale SCR units requires optimal distribution of the flow rates as well as the local NH3 to NOx molar ratios [86]. Usually, the deviation is specified to be below 5% of the mean. During fine-tuning of the ammonia injection grid at the start-up of full-scale units, tests of NOx , NH3 and flow distribution are used for feedback to the grid adjustment. This requires traversal of the flue gas duct with long probes. Multipoint sampling and analysis may reduce the time required [87]. Usually, SCR catalyst testing at the laboratory or bench scale is conducted at a NH3 : NO molar ratio of 1, or with excess NH3 . In practical use, the SCR reactors are designed for NH3 conversion at NH3 : NO molar ratios less than unity. Furthermore, due to the high levels of NOx conversion over SCR catalysts, low concentrations 11.3.3.7.4

0.8

500 450

0.7

400 0.6 350 0.5

300

0.4

250 200

0.3

NH3 slip / ppm

11.3.3.7.2 Pilot-Scale Testing Certain tests are best performed under realistic conditions with several catalyst beds, as in the industrial reactor. Figure 27 illustrates schematically the process diagram of a pilot-scale SCR test unit, together with a photograph of a set-up used for high-dust tests with customer fly ash. The pilot scale will normally involve real industrial elements and a cross-section of at least 15 × 15 cm2 . Pilot-scale testing offers greater flexibility than that of slip stream reactors (see below).

in high-dust installations; a minimum maintenance is normally desired of such unmanned pilot reactors. Care must be taken to conclude from tests where significant catalyst clogging is observed. Slip-stream reactors may even be used for several catalyst elements at the same time [84]. Recently, an in-situ activity test for continuous evaluation of the NO removal potential of the installed catalyst was presented [85].

NOx conversion

is, the wetted catalyst channel perimeter multiplied by the catalyst length. Xi is NH3 or NOx conversion; therefore, it is important to know the geometric area of the catalyst tested. At NH3 : NOx below 1, the NOx conversion increases linearly with increasing ratio and the rate depends on the concentration of NH3 . The conversion and activity at stoichiometric NH3 : NOx are determined by interpolation [81]. Figure 26 shows, graphically, the typical effect of increasing the NH3 : NO ratio on NO conversion and NH3 slip. The activity relative to that of the fresh catalyst k/k0 is normally plotted against the operating time to evaluate the actual deactivation relative to the projected one. Both micro-scale and bench-scale reactors are used for tests of deactivated catalysts. If sections are cut from elements extracted from the full-scale reactor, it is important to note that the distribution of accumulated poisons is not even throughout the length of the element. The rate constant of samples from the inlet section will be lower than that of the outlet section [2]. For testing the conversion of SO2 or Hg0 over SCR catalysts it is very important that the reaction in the reactor set-up (e.g., on hot steel surfaces) and condensation or adsorption in the sample lines are avoided. Glass-lined pipes and heated sample lines are commonly used to overcome these issues. The VGB guideline specifies a 72-h period of stabilization prior to testing the SO2 conversion rate; this is necessary due to the ability of the titania support to bind sulfate.

150

Slip-Stream Reactors The tests described above require that a catalyst sample or whole element is retracted from the full-scale reactor in order to evaluate the activity. Whilst small sample coupons may be removed during the operation, the retrieval of whole elements requires a complete shut down. Alternatively, slip-stream reactors may be installed in the flue gas duct. This enables a fraction of the flue gas to be taken out, passed through a bed of test SCR catalyst elements, and led back into the duct [83]. Normally, the flue gas is pumped using a pressurized air ejector. Care should be taken to avoid clogging due to the deposition of fly ash when used

11.3.3.7.3

0.2 100 0.1

50

0 0

0.5 1 NH3:NOx feed ratio

0 1.5

NO conversion and NH3 slip plotted against the NH3 : NO ratio for a deactivated catalyst sample at 200 ◦ C (
) and 300 ◦ C (•). (Adapted from Ref. [2].)

Fig. 26

References see page 2382

2370

11.3 Flue Gases from Stationary Sources

To analyzers

Mixer To analyzers

Catalyst layer

NH3

CH2

Fly NH3 CO HCl SO2 ash Cooler

To analyzers

Boiler Stack

(a)

(b)

(a) Schematic process diagram of pilot-scale SCR test set-up (Adapted from Ref. [2].) (b) An SCR pilot test for the study of catalyst clogging and erosion by fly ash at high temperature and in the presence of water and SO2 /SO3 . The dimensions are 15 × 15 cm2 in cross-section, with an option of three catalyst layers (50 cm each).

Fig. 27

of the gas compounds must be detected. It is not always possible to control unit load (e.g., temperature and flow) and NOx conversion by adjusting NH3 addition. Correct gas sampling is carried out using an isokinetic nozzle which includes a filter to collect a representative ash sample. An analyzer, a liquid or solid sorption system is used to collect a particulate-free gas sample after sample conditioning, generally according to the EPA’s methods [88–90]. Inertial separation probes make it possible to sample from a high-dust gas stream without the use of a filter [91]. If gas compounds are able to react with or adsorb onto the fly ash surface, the full-scale measurements in high-dust flue gas ducts can be quite complicated [91–93]. This is true in particular for batch sampling methods, where several hours of sampling are required. This was found for Hg sampling [92] as well as for SO3 sampling in high-dust units using the controlled condensation method [93]. For example, Nielsen showed that the analyzed SO3 concentration decreased as the sampled gas volume increased due to adsorption of SO3 onto the fly ash [93].

is based on the oxidation of NO to NO2 with ozone, and requires that NH3 and H2 O are removed from the gas stream, for example by cooling and condensation, followed by scrubbing with oxalic acid or phosphoric acid [94]. Gas chromatography and mass spectroscopy are also used. Many components are analyzed with infrared (IR) spectroscopy, for example CO2 , NH3 , N2 O and SO2 , often combined with accurate wet chemical methods of calibration. In this context it is important to be aware of interference by certain gas components on the analysis in question. For example, SO2 interferes with the wet chemical analysis of NH3 [94]. New techniques such as tunable laser spectroscopy may be used for O2 and NH3 analysis. SO3 must be sampled by controlled condensation of the aerosols using heated glass or Teflon-lined sample lines [95, 96]; a typical set-up is shown schematically in Fig. 28. Sampling from high-dust gas ducts requires removal of the fly ash by filtration, with or without the combination of inertial separation by such cyclones or inertial separation probes [91]. 11.3.4

Gas Sampling and Analysis The parameters normally controlled and measured include NO, NO2 , N2 O, NH3 , SO2 , SO3 , CO2 , H2 O, O2 , temperature, and gas flow rate. There is a rapid and continuous development of on-line analysis methods in particular. Traditionally, chemiluminescence has been used for the analysis of NO and NO2 over a wide concentration range. The method 11.3.3.7.5

Other Emissions 11.3.4.1

Sulfur Removal

11.3.4.1.1 The Wet Sulfuric Acid (WSA) Process and Combined WSA and DeNOx (SNOX) The Topsoe WSA process recovers up to 99.5% of the sulfur in sulfurous

11.3.4 Other Emissions

Glass ‘‘mouse’’, packed with quartz wool

Flue gas

Heated SO3 sampling probe (Quartz tube, 180–200 °C)

2371

Heated sample line (Teflon-coated, 180 °C)

Midget impingers 30 mL 3% H2O2

Water 70–80 °C

Pump dry gas meter thermocouple

SO3 condenser coil G4 glass fritt

Sample collecting flask

Silica gel dryer

The set-up for sampling of SO3 by controlled condensation and SO2 by absorption and oxidation in H2 O2 . The system is developed from the ASTM D-3226-73T standard [93].

Fig. 28

off-gases as concentrated sulfuric acid from gases with up to 30% H2 O and up to 6% SOx , with no lower limit with regard to the SOx concentration and without drying of the gas. The sulfur compounds in the feed gas are first oxidized to SO3 which is then hydrated in gas phase, condensed and recovered as 93–98.6% concentrated sulfuric acid by cooling of the gas in air-cooled glass tubes. The WSA process is applied in various versions, depending on the application and the temperature and composition of the feed gas to be treated. The first WSA plant began operation in 1986 on a molybdenum roasting plant in Sweden, and several plants have been installed for diverse applications. These comprise the treatment of off-gases with 0.5–6% SO2 from mineral roasting, or off-gases with H2 S and other combustibles from Claus plants, viscose plants, purification of natural gas, gasification processes and hydrodesulfurization (HDS) processes in refineries [97–99]. WSA/SAR plants for the recovery of concentrated acid from the incineration of spent sulfuric acid, and WSA/SNOX plants for purification of flue gas from boilers and power plants are other important applications of the WSA process [100]. The main steps of the WSA process are as follows. Upstream of the WSA plant the gas is conditioned to remove other acidic components such as HCl, HF, H2 S, COS, CS2 , organic sulfur compounds, arsenic, and dust. This is typically carried out in a bag filter using alkaline sorbent and by catalytic or thermal oxidation. The gas is preheated to about 400 ◦ C prior to the catalytic oxidation

of SO2 to SO3 on a suitable vanadium catalyst: − SO3 , SO2 + 12 O2 ← −−− −− → ◦

−
H0 = +3.09 MJ kg−1 sulfur at 25 C. The gas is subsequently cooled to 240–295 ◦ C, or minimum 16 ◦ C above the H2 SO4 dew point of the gas in a boiler, whereby most of the SO3 is hydrated to H2 SO4 vapor by: ← − H2 SO4 (vap), SO3 + H2 O − −− −− →
H0 = +3.05 MJ kg−1 sulfur. Finally, the gas is cooled to about 100 ◦ C in vertical, air-cooled glass tubes of the ‘‘WSA condenser’’ where remaining SO3 is hydrated and the H2 SO4 vapor condensed as concentrated sulfuric acid that is drained from the bottom of the WSA condenser at a temperature close to the H2 SO4 dew point of the gas: ◦

H2 SO4 (vap) → 96%H2 SO4 (200 C liq), −
H0 = +1.8 MJ kg−1 sulfur. More than 90% of the heat of formation of sulfuric acid from SO2 is utilized for air preheating and steam production [101]. The formation of acid mist is suppressed by heterogeneous nucleation control, and by special internals installed in the glass tubes, as shown in Fig. 29 [102]. The arrangement of the bundles of glass References see page 2382

2372

11.3 Flue Gases from Stationary Sources

Clean gas outlet

Cooling air inlet

Hot air outlet

70%

Fig. 29

Acid strength (%)

98% Acid gas inlet

Sulfuric acid (to acid cooling system)

Acid mist control in the glass tubes of the WSA process.

Clean gas outlet

Cooling air outlet

Hot air outlet

Acid gas inlet Product acid outlet

Fig. 30

Schematic representation of the glass tube bundles in a WSA condenser.

tubes inside a WSA condenser is shown schematically in Fig. 30. Actual examples of applications of the WSA process are shown in Fig. 31, which demonstrates the WSA process layout for the treatment of off-gas from a

viscose plant in Austria. The layout of the WSA/SNOX plant purifying up to 1 150 000 Nm3 h−1 flue gas from the combustion of petroleum coke in the power plant of the Agip refinery in Gela, Italy [103], is shown in Fig. 32.

11.3.4 Other Emissions

2373

Steam BFW preheater

BFW

Steam drum

Peroxide Tail gas unit

Hot air from condenser

WSA condenser

Lean gas Burner

Fuel Sulfur

SO2− converter

Waste heat boiler

Rich gas

Hot air to BFW Preheater

Sulfuric Acid

Cooling water

Fig. 31

WSA process lay-out for the treatment of off-gases from a viscose plant in Austria.

CH4

SNOX plant

3 Boilers

NH3

Air 3

1

m 2

3

1 000 000 Nm /h 0.3% SOx 320 ppm NOx from 3 boilers

4

5

~130°

11 8

15 18

395° 102 °C

380 – 385° 405°

12 14

7

6

220 °C

10

9

13

360° 50 mbar 18

Combustion air 160 °C

17

16

190 °C C.W.

1 Boiler 2 Air preheaters 3,5 Dust precipitators 4 Boiler flue gas fans 6 Flue gas fans 7 Gas-gas heat exchanger

Fig. 32

8 Gas heater 9 NH3 injection grid 10 SCR deNOx reactor 11 SO2 SO3 reactor 12 Catalyst cleaning system 13 WSA sulphuric acid condenser

The combined WSA and DeNOx process known as SNOX [103].

Hydrogen Sulfide Hydrogen sulfide (H2 S) is a smelly, corrosive, highly toxic gas which is commonly found as a component of natural gas and is also generated at oil refineries during the hydroprocessing of sulfur-containing crude oil. Usually, it is converted to non-toxic and useful elemental sulfur at most locations by the Claus Sulfur Recovery process. 11.3.4.2

13 t/h 95% H2SO4

14 Air fan 15 Stack 16 Acid cooling system 17 Air cooler/boiler 18 Excess air to stack

Gases with a H2 S content greater than 25% are suitable for the Claus process. These gases, which may also contain HCN, HC, SO2 or NH3 , mainly originate from physical and chemical gas treatment units in refineries, gasification or synthesis gas plants. The main reaction References see page 2382

2374

11.3 Flue Gases from Stationary Sources

equation is:

Furnace

Catalytic section

← − 2S + 2H2 O 2H2 S + O2 − −− −− → First, H2 S is separated from the gas stream using amine extraction and fed to the Claus unit, where it is converted in two steps: (i) a thermal step; and (ii) a catalytic step. In the thermal step, H2 S is partially oxidized above 850 ◦ C so that elemental sulfur precipitates in the downstream cooler. Claus gases (acid gas) with no further combustible contents apart from H2 S are burned in lances surrounding a central muffle. Gases containing ammonia, sour wet stripper gas (SWS gas), or hydrocarbons are converted in the burner muffle. Sufficient air is injected into the muffle for the complete combustion of all hydrocarbons and ammonia. The air to acid gas ratio is controlled so that one-third of the H2 S is converted to SO2 : ← − 2SO2 + 2H2 O 2H2 S + 3O2 − −− −− → Usually, 60 to 70% of the total amount of elemental sulfur produced in the process is obtained in the thermal process step. The process extensively utilizes integrated heat exchange and hot gas bypass. In the catalytic step the remaining H2 S is reacted catalytically with the SO2 at lower temperatures (ca. 200–350 ◦ C) over titania- or alumina-based catalysts to produce more sulfur. H2 S reacts with the SO2 formed during combustion in the reaction furnace, and this results in gaseous, elemental sulfur. This is known as the Claus reaction (see also Chapter 12.4): ← − 3 S2 + 2H2 O 2H2 S + SO2 − −− −− → 2 Two or three catalytic stages are normally used to obtain sufficient conversion, with sulfur being removed by condensation between the stages (Fig. 33). Small amounts of H2 S remain in the tail gas and are treated in a tail gas unit; this results in overall sulfur recoveries of up to 99.8%. Where an incineration or tail gas treatment unit is added downstream of the Claus plant, only two catalytic stages are usually installed. The first process step in the catalytic stage is the process gas heating, this being necessary to prevent sulfur condensation in the catalyst bed, which can lead to catalyst fouling. Typically, the operating temperature of the first catalyst stage is 315 ◦ C to 330 ◦ C in order to hydrolyze COS and CS2 , which is formed in the furnace and would not otherwise be converted in the modified Claus process. Catalytic conversion is maximized at lower temperatures, and care must be taken to ensure that each bed is operated above the dew point of sulfur. Operating temperatures of the subsequent catalytic stages are typically 240 ◦ C for the second stage and 200 ◦ C for the third stage. In

H2S Air

Tail gas

Liquid sulfur Fig. 33

H2 S conversion to elemental sulfur in the Claus process.

the sulfur condenser, the process gas coming from the catalytic reactor is cooled to between 150 and 130 ◦ C. The condensation heat is used to generate steam at the shell side of the condenser. Before storage and downstream processing, liquid sulfur streams from the process gas cooler, the sulfur condensers and from the final sulfur separator are routed to the degassing unit, where the gases (primarily H2 S) dissolved in the sulfur are removed. Many improvements have been made to the Claus process, including SUPERCLAUS, where iron and chromium oxides on α-Al2 O3 carrier catalysts in the last reactor oxidizes H2 S selectively to sulfur, avoiding the formation of SO2 [104]. Also, Oxygen Claus is used, where the combustion air is mixed with pure oxygen. This reduces the amount of nitrogen passing through the unit, making it possible to increase throughput. Furthermore, better catalysts with higher surface areas and macro porosity have improved performance. Alternatives to the chromia/alumina catalysts have been identified. Ce-V mixed oxides as well as Fe-Mn-Zn-Ti-O mixed oxide catalysts have been proposed [105, 106]. Carbon Monoxide Carbon monoxide in a flue gas stream is normally oxidized to carbon dioxide by catalytic means. For the oxidation of CO, platinum is often preferred over other noble metals such as palladium or rhodium, due to its low light-off temperature [107]. The noble metals are normally supported on alumina for dense-packed granular beds or monoliths. 11.3.4.3

Carbon Dioxide Today, substantial effort worldwide is being applied to the research and development of processes for reducing CO2 emissions. While CO2 itself is not toxic, the amount 11.3.4.4

11.3.4 Other Emissions

2375

Cleaned air 10 mg / Nm3 HC 130

Recovered heat 260 000 kcal h−1

Off-gas 10 000 Nm3 h−1

Support fuel 0 kg h−1

280

150 340

210

Drying furnace Fig. 34

280

°C

Process layout for catalytic oxidation process with preheating in feed/effluent heat exchanger and support burner, the CATOX

process.

Clean air

45

Off-gas

30

Burner

300

Blower

Natural gas

°C Heat exchange material Fig. 35

Catalyst

Simplified process scheme for catalytic oxidation using regenerative heat exchange beds, the Topsøe REGENOX process.

emitted worldwide, and the effect on global warming, cause much concern. Although most projects consider using a ‘‘CO2 -neutral’’ biomass or CO2 sequestration in the sea or underground, CO2 can also be reduced to CO catalytically by reaction with hydrogen over, for example, tungsten sulfide or transition metal on zinc oxide catalysts [108, 109]. A high-pressure, liquid phase process (300 ◦ C, 107 Pa, hydrochloric acid) using an iron catalyst is reported to produce hydrocarbons at reasonable yields [110].

Hydrocarbons A diversity of volatile organic carbon (VOC) compounds is emitted from a large number of sources. These gas streams can be treated by thermal incineration, chemical scrubbing or adsorption, for example on active carbon. Furthermore, catalytic oxidation is widely used. In this process the flue gas is generally preheated to be at a minimum temperature where the catalyst is active. Above 11.3.4.5

References see page 2382

11.3 Flue Gases from Stationary Sources

a certain temperature, the catalyst may be thermally deactivated due to sintering. If the concentration of combustibles is high, the adiabatic temperature rise over the catalyst bed must be taken into account. Catalytic oxidation is particularly suited to the treatment of low concentrations of VOCs (in the ppm range), and feed–effluent heat exchange can be used for better energy efficiency. The operating temperature will depend on the catalyst type and concentration, the linear gas velocity, operating pressure, catalyst geometry, bed length, and concentration of the VOCs [111]. The process can be designed as a single fixed bed with preheat (see Fig. 34) or as a regenerative process operating two beds in cyclic mode, in which one bed functions as the preheater, whereas the other works as the catalytic oxidation reactor. This scheme is illustrated in Fig. 35. The highly exothermic nature of the oxidation reactions may require careful process design and control. For details regarding catalytic oxidation reactor modeling, the reader is referred elsewhere (e.g., Refs. [112, 113]). Usually, light-off curves showing the conversion versus the temperature are used to describe the processes. The temperatures of 50% and 90% conversion, T50 and T90 , are used to visualize the relative performance of the catalysts. As the conversion can be highly sensitive towards the conditions, care should be taken when these temperatures are compared. Deactivation of the catalyst occurs by poisoning, surface fouling, and sintering. Normally, the catalyst bed is designed to accommodate the deactivation so that a guaranteed end-of-run conversion is met. This means that the zone with the highest conversion shifts downstream as the catalyst eventually deactivates with a final breakthrough of the components to be oxidized. Guard beds may also be installed to capture trace amount of poisons, such as chlorine or sulfur. Noble metal catalysts, Pt or Pd, have been used extensively both in monolithic and granular forms for the oxidation of VOCs and CO. The noble metals are normally supported on high-surface-area alumina, which may be promoted for better thermal stability. For better utilization of the diffusion-limited catalysts, eggshell impregnation – where the active components are located in the outermost part of the catalyst particles or monoliths – is often used. Due to the high cost of noble metals, transition metal oxide catalysts have also been investigated and used extensively. For example, CuO, V2 O5 , NiO, MoO3 , Cr2 O3 , MnOx , zeolites and mixed oxides such as perovskites all exhibit catalytic oxidation activity [111, 114]. Lou and Tu [115] found ferrospinel MnFeO2 O4 to be most active for the oxidation of isopropyl alcohol among several iron-based catalysts prepared from the ferrite route. An example of light-off curves for n-hexane obtained at different space velocities over a metal oxide catalyst

supported on alumina is shown in Fig. 36. Here, it is seen that the temperature of 90% conversion, T90 , varies between 170 and 190 ◦ C depending on the conditions. Everaert and Baeyens [114] proposed a first-order reaction for the oxidation of a wide range of VOCs in the range of 260 to 340 ◦ C, the apparent rate constant being a combination of kinetics and mass transfer. In a novel V2 O5 –WO3 /TiO2 catalyst coated onto metal-fleece, these authors found few mass transfer restrictions, and the rate to zero order in oxygen at 300 ◦ C. Moreover, multiple chlorine substitution in the molecule was seen to enhance reactivity with this catalyst. Dioxins Two chemical types of compound are referred to as dioxins, namely polychlorinated dibenzo-para-dioxins (PCDDs) and polychlorinated dibenzofurans (PCDFs) (Fig. 37) [115]. In total, there are 419 related compounds of which 30 are regarded as significantly toxic; the most toxic of these compounds is 2,3,7,8-tetrachlorodibenzopara-dioxin (TCDD). Dioxins are stable and do not decompose under normal circumstances in Nature. The World Health Organization (WHO) has estimated the average half-life for dioxins in the human body to be seven years. Dioxins are also soluble in the body fat, and hence will be accumulated in animals to the 11.3.4.6

1.0 0.8 Conversion

2376

0.6 Increasing space velocity

0.4 0.2 0.0

0

100

160 140 Temperature/°C

120

180

220

200

Conversion of n-hexane over a metal oxide on alumina catalyst at different space velocities.

Fig. 36

9 8 7 Clx

6

O O PCDD

1

4

9 2

8

3

7 Cly

Clx

1 2

6

O

3 4

Cly

PCDF

Fig. 37 The molecular structures of polychlorinated dibenzo-paradioxins (PCDDs) and polychlorinated dibenzofurans (PCDFs).

11.3.4 Other Emissions

top of the food chain – so-called ‘‘bio-accumulation’’. The main toxic effects of dioxins are short-term skin lesions and reduced liver function, but in the long term they produce impairment of the immune, nervous, endocrine and reproductive systems; moreover, they may also be regarded as carcinogenic. Dioxins are normally destroyed in an incinerator and occur during complicated reformation reactions, when chlorine, organic compounds and a material that can perform a catalytic reaction are present in a gas [116, 117]. The details of published mechanisms suggest that there are two reaction pathways – one via the polycondensation of aromatic precursors and another (called de-novo synthesis) where solid carbon, oxygen and chlorides are needed for the reaction to occur. It has been suggested by some authors that the de-novo synthesis is the predominant reaction pathway [118]. The major sources, which account for about 62% of the total emissions, appear to be waste incineration, followed by metal sinter plants, incinerators for medical waste and facilities of the non-ferrous metal industry. The remainder is produced from other industrial sources, domestic heating (wood combustion), fires, and road traffic. A number of methods have been proposed for dioxin abatement, such as wet scrubbing techniques, and adsorption onto active carbon in combination with a bag filter [119]. Moving-bed or fixed-bed reactors with activated carbon are effective alternatives, and provide a removal of >95% between 120 and 150 ◦ C [119–121]. V2 O5 /TiO2 catalysts have been shown as effective for the oxidation of chlorinated VOCs, both chlorinated and non-chlorinated [114, 116]. Everaert and Baeyens [114] reported the activation energy for oxidation of trichloroethene over a V2 O5 –WO3 /TiO2 catalyst to be significantly lower than for Pt or PdO on γ -Al2 O3 , which illustrates the affinity towards chlorine. As mentioned above, these authors found the reaction order to be zero in oxygen at 300 ◦ C, whilst Stoll et al. [116] found a reaction order of about 0.3 in oxygen at lower temperatures (180–240 ◦ C) to be more relevant to the tail-end operation of waste incineration plant SCR reactors. Stoll et al. also identified a very negative influence of water on the conversion of mono- and 1,2 di-chlorobenzene model compounds under these conditions. SO2 , HCl and CO2 were not found to have negative influence. The results were fitted to a Langmuir–Hinshelwood model [116]. Higher temperatures (>300 ◦ C) may be required (by reheating the flue gas) to obtain almost complete destruction of dioxins. The SCR catalyst is normally very stable, and no significant loss of activity for dioxin removal over 50 000 h of operation has been reported [122]. Extruded catalysts can also be used [123]; the catalytic oxidation efficiency of these types of catalyst begin between 100 and 150 ◦ C, and they adsorb dioxin at lower temperatures.

2377

It is also possible to cover soot or dust filters with a catalytic coating which consists of V2 O5 /WO3 –TiO2 on fibrous filters [124]. Tests with these types of filter have shown that it is possible to destroy >99% of dioxins, even at temperatures down to 200 ◦ C. Hg Oxidation and Capture Most waste incineration plants are of the stoker design, which is flexible towards the fuel. Flue gas cleaning is normally performed by a bag house filtration, followed by a scrubber to remove acid gases. A tail-end SCR can also be installed for NOx removal. Sorbents for capturing heavy metals may be injected upstream of the bag house. In recent years, attention has focused increasingly on reducing the emissions of toxic metals from coal-fired power plants; in particular, mercury emissions have received significant interest as the most important of the volatile metals. About 40 states in the US have issued fish consumption advice based on the bio-accumulation and neurotoxicologic effects of mercury. The highest emitters of mercury to the air include coal-burning power plants, and combustors or incinerators of municipal, medical, and hazardous waste. This resulted in the EPA issuing, in May 2006, the final rule regarding the regulation of mercury emissions from utilities, the Clean Air Mercury Rule (CAMR) [125]. The rule is the first of its kind, and is projected to reduce annual emissions of mercury from 48 tons to 31 tons, beginning in 2010, followed by a second phase cap of 15 tons when the cap-and-trade program is fully implemented. Mercury emissions from municipal, medical and hazardous waste combustion and incineration will be reduced by 50 to 94% compared to the 1990 level. From 2009, power plants will be required to install continuous Hg monitoring equipment in their stacks. At the power plant, mercury enters with the fuel and leaves through the boiler bottom ash, fly ash from the filters, from the scrubber byproducts and waste, and through the stack. Three mercury species are normally considered: elemental mercury (Hg0 ), oxidized mercury, usually bivalent Hg2+ , and particle-bound mercury Hgp . The speciation in the flue gas is important to understand due to the different physical and chemical properties, and in order to optimize the gas cleaning equipment downstream of the boiler. Hg exits in the furnace mainly as reduced, metallic Hg0 , which may pass unchanged through the flue gas treatment system, be oxidized (e.g., to HgCl2 ) or be adsorbed onto the fly ash as particle-bound Hgp . The oxidized species are hydrophilic and can be captured with a high removal efficiency in conventional wet scrubbers, whereas elemental Hg0 is 11.3.4.7

References see page 2382

2378

11.3 Flue Gases from Stationary Sources

Stack: 3 – 34% (25%)

Boiler

Coal

SCR

Slag: 0.2 – 20% (1%)

ESP

FGD

Fly ash: 27– 92% (49%)

Gypsum: Waste: 4 – 51% 2 – 10% (17%) (9%)

Ranges of the distribution of Hg emissions from Dutch coal-fired power plants. Data from Meij and te Winkel [128]. The numbers in parentheses are average values.

Fig. 38

11.3.4.7.1 Hg Oxidation Thermodynamics The oxidation of Hg0 to HgX2 , where X is a halogen, is thermodynamically favored by a low temperature and a high concentration of the halides. Above 300 ◦ C, chlorine in the gas phase is mainly present as HCl(g). Bromine is mainly present as Br2 (g) up to 700 ◦ C. On a thermodynamics basis [130–136], it can be calculated that bromine is a much more powerful oxidizing agent than chlorine, which is illustrated in Fig. 39. Davidson [137] notes that the experimental and plant data do not support a strong effect of coal chlorine on the speciation and capture of mercury. Although coal chlorine is one of the factors that affect mercury oxidation, it does not seem to be the determining factor, possibly due to reactions of chlorine with the fly ash. Ash minerals may also bind the acidic gas components and influence the flue gas equilibrium composition. There is strong evidence that unburned carbon in the fly ash plays a significant role in the Hg capture. Furthermore, other species may form, such as Hg(NO3 ) · H2 O [7, 137]. 11.3.4.7.2 Hg Removal from Coal-Fired Power Plants A Conventional Flue Gas Cleaning Processes Several reviews on the behavior of Hg in coal-fired power plants

120 100 Oxized fraction / %

hardly captured [126, 127]. Figure 38 shows an example of how Hg emissions may be distributed in a power plant. Meij and te Winkel [128] provide a range of separation factors for a number of Dutch power plants operating with a range of coal compositions. While these authors observed a beneficial effect of SCR on total Hg removal, no difference was seen between pure coal combustion and cofiring with biomass. Hg may be re-emitted from the solids (fly ash, active carbon, scrubber sludge) and liquids. For example, Xin et al. [129] found that wet FGD samples emitted Hg, whereas dried samples acted as sinks for atmospheric Hg. Similarly, a high carbon content of ash samples is beneficial for avoiding the re-emission of Hg.

80 60 40 20

20 ppm HCI 20 ppm HBr 2 ppm HCI 2 ppm HBr

0 200 250 300 350 400 450 500 550 600 650 700 Temperature / °C Fig. 39 Equilibrium oxidized fraction of gaseous Hg-species in the presence of HBr or HCl. 3.5% O2 and 8% H2 O and 1 ppb Hg(g) were used for the equilibrium calculations using HSC chemistry [136].

and flue gases have been published [127, 138]. Although primary measures can be taken to control Hg emissions, there is substantial evidence that conventional flue gas cleaning processes will assist in removing only part of the mercury in the coal (see Table 7) [139]. Even with a fabric filter (FF), significant Hg removal is achieved, probably due to the long residence time, the reaction with residual carbon on the fly ash, and adsorption into the filter cake. Significant oxidation over the air preheater due to the residence time and cooling to a lower temperature can also be observed, depending on the coal type. A recent comprehensive field test series demonstrated a pronounced beneficial effect of SCR and FGD on Hg removal from bituminous coal firing [140]. For plants without SCRs, 50 to 80% of the coal Hg has been observed to be captured by the FGD, while the range was 70 to 97% when SCR was installed, even at low chlorine content. An increase in the fraction of Hg2+ over the SCR was demonstrated for several full-scale units,

11.3.4 Other Emissions

2379

EPA data showing average mercury capture by coal rank and air pollution control configuration

Tab. 7

Control Technology Configuration Particulate

NOx

SO2

CS-ESP CS-ESP CS-ESP CS-ESP HS-ESP HS-ESP HS-ESP HS-ESP PS PS FF FF FF FF FF

None None None SCR None None None SCR None None None None None SCR SCR

None Wet FGD SDA Wet FGD None Wet FGD SDA Wet FGD None Wet FGD None Wet FGD SDA Wet FGD SDA

Bituminous

Sub-bituminous

Lignite

36 66 36 90 10 42 40 90 n.a. 12 89 97 95 90 98

3 16 35 66 6 20 15 25 9 0 73 73 25 85 n.a.

0 44 n.a. n.a. n.a. n.a. n.a. n.a. n.a. 33 n.a. 0 0 n.a. n.a.

CS-ESP = cold-side ESP; ESP = electrostatic precipitator; FF = fabric filter; FGD = flue gas desulfurization; HS-ESP = hot-side ESP; PS = particulate scrubber; SDA = spray dryer adsorber. n.a. = data not available.

where Hg removal was improved by SCR catalysts, in particular for bituminous coals [141]. The negative effect of NH3 on Hg oxidation seen in laboratory and pilot tests does not seem to be as important in the full-scale testing [92, 142]. In a slip-stream reactor test, different SCR catalysts showed about 90% NOx reduction and 25 to 65% oxidation of Hg0 at typical full-scale space velocities in flue gas from an 87 : 13 blend of subbituminous : bituminous coals. After 2200 h of exposure to the flue gas, the catalysts appeared to lose activity for NOx and Hg conversion [142]. The significant variation in Hg content of the coals, the dynamic operation of the power plants, the low concentrations, and the interaction with the fly ash all make the analysis and test of Hg removal a major challenge [143]. The re-emission of Hg from scrubbers due to sulfite oxidation can, at least to some extent, be controlled by adjusting the operating conditions. With correct control it is possible to determine whether the Hg will be in the sludge or in the liquid phase [144].

types seem to be brominated carbons, where up to 95% mercury removal may be obtained [145]. The amount of carbon injected corresponds to an increase of the fly ash carbon by 1 to 3 wt.%, which may make the fly ash unsellable for cement production. The TOXECON process, which is designed to overcome this issue, is based on treatment of the gas out of the electrostatic precipitator (ESP) with sorbent, followed by a FF upstream of the FGD [146]. Other sorbents are silicates or the manganese oxide-based Pahlmanite . An alternative to sorbent injection for low-rank fuels is coal blending, where up to 15% of high-chlorine bituminous coal (100 ppm Cl) with PRB coal (80% Hg capture when used in combination with either standard carbon injection or high residual carbon on the fly ash [148].

B Sorbent Injection, Coal Additives, and Coal Blending The injection of activated carbon up front of a filter has been used for incineration plants in the capture of toxic substances. This technology has also been extensively tested for the capture of mercury from coal-fired power plants, especially for low-rank fuels where it is effective in combination with bag filters [92]. The most efficient

11.3.4.7.3 Hg Oxidation Mechanism and Catalysis Thermodynamic data on Hg oxidation may be inadequate, as they tend to underestimate Hg oxidation under conditions where it should be thermodynamically limited [137]. However, this is most likely caused by heterogeneous References see page 2382

2380

11.3 Flue Gases from Stationary Sources

reactions with the fly ash, especially if the carbon level is high. Although field measurements confirm the general trend from thermodynamics, models based solely on homogenous chemistry have failed. The introduction of heterogeneous models improved the prediction somewhat [126]. Niksas and Fujiwara [149] have proposed an Eley–Rideal model for oxidation of Hg0 over SCR catalysts based on the EPA power plant data. The mass balance gives: S S ) = kHg · CHg · Cl km,Hg (CHg − CHg

= kHg

0 KH Cl CH Cl S 0 1 + KN H3 CN H3 + KH Cl CH Cl

S CHg

where km,Hg is the film transfer coefficient as determined from the Gratz–Nusselt analysis, subscripts S and 0 denote surface and bulk conditions, respectively, ki is the surface reaction rate,  the surface coverage, and Ki the adsorption equilibrium constant. The model includes the NH3 /NOx ratio, the HCl and Hg0 concentrations, catalyst geometry and space velocity. In this model, HCl competes with NH3 for the surface sites, and Hg0 combines with the adsorbed HCl. This explains the observed inhibitory effect of NH3 on Hg oxidation, and the model confirms the observation that smaller catalyst channels promote the Hg conversion. A mechanism for catalytic Hg oxidation and subsequent capture on carbon, where oxidation and capture are decoupled, has been proposed [150]. It is suggested that Hg2+ as a Lewis acid competes with, for example H2 SO4 , for the basic sites on carbon. HCl was found initially to promote the capture kinetics, while not acting as an oxidizing agent. The hypothesis is that HCl binds to the basic sites forming a carbenium ion (Fig. 40). It was observed that, while no further

SO2

Hg0(g)

HCl HSO−4(5)

Hg0(s)

Hg(II) − Base Cr ne−(s) NO2(g)

Hgn+

Hg(NO3)2(s)

Alternative Processes Dedicated Hgoxidation catalysts installed at the outlet of the ESP have shown promise for the oxidation of Hg0 when installed after an ESP [127]. In particular, Pd and carbon-based catalysts have been shown to be active for Hg oxidation in a pilot-scale test. The photocatalytic oxidation of Hg over titania, zinc oxide, tin oxide, and cerium oxide was reported as early as 1971 [154]. The photooxidation is apparently caused by OH radicals, arising from the basic OH− groups, together with a chemisorbed oxygen species, very likely O2− . Recently, this concept has been revitalized. Up to 90% Hg oxidation was demonstrated in a benchscale test at an (un-optimized) energy consumption of 164 kWh Nm−3 flue gas [155]. Multipollutant processes such as the electrocatalytic ECO process have also been demonstrated on a large scale [156]. 11.3.4.7.4

HgCl2

(s)

NO2(g)

adsorption took place (a saturated carbon surface was detected at a constant Hg concentration in the gas leaving the carbon bed), oxidation still took place over carbon. Consequently, the carbon acts as a catalyst for oxidation of Hg. Olson et al. [150] found strong competition with SO3 /H2 SO4 on carbon, which was proposed as a ‘‘deadend’’ blocking the active oxidation sites due to formation of sulfate. The effect of NO2 and HCl as promoters for oxidation and capture on carbon was found to be additive. HCl, NO, and NO2 have each been observed to promote oxidation and capture, both individually and in combination. However, the combination of SO2 with NO2 greatly reduces the capture of Hg0 on activated carbon, whereas oxidation continues on the solid surface [127]. Zhao et al. [151] found that NO and SO2 inhibit Hg0 oxidation under homogeneous conditions at high temperatures (>500 ◦ C) when water was present. The catalytic oxidation of Hg2+ and, to some extent, the subsequent capture take place on the fly ash, in particular due to the content of iron oxides and unburned carbon [126, 152, 153]. No capture of elemental mercury was observed without accompanying oxidation – that is, the catalytic effect is determining the performance of the adsorbent.

11.3.5 Hg(NO3)2

NO−2(g) Carbon Fig. 40 The proposed mechanism for Hg oxidation on carbon. (Adapted from Ref. [150].)

Multipollutant Processes The DESONOX Process The DESONOX process is based on a similar process scheme as the above-mentioned WSA process [157]. An electrostatic precipitator reduces the dust to about 20 mg Nm−3 before the flue gas is led to a reactor containing two catalyst types. NH3 is added to the 450 ◦ C 11.3.5.1

11.3.5 Multipollutant Processes

hot flue gas for NOx removal in the first SCR catalyst. The second catalyst oxidizes SO2 to SO3 . The issues have been to obtain a sufficiently concentrated sulfuric acid during cooling and condensation, and to eliminate emissions of the SO3 and H2 SO4 aerosols. Three-stage scrubbing and wet electrostatic precipitator were applied for this final clean up.

2381

while those from NOx reduction by NH3 are gaseous. A cyclic thermal regenerative process of the coke is required [162]. The process has been in commercial operation since 1984, and can also include the removal of dioxin and Hg by adsorption onto the active coke [163]. Catalytic Filters In recent years, the development of catalytic hightemperature filters has resulted in the potential advantage of combined dust removal and catalytic flue gas cleaning from compounds such as VOC, NOx , SO3 , and dioxin. Future legislation for fine particle emissions could call for this technology as an efficient combined process. Filter materials, catalysts and process concepts were reviewed by Fino et al. [164]. The same group obtained high removal rates of NOx and VOC removal at 200–210 ◦ C by passing the flue gas through a high- temperature-resistant polymeric filter bag enclosing a ceramic foam carrier with combinations of MnOx /CeO2 and V2 O5 /WO3 /TiO2 catalysts [165]. The superficial velocity was in the range of 10 to 60 Nm3 m−2 h−1 . This catalyst system was found to be sensitive to SO2 . Figure 41 shows the concept of a high-temperature, fiber-based catalytic filter. The dust is collected on a barrier layer, which protects the catalyst from particulate poisons. The lateral flow also improves the catalyst effectiveness, since diffusion and film transfer restrictions are significantly reduced [166, 167]. The integrated layout for a combined process is shown in Fig. 42. Catalytic filters are now commercially available for dioxin removal (Goretex Remedia from Gore) or combined removal of dioxin and NOx (Cerafil TopKat from Madison and Haldor Topsøe A/S) [166]. 11.3.5.3

Other Combined Processes Liu et al. [158] presented details of a catalyst based on CuO and Na2 O supported on Al2 O3 -coated cordierite ceramic monoliths for the combined removal of SO2 and NOx in a cyclic redox process. CuO and CuSO4 are active for the SCR reaction; therefore, the injection of NH3 into the Shell FGD process mentioned above can result in the combined removal of SO2 and NOx . Jeong and Kim [159] carried out experiments for the simultaneous removal of SO2 /NO by CuO/γ -Al2 O3 sorbent/catalysts in a fluidizedbed reactor. The optimum temperature ranges for NO reduction over fresh and sulfated CuO/γ -Al2 O3 catalyst are found to be 250–300 ◦ C and 300–450 ◦ C, respectively, with the NO removal efficiency over the sulfated catalyst being somewhat higher than that with the fresh catalyst. The optimal Cu/S ratio for the simultaneous removal of SO2 and NO was found to be 1.5. Over the temperature range of 350 to 400 ◦ C, SO2 and NO removal efficiencies were high (>90%) at reasonably low N2 O formation (90% implied that minimal coke formation was occurring under these oxidative conditions. The authors noted, however, that the chlorine balance was in the range 75 to 90%, suggesting that there is significant chlorine retention by the zeolite, which may lead to longer-term catalyst deactivation. The acid zeolite catalysts offer the advantage over noble metal catalysts of producing no Cl2 during reaction. The performance of these catalysts for trichloroethene oxidation was less impressive, with less than 50 K difference between the temperature required for 50% conversion in the catalyzed and uncatalyzed combustion. Zeolite catalysts also performed poorly relative to amorphous silica–alumina for the oxidative dechlorination of chloropropanes [21], but this was attributed by the authors to the formation of coke within the zeolite pores. There is a considerable body of evidence to show that zeolites are dealuminated on exposure to CFCs, Halons or chlorocarbons at elevated temperatures. Recently, Kiricsi and Nagy [13] reviewed the chemistry of interaction of C-1 CFCs with oxides and zeolites. As an example, Fig. 2 shows a generalized reaction scheme for the reaction of CCl2 F2 with a zeolite. Aluminum is removed from the zeolite framework as AlF3 , but both chlorine and fluorine are also incorporated into the zeolite. The CCl4 formed as an initial product then reacts further with oxide ions of the zeolite to produce phosgene (COCl2 ), which can cause further dealumination. The

+ M+ O O M O Si Al− Al− Si O O O O O O O O Si Si Si

O

O

+ 3 CCl2F2 MCl + AlF3 + 2 CO2 + CCl4 + O

Cl

F

M+

O

F Si

O

F Si

O

Al−

Si

Si O

O O

O

O

Si

Reaction of CCl2 F2 with a zeolite. (Reproduced from Ref. [13], with permission.)

Fig. 2

various species involved have been identified by solidstate NMR, FTIR and chemical analysis. The phosgene intermediate in particular is believed to be the key to structural damage of the zeolite. Very analogous chemistry occurs when Halon 1301 (CF3 Br) interacts with zeolites at elevated temperatures. Surface analysis with X-ray photoelectron spectroscopy (XPS) of ZSM-5 zeolites exposed to CF3 Br at elevated temperatures shows that the surface is depleted in aluminum but enriched in fluorine. Although the zeolite remains crystalline to X-ray diffraction, the acid sites are either diminished or completely removed [12]. No bromine can be detected, which might suggest that aluminum is lost as AlBr3 , but Al K-edge X-ray absorption near edge structure (XANES) spectra show the characteristic signature of AlF3 in the zeolite. Transition-metal cations such as Ni2+ ion-exchanged into the zeolite appear to reduce the extent of dealumination, possibly by forming nickel fluoride or oxyfluoride species. Transition-metal exchanged ZSM-5 zeolites are able to sustain a steady-state hydrodebromination reaction of CF3 Br with CH4 for periods of up to 20 h on stream before gradual deactivation occurs. Catalyst activity could be regenerated by treatment in hydrogen, but not in oxygen, which suggested that the deactivation was due to the build up of halogen (fluoride) in the zeolite rather than to coke deposition. Problems associated with coke formation or zeolite dealumination caused by exposure to halocarbons at elevated temperatures should not be present in lowtemperature hydrodehalogenation in aqueous solution. Real wastewater samples will, however, contain many different species other than chlorocarbons. In field applications, supported palladium catalysts which work well for hydrodehalogenation in the laboratory suffer from poisoning by sulfite ions or other reduced ionic species. For example, McNab et al. used a palladium–alumina catalyst operated in situ to degrade chlorinated ethenes in a reactive well approach [47]. The unit was operated for 4 to 8 h per day, and the catalyst regenerated between use by flushing with clean water and exposure to air. The frequent periodic regeneration was found to be necessary to maintain performance, and the authors warned that the method may not be useful in groundwaters containing reduced sulfur species. A novel solution to this problem, utilizing a hydrophobic zeolite support, is described by Sch¨uth et al. [48, 49]. These authors found, in laboratory studies, that palladium supported in a dealuminated zeolite Y maintained activity for hydrodechlorination of 1,2-dichlorobenzene in water at room temperature even in the presence of 30 mg L−1 Na2 SO3 , whereas alumina and other supported palladium catalysts were immediately deactivated. They argued that the hydrophobicity of the dealuminated zeolite prevented aqueous sulfite ions from entering the

Total CHC concentration n / µmol·L−1

References

2393

250 Inflow sampling port 1 Outflow sampling port 3

200 150 100 50 0 0

150

300

450

600

750

Time / days

Total chlorocarbon concentrations at inlet and outlet during wastewater treatment with a palladium on hydrophobic zeolite catalyst. (Reproduced from Ref. [49], with permission.)

Fig. 3

pores of the zeolite and poisoning the palladium sites for hydrodechlorination. The approach was then demonstrated with 20 kg of catalyst in a pilot-scale water treatment plant. Figure 3 shows, graphically, the inlet and outlet total chlorocarbon concentrations over the 2-year test period. In this system also, hydrogen sulfide must be avoided, and the system was flushed every 42 h with H2 O2 for 10 min to inhibit sulfate-reducing bacterial growth (the occasional spikes in system performance seen in Fig. 3 were due to failure of the H2 O2 injection system). No catalyst regeneration was required. This study demonstrates that the concept of a hydrophobic zeolite support for a hydrodehalogenation catalyst shows great promise for the treatment of chlorocarbon-contaminated groundwaters. References 1. P. Hohener, D. Werner, C. Balsiger, G. Pasteris, Crit. Rev. Environ. Sci. Technol. 2003, 33, 1. 2. L. Manzer, M. J. Nappa, Appl. Catal. A: General 2001, 221, 267. 3. F. Alonso, I. P. Beletskaya, M. Yus, Chem. Rev. 2002, 102, 4009. 4. V. I. Kovalchuk, J. L. D’Itri, Appl. Catal. A: General 2004, 271, 13. 5. R. F. Howe, Appl. Catal. A: General 2004, 271, 3. 6. E. Kemnitz, D.-H. Menz, Prog. Solid State Chem. 1998, 26, 97. 7. J. M. Berty, Ind. Eng. Chem. Res. 1997, 36, 513. 8. G. Aylward, T. Findlay, S. I. Chemical Data, 4th Ed., Wiley, Brisbane, 1999, p. 115. 9. E. Shin, M. Keane, Catal. Lett. 1999, 58, 141. 10. T. N. Kalnes, R. B. James, Environ. Prog. 1988, 7, 185. 11. R. Tran, E. Kennedy, B. Dlugogorski, Ind. Eng. Chem. Res. 2001, 40, 3139. 12. K. Li, F. Oghanna, E. Kennedy, B. Dlugogorski, A. Fazeli, S. Thomson, R. F. Howe, Microporous Mesoporous Mater. 2000, 35-36, 219; R. F. Howe, S. Thomson, Y. Yang, K. Lee,

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13. 14. 15. 16. 17. 18.

19. 20. 21. 22. 23. 24.

25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35.

36. 37. 38. 39. 40. 41. 42. 43. 44.

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E. Kennedy, B. Dlugogorski, J. Mol. Catal. A: Chemical 2002, 181, 63. I. Kiricsi, J. Nagy, Appl. Catal. A: General 2004, 271, 27 and references therein. Y. Sun, S. M. Campbell, J. H. Lunsford, G. E. Lewis, D. Palke, L. M. Tau, J. Catal. 1993, 143, 32. Y. Ukisu, T. Miyadera, Appl. Catal. A: General 2004, 271, 165. F. Kopinke, K. Mackenzie, R. Koehler, A. Georgi, Appl. Catal. A: General 2004, 271, 119. R. L. Calhoun, K. Winkelman, G. Mills, J. Phys. Chem. B 2001, 105, 9739. K. Li, E. Kennedy, B. Dlugogorski, R. F. Howe, Chem. Commun. 1999, 709; K. Li, E. Kennedy, B. Dlugogorski, R. F. Howe, Catal. Today 2000, 63, 655. R. L´opez-Fonseca, J. I. Guti´errez-Ortiz, J. R. Gonz´alez-Velasco, Appl. Catal. A: General 2004, 271, 39. C. Pistarino, E. Finocchio, M. A. Larrubia, B. Serra, S. Braggio, G. Busca, Ind. Eng. Chem. Res. 2001, 40, 3262. E. Finocchio, C. Pistarino, S. Dellepiane, B. Serra, S. Braggio, M. Baldi, G. Busca, Catal. Today 2002, 75, 263. M. Hegedus, A. Dombi, Appl. Catal. A: General 2004, 271, 177, and references therein. W. Chu, C. C. Wong, Ind. Eng. Chem. Res. 2004, 43, 5027. M. Tajima, M. Niwa, Y. Fujii, Y. Koinuma, R. Aizawa, S. Kushiyama, S. Kobayashi, K. Mizuno, H. Ohuchi, Appl. Catal. B: Environmental 1996, 9, 167. N. Coute, J. D. Ortega, J. T. Richardson, M. V. Twigg, Appl. Catal. B: Environmental 1998, 19, 175. F. J. Urbano, J. M. Marinas, J. Mol. Catal. A: Chemical 2001, 173, 329. T. Mori, J. Kubo, Y. Morikawa, Appl. Catal. A: General 2004, 271, 69. M. A. Keane, Appl. Catal. A: General 2004, 271, 109. S. Chandra Shekar, J. Krishna Murthy, P. Kanta Rao, K. S. Rama Rao, Appl. Catal. A: General 2004, 271, 95. M. Legawiec-Jarzyna, A. Srebowata, W. Jusczyk, Z. Karpinski, Appl. Catal. A: General 2004, 271, 61. V. de Jong, R. Louw, Appl. Catal. A: General 2004, 271, 153. V. Y. Borovkov, D. R. Luebke. V. I. Kovalchuk, J. L. D’Itri, J. Phys. Chem. B 2003, 107, 5568. B. Heinrichs, J. P. Schoebrechts, J. P. Pirard, J. Catal. 2001, 200, 309. C. C. Chang, C. M. Reo, C. R. F. Lund, Appl. Catal. B: Environmental 1999, 20, 309. A. Malinowski, W. Juszczyk, M. Bonarowska, M. Wojciechowska, Z. Kowalczyk, Z. Karpinski, Reaction Kinetic Catal. Lett. 1999, 68, 53. M. Martino, R. Rosal, H. Sastre, F. V. Diez, Appl. Catal. B: Environmental 1999, 20, 301. B. Coq, F. Figu´eras, S. Hub, D. Tournigant, J. Phys. Chem. 1995, 99, 11159. S. Deshmuk, J. L. D’Itri, Catal. Today 1998, 40, 377. K. Early, V. I. Kovalchuk, F. Lonyi, S. Deshmukh, J. D’Itri, J. Catal. 1999, 182, 219. P. P. Kulkarni, V. I. Kovalchuk, J. L. D’Itri, Appl. Catal. B: Environmental 2002, 36, 299. B. Heinrichs, F. Noville, J. P. Schoebrechts, J. P. Pirard, J. Catal. 2003, 220, 215. H. M. Roy, C. M. Wai, T. Yuan, J. K. Kim, W. D. Marshall, Appl. Catal. A: General 2004, 271, 137. C. B. Wang, W. X. Zhang, Environ. Sci. Tech. 1997, 31, 2154. X. Xu, M. Zhou, P. He, Z. Hao, J. Hazardous Mater. B 2005, 123, 89.

45. E. J. Creyghton, M. Burgers, J. C. Jansen, H. van Bekkum, Appl. Catalysis A: General 1995, 128, 275. 46. J. R. Gonz´alez-Velasco, R. L´opez-Fonseca, A. Aranzabel, J. I. Guti´errez-Ortiz, P. Steltenpohl, Appl. Catal. B: Environmental 2000, 24, 233. 47. W. W. McNab, R. Ruiz, M. Reinhard, Environ. Sci. Technol. 2000, 34, 149. 48. C. Sch¨uth, S. Disser, F. Sch¨uth, M. Reinhard, Appl. Catal. B: Environmental 2000, 28, 147. 49. C. Sch¨uth, N. A. Kummer, C. Weidenthaler, H. Schad, Appl. Catal. B: Environmental 2004, 52, 197.

11.5

Solid Catalysts for the Oxidation of Volatile Organic Compounds James J. Spivey∗

11.5.1

Introduction Significance of the Problem Volatile organic compounds (VOCs) consist of a wide range of relatively low molecular-weight compounds that are emitted from a wide range of industrial processes. Over 300 chemicals are designated as VOCs by the U. S. Environmental Protection Agency [1]. VOCs are linked to photochemical smog [2], stratosphereic ozone depletion [3, 4], and the formation of tropospheric ozone [5]. In addition, they may have inherent toxicity or carcinogenicity [6]. For these reasons, VOCs are carefully regulated, and a number of processes are used to control the emissions of these compounds [7–11]. These include adsorption [12, 13], thermal oxidation [14, 15], wet scrubbing [16, 17], photocatalysis [18, 19], biofiltration [16, 20, 21], plasma oxidation [22, 23], plasma-catalytic oxidation [24–26], and catalytic oxidation [7, 10, 27]. In order to reduce costs, these processes are sometimes linked; for example, adsorption can be used to concentrate the oftendilute VOC emissions so that the size and cost of subsequent processes such as catalytic oxidation are minimized [28, 29]. 11.5.1.1

Applications of Catalytic Oxidation for VOC Control VOCs are generated by a wide range of industrial operations, such as chemical processing [30, 31], semiconductor manufacture [32, 33], polymer synthesis and processing [34], wastewater treatment [35], and coating 11.5.1.2

∗

Corresponding author.

11.5.2 Reactors and Process Configurations for Catalytic Oxidation Systems

operations [36–38]. They are also of concern in controlling indoor air quality [39–41]. Catalytic oxidation has been applied to VOC-containing gas streams from each of these types of operation. The VOC concentration, total flow rate, gas temperature, and gas composition of the VOC-containing air streams from these operations will differ significantly, but they are all characterized by relatively dilute levels of VOCs (102 to 104 ppm; well below the explosive limits), high flow rates, inlet temperatures at or near ambient, and mixtures of VOCs rather than single components. Typically, the most cost-effective applications are those with: • VOC concentrations and inlet temperatures sufficiently high to minimize the cost of preheat (roughly 500–5000 ppm total VOC concentration, and temperatures well above ambient) • Reasonably consistent concentrations, flow rates, and temperatures to avoid the need to over-design the unit • Low levels of catalyst poisons (e.g., heavy metals or phosphorus-containing compounds) and particulates (e.g., silica) to minimize the rate of catalyst deactivation. In addition to the control of VOCs by ‘‘end-ofpipe’’ technologies such as catalytic oxidation, significant efforts are being directed at developing processes that minimize or eliminate VOCs, such as water-based coatings and paints [42–44]. These approaches require significant development time, while product quality (e.g., automotive paint finish appearance and durability) must be maintained and added cost kept to a minimum. However, this type of alternative to processes such as catalytic oxidation must be considered in the evaluation of VOC control options. Catalytic Oxidation Catalytic oxidation is one of the most widely used techniques to control VOC emissions. In many cases, the lower operating temperature compared to thermal oxidation helps to reduce auxiliary fuel costs, which can account for 40 to 70% of the total operating costs of these systems [45–48]. Catalytic oxidation also eliminates NOx emissions associated with thermal oxidation, because the reaction temperatures are well below those at which NOx are formed [49]. Compared to adsorption processes alone, catalytic oxidation converts the VOCs to oxidized products rather than simply transferring them from one phase to another. However, if the VOCs can be recovered in pure enough form to be recycled, then adsorption or other processes that do not destroy them may be more cost-effective. There are, however, significant challenges to the use of catalytic oxidation for the removal of VOCs: 11.5.1.3

2395

• the catalyst can add significantly to the cost of the system • catalyst deactivation is inevitable, though it can be greatly minimized by appropriate choice of the catalyst and proper operation of the unit • it is possible (as with other chemical reactions used for VOC removal) to form byproducts that are as toxic or harmful as the original VOC • the reaction of a multicomponent VOC mixture cannot typically be predicted from studies on single components, so that the high conversion required by regulations must be experimentally verified for the specific mixture to be controlled • as with any emission control process, changes in flow rate or composition can change the overall VOC conversion, so that reactor design and process control become critical in minimizing costs. Scope of this Chapter This chapter focuses on the catalytic oxidation of VOCs, specifically the oxidation of low concentrations of these compounds in air. Although catalytic oxidation can be integrated into processes such as adsorption [28] and particulate filtration [50–52], these types of combined processes, various reactor designs for catalytic oxidation systems [53], or the analysis of mass- and heat-transport phenomena [6, 54] are not discussed in this chapter. Instead, the discussion here focuses on the catalytic processes involved in the oxidation of VOCs over supported noble metal and metal oxide catalysts at conditions typical of VOC control systems: 11.5.1.4

• dilute concentrations of VOCs in air (typically 102 to 104 ppm) • relatively high space velocities (104 to 105 h−1 ) • operating temperatures of 200 to 500 ◦ C • ambient pressure (0.1 MPa). 11.5.2

Reactors and Process Configurations for Catalytic Oxidation Systems

Though not discussed in detail here, the literature shows that a wide range of reactor designs and process configurations have been developed for catalytic oxidation systems. A distinction is made here between the reactor design, which includes only the portion of the process in which the active catalyst is located, and the process configuration, which includes the overall arrangement of unit operations (including the catalytic reactor). The general objective of both is to minimize overall References see page 2408

2396

11.5 Solid Catalysts for the Oxidation of Volatile Organic Compounds

Monolith

Fig. 1

Washcoat on the walls of the monolith

Active components on the washcoat

Catalytic monolith for volatile organic compound (VOC) oxidation [57].

costs. A brief overview of reactor designs and process configurations is presented below. Reactors Reactors must be designed to deal with two primary constraints. First, because of the high volumetric flow rates in most catalytic oxidation systems, reactors must be designed to minimize pressure drop. Second, emission regulations typically require very high conversions of the VOCs in the gas stream. The result of these two constraints is that fixed beds of catalyst particles are not generally suitable for practical applications [54]. Several reactor designs have been developed to meet these constraints. Perhaps the most widely applied is the coated ceramic monolith [55, 56], which consists of parallel, non-intersecting channels of an inert oxide (e.g., cordierite; 2MgO · 2Al2 O3 · 5SiO2 ) which is coated with a thin layer of active, microporous catalyst (Fig. 1) [57]. Structurally similar reactors (e.g., with non-intersecting parallel channels), in which a metal is used in place of the inert oxide, are also reported [58, 59]. In this case, the active catalyst is coated using methods such as anodic oxidation [58, 60] or deposition of organic precursors on calcined metal [61]. One example of these structures is shown in Fig. 2 [62]. Details of the design, analysis, and preparation of these reactors are provided elsewhere [63]. 11.5.2.1

Other reactor designs have been reported; these include fluid beds [64], coated wire gauze [65], and catalytic ‘‘filters’’, which are designed to remove both particulates and to oxidize VOCs [66, 67]. An example of a catalytic filter is shown in Fig. 3 [67]. Process Configurations In general, the configuration for all VOC oxidation processes can be represented as shown in Fig. 4. The VOC-containing gas is first preheated to bring it to a temperature at which the catalytic oxidation reaction rate is sufficiently high. This heat exchange can be accomplished using either indirectly (e.g., with a conventional heat exchanger), or directly (e.g., with a heated bed of ceramic beads). For hydrocarbons and oxygenates, the catalytic reactor operates at roughly 250–400 ◦ C, depending on the catalyst and the chemical nature of the VOC. For chlorinated VOCs, a higher temperature is required, usually above 400 ◦ C. 11.5.2.2

Filter

Purified gas

Flue gas

Catalytic abatement of gaseous pollutants

Dust cake

2b

2b

Pore

Support

Corrugated metal monolith reactor showing direction of gas flow. (From Ref. [62].)

Fig. 2

Catalytic layer

Schematic of a catalytic filter for particulate removal and VOC oxidation. (From Ref. [67].)

Fig. 3

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11.5.3 General Considerations

After catalytic oxidation in the reactor, the exit gas is used to preheat the incoming VOC-containing gas before being vented. It is usually necessary to burn auxiliary fuel to raise the gas temperature to the desired value. 11.5.3

General Considerations 11.5.3.1 Specific Problems in Chlorinated VOC (CVOC) Oxidation Several challenges specific to the oxidation of chlorinated VOCs (CVOCs) have been identified, including: deactivation; byproduct formation; and the fact that higher temperatures are required.

Catalyst Deactivation in CVOC Oxidation Many widely used VOCs of industrial interest are C1 −C2 chlorocarbons. These compounds pose several problems that are not associated with the catalytic oxidation of hydrocarbon VOCs. For example, they tend to deactivate both noble metal and metal oxide catalysts because chlorine (or byproducts of its oxidation) adsorb strongly onto the catalyst. This has been addressed in one interesting catalyst formulation: NaCO3 is used as part of the catalyst support to capture chlorinated products, primarily HCl, and to minimize deactivation [68]. Although it eventually requires regeneration or disposal, this is one way of minimizing the rate of deactivation. 11.5.3.1.1

11.5.3.1.2 Byproducts There is a possibility of forming byproducts that are more hazardous than the CVOC reactant; for example, phosgene (a toxic gas) can be formed from the catalytic hydrolysis of CCl4 [56]. Elemental chlorine (also toxic) can be formed from the Deacon reaction in the presence of the typically large excess of oxygen:

(Note: Even if this reaction does not occur, HCl is formed from CVOC oxidation and must be removed from the product gas.) This reaction shifts toward HCl at higher temperatures. As an example, at 420 ◦ C, for a starting composition of 2000 ppm Cl atoms, 3.3% O2 and 1.3% H2 O, the equilibrium composition is 240 ppm Cl2 and 1520 ppm HCl [69]. For some CVOCs, there is insufficient hydrogen in the reactant to form HCl as the only chlorinated oxidation product; an example is the catalytic oxidation of trichloroethene forms some elemental chlorine: ClCH=CCl2 + 2O2 −−−→ 2CO2 + HCl + Cl2

(2)

Equilibrium calculations for this reaction show that, although complete conversion is possible at typical conditions, some elemental chlorine is present at equilibrium [54]. (The reverse Deacon reaction can produce HCl at high temperatures if sufficient water vapor is present.) Figure 5 shows the equilibrium composition of a mixture of 1000 ppm trichloroethene in air at typical conditions (including ambient levels of humidity). HCl is the thermodynamically favored product at high temperatures, but significant elemental chlorine is present at all temperatures above about 200 ◦ C. This suggests that for CVOCs with low H/Cl ratios, the potential for toxic byproducts such as Cl2 must be recognized. Olefins are another potential byproduct of CVOC oxidation. This is because paraffinic CVOCs often react via a hydrodechlorination path in oxidative conditions

0.020

mol-%

0.015

2HCl(g) + 12 O2(g) −−−→ H2 O(g) + Cl2(g)
G298 = −38.0 kJ mol−1 ;
H298 = −57.2 kJ mol−1

H2O (g)

0.010

(1) 0.005

VOC-containing gas

0.000

Reactor

Heat exchanger

0

HCl (g)

CO2 (g)

Cl2 (g)

100

200

300

400

500

600

700

Temperature / °C Vent

General process configuration for VOC catalytic oxidation systems.

Equilibrium composition for a trichloroethylene (TCE)/ humid air mixture; starting composition: 1000 ppm TCE; 21% O2 , 78% N2 , 1% H2 O. (Data calculated using HSC Chemistry 4.1; Outokumpu Oy, Pori, Finland.) CO and CO2 concentrations are negligible at all temperatures.

Fig. 5

Fig. 4

References see page 2408

2398

11.5 Solid Catalysts for the Oxidation of Volatile Organic Compounds

(and in the presence of acidic supports such as Al2 O3 or silica/alumina) to produce HCl and the corresponding olefins [70, 71]. These olefins, though not unusually toxic, can oligomerize to produce coke on the catalyst surface.

Mixture Effects Many VOC-containing gases of practical interest are multicomponent mixtures. Because of competitive adsorption and temperature effects, the rates and selectivities for individual compounds are not generally predictable from single-component studies. The rate of oxidation for a specific VOC can be either enhanced [6] or inhibited [73] in a mixture. For example, Taylor et al. showed that benzene oxidation over a U3 O8 catalyst was enhanced in a mixture with chlorobenzene [6]. At >350 ◦ C, benzene conversion in a 1 : 1.3 mixture with chlorobenzene was more than 86%, whereas the conversion of benzene as a single component at these same conditions was below detection limits. This is attributed to an increase in temperature in the catalyst bed once chlorobenzene oxidation was initiated. In contrast, Gangwal et al. showed that n-hexane conversion is suppressed by benzene in a binary mixture over 0.3% Pt–3% Ni/alumina when both n-hexane and the benzene +n-hexane inlet concentrations are the same (thereby minimizing temperature differences between these results), as shown in Fig. 6 [60]. However, benzene conversion was not affected by the presence of n-hexane. This appears to be due to preferential adsorption of benzene on the catalyst. The results are interpreted using a multicomponent Mars–van Krevelen model in which adsorption of the two reactants occurs on different sites. Similar anomalous mixture effects have been reported: for example, benzene oxidation is suppressed in the presence of butanol on 0.3% Pt/γ -Al2 O3 [74], 2-propanol preferentially displaces methyl ethyl ketone but does not affect toluene oxidation on a Pt-coated metal monolith [75], CO in the reactant gas inhibits oxidation of benzene:toluene:1-hexene mixtures on Pt, but preferentially enhances oxidation of the aromatics on Rh [76]. Together, these results demonstrate the importance of experimental study of VOC mixtures typical of industrial practice.

100 90 Percent conversion of n-hexane

Higher Temperatures Another problem with CVOCs is that higher temperatures are normally required for complete conversion than for corresponding hydrocarbon or oxygenated VOCs. In addition to the possibility of byproduct formation and deactivation, these high temperatures require more auxiliary fuel to heat the gas stream [72], adding significantly to the operating costs of the control system [45–48]. 11.5.3.1.3

110

80 70 60 50 40 30 20

11.5.3.2

10 0 140

180

220

260

300

340

380

Temperature / °C

Effect of benzene on n-hexane conversion. Reaction conditions: 209 WHSV, concentrations in air: () 410 ppm n-hexane; (•) 190 ppm n-hexane plus 189 ppm benzene. (From Ref. [60].)

Fig. 6

Deactivation Catalyst deactivation in the oxidation of VOCs has been reviewed [77], and is of critical practical importance for several reasons. First, because deactivation leads to loss of overall VOC conversion with time, eventually emissions will exceed design (and perhaps regulatory) limits and the catalyst must be replaced. Replacement costs for the catalysts, particularly noble metal-based catalysts, can be significant. On these noble metals, deactivation may be more rapid for smaller metal clusters [10], which are otherwise desirable to lower the total metal loading and therefore the cost. The second reason is that toxic byproducts can sometimes be formed even when the overall conversion of the target VOC is complete. An example is the formation of toxic phosgene and elemental chlorine from C2 -chlorinated VOCs [56]. The third reason is that because the operating temperature is often raised to compensate for deactivation, the processes leading to deactivation can be unintentionally accelerated if the activation energy for the deactivation process is greater than that of the main reaction(s) (e.g., oxidation). Butt et al. have provided a simple analysis of this phenomenon [78]. Assuming the deactivation rate is given by: 11.5.3.3

−rd =

ds = kd s n dt

(3)

11.5.4 Noble Metal Catalysts

where rd = the rate of deactivation (in s−1 ), s is the activity variable (rate of main reaction at time t/rate of main reaction at time zero), n is the order of deactivation reaction, kd is the deactivation rate constant (in s−1 = Ad exp(−Ed /RT )), and Ed is the activation energy of deactivation (in kJ mol−1 ). If the temperature is raised so that the rate of the reaction for the target VOC is kept constant, the rate constant for the deactivation reaction becomes: kd = (ko /s) = Ae(−E/RT )

(4)

where ko is the rate constant at t = 0 (and T = To ) (in s−1 ), A is the pre-exponential factor for the main reaction (in s−1 ), and E is the activation energy for the main (e.g., oxidation) reaction (in kJ mol−1 ). From the above, and assuming that n = 1, it can be shown that the time t corresponding to a particular activity s(s < 1) is given by: t = (E/Ad Ed )e(Ed /RTo ) [1 − s (Ed /E) ]

(5)

This expression shows that if E/Ed < 1, the time for the activity to decrease to a particular activity s will be less than if E/Ed > 1; that is, the catalyst will deactivate more rapidly. This increase in temperature also adds to the cost of fuel needed to heat the incoming gas stream. This auxiliary fuel cost is typically a large portion of the annualized operating costs, perhaps 40 to 70% [45–47]. An energy balance has been used to develop an expression for the auxiliary fuel flow rate required to maintain constant conversion as the catalyst deactivates as a function of time [45]. Deactivation is discussed in more detail below for specific catalysts. 11.5.4

Noble Metal Catalysts

Catalysts for the oxidation of VOCs can be divided into supported noble metals and metal oxides. As a general rule, the literature shows that supported noble metals show higher activity and selectivity to carbon oxides (for hydrocarbon VOCs) than metal oxides [6, 10]. However, these catalysts are significantly more expensive and tend to deactivate more rapidly in the presence of chlorinated VOCs [6, 10, 54, 79]. The two supported noble metals that have been most widely studied for VOC oxidation are Pt and Pd. Both have inherently high total oxidation activity, but studies show that the synthesis conditions and the nature of the support can affect the activity.

11.5.4.1

2399

Pt catalysts

Effect of Synthesis Conditions Typically, these catalysts are prepared by impregnation followed by calcination. The resulting metal oxide is then reduced before testing which, as studies have shown, results in a more active [80] and dispersed [81] catalyst. This reduction can be complete at temperatures as low as 160 ◦ C [68]. However, the working catalyst may not be in the reduced state. Studies using X-ray photoelectron spectroscopy (XPS) have shown that PtIV is likely the active site for the oxidation of toluene in air [68]. The catalyst precursor can also affect the activity. Chloroplatinic acid, H2 PtCl6 , is often used to prepare supported Pt catalysts. However, chlorine removal during calcination is not always complete. This residual chlorine has been shown to inhibit oxidation activity [61], suggesting that other precursors should be considered for supported Pt catalysts used for VOC oxidation. Although the reduced Pt is usually well dispersed following reduction, the particle size may change with time on stream and this, in turn, can affect the activity. In one study, an increase in activity with time on stream was attributed to a corresponding increase in Pt cluster size in the oxidation of o- and m-xylene on Pt/carbon aerogel [82]. 11.5.4.1.1

11.5.4.1.2 Support Effects Support effects are important, especially hydrophobicity of the support. In a typical VOC-air stream, water vapor is present in concentrations found in ambient air (∼104 ppm), which is often 10- to 100-fold greater than the concentration of the VOC. If the support is hydrophilic, high surface concentrations of water vapor may spill over to the metal, limiting the activity of the catalyst [83]. Water is also produced during the oxidation reaction, and this provides another source of undesirable water on the catalyst surface. The effect of surface water is accentuated at low temperatures, which otherwise are desirable for economic reasons. Chuang et al. [84] compared the oxidation of benzene/toluene/xylene, methanol and formaldehyde over Pt supported on hydrophilic and hydrophobic supports. These authors found that complete oxidation could be achieved at lower temperatures on hydrophobic supports. Similar results were found by Wu and Chang for the oxidation of toluene in air over Pt supported on hydrophobic styrene-divinyl benzene and hydrophilic activated carbon [68]. These authors propose a Mars–van Krevelen mechanism to account for their results, and use it to explain the effect of support hydrophobicity. Specifically, the mechanism can be summarized in two steps References see page 2408

2400

11.5 Solid Catalysts for the Oxidation of Volatile Organic Compounds

(HC = hydrocarbon):

800 Pt/TiO2 k0

O2 + reduced sites −−−→ oxidized sites

(6)

600 ppm

kHC

HC + oxidized sites −−−→ reduced sites + H2 O + CO2

13

CO2

CO2

400

(7)

11.5.4.1.3 Pt-Based Catalysts: Catalyst Deactivation Although not normally used for the oxidation of CVOCs because of concern about deactivation, Pt-based catalysts with a relatively high loading (2.15% Pt), supported on an acid-resistant support (α-Al2 O3 ) have been shown to be effective for oxidation of cyanogen chloride (CNCl), a

Tab. 1

Kinetic constants for toluene oxidation [68]

Temperature/ ◦ C

130 140 150

Pt/styrene-divinyl benzene

Pt/carbon

kHC

k0

kHC /k0

kHC

k0

kHC /k0

19.0 24.4 26.3

0.023 0.042 0.065

826 573 399

7.43 8.55 19.4

0.0197 0.0334 0.0592

377 256 327

200 13

(a)

EAc

0 400

Pt/TiO2(W4+) 13

300 ppm

From their data, values of k0 and kH C can be estimated for the two catalysts: Pt/styrene-divinyl benzene and Pt/carbon. The data listed in Table 1 show that the absolute rate of hydrocarbon oxidation, kH C , and the ratio kH C /k0 is greater for the hydrophobic styrenedivinyl benzene support. These authors attribute this to the rapid desorption of water vapor, presumably spilled over from the metal site in Eq. (7), from the styrene-divinyl benzene surface. This effect is especially noticeable at low temperatures. However, the Pt cluster size increases significantly with time on stream for the Pt/styrene-divinyl benzene, but not for Pt/carbon. As larger clusters may be more active [69], the higher activity of the Pt/styrene-divinyl benzene may not be due entirely to the hydrophobicity of the support. There is apparently an optimal acidity: high acidity leads to inhibition by water vapor [70, 85], and low acidity can limit adsorption of some VOCs onto the support, which lowers oxidation activity. This is the case if the mechanism is Langmuir–Hinshelwood – that is, if both oxygen and the VOC adsorb on the catalyst and then react. Figure 7 compares Pt supported on (a) TiO2 and (b) Wpromoted TiO2 (which has the greater density of acid sites, as measured by NH3 TPD) [86]. The oxidation of ethyl acetate to CO2 was complete at a lower temperature on the W/TiO2 , with this higher activity attributed to the higher density of acid sites.

EAc

13

CO2

CO2

200 Acetaldehyde 100 0

0

100

(b)

200

300

400

500

600

Temperature / °C

Temperature-programmed oxidation following adsorption of labeled ethyl acetate, 13 CH13 3 COOCH2 CH3 , showing oxidation to CO2 at lower temperatures on the more acidic Pt/TiO2 (W4+ ). (From Ref. [86]; Note different scales on y-axes.)

Fig. 7

chemical warfare agent [87]. Catalytic hydrolysis is a major reaction pathway [88]; conversion at 375 ◦ C was 98% in the presence of ambient levels of water vapor, but was only 20% at 440 ◦ C without water vapor. There was also a difference in selectivity: in the presence of water vapor, the products were CO2 and HCl, whereas in dry conditions, significant concentrations of CO and Cl2 were produced. 11.5.4.2

Pd catalysts

11.5.4.2.1 Effect of Reduction Several studies of Pdbased catalysts suggest that reduced Pd0 is more active for VOC oxidation than PdO [89], despite the fact that PdO is generally recognized as being more active for the oxidation of methane [90]. As an example, Ihm et al. studied the oxidation of 250 ppm n-hexane in air over a 5% Pd/Al2 O3 catalyst pretreated in (a) hydrogen, and (b) air [91]. The reduced catalyst was primarily Pd0 , and the air-treated catalysts contained mostly PdO. The reduced catalyst was far more active (Fig. 8), and did not oxidize with time on stream at 180 ◦ C, as shown by post-run XPS and X-ray diffraction (XRD). Thus, reduced Pdo is the active species during the reaction.

11.5.4 Noble Metal Catalysts

2401

100

80

60

40

5% Pd(R)/Al2O3 5% Pd(O)/Al2O3

Intensity of CO2 / A.u.

Conversion / %

TR = 180 °C

20

TR = 280 °C

0 160

200

240

280

320

360

Temperature / °C

Conversion versus temperature for reduced (at 500 ◦ C in H2 [5% Pd(R)/Al2 O3 ]) and oxidized (at 500 ◦ C in air [5% Pd(O)/Al2 O3 ]) catalysts for the oxidation of 250 ppm n-hexane in air. (From Ref. [88].)

100

200

Fig. 8

These results are consistent with those of Cordi and Falconer [76], who also showed that, at temperatures above 330 ◦ C, Pd supported on Al2 O3 goes through a redox cycle in which oxygen is adsorbed into the Pd lattice and is then reduced by the VOC via a Mars–van Krevelen mechanism. At lower temperatures, the VOCs adsorb on the Al2 O3 support and diffuse to the Pd site, where they are oxidized by PdO. These lower reaction temperatures also lead to carbon deposition. Figure 9 shows the TPO of the reduced Pd catalyst after n-hexane oxidation at 180 ◦ C and 280 ◦ C: substantial carbon is formed at 180 ◦ C, but none at 280 ◦ C [78]. Although the effect of water vapor was not specifically studied, it can be postulated that water vapor would compete with the VOCs for adsorption sites at lower temperatures, inhibiting the reaction. 11.5.4.2.2 Effect of Promoters Promoted Pd catalysts have also been reported. Compared to umpromoted Pd/TiO2 and V/TiO2 , vanadium-promoted Pd/TiO2 had the highest oxidation activity for benzene and naphthalene [92]. However, there is a difference in the reactivity of these two VOCs on the promoted catalyst. Only CO2 was produced in benzene oxidation, but the oxidation of naphthalene was incomplete at temperatures below about 300 ◦ C. These differences are attributed to the difference in nucleophilic vanadium sites and electrophilic Pd sites, with the latter leading to total oxidation products.

300

400

500

600

Temperature / °C

Temperature-programmed oxidation of Pd/Al2 O3 after n-hexane oxidation at reaction temperatures (TR ) of 180 ◦ C and 280 ◦ C. (From Ref. [89].)

Fig. 9

11.5.4.2.3 Pd Catalysts: CVOC Oxidation Deactivation is also a concern for Pd-based catalysts as it is for Pt-based materials in CVOC oxidation [93, 94]. In general, results show that Pd is more active than Pt but less selective to CO2 [95–98], although no long-term studies are available on Pd-based catalysts. Available studies also show the importance of support acidity. For example, L´opez-Fonseca et al. claim that support acidity is the primary catalyst property in predicting the oxidative decomposition of CVOCs [59], and this has been widely confirmed in other studies [99–104]. In particular, high acidity seems to inhibit selectivity to the undesirable product Cl2 in favor of HCl [105]. A direct comparison of Pt and Pd, both at 0.7 wt.% loading, and both supported on (a) H-ZSM-5 and (b) HBETA for the oxidation of dichloromethane (DCM) and trichloroethene (TCE), showed significant differences among the two metals and two supports [106]. H-ZSM5 had slightly lower total acidity (as measured by NH3 TPD) than H-BETA (0.49 versus 0.63 mmol NH3 g−1 ), but more strong acid sites (as measured by NH3 TPD above 250 ◦ C). These are strong adsorption sites for CVOCs [107]. Complete conversion of DCM and TCE required temperatures above 450 ◦ C on all catalysts. Because of the low H/Cl ratio in TCE, HCl selectivity was always lower than for DCM at comparable conditions, and decreased for all catalysts as the temperature increased References see page 2408

2402

11.5 Solid Catalysts for the Oxidation of Volatile Organic Compounds

from 350 to 550 ◦ C. Selectivity to CO2 was 100% within experimental error for all Pt catalysts, but always less than complete on the Pd catalysts. The results suggest that the strong acid sites of H-ZSM-5 increase selectivity to HCl for both Pt and Pd. Interestingly, toxic tetrachloroethene was formed from TCE on all catalysts (presumably from the direct chlorination of TCE) in concentrations up to 420 ppm. This had been observed previously with similar catalysts by Shaw et al. [108]. 11.5.5

Metal Oxide Catalysts

Metal oxide catalysts are, in general, less active than supported noble metals, but are more resistant to poisoning, especially by chlorine [10]. The most common oxidation catalysts are based on V [79] Ce [104, 105], Mn [109], Cr [66], Cu [110], Co [111, 112], and results are also reported for Ni [113], W [114], and U [6]. Vanadium-Based Catalysts Catalysts containing vanadium alone are not typically active for VOC oxidation. They tend to produce partial oxidation products at temperatures of interest here [79]. Figure 10 shows results for four V/TiO2 catalysts, prepared with different vanadium loadings, for the 11.5.5.1

Conversion C8H10 / %

120 100

60 40 20 0 100

150

(a)

Ceria-Based Catalysts Although ceria is often used as a support, it has been studied for VOC oxidation without added metals [117, 118]. In fact, ceria can be more active than noble metals in the oxidation of mono-aromatics such as toluene [119]. Garcia et al. [104] compared non-crystalline ceria, prepared using various methods, with 0.5% Pt/Al2 O3 and MnO2 , both of which are active oxidation catalysts. Ceria

200 250 Temperature / °C

300

80 70 60 50 40 30 20 10 0

1.0 Destruction efficiency / h

Yield CO2 / %

11.5.5.1.1 V-Based Catalysts: CVOC Oxidation A recent review of the literature on vanadia-based catalysts for the oxidation of CVOCs shows that most studies deal with SCR-type V2 O5 /WO3 /TiO2 , which appears to be relatively stable to HCl and Cl2 [54]. This catalyst has been used for oxidation of chlorocarbons [54], and for the oxidation of chlorinated dioxins in incinerator waste gases [80, 115]. In this latter case, the V2 O5 /WO3 /TiO2 SCR catalysts perform two functions: the reduction of NOx and the oxidation of dioxins [54]. There are clear differences in the reactivity of various CVOCs over these catalysts. The ‘‘destruction efficiency’’ (equal to conversion) for the aromatics increases with the degree of chlorine substitution at a given temperature [54]. This is consistent with kinetic analysis, which showed a decrease in activation energy and an increase in adsorption with chlorine content [116]. For perhaps the same reasons, chlorinated alkanes are more readily oxidized on these catalysts than the corresponding chlorinated alkenes, as shown in Fig. 11 [54]. 11.5.5.2

80

100 (b)

oxidation of naphthalene [79]. At conditions where naphthalene conversion is complete, the CO2 yield is far from complete. Analysis of the byproducts showed that a range of aromatics and oxygenates were present in the gaseous product and were also deposited on the catalyst.

150

200

250

300

Temperature / °C

Conversion and CO2 yield for oxidation of naphthalene over V/TiO2 with various vanadium loading:
, 0.5%; , 1.0%; , 1.5%, •, 3.0%. (From Ref. [79].)

0.8

TCA TCE

0.6 0.4 0.2 0.0 250

300 Temperature / °C

350

Fig. 10

Fig. 11 Destruction efficiency of trichloroethylene (TCE) and 1,1, 1-trichloroethane (TCA) over V2 O5 /WO3 /TiO2 . (From Ref. [54].)

11.5.5 Metal Oxide Catalysts

prepared by precipitation of (NH4 )2 Ce(NO3 )6 with urea was more active for the oxidation of naphthalene than either 0.5% Pt/Al2 O3 or MnO2 , and more selective to CO2 . Other methods of preparing ceria (e.g., precipitation of cerium nitrate) did not yield a similarly active catalyst. The authors ascribed the difference to the higher naphthalene adsorption for the non-crystalline, urea-precipitated sample. At temperatures below those required for total oxidation, ceria can adsorb naphthalene on the surface at low temperatures, without oxidizing it. Garcia et al. point out that this may be important in the operation of a practical VOC oxidation system [104]. For example, ceria can store the naphthalene during start-up (or periods of time when the concentration fluctuates) and then oxidize it simply by raising the temperature. It is not clear that this mechanism would work for other VOCs, or whether water vapor would adversely affect the storage process. However, this type of practical consideration is important in evaluating catalysts for VOC removal. Manganese-Based Catalysts Mn-based materials have long been recognized as active catalysts for the oxidation of CO and VOCs [120]. They can be alloyed with other metals such as Cu [137], and incorporated into structures such as perovskites [121–123], although supported MnO2 is also an active VOC oxidation catalyst [124]. In general, Mn-based catalysts tend to be structuresensitive for VOC oxidation [125]. For example, a direct comparison of bulk Mn2 O3 and an Mn-containing perovskite (SmMnO3 ) showed that the perovskite was more active for VOC oxidation and more resistant to deactivation [108]. Lamaita et al. [126, 127] prepared substoichiometric MnOx (x < 2) by two different methods: decomposition of MnCO3 and oxidation of MnSO4 . Compared to MnO2 and Mn2 O3 , they found that the higher activity of these materials for ethanol oxidation was due to the presence of both Mn3+ and Mn4+ ions and OH− species generated by Mn4+ vacancies. Mn oxides can also form molecular sieve-type structures. One such material, cryptomelane, forms one-dimensional channels similar to mordenite. Interestingly, this material has labile lattice oxygen that makes it active for the oxidation of VOCs, such as ethanol and benzene [128]. It is also hydrophobic, and results show that it resists inhibition in the presence of water vapor, preferentially adsorbing benzene in the presence of ambient water vapor concentrations [115]. 11.5.5.3

11.5.5.3.1 Mn-Based Catalysts: CVOC Oxidation Two forms of Mn-based catalyst have been reported for CVOC oxidation: Mn-based perovskites [56, 129], and zeolitesupported catalysts [117]. Both show complete conversion

2403

of C1 −C2 CVOCs between 450 and 550 ◦ C, and these studies show the importance of water vapor in the reactant gas and the support acidity. Effect of Water Vapor The effect of water vapor on the reaction pathway has been studied for a series of three chloromethanes with increasing chlorine content (CH2 Cl2 , CHCl3 , CCl4 ) on LaMnO3 perovskite [56]. With ambient levels of water vapor in the reactant gas (1.3%), the reaction pathway for CH2 Cl2 and CHCl3 appears to be both hydrolysis and oxidation, whereas CCl4 reacts only by hydrolysis. CCl4 formed only CO2 and HCl whether there was oxygen in the reactant gas or not, presumably via the following reaction:

CCl4 + 2H2 O −−−→ CO2 + 4HCl

(8)

However, CH2 Cl2 conversion shows a positive dependence on both water and oxygen concentrations, suggesting that both hydrolysis and oxidation reactions are significant. CHCl3 behavior is intermediate between CH2 Cl2 and CCl4 . Byproducts Byproducts of the oxidation of these CVOCs include CO, Cl2 , and traces of phosgene (COCl2 ). (Interestingly, the reaction of CHCl3 also produces CCl4 .) High selectivity to CO (versus CO2 ) is observed in the oxidation of these compounds at temperatures where the conversion of the CVOC is incomplete; these results suggest that HCl is the primary product of the reactions of CH2 Cl2 and CHCl3 , and that Cl2 is formed by the Deacon reaction (in agreement with Guit´errez-Ortiz et al. for Mn–H-ZSM-5 [130]):

2HCl(g) + 12 O2(g) −−−→ H2 O(g) + Cl2(g)
G298 = −38.0 kJ mol−1 ;


H298 = −57.2 kJ mol−1

Phosgene, COCl2 , is a toxic byproduct and is formed in trace amounts in the oxidation of CCl4 only. The presence of water vapor greatly limits its concentration in the product gas, presumably since it is readily hydrolyzed: COCl2 + H2 O −−−→ 2HCl + CO2

(9)

The presence of water vapor greatly affects the product distribution. Figure 12 [56] shows that CCl4 is actually produced in the catalytic oxidation of CH2 Cl2 with no added water vapor in the reactant gas, but not when water vapor is present. In this same experiment [56], CHCl3 is produced from CH2 Cl2 whether water vapor is present, or not, although much greater concentrations are seen with no added water References see page 2408

2404

11.5 Solid Catalysts for the Oxidation of Volatile Organic Compounds

do not permit a direct comparison of LaMnO3 [56] and 4.3% Mn−HZSM-5 [117], results for the oxidation of 1000 ppm CH2 Cl2 suggest that these catalysts have comparable activity. However, LaMnO3 produces up to 150 ppm CHCl2 and 60 ppm CCl4 , while no such products are reported for the 4.3% Mn−HZSM-5. In both cases, Mn shows significant activity for the Deacon reaction, leading to Cl2 in the product gas.

180 160 ppm by products

140 120 100 80 60 40 20 0 200

Chromia-Based Catalysts Most of the literature on chromia-based catalysts deals with the oxidation of CVOCs, because these catalysts tend to be resistant to deactivation. The relatively few studies on non-chlorinated VOC oxidation show that some forms of chromia have remarkably high activity. For example, chromium oxide nanocrystals synthesized as bulk xerogels and aerogels have shown higher activity than supported Pt. Specifically, the activity of the chromia aerogel (α-CrOOH, 630 m2 g−1 ) was fourfold higher than 0.5% Pt/Al2 O3 and 30-fold higher than 30% Cr2 O3 /SiO2 for the oxidation of ethyl acetate [131]. Cr2 O3 /Fe2 O3 also showed higher activity than bulk Cr2 O3 in the oxidation of acrylonitrile, acetic acid, and cyanic acid in a reverse-flow reactor [132]. 11.5.5.4

250

300

350

400

450

500

550

Temperature / °C

Concentration (ppm) of CCl4 ( ) and CHCl3 () byproducts versus temperature in the catalytic oxidation of 1000 ppm CH2 Cl2 in absence (- - - - - -) and presence of 1.3% ( – ) water on LaMnO3+δ . (From Ref. [54].)

Fig. 12

vapor. These results suggest that CHCl3 is formed first (note the lower temperature at which it is formed in the absence of water), and that both of the CCl4 and CHCl3 byproducts are subsequently converted to final products (HCl and CO2 ) at temperatures above ∼525 ◦ C. Oxidation of CH2 Cl2 over Mn−ZSM-5 in the absence of added water vapor produced only trace levels of methyl chloride, CH3 Cl, as a byproduct, suggesting a different reaction path on this catalyst [117]. Effect of Support Acidity Oxidation of CH2 Cl2 on 4.3% Mn−HZSM-5, 4.3% Mn/Al2 O3 , and 4.3% Mn/SiO2 shows that the oxidation activity for a series of Mnsupported catalysts correlates directly with support acidity and metal dispersion [117]. The results show that CH2 Cl2 adsorbs on the acid sites of the support and oxygen is activated by Mn, with more dispersed clusters of Mn showing greater oxygen chemisorption capacity. Chlorine Retention Deactivation due to chlorine retention on the catalyst is always possible in CVOC oxidation. Sinquin et al. showed that chlorine is not incorporated into LaMnO3 during CH2 Cl2 oxidation, but that it is retained within the structure of a directly comparable LaCoO3 perovskite, especially at temperatures where conversion is incomplete, ∼350 ◦ C in this case [56]. Interestingly, the presence of water vapor seems to limit chlorine deactivation on both catalysts, presumably by removing it in the form of HCl. Zeolite-Supported Mn versus Mn Perovskites Although different experimental conditions (e.g., space velocities)

11.5.5.4.1 Cr-Based Catalysts: CVOC Oxidation Catalysts based on supported chromium oxides have been shown to be active for the oxidation of a series of C1 −C2 CVOCs and chlorobenzene [133–143]. Various forms of supported chromium oxide are used commercially for the destruction of CVOCs [122, 129]. A comparison of supported Cr catalysts to other supported metals suggests that chromium oxides are among the most active catalysts for CVOC oxidation. For example, CrOx showed higher trichloroethylene conversion than MnOx , Pt, CoOx , CuOx , FeOx , or Ni when all were supported on TiO2 [144]. However, several studies point out the potential for the formation of toxic Cr2 O2 Cl2 [127, 130] or vinyl chloride [145], suggesting that care must be taken to quantify even trace levels of byproducts.

Effect of Support As with other supported catalysts for CVOC oxidation, the support plays a crucial role. For Cr-based catalysts, this appears to be most closely related to the achievable loading on a given support. Activity increases with loading up to ∼12 wt% [121, 146], but activity decreases at higher loadings because of the formation of an inactive bulk Cr2 O3 phase [121, 128]. Strong metal–support interaction on supports such as TiO2 and Al2 O3 seems to stabilize the active Cr(III) and Cr(IV) ions, however [121, 147]. Recent studies point out the advantages of a bulk (rather than supported) Cr-based catalyst for CVOC

11.5.5 Metal Oxide Catalysts

oxidation: less potential pore blocking of the support at the high metal loadings required for an active catalyst (12–20%), and the necessity for strong metal support interaction to stabilize the active phase of Cr [148, 149]. Rotter et al. prepared an unsupported, bulk nanostructured CrOOH aerogel (500 m2 g−1 ) and tested it for the oxidation of dichloroethane and chlorobenzene, and compared it to this same material containing Pt and other metals, as well as to supported Cr2 O3 and Pt/Al2 O3 [66]. The unpromoted CrOOH was more active for dichloroethane oxidation than Pt/Al2 O3 and supported Cr2 O3 , and comparable to Pt/CrOOH. These results suggest the advantages of bulk aerogels as a means of improving specific activity of Cr-based catalysts. Figure 13 [150] shows that there are substantial differences in CVOC oxidation rates on Cr-based catalysts. CH2 Cl2 was the most refractory on a commercial 9% Cr2 O3 /alumina catalyst [137]. The apparent activation energies of all four CVOCs in this study were similar (85–105 kJ mol−1 ), consistent with other results [48], and suggesting no significant difference in mechanism. The lower rate for CH2 Cl2 can therefore be attributed to a lower surface coverage in this four-component mixture, which also contained 1.6% water vapor. Deactivation Tests on a commercial Cr-alumina catalyst show that >99% total conversion of mixture of C2 chlorocarbons (total concentration 500 ppm) could be maintained for 3600 h at 350 ◦ C [137, 151]. However, the conversion of CH2 Cl2 dropped from >99% to 93% and CO selectivity increased from 32 to 54% over this same time period, suggesting a change to the catalyst over this extended period of time. XPS analysis showed no change

100

Conversion / %

80 60 40 20 0 225

250

275

300

325

350

375

Temperature / °C Fig. 13 Conversion of individual CVOCs in a mixture over 9% Cr2 O3 versus temperature. Total CVOC concentration = 200 ppm in humid air (dew point = 21 ◦ C), 23 970 h−1 . •, dichloromethane; , 1,2-dichloroethane; ◦, trichloroethylene; , 1,1-dichloroethylene; , total. (From Ref. [150].)

2405

in the Cr 2p binding energy, but there was a decrease in bulk Cr content, from 9.2% to 8.3%, suggesting loss of Cr, perhaps as Cr2 O2 Cl2 [127, 130]. Copper-Based Catalysts Supported Cu is not typically considered to be an active total oxidation catalyst. Although often used as a promoter or cocatalyst with Pt [152], Fe [139], or Mn [153], recent studies have shown that Cu can be used alone for the total oxidation of VOCs [154]. Studies comparing supported Cu, Mn, Fe, V, Mo, Co, Ni, and Zn (each at 15 wt.% loading) for the complete oxidation of toluene [155] and a benzene/toluene/xylene mixture [141] showed that 5% Cu/γ -Al2 O3 was the most active. Smaller, more disperse Cu clusters were formed on this support compared to TiO2 and SiO2 . However, these supported Cu catalysts are far less active than Pt for the oxidation of toluene: complete oxidation was reached at ∼120 ◦ C for Pt/styrene-divinyl benzene [68], but only at ∼260 ◦ C for Cu/γ -Al2 O3 [141]. 11.5.5.5

11.5.5.5.1 Cu-Based Catalysts: CVOC Oxidation There are few reports of Cu-based catalysts for the oxidation of CVOCs. CuCl has been shown to be an active catalyst for the oxidation of CH2 Cl2 at temperatures as low as 400 ◦ C [156]. However, it is hydrolyzed in the presence of water vapor and deactivates rapidly, which makes it impractical for commercial application.

Cobalt-Based Catalysts A direct comparison of Co to Mn, Cu, and Zn for n-C4 and n-C6 oxidation at a loading of 5% on an activated carbon and two ion-exchange resins showed that the activities of supported Co and Mn were comparable, and both were more active than Cu or Zn [98]. However, the comparison between Mn and Co was dependent upon the support. On an activated carbon support, cobalt was more active for n-C4 oxidation: at 250 ◦ C, n-C4 conversion of 5% CoO/carbon was 99 + % while on 5% MnO/carbon it was 94%. However, the carbon support was apparently not stable at these conditions. It seems to have oxidized, losing 12.5% weight over 45 h at 250 ◦ C. However, on Ambersorb supports (ion-exchange resins), the reactivity of Mn and Co are reversed. For n-C6 oxidation, 99.9% conversion is reached at 175 ◦ C on the 5% MnO/Ambersorb catalyst, but a temperature of 200 ◦ C is required to reach 99.1% conversion on 5% CoO/Ambersorb . The water vapor content of the reactant gas is not provided, and it is possible that water vapor adsorbed on the more hydrophilic Ambersorb support affects the oxidation on the cobalt catalyst more than for the manganese catalyst. 11.5.5.6

References see page 2408

11.5 Solid Catalysts for the Oxidation of Volatile Organic Compounds

In this same study, longer-term runs showed substantial deactivation of the 5% CoO/Ambersorb ; the conversion of n-C4 at 200 ◦ C dropped from 99.1% to 78% over 50 h. A 368-h run on this catalyst showed that the activity for n-C6 oxidation could be restored by treatment in air at 300 ◦ C, although a gradual decline in activity appears inevitable. 11.5.5.6.1 Co-Based Catalysts: CVOC Oxidation Cobalt has been used both as a catalyst (supported on Al2 O3 ), and a cocatalyst (Au supported on Co3 O4 ), for the oxidation of CH2 Cl2 [112]. A coprecipitated 0.2% Au/Co3 O4 catalyst was claimed to consist of small Au particles on cobalt oxide, though no particle size characterization is reported. The addition of Au to cobalt oxide increased the activity significantly, as shown in Fig. 14, but there was little effect of Au loading above 0.2%. Complete conversion in humid air (0.6% H2 O) to CO2 and HCl is reported. The activity was independent of water vapor pressure over the range of 0.3 to 1.5%. However, even at 0.3%, the inlet water vapor concentration was 60-fold that of CH2 Cl2 (which was 500 ppm inlet), so hydrolysis may still have been important. A longer-term study (140 h) showed that the 5% Au/Co3 O4 catalyst was more active and more stable than supported Pt or Pd catalysts (Fig. 15). (No similar study was carried out for the 0.2% Au/Co3 O4 catalyst, so it is not known whether this lower-loaded catalyst would have similar stability.) During this test, no products other than CO2 and HCl were observed.

Nickel-Based Catalysts Like Cu, Ni is often used with other metals in VOC oxidation [157]. A direct comparison of Ni with Cu, Mn, Fe, V, 11.5.5.7

100

100

5% Au/Co3O4

80

Conversion / %

2406

60

0.5% Pt/Al2O3

40

20

0.5% Pd/Al2O3

T = 350°C GHSV= 90,000 h−1 100 ppm CH2Cl2 0.6 wt% H2O Balance air

0 0

30

60

90

120

150

Time / h Conversion versus time for 5% Au/Co3 O4 compared to supported Pt and Pt catalysts. (From Ref. [112].)

Fig. 15

Mo, Co and Zn (all at 15% loading on γ -Al2 O3 ) shows that Ni has very poor activity as the sole supported metal [141]. However, when alloyed with other metals, Ni can be an important component of a VOC oxidation catalyst. For example, Ni has been investigated as a cocatalyst with cobalt in spinel structures. Modification of NiCo2 O4 spinel with aluminum and potassium was investigated for the oxidation of toluene, acetone, ethanol, and acetic ether [100]. Results show that potassium addition increased activity, while aluminum addition decreased activity. The VOC oxidation activity correlated well with particle size, with smaller NiCo2 O4 particles being more active. Potassium also stabilized the spinel structure, making it more active than either the Al-modified or the parent spinel.

5% Au/Co3O4

Conversion / %

60

0.2% Au/Co3O4

1% Au/Co3O4

Co3O4

10% Au/Co3O4 40 GHSV=15.000 h−1 500 ppm CH2Cl2 0.6 wt% H2O Balance air

20

0 150

200

250

300

350

400

Temperature / °C Effect of Au loading on CH2 Cl2 oxidation. (From Ref. [112].)

Fig. 14

Tungsten-Based Catalysts Tungsten is perhaps most often used as a modifier for SCR-type V2 O5 /WO3 /TiO2 catalysts. WOx (x < 3) supported on activated carbon (AC) and TiO2 has been studied for the oxidation of toluene [114]. Tungsten oxide has acidic character [158], and forms a wide range of non-stoichiometric oxides depending on temperature and redox atmosphere [158–163]. The formation of a nonstoichiometric oxide was shown to be essential to the oxidation of toluene in air [114]. WOx supported on activated carbon was found to result in a more active catalyst than WOx on SiO2 , Al2 O3 or zeolite Y. This was attributed to the hydrophobicity of the AC support, which prevents water vapor adsorption. These catalysts are restricted to temperatures below about 400 ◦ C by the potential for gasification of the carbon. 11.5.5.8

80

11.5.6 Kinetics

Uranium-Based Catalysts Although not commonly thought of as an oxidation catalyst, uranium oxide has been studied for the oxidation of VOCs [6] and CO [164]. In particular, U3 O8 has uranium present in mixed oxidation states with facile transition between states, making it an active oxidation catalyst [6]. In experiments on the oxidation of a wide range of hydrocarbon VOCs, supported and bulk U3 O8 showed activity comparable to Co3 O4 [6]. Although conversions of a wide range of VOCs was >99% at temperatures from 300 to 600 ◦ C, selectivities to CO were relatively high (47% for the oxidation of methanol), suggesting that U3 O8 may require a cocatalyst for CO oxidation to be useful for complete VOC oxidation. Temporal analysis of products (TAP) experiments showed that oxidation takes place in the absence of gas-phase oxygen, suggesting that lattice oxygen from U3 O8 is active during the catalytic cycle. 11.5.5.9

Langmuir–Hinshelwood (L–H) Models Several studies on hydrocarbon VOCs [74, 115, 166] and chlorinated VOCs [54] show kinetic results that are consistent with a L–H mechanism. For example, Barresi and Baldi [166] developed a model based on the following simplified L–H mechanism for a series of aromatic hydrocarbon VOCs: 11.5.6.1

11.5.6

Kinetics

In addition to simpler power law fits of the data [54, 56, 61], kinetic models for VOC oxidation have been developed based on Langmuir–Hinshelwood [61, 102] and Mars–van Krevelen [68] mechanisms. Although some results have been interpreted using the Eley–Rideal mechanism [54], it has not been widely investigated for VOC oxidation and is not considered further here. In most of these models, the relatively high concentrations of oxygen in ambient air compared to the VOC lead to models that are zero-order in oxygen concentration [56, 61, 102], although at oxygen concentrations of a few percent this may not be the case [166]. Dependence on VOC concentration is generally positive order, although competitive adsorption effects in gases containing VOC mixtures can make this dependence near zero for strongly adsorbed species.

f ast

ko

O2 + ( )−−−→(O2 )−−−→2(O)

(10)

VOC + [ ] −−−→ [VOC]

(11)

kr

[VOC] + (O)−−−→P + ( ) + [ ]

(12)

where VOC is a hydrocarbon and P is the product (CO2 + H2 O in this case). The rate-determining step is the surface reaction [Eq. (12)]. For a strongly adsorbed hydrocarbon VOC, the model becomes: −rV OC =

11.5.5.9.1 U-Based Catalysts: CVOC Oxidation Studies show that, for chlorobenzene, U3 O8 activity is comparable to supported Cu, Pt, Co, and WO3 and showed little deactivation over 400 h on stream [6, 165]. In TAP experiments on CVOCs, no chlorine-containing products were observed, so that chlorine is likely retained in the catalyst. Although this might be expected to lead to eventual deactivation, conventional flow experiments showed that HCl is present in the gas-phase product and no deactivation was observed. Chlorine apparently accumulates on the surface until a steady state is reached. This steady-state level is not sufficient to cause observable deactivation over the time scale of these experiments, although eventual deactivation might be expected.

2407

k0 kr PO2 ϑV OC k0 PO2 + vkr ϑV OC

(13)

where PO2 is the partial pressure of O2 , ϑV OC is the fractional surface concentration of the VOC, and v is the stoichiometric coefficient of oxygen. For a strongly adsorbed VOC, ϑV OC = 1 and the rate becomes apparent zero order in the VOC. For high oxygen concentrations (as is true in most applications), the rate becomes first order in ϑV OC . This approach can be extended to n-component mixtures, with the following result [166]: ri,mix = 




ko ki PO2 n 

(bj Pj /bi Pi ) + vi ki  ko PO2 1 +  j =1;j =i   n   + vj kj (bj Pj /bi Pi )

     

j =1;j =i

(14) where pi is the partial pressure of the ith VOC, and bi and bj are the adsorption constants for the ith and j th VOC. Barresi and Baldi show that this model accounts for the competitive adsorption among five aromatic VOCs (benzene, toluene, xylene, ethyl benzene, and styrene), as well as the oxidation of individual components [153]. The effect of water vapor in the reactant gas was modeled by Stoll et al. for the oxidation of chlorobenzenes using an L–H model of the form [115]: −rV OC =

a pb ApO pc 2 V OC H2 O

(1 + BpO2 + CpH2 O )d

(15)

where A, B, and C are constants. Their results showed zero-order in water vapor (c = 0) over a range of 2 References see page 2408

2408

11.5 Solid Catalysts for the Oxidation of Volatile Organic Compounds

to 30%, suggesting strong adsorption of water vapor over this range. Interestingly, the reaction was first order in the VOC (b = 1), and 0.3 order in oxygen (B = 0 and a = 0.3 from 1 to 15%), suggesting that water vapor preferentially adsorbs compared to oxygen on their V2 O5 /WO3 /TiO2 catalyst, and inhibits the reaction significantly (d = 1.5–2.0). Mars–van Krevelen (M–vK) Model The M–vK model [167] has been widely applied to the oxidation of VOCs [81, 168–170]. It is conceptually based on elementary steps involving a redox cycle on the catalyst surface [81]:

Although this mechanism is based on binary mixtures, the concept has general applicability and the resulting expression for oxidation of mixtures of VOCs is as follows (vi is the stoichiometric coefficient for the oxidation of the ith VOC) [171]: ri,mix =

11.5.6.2

ko O2 + (reduced sites)−−−→(oxidized sites)

(16)

kV OC V OC + (oxidized sites)−−−→(reduced sites) + H2 O + CO2

(17)

The fundamental difference between this and the L–H mechanism is that the VOC reacts directly from the gas phase with an oxidized site, rather than adsorbing onto the catalyst and then reacting. The resulting rate expression is as follows (v is the stoichiometric coefficient of oxygen) [81, 170]: −rV OC =

ko kV OC pO2 pV OC ko pO2 + vkV OC pV OC

(18)

Since pO2 can be considered a constant (pO2  pV OC ), this equation reduces to [81, 170]: −

1 v 1 = + rV OC kV OC pV OC ko pO2

(19)

This model is consistent with experimental results on the oxidation of toluene, and could be used to explain support effects for supported Pt catalysts [81]. Specifically, a more hydrophobic support is claimed to enhance kV OC in Eq. (18) by removing water from the catalyst surface. The M-vK model has been extended to multicomponent mixtures, with the assumption that the catalyst is reduced to different surface species by different VOCs [171]: k1 V OC1 + (oxidized site)−−−→(reduced site 1) + H2 O + CO2 ko1 O2 + (reduced site1)−−−→(oxidized site)

(19a) (20)

k2 V OC2 + (oxidized site)−−−→(reduced site2) + H2 O + CO2 ko2 O2 + (reduced site2)−−−→(oxidized site)

(21) (22)

koi ki PO2 PV OC,i , i = 1, 2, 3 . . . . n k  oi koi PO2 + vi kj PV OC,j j =1 koj (23)

11.5.7

Summary

The oxidation of VOCs using solid catalysts has been shown for a wide range of catalysts and VOCs. Generally, supported noble metals are more active than metal oxides. Although these catalysts have been studied for the oxidation of chlorinated VOCs (CVOCs), they tend to deactivate at conditions of practical interest. For metal oxide catalysts, the surface properties of the support (acidity and hydrophobicity) play a significant role in the reaction rate, deactivation, and product distribution. Byproducts that are more toxic and hazardous than the reactant(s) can be formed from relatively innocuous VOCs, so the reaction products must be carefully monitored. Competitive adsorption in multicomponent mixtures can cause unexpected effects, and the results of single-component studies cannot generally be extended to multicomponent mixtures. Kinetic models based on Langmuir–Hinshelwood and Mars–van Krevelen mechanisms have been used to model reaction data. Under typical conditions, the high oxygen concentrations generally result in zero-order kinetics in oxygen, while dependence on VOC concentration can be from first to fractional orders. Water vapor generally inhibits the reaction, pointing to the need for hydrophobic supports. References 1. N. Mukhopadhyay, E. C. Moretti, Current and potential future industrial practices for controlling volatile organic compounds, Center For Waste Control Management, 1993. 2. M. S. Jennings, M. A. Palazzolo, N. E. Krohn, R. M. Parks, R. S. Berry, K. K. Fidler, in. Noyes (Ed.), Catalytic Incineration for the Control of Volatile Organic Compound Emission, Pollut. Technol. Rev. 121, 1985. 3. M. J. Molina, F. S. Rowland, Nature 1974, 249, 810. 4. A. Musialik-Piotrowska, B. Mendyka, Catal. Today 2004, 90, 139. 5. Chemistry in Britain, February, 1997. 6. S. H. Taylor, C. S. Heneghan, G. J. Hutchings, I. D. Hudson, Catal. Today 2000, 59, 249–259. 7. P. Hunter, S. T. Oyama, Control of Volatile Organic Compound Emissions, Wiley, New York, 2000.

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11.6

Catalytic Combustion Pio Forzatti∗ , Gianpiero Groppi, and Cinzia Cristiani

11.6.1

Introduction

Catalytic combustion can achieve stable and effective combustion over a wider concentration range and at lower temperatures than in conventional burners, based on gasphase radical chemistry. Because of this it is commonly used for the abatement of gaseous emissions in many manufacturing and energy-conversion processes. During the past few decades, catalytic combustion has also been explored as a primary control method in the production of heat and energy with ultra-low emissions of NOx , CO, and unburned hydrocarbons (UHC). Catalytic combustion in gas turbines (GT) has attracted particular interest, and today such turbines represent the preferred energy conversion technology in medium- and largescale power stations. Likewise, major opportunities are foreseen in the small-scale distributed power generation sector. NOx emissions are a major concern of GT systems. In the United States, emission regulations are based mainly on air quality standards fixed by federal laws. Although NOx standards are now usually attained, NOx is deemed to be a precursor of ozone, the standards for which are not agreed in several areas. In order to cope with these problems, a permit strategy is applied by local authorities which requires the adoption of a Best Available Control Technology (BACT) in attainment areas, or a Lowest Achievable Emission Rate (LAER) technology in non-attainment areas, typically in combination with emission fee/credit systems. A similar regulation system is currently applied in the European Union (EU). To date, several NOx control technologies for GTs have been developed, based on both primary and secondary methods. Dry low-NOx (DLN) burners based on lean premixed combustion guarantee NOx emission levels of References see page 2424 ∗ Corresponding author.

2412

11.6 Catalytic Combustion

between 20 and 25 ppm, though several GT manufacturers claim that the system can be designed and operated down to 9–10 ppm. Any further reduction may be precluded by flame-stability problems. In order to meet the most stringent emission regulations, many installations include a selective catalytic reduction (SCR) unit which allows a further reduction in NOx emission levels. Catalytic combustion has been demonstrated, on a commercial basis, to reduce NOx emissions below 3 ppm while maintaining CO and UHC emissions below 10 ppm, without the need for expensive exhaust cleanup systems. In addition, a catalytic combustor reduces typical DLN problems such as risk of blow-out and flame instability. A cost analysis commissioned by DOE [1] considered GT of small (5 MW), medium (25 MW) and large (150 MW) sizes, and compared the following technologies: water/steam injection; DLN; conventional, low-T and high-T SCR; SCONOX (a secondary control method based on NOx adsorbers); and catalytic combustion. The results clearly indicate the economic advantage of primary methods, including catalytic combustion as opposed to secondary clean-up measures (SCR and SCONOX). The disparity of cost impact is particularly large for small-scale gas turbines that are deemed most suitable for the distributed generation market which, in turn, is threatened by the most strict environmental regulations. Although the potential of catalytic combustion has been recognized for more than 30 years, only recently has this technology been proven commercially viable and finally marketed, albeit to a limited level. This chapter provides a review of the status and of the perspectives of catalytic combustion for GTs. Initially, details of the base concepts associated with the system requirements are presented, after which the design approaches based on lean and rich catalytic combustion are illustrated, together with reports of performances demonstrated in full-scale and field tests. Details of the most relevant characteristics of PdO-supported catalysts and

Fuel 300–400 °C

of transition metal-substituted hexaaluminates that have been most extensively considered for lean combustion applications are outlined, along with those of the noble metal catalysts adopted in rich combustion systems. Finally, brief details are provided of the use of mathematical modeling as a tool for the design and analysis of catalytic combustors, and the perspectives for this technology outlined. 11.6.2

Base Concepts and System Requirements

As illustrated in Fig. 1, in a conventional system a fraction of the air delivered by the compressor is mixed with the fuel, typically natural gas. The mixture is then combusted in a flame, such that the hot gas produced expands and drives the turbine. Flame stability requires adiabatic combustion temperatures to be as high as 1600–1800 ◦ C, but these must be reduced to 1100–1450 ◦ C, by means of cooling bypass air, before the hot compressed gas is delivered to the turbine, in order to avoid damage to the inlet blades. At such temperatures, and within the tens of milliseconds residence time required for complete burn-out of fuel and CO, significant amounts of NOx are produced, mostly via the Zeldovich thermal mechanism [2]. In a catalytic burner, the combustion is ignited and stabilized under ultra-lean conditions, which result in adiabatic temperatures close to those allowed for delivering the hot compressed gas to the turbine. Thus, the need for bypass air is minimized and the formation of thermal NOx is almost prevented due to an absence of a hot combustion zone. The reduction in NOx emission has been reported as being even larger than expected from the lower combustion temperature if a significant fraction of the fuel is oxidized on the catalyst surface [3]. This effect has been attributed either to a reduction in the formation of prompt NOx , in view of the decrease in CH radicals in the gas phase due to fuel complete oxidation on the Fuel

1600–1800 °C HC

1100–1400 °C

300–400 °C

By pass C

Catalyst segments 1100–1400 °C CC

T

C Output

Air (a) Flame combustion

Exhaust

T Output

Air

Exhaust

(b) Catalytic combustion

Comparison between conventional (a) and catalytic (b) gas-turbine systems. HC: homogeneous combustor, CC: catalytic combustor, C: compressor, T: turbine.

Fig. 1

11.6.3 Design Approaches

catalyst surface [4], or to a reduction in the formation of NOx due to the release of H2 O produced at the catalyst surface. Indeed, a beneficial effect of increased H2 O concentration on the reduction of NOx emission under typical GT operating conditions has been reported [5]. The operating constraints of GT systems (Table 1) pose severe requirements to the catalytic combustor. Air is delivered by the compressor at temperatures which typically range from 300 ◦ C to 450 ◦ C, depending on load conditions and the nominal pressure ratio of the machine. Upon accurate fuel/air mixing, ignition must occur readily and fuel conversion should proceed rapidly to completion, while the gas reactants heat up to the adiabatic combustion temperature. Due to high gas velocity (10–30 ms−1 at the combustor inlet) and size constraints of the combustion chamber, the process must be completed within a few tens of milliseconds, while the overall pressure drop (including mixing) must be kept below 5% of the turbine inlet pressure in order to prevent significant energy efficiency losses. These characteristics of GT operations result in the following requirements: • Highly active catalysts able to ignite the combustion of natural gas at temperatures as close as possible to those at the compressor discharge must be employed. Indeed, to fill the gap between compressor discharge temperature and catalyst ignition temperature, a homogeneous preburner is needed, but this may produce significant amounts of NOx . • Materials with high thermal stability able to hinder catalyst deactivation by sintering, phase transformation and volatilization, as well as to secure mechanical integrity upon thermal shocks, are required. In fact, strong temperature excursions are experienced during start-up and shut-downs, and particularly during the load trip of the turbine. To prevent overspeeding and destruction of the turbine in this case, the fuel feed is Design criteria and operating conditions of gas turbine combustors

Tab. 1

Design criteria Emission targets

Pressure drop Catalyst durability

NOx < 5 ppm CO < 10 ppm UHC < 10 ppm 90% AlF3 ) alumina is used [8]. A convenient method of preparing HFC-152a, useful in refrigerant blends, is by the reaction of vinyl chloride (VCl) with HF [Eq. (2)]. An 86% conversion of vinyl choride with

11.7.2 Catalytic Transformations

a 99% selectivity to HFC-152a has been reported for the vapor-phase reaction of vinyl chloride with HF using an alumina plus transition metal catalyst [9]. HF + CH2 =CHCl −−−→ CH3 CHF2 VCl 152a

(2)

HCFC-141b, HCFC-142b and HFC-143a can all be prepared by the reaction of HF with vinylidene chloride [VCl2 ; Eq. (3)]. In the vapor phase using an HFactivated Al2 O3 catalyst, the product selectivity can be dramatically altered using bismuth as a dopant. An HF-activated Al2 O3 catalyst affords a >99% selectivity to HCFC-141b [10], whereas the same catalyst doped with bismuth affords a >99% selectivity to HFC-143a [11]; both conversions are essentially quantitative. HF + CH2 =CCl2 −−−→ CH3 CCl2 F VCl2 141b + CH3 CClF2 + CH3 CF3 142b 143a

(3)

HFC-227ea, which has been proposed as a replacement for Halon fire extinguishants [12], can be prepared by the addition of HF to hexafluoropropene [HFP; Eq. (4)] over either an activated carbon catalyst [13] or a CrO2 F2 catalyst [14]. HF + CF2 =CFCF3 −−−→ CF3 CHFCF3 HFP 227ea

(4)

Halogen Exchange Halogen exchange is the operation whereby a C−Cl bond is converted to a C−F bond using a catalyst and HF as the fluorine source and HCl as the chlorine product (not shown in the equations below). When an alkene is the starting material, the initial step is HF addition to the double bond. In many cases this is followed by halogen 11.7.2.2

Tab. 1

CFC substitutes under development or in production

Market Refrigerants

Current CFC

CFC alternative

CFC-12 (CF2 Cl2 )

HFC-134a (CF3 CFH2 ) HCFC-22 (CHF2 Cl) HFC-32 (CH2 F2 ) HFC-125 (CF3 CF2 H) HCFC-124 (CF3 CHFCl) HFC-152a (CH3 CHF2 )

Blowing agents

CFC-11 (CFCl3 )

HCFC-141b (CH3 CFCl2 ) HCFC-123 (CF3 CHCl2 ) HCFC-22 (CHF2 Cl)

Cleaning agents

CFC-113 (CF2 ClCFCl2 )

blends/azeotropes new compounds hydrocarbons

2427

exchange to make more highly fluorinated analogs. In the previous section, HF was added to perchloroethene to make HCFC-121, which was then converted directly to HCFC-123, HCFC-124, and HFC-125 by halogen exchange. It is also possible to start with HCFC-121 or HCFC-122 [Eq. (5)]. CHCl2 CCl2 F −−−→ CHCl2 CClF2 −−−→ CHCl2 CF3 122 123 121 (5) This has the advantage of less catalyst deactivation since there is a lower alkene concentration in the reactor. A combination of fluorinated chromium and magnesium oxide shows high selectivity for HCFC-123 (64%) and HCFC-124 (15%) even after 6 months of using HCFC-122 as the feedstock [15]. Other catalysts, such as CrCl3 [16], CoCl2 /CeCl3 [17], NiCl2 [18], and CoCl2 /MgCl2 [19], all supported on γ -Al2 O3 give a high selectivity for HCFC123 (83–86%). Although it is possible to start with perchloroethene, HCFC-121, or HCFC-122 and convert these to HCFC123 in the liquid phase using a TaX5 [20] or SbCl5 [21] catalyst it is only in the vapor phase over a heterogeneous catalyst that the conversion to the more highly fluorinated HCFC-124 or HFC-125 [Eq. (1)] occurs to any appreciable extent. The rate of fluorination decreases as the fluorine substitution increases in the series HCFC-121 to HFC125. The rate-determining step in the formation of HFC-125 is the fluorination of HCFC-124, and Coulson has shown that the fluorination of HCFC-123 to HFC125 using Co/γ -A12 O3 occurs sequentially on the catalyst surface [22]. Other metals on γ -A12 O3 have been claimed to fluorinate HCFC-123 to HFC-125, such as Zn and Cr [23] as well as a coextrudate of Cr2 O3 and A12 O3 [24]. Chromium oxide on a carbon support [25], and a catalyst made from Cr2 O3 and MgO or Mg(OH)2 [26] show high selectivity for HFC-125 in the fluorination of HCFC-123. Chromium oxide low in alkali metals is known to be highly active for converting HCFC-123 to HCFC-124 and has long catalyst life [27]. The conversion of CHCl3 to HCFC-22 [Eq. (6)] has been practiced for many years since HCFC-22 is a precursor to tetrafluoroethene, a valued fluoromonomer. During the conversion a small amount of HFC-23 is formed. CHCl3 −−−→ CHCl2 F −−−→ CHClF2 −−−→ CHF3 20 22 23 21 (6) Milks reported the conversion of CHCl3 to predominantly HFC-23 using Ni- and Co-activated Al2 O3 (73% and 84%, respectively) [28]. In related work, Cu supported on Al2 O3 , gave 90% HFC-23 [29]. Ruh and Davis References see page 2433

2428

11.7 Catalytic Routes to Hydro(chloro)fluorocarbons

subsequently reported that hydrated CrF3 activated by reaction with oxygen, gives a high selectivity for HFC-23 (91%) [30]. Similarly, a process for producing fluorinated hydrocarbons using an improved chromium oxide catalyst made by treating a chromium hydroxide paste with water or steam before being dried and calcined, gives predominantly HFC-23 (37%) [31]. While carbon alone can catalyze the partial fluorination of CHCl3 , FeCl3 /C is highly active and selective for HFC-23 (87%) [32]. HCFC-22 can also be converted to HFC-23 by disproportionation (Section 11.7.2.3) which has the advantage of lower temperature and no HF is used. However, HCFC21, which is a product of the reaction, will have to be recycled back to HCFC-22. For example, Guo and Cai reported the formation of HFC-23 in high yield (>99%) using AlF3 , Bi/La, and Co-activated AlF3 [33]. HFC-134a has been found to be especially useful as a substitute for CFC-12 in automobile air conditioners. A fairly direct preparation method is shown in Eqs. (7) and (8); TCE = trichloroethene. The reactions shown are typically conducted separately in several reaction zones. HCFC-133a can be prepared using catalysts such as Cr/Mg [34], Cr2 O3 /AlF3 [35], AlClx Fy [36], Zn/fluorinated alumina [37], and Zn/Cr [38, 39]. CHCl=CCl2 + HF −−−→ CH2 ClCF3 + 2HCl TCE 133a CH2 ClCF3 + HF −−−→ CH2 FCF3 + HCl 133a 134a

(7) (8)

This conversion is equilibrium-limited [40], and a large excess of HF is required to achieve a high conversion of HCFC-133a to HFC-134a. Using a Cr2 O3 catalyst and increasing the HF: 133a ratio from 3.3 to 4, increases the HFC-134a selectivity from 16% to 18.2%, while decreasing the quantity of alkenes formed [41]. It is believed that alkenes at low HF ratios lead to carbon deposits and deactivation of the catalyst by blinding active sites. The cofeeding of chlorine to a chromium oxide-catalyzed reaction has been used to extend catalyst life, presumably by removing this carbon coating [42]. In lieu of chlorine, oxygen can be added. The oxygen reacts with HCl in a Deacon-like reaction to liberate Cl2 and H2 O, and the Cl2 functions as described above [43]. Removal of the HCl has the additional benefit of driving the equilibrium reaction to form more HFC-134a. A commercial process using a high HF ratio, needed to drive the equilibrium, will result in high energy costs, and the use of multiple distillation columns and complex recycle loops can effectively circumvent this [44]. Deactivation occurs via coking, as described above, and potentially by catalyst over-fluorination. In lieu of cofeeding oxygen (as described previously), air can be used

in a separate step to reactivate used catalyst [45]. To further remove coke or to reduce the fluorine content, a treatment with water vapor is also useful [46]. Many of the catalysts capable of fluorinating HCFC-133a are chromium-based, such as the Cr2 O3 examples described above. Many metals in combination with Cr are also effective such as Al, Co, Mn, Mg [47], and Zn [48]. Alumina or aluminum fluoride catalysts can also be used [49]. The conversion of TCE to HFC-134a using a single reactor zone for both reaction stages, which utilizes recycle of the HFC-133a reaction product using catalysts based on trivalent chromium, has been disclosed [50]. As in the synthesis of HFC-134a from HCFC-133a, the conversion of CH2 Cl2 to HFC-32 (CH2 F2 ) [Eq. (9)] using HF is an equilibrium-limited reaction, and NaF has been used with Cr2 O3 to remove HCl and drive the equilibrium to form more HFC-32 [51]. CH2 Cl2 −−−→ CH2 F2 30 32

(9)

Many different catalysts have been reported for this transformation, ranging from CrCl3 /C, CrCl3 /Al2 O3 , NiCl2 /C (all with aqueous HF [52]), Co/Al2 O3 [53], Bi2 O3 –CrCl3 /AlF3 , Bi2 O3 –CoCl3 /AlF3 , Bi2 O3 –MnCl2 / AlF3 [54], AlF3 , CrF3 /AlF3 , FeCl3 /C, and CrF3 /C [55]. In these latter cases, the most important factor was the HF: CH2 Cl2 ratio, not the catalyst or the temperature, suggesting that the reaction is equilibrium, not kinetically, limited with very little temperature dependence on the equilibrium. Isomerizations CFC-113 (CCl2 FCClF2 ) had been a large-volume fluorochemical used in solvent applications. The successful isomerization of CFC-113 to a useful intermediate, CFC113a (CF3 CCl3 ), has enabled some producers to utilize existing facilities to rapidly develop HCFC and HFC alternatives. A common catalyst for the isomerization of CFC-113 was described by Miller et al. [56], but the heterogeneous reaction between CFC-113 and AlCl3 , which undoubtedly made some CFC-113a, was originally reported by Henne and Newman [57]. Henne described a reaction in which AlCl3 reacts with CFC-113 to give CFC112a, by extracting fluorine from the organic reactant [Eq. (10)]. 11.7.2.3

2CClF2 CCl2 F + AlCl3 −−−→ 2CClF2 CCl3 + AlClF2 112a 113 (10) An active catalyst is formed in this step which can also disproportionate the CFC-112a to CFC-111, CFC-113a,

11.7.2 Catalytic Transformations

[Eq. (14)], and many byproducts are formed.

hexachloroethane, and CFC-114a [Eq. (11)]. CClF2 CCl3 −−−→ CCl2 FCCl3 + CF3 CCl3 112a 111 113a + CCl3 CCl3 + CF3 CFCl2 110 114a

CF2 ClCHF2 −−−→ CF3 CHClF 124a 124 (11)

The activation and isomerization reactions were examined in much detail by Okuhara [58] and Paleta et al. [59]. Okuhara made some pertinent observations such as: (i) the isomerization did not occur with partially deteriorated aluminum chloride; (ii) the rate of isomerization is sensitive to stirring efficiency; and (iii) the rate of isomerization is retarded by highly chlorinated organics (e.g., CCl4 ). Paleta et al. reported that the behavior of the catalyst in both activations and isomerizations depends on the type of commercial aluminum chloride used. This has recently been supported by claims that catalyst life and efficacy are related to the surface area of the AlCl3 catalyst precursor, which varies from batch to batch of commercial material [60]. The isomerization of CFC-114 to CFC-114a [Eq. (12)] is more difficult than the isomerization of CFC-113 since disproportionation of CFC-114a to CFC-113a and CFC-115 competes with isomerization, leading to yield losses. CClF2 CClF2 −−−→ CF3 CCl2 F 114 114a

(12)

AlX3 , activated with trace amounts of Cr and Mn powders, isomerizes CFC-114 in the liquid phase under pressure to yield 73% C2 Cl2 F4 (114a : 114 = 9.6) [61]. A large fraction (27%) of the product was CFC-113a, coming from the disproportionation of CFC-114a. In the vapor phase, an alumina catalyst activated with Cr and Mg was used to disproportionate CFC-114 to CFC-113a and CFC-115, while simultaneously feeding Cl2 and HF to fluorinate the CFC-113a to CFC-114a [62]. The net result is the conversion of CFC-114 to CFC-114a and CFC-115. Aluminum halide catalysts are not effective isomerization catalysts for hydrogenated chlorofluorocarbons because of HCl elimination to form alkenes which are catalyst poisons. HF elimination from HFCs is not as facile, thus AlF3 [63] and AlClx Fy Oz (x + y + 2z = 3) [64] are effective catalysts for the isomerization of 134 to 134a [Eq. (13)]. CHF2 CHF2 −−−→ CF3 CH2 F 134 134a

2429

(13)

Prefluorinated Cr2 O3 , an effective fluorination catalyst, also isomerizes 134 to 134a in the vapor phase [65]. While effective at isomerizing 134 to 134a, Cr2 O3 is too active for the selective isomerization of 124a to 124

(14)

However, γ -Al2 O3 fluorided with CCl2 F2 selectively isomerizes 124a to 124 in a heterogeneous liquid phase reaction [66]. In the vapor phase, γ -Al2 O3 fluorided with CCl4 will interconvert 124a and 124, depending on the temperature of reaction [67]. Just as in the isomerization of 124a, aluminum oxyhalides are effective in the isomerization of 123a [Eq. (15)]. CF2 ClCHF2 −−−→ CF3 CHClF 124a 124

(15)

AlF0.2 O1.6 , made by calcining AlF3 · H2 O containing TFE fibrils, isomerizes HCFC-123a at 40 ◦ C with 95% conversion after 1 h and 88% conversion after 3 h [68]. Cr2 O3 or CrCl3 /Al2 O3 isomerizes HCFC-123a to HCFC123 with higher amounts of disproportionation products, HCFC-133a, and CFC-113a than Al2 O3 , but with less alkene formation due to HX elimination [69]. Disproportionations/Conproportionations Disproportionation or dismutation is the reaction between two of the same molecules resulting in an exchange of atoms and the formation of at least two new molecules. Conproportionation is the reverse of this, where at least two different molecules react to form two molecules, which may be of the same kind. In conproportionations, it is common to view one of the reactants as a substrate and the other as a fluorinating agent. Chromium catalysts are proficient at disproportionation, and this feature has been exploited in the synthesis of CFC114a, HCFC-123, HCFC-124, and HFC-125 [Eqs. (16) and (17)]. An MgO or Mg(OH2 )/Cr2 O3 catalyst [70] catalyzes the conproportionation reaction between CFC113a and HCFC-133a to make CFC-114a, with byproduct formation of trichloroethene. 11.7.2.4

CCl3 CF3 + CH2 ClCF3 −−−→ CCl2 FCF3 + CHCl=CCl2 113a 133a 114a TCE (16) CHCl2 CF3 + CH2 ClCF3 123 133a (17) −−−→ CHClFCF3 + CHCl=CCl2 124 TCE HCFC-123 reacts with HCFC-133a over the Mg/Cr catalyst to give HCFC-124, also with TCE formation. At different feed ratios and over a Cr2 O3 catalyst, CFC-113a and References see page 2433

2430

11.7 Catalytic Routes to Hydro(chloro)fluorocarbons

HCFC-133a react to form HCFC-123 as the primary product [71] [Eq. (18)]. CF3 CCl3 + CF3 CH2 Cl −−−→ CF3 CHCl2 113a 123 133a

(18)

Whereas γ -Al2 O3 and FeCl3 /γ -Al2 O3 were not effective, Cr2 O3 was also shown to catalyze the conproportionation of HCFC-123 and HFC-125 to make HCFC-124 [72, Eq. (19)]. CF3 CHCl2 + CF3 CHF2 −−−→ 2CF3 CHClF 123 125 124

(19)

Cr2 O3 will also catalyze the approach to equilibrium from the reverse direction; HCFC-124 will disproportionate over Cr2 O3 to give a mixture of HFC-125 and HCFC-123, free of the a and b isomers [73] . HX Elimination Reactions HFC-245eb can be prepared from hexafluoropropene (HFP) by the sequence shown in Eq. (20), in which the HFP is hydrogenated over a Pd/C catalyst to HFC-236ea. This is dehydrofluorinated over carbon to HFC-1225ye, which is then hyrogenated over a Cu/Pd/C catalyst to HFC-245eb [74]. 11.7.2.5

CF3 CF=CF2 −−−→ CF3 CHFCHF2 HFP 236ea −−−→ CF3 CF=CHF −−−→ CF3 CHFCH2 F 1225ye 245eb

(20)

Haloalkenes can be prepared by dehydrohalogenating saturated hydrogen-containing polyhalocarbons using liquid alkali metal acid fluoride and/or alkali metal fluoride compositions [75]. HCFC-133a can be converted to CF2 =CHCl using these catalyst systems as shown in Eq. (21). CF3 CH2 Cl −−−→ F2 C=CHCl 133a 1122

(21)

Hydrodehalogenations/Dehydrohalogenations The substitution of chlorine with hydrogen by catalytic hydrodechlorination (HDC) is a well-known reaction. Earlier work by Bitner et al. [76] and Gervasutti et al. [77] over Pd/C using CFC-114a showed that the reaction proceeded primarily to HFC-134a although a small amount of HCFC-124 was also obtained [Eq. (22)]. 11.7.2.6

CF3 CFCl2 + H2 −−−→ CF3 CFH2 + CF3 CFHCl + HCl 114a 134a 124 (22) Conversion of the HCFC-124 to HFC-134a also occurred, but at much higher temperature. This has been recently

studied and a proposed mechanism for the direct formation of HFC-134a involves a surface carbene intermediate, CF3 CF: [78]. A significant problem with the HDC reaction, using Pd/C is catalyst deactivation, resulting from sintering due to the HCl produced in the reaction [79]. Kellner et al. [80] have found that the sintered Pd can be regenerated by treating the deactivated catalyst with a CFC or HCFC at relatively mild conditions. Morokawa et al. [81], have disclosed a series of multimetallic catalysts based on a Group VIII metal and other metal additives. They have interpreted the extended life of these catalysts to chemical adsorption energy and geometric factors. The effect of catalyst support has also been reported to have a pronounced effect. Kellner and Rao [82] reported 100% selectivity to HFC-134a from HCFC-124, using the unconventional support AlF3 . Acid washing of the carbon to remove trace quantities of metal impurities was reported by Rao [83] to increase conversion of the CFC-114a and to significantly increase the HFC134a : HCFC-124 ratio. Catalytic HDC has been reported for the synthesis of several other CFC alternatives or their intermediates. Particularly interesting is the report by Ichikawa and coworkers [84] that Pd catalysts promoted with additives such as Tl, Bi, Sn, Cu, or In result in dehydrochlorination (DHC) to give fluoroalkenes [Eq. (23)]. CF2 ClCFCl2 + H2 113 −−−→ CF2 =CFCl + CF2 =CFH + HCl 1113 1123

(23)

It is reported that a Pd/Bi catalyst gave excellent selectivity to the desired 1123, which is important since addition of HF readily provides HFC-134a [85] using Crbased catalysts. Pd modified with Au, Te, Sb, and As [86], and Re catalysts modified with Group VIII metals [87] have also been reported. The hydrodechlorination of CFC-115 to HFC-125 [Eq. (24)] has been carried out using a variety of Group VIII metal catalysts [88]. CF3 CF2 Cl + H2 −−−→ CF3 CF2 H + 2HCl 115 125

(24)

Bitner et al. reported that the HDC of CFC-12 (CF2 Cl2 ) over Pd/C yields primarily methane. More recently, it has been reported that the desired HFC-32 (CH2 F2 ) is formed in about 5% yield using similar catalysts based on Pd/C [89]. These same workers report that higher yields can be obtained starting with HCFC22 (CHClF2 ), although methane is still the major product.

11.7.2 Catalytic Transformations

Oxygenates as Reactive Intermediates With world regulatory pressures on the use of chlorinated compounds increasing, the synthesis of fluorinecontaining hydrocarbons from oxygen-containing substrates has received much attention in recent years. Reacting simple molecules such as formaldehyde and carbonyl fluoride over an activated carbon catalyst [90] gives an 86% yield of HFC-32 [Eq. (25)]. 11.7.2.7

H2 CO + COF2 −−−→ CH2 F2 + CO2 32

(25)

Separating the steps and reacting bisfluoromethyl ether, made from formaldehyde and HF over a Cr2 O3 catalyst, gives 57% HFC-32 and some methyl fluoride. The approach of decomposing ethers over heterogeneous catalysts has been extended to make other fluorocarbons of interest such as HFC-134a, HFC-125, etc. [91]. γ -Al2 O3 was used to catalyze the conversion of fluorinated alcohols (Rf CH2 OH) to fluorinated alkanes (Rf CH2 F). For example, the reaction of CF3 CH2 OH with COCl2 and HF [Eq. (26)] gives HFC-134a in a single step over γ -Al2 O3 [92]. CF3 CH2 OH + COCl2 + 2HF −−−→ CF3 CH2 F + CO2 + 2HCl 134a

(26)

HFC-245ca (CH2 FCF2 CHF2 ), HFC-449pccc (CHF2 CF2 CF2 CF2 CH2 F), and HFC-338q (CF3 CF2 CF2 CH2 F) were also made this way via the corresponding alcohol. This catalyst will also convert corresponding carbonate esters to fluorinated alcohols. (CF3 CF2 CH2 O)2 CO/COCl2 / HF is converted to HFC-236cb (CF3 CF2 CH2 F) in this manner. It is also possible to carry this out in two steps. The alcohol is first reacted with either SO2 Cl2 or COCl2 to make Rf CH2 O-SO2 Cl or Rf CH2 OCOCl, which is then reacted with HF, catalyzed by either γ -Al2 O3 or AlF3 to give the fluoroalkane [93]. Coupling Polyfluoroalkenes can be easily prepared by the addition of polyfluoroallylic fluorides to fluoroethenes in the presence of catalysts of the structure AlX3 where X is one or more of Br, Cl or F, provided that X cannot be entirely F. Of particular interest are reaction products of hexafluoropropene (HFP), where R1 , R2 , R3 , and R4 of (1) in Eq. (27) are all F, and tetrafluoroethene (TFE) where R5 of (2) is F, which are reacted in the presence of an AlFx Cly catalyst. 11.7.2.8

R1 R2 C=C(R3 )CF2 R4 + R5 FC=CF2 (1) (2) −−−→ R1 R2 C=C(R3 )CF(R4 )CF(R5 )CF3

(27)

2431

F(CF2 )2 CF=CFCF3 is obtained by reacting HFP : TFE in a 1 : 1 molar ratio. The major products obtained in 32% and 39% yields from the reaction of HFP : TFE in a 1 : 2 molar ratio are F(CF2 )2 CF=CFCF3 and F(CF2 )3 CF=CF(CF2 )2 F, respectively [94]. These alkenes can then be reduced with hydrogen over a palladium catalyst to afford HFCs which can replace CFCs as cleaning liquids [95] or as solvents for fluoromonomer polymerizations [96]. HCFC-225aa, HCFC-225ca and HCFC-225cb are solvents which are useful for the cleaning of electronic circuit boards, and can provide a temporary replacement for CFC-113 which has been extensively used for this application. The 225 isomers are readily prepared by the addition of HCFC-21 to TFE using an AlClx Fy catalyst [97] as shown in Eq. (28). CHCl2 F + CF2 =CF2 −−−→ CHCl2 CF2 CF3 TFE 225ca 21 + CHClFCF2 CClF2 + CHF2 CCl2 CF3 225cb 225aa

(28)

Chlorinations/Chlorofluorinations Chlorination of C−H bonds in readily available precursors, provides yet another route to CFC alternatives. Conversion of either HFC-143a or HCFC-133a to the desired HCFC-123 [Eq. (29)] has been reported using activated carbon or chromium oxide [98] and photochemical techniques [99]. The first chlorination is the most difficult, thus progressive chlorinations become faster, so that selectivity is best achieved at low conversion to minimize the formation of CFC-113a. 11.7.2.9

CF3 CH3 + Cl2 −−−→ CF3 CH3 Cl 143a 133a −−−→ CF3 CHCl2 −−−→ CF3 CCl3 (29) 123 113a A series of patents from DuPont describe the chlorofluorination of propane and propene [100] directly to perhalogenated products such as perfluorocarbon PFC218 and CFC-216aa [Eq. (30)]. PFC-218 is a potential etching agent for silicon and CFC-216aa is an intermediate, through hydrodechlorination [Eq. (31)], to HFC-236fa, a potential foam-blowing and fire extinguistant agent. It is also a precursor to fluoromonomers such as hexafluoropropene. The catalysts used for chlorofluorinations are AlF3 , Cr2 O3 and CrCl3 /C. CH3 CH = CH2 + Cl2 + HF −−−→ CF3 CCl2 CF3 216aa (30) + CF3 CF2 CF3 + HCl 218 References see page 2433

2432

11.7 Catalytic Routes to Hydro(chloro)fluorocarbons

CF3 CCl2 CF3 + H2 −−−→ CF3 CH2 CF3 + HCl (31) 216aa 236fa The chlorofluorination reaction typically gives perhalogenated products, although it is now possible through new catalysts to prepare HFCs and HCFCs such as the HCFC-225s [101], and HFC-125 [102] [Eqs. (32) and (33)]. CH3 CH2 =CH2 + Cl2 + HF −−−→ C3 HF5 Cl2 225 CF3 CH3 + Cl2 + HF −−−→ CF3 CF2 H 125

(32) (33)

11.7.3

CFC Destruction

A variety of methods have been proposed for removing CFCs from the environment. These include combustion in air, oxygen, ammonia, or water atmospheres. CFCs can also be reacted with HCl to afford starting materials for the preparation of some of the HFCs and HCFCs discussed in this chapter. The combustion of CCl2 F2 in oxygen at 500 ◦ C was studied in the presence of more than 20 different catalysts. The combustion products are CO, CO2 , and F2 . Because of fluorine’s tremendous reactivity, metallic and silicon-based catalysts degraded rapidly. The most durable catalyst was found to be BPO4 ; however, it also did not have a satisfactory life because of fluorine’s reaction with boron [103]. CFC-12 was reacted with NH3 [Eq. (34)] over metal catalysts on LaF3 and activated carbon to form HCN, HF, HCl, and N2 [104]. CF2 Cl2 + NH3 −−−→ HCN + HF + HCl + N2 12

(34)

CFCs were decomposed to HCl, HF, and CO2 at 150 ◦ C to 350 ◦ C by the reaction of H2 O over amorphous alloy catalysts consisting of at least one element selected from the group of Ni and Co, at least one element selected from the group Nb, Ta, Ti, and Zr, and at least one element selected from the group Ru, Rh, Pd, Ir, and Pt. The alloys were activated by immersion in HF [105]. CFCs are decomposed by the reaction of water vapor at temperatures above 300 ◦ C in the presence of iron oxide supported on activated carbon [106]. They are also decomposed by steam in the presence of catalysts comprising alumina or alumina/silica at 350 to 1000 ◦ C [107]. The fluorine content of CFCs can be reduced by reacting the CFC with HCl in the vapor phase in the presence of a vapor-phase fluorination catalyst (e.g., aluminum fluoride and trivalent chromium compounds) [108]. A CFC compound such as CClF2 CClF2 can be converted

to CCl3 CF3 (CFC-113a) and CCl2 =CCl2 (PCE). CFC113a and PCE can be used as starting materials for the preparation of HFC-134a [Eqs. (7) and (8)]. 11.7.4

Purification Techniques

Many fluoroalkenes have been shown to be toxic. Fluoro alkenes, as well as other types of alkene, can be removed from HFCs, by hydrogenolysis in the presence of a hydrogenation catalyst [109]. Alkenes can be removed from HCFCs, in the presence of oxygen-containing phase-transfer catalysts with complex hydrides and/or strong bases [110]. A more convenient method involves hydrogenolysis of HCFCs, containing alkenes over a hydrogenation catalyst [111]. 11.7.5

Opportunities for the Future

Fluorocarbon catalysis is still largely dominated by halogen exchange reactions involving initial formation of a C−Cl bond and halogen exchange with HF. These reactions lead to very large quantities of HCl, as a byproduct, which is under severe environmental scrutiny. New methods are needed to prepare C−F bonds selectively, without the use of chlorine. Although electrochemical fluorinations have been practiced commercially for many years, the main products of the reactions are typically perfluorocarbons since C−H bonds rarely survive. There is a need to develop new techniques which are chlorine-free. An intriguing reaction was published many years ago [112] involving the oxidative fluorination of ethene to vinyl fluoride [Eq. (35)] using Pd/Cu catalysts. The reaction probably involves Wacker chemistry, proceeding through acetaldeyde, so high degrees of fluorination are unlikely. However, it remains one of the few reported examples with which to make C−F bonds, from a hydrocarbon, without going through a chlorinated intermediate. C2 H4 + O2 + HF −−−→ CH2 =CHF + H2 O

(35)

Another example that has been reported in the patent literature [113] is the reaction of fluorspar with CO and SO3 [Eqs. (36) and (37)]. The mechanism and viability of this process remains to be determined. CaF2 + SO3 −−−→ Ca(SO3 F)2

(36)

Ca(SO3 F)2 + CH3 CHO −−−→ CH3 CHF2 + CaSO4 (37)

References

11.7.6

Conclusions

Heterogeneous catalysis has played a key role in the synthesis of CFC alternatives. However, for the process to be environmentally safer, the HCl must be recycled or sold. Although it is beyond the scope of this chapter, it is important to recognize that new catalytic processes based on Deacon chemistry [Eq. (38)] have been commercialized by Mitsui-Toatsa [114] using Cr-based catalysts. HCl + O2 −−−→ Cl2 + H2 O

(38)

For practitioners of heterogeneous catalysis, the challenges are to bring some mechanistic understanding and science to the many new reactions that have been reported over the past years for the development of the CFC alternatives – HFCs and HCFCs. The challenge is not over. Much work is in progress in industrial and government laboratories to develop environmentally safer ‘‘third-generation’’ CFC alternatives. References 1. F. Swartz, Bull. Acad. R. Belg. 1892, 24, 309. 2. L. E. Manzer, Science 1990, 31. 3. M. O. McLinden, D. A. Didion, ASHRAE J. 1987, December 3. 4. L. E. Manzer, V. N. M. Rao, Adv. Catal. 1993, 39, 329. 5. S. Morikawa, S. Samejima, M. Yoshitake, N. Tatematsu, Jpn Patent 2-178 237 (Asahi Glass.Co.) 1990, (Chem. Abstr. 113, 151 815). 6. P. Cuzzato, A. Masiero, EU Patent 408 004 (Ausimont) 1991, (Chem. Abstr. 114, 228 352). 7. D. R. Corbin, V. N. M. Rao, US Patent 5 300 711 (DuPont) 1994 (Chem. Abstr. 118, 105 304). 8. L. E. Manzer, V. N. M. Rao, US patent 4 766 260 (DuPont) 1988 (Chem. Abstr. 110, 97 550). 9. A. Akramkhodzhaev, T. S. Sirlibaev, Kh. V. Usmanov, Uzb. Khim. Zh. 1980, 29 (Chem. Abstr. 92, 215 810). 10. S. H. Swearingen, J. F. Wehner, M. G. Ridley, US Patent 5 105 033 (DuPont) 1992 (Chem. Abstr. 114, 101 118). 11. N. Schultz, H. J. Vahlensieck, R. Gebele, US Patent 3 904 701 (Dynamit-Nobel) 1975 (Chem. Abstr. 74, 3310). 12. R. E. Fernandez, US Patent 5 084 190 (DuPont) 1992 (Chem. Abstr. 116, 155 038). 13. P. Hopp, U. Wirth, EU Patent 562 509 (Hoechst) 1993 (Chem. Abstr. 119, 270 624). 14. S. P. v. Halasz, Ger Patent 2 712 732 (Hoechst) 1978 (Chem. Abstr. 89, 214 885). 15. S. Morikawa, S. Samejima, M. Yoshitake, S. Tatematsu, Jpn Patent 2-172 932 (Asahi Glass), 1990. 16. K. Yangii, S. Yoshikawa, K. Murata, Jpn Patent 2-157 235 (Central Glass), 1990. 17. S. Morikawa, S. Samejima, M. Yoshitake, S. Tatematsu, Jpn Patent 4-029 940 (Asahi Glass), 1992. 18. S. Morikawa, S. Samejima, M. Yoshitake, S. Tatematsu, Jpn Patent 4-029 942A2 (Asahi Glass), 1992. 19. S. Morikawa, S. Samejima, M. Yoshitake, S. Tatematsu, Jpn Patent 4-029 943A2 (Asahi Glass), 1992.

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20. W. H. Gumprecht, W. G. Schindel, V. M. Felix, US Patent 4 967 024 (DuPont), 1990; V. N. M. Rao, US Patent 5 015 791 (DuPont), 1991. 21. G. Fernschild, C. Brosch, DE (German patent application) 4 005 944A1 (Kali-Chemie) 1990. 22. D. R. Coulson, J. Catal. 1993, 142, 289. 23. D. R. Corbin, V. N. M. Rao, WO92/16482 (DuPont), 1992. 24. H. S. Tung, A. M. Smith, US Patent 5 155 082 (Allied Signal), 1992. 25. B. Cheminal, E. Lacroix, A. Lantz, AU (Australian patent application) 76053/91 (Atochem), 1991. 26. M. Hoeveler, M. Schnauber, M. Schott, CA (Canadian patent application) 2 068 832 (Hoechst), 1992. 27. V. M. Felix, W. H. Gumprecht, B. A. Mahler, US Patent 5 334 787 (to DuPont) 1994. 28. W. N. Milks, US Patent 2 744 147 (Dow), 1956. 29. R. P. Ruh, R. A. Davis, US Patent 2 744 148 (Dow), 1956. 30. R. P. Ruh, R. A. Davis, US Patent 2 745 886 (Dow), 1956. 31. R. Firth, G. Toll, US Patent 3 755 477 (ICI), 1973. 32. L. J. Belf, US Patent 2 946 827 (National Smelting, LTD.), 1960. 33. X. Guo, Z. Cai, CN (Chinese patent application) 85/105080 (Zhejiang Chemical Industry Research Institute), 1985. 34. W. Wanzke, G. Siegemund, T. M¨uller, EU Patent 407 961 (Hoechst), 1991 (Chem. Abstr. 114, 121 449). 35. P. Cuzzato, EU Patent 583 703 (Ausimont), 1994 (Chem. Abstr. 120, 269 636). 36. H. W. Swidersky, W. Rudolph, T. Born, EU Patent 492 386 (Kali-Chemie), 1992 (Chem. Abstr. 117, 150 539). 37. D. R. Corbin, V. N. M. Rao, PCT International Application WO 92/16480 (DuPont), 1992 (Chem. Abstr. 117, 253 904). 38. J. D. Scott, M. J. Watson, US Patent 5 281 568 (ICI), 1994 (Chem. Abstr. 117, 236 287). 39. D. R. Corbin, V. N. M. Rao, US Patent 5 321 170 (DuPont), 1994 (Chem. Abstr. 117, 253 903). 40. J. Barrault, S. Brunet, B Requieme, M. Blanchard, J. Chem. Soc., Chem. Commun. 1993, 374. 41. S. L. Bell, US Patent 4 129 603 (ICI), 1978. 42. W. H. Gumprecht, US (statutory invention application) H1129, 1993. 43. L. E. Manzer, US Patent 5 051 537 (DuPont), 1991. 44. H. Ohno, T. Arai, K. Muramaki, T. Ohi, H. Nakayama, Y. Shoji, US Patent 5 276 223 (Showa Denko), 1994. 45. P. K. Dattani, J. D. Scott, EU Patent 475 693 A1 (ICI), 1990. 46. J. D. Scott, A. Oldroyd, WO (PCT International application) 93/10898 (ICI), 1993. 47. S. Morikawa, S. Samejima, M. Yoshitake, S. Tatematsu, Jpn Patent 2-179 223 (Asahi Glass), 1990. 48. J. D. Scott, M. J. Watson, US Patent 5 281 568 (ICI), 1994. 49. L. E. Manzer, US Patent 4 922 037 (DuPont), 1990; D. R. Corbin, V. N. M. Rao, WO 92/16481 (DuPont), 1992. 50. L. E. Manzer, US Patent 5 185 482 (DuPont), 1993 (Chem. Abstr. 114, 810 16). 51. W. R. Buckman, US Patent 3 644 545 (Allied), 1972. 52. T. R. Fiske, D. W. Baugh, Jr., US Patent 4 147 733 (Dow), 1979. 53. M. Yoshitake, S. Tatematsu, S. Morikawa, Jpn Patent 5-339 179A2 (Asahi Glass), 1993. 54. X. Guo, Z. Ye, X. Ma, CN (Chinese patent application) 1 076 686 A, (Zhejiang Chemical Research Institute), 1993. 55. S. Takayama, H. Nakayama, H. Kawasaki, N. Hashimoto, N. Kawamoto, EU Patent 128 510 (Showa Denko), 1984. 56. W. T. Miller, Jr., E. W. Fager, P. H. Griswold, J. Am. Chem. Soc. 1950, 72, 705.

2434

11.8 Heterogeneous Catalysis in the Troposphere

57. A. L. Henne, M. S. Newman, J. Am. Chem. Soc. 1938, 60, 1697. 58. K. Okuhara, J. Org. Chem. 1978, 43, 2745. 59. O. Paleta, L. Frantisek, A. Posta, V. Dedek, Collection Czechoslov. Chem. Commun. 1980, 45, 104. 60. W. H. Gumprecht, W. J. Huebner, M. J. Nappa, WO (PCT International application) 94/03417 (DuPont), 1994. 61. R. C. Zawalski, US Patent 5 017 732 (Dixie Chemical), 1991. 62. S. Morikawa, M. Yoshitake, S. Tatematsu, Jpn Patent 1-172 347 (Asahi Glass), 1989. 63. L. E. Manzer, V. N. M. Rao, US Patent 4 902 838 (DuPont), 1990. 64. S. Morikawa, S. Samejima, M. Yoshitake, S. Tatematsu, T. Tanuma, Jpn Patent 2-115 135 (Asahi Glass), 1990. 65. G. J. Moore, H. M. Massey, US Patent 5 091 600 (ICI), 1992. 66. S. Morikawa, M. Yoshitake, S. Tatematsu, Jpn Patent 2-040 332 (Asahi Glass), 1990. 67. W. H. Manogue, V. N. M. Rao, US Patent 5 030 372 (DuPont), 1991. 68. S. Okazaki, M. Ogura, Y. Mochizuki, EU Patent 450 467 (Du Pont-Mitsui Fluorochemicals), 1991. 69. P. Cuzzato, A. Masiero, EU Patent 537 760 A2 (Ausimont), 1993. 70. T. Muller, G. Siegemund, US Patent 4 547 483 (Hoechst), 1992. 71. R. F. Sweeney, B. F. Sukornick, US Patent 4 192 822 (Allied), 1980. 72. L. E. Manzer, F. J. Weigert, WO (PCT International application) 92/02476 A1 (DuPont), 1992. 73. P. Cuzzato, EU Patent 569 832 A1 (Ausimont), 1993. 74. H. Aoyama, E. Seki, WO PCT International Application 93/25510 (Daikin), 1993 (Chem. Abstr. 120, 298 045). 75. R. E. Fernandez, R. B. Kaplan, US Patent 5 180 860 (DuPont), 1993 (Chem. Abstr. 118, 212 469). 76. J. L. Bitner, et. al., U.S. Dept. Comm. Off. Tech. Serv. Rep. 136 732 1958, 25. 77. C. Gervasutti, L. Marangoni, W. Marra, J. Fluorine Chem. 1981, 19, 1. 78. E. D. Boyes, D. R. Coulson, G. W. Coulston, M. P. Diebold, P. L. Gai, G. A. Jones, C. S. Kellner, J. J. Lerou, L. E. Manzer, et al., Proceedings of American Chemical Society, Division of Petroleum Chemistry 1993, 38, 847. 79. J. I. Darragh, UK Patent 1 578 933, 1980. 80. C. S. Kellner, J. J. Lerou, K. G. Wuttke, V. N. M. Rao, U.S. Patent 4 980 324, 1990. 81. S. Morikawa, S. Samajima, M. Yoshitake, S. Tatematsu, EU Patent 317 981, 1989 (Chem. Abstr. 111, 194 096). 82. C. S. Kellner, V. N. M. Rao, US Patent 4 873 381, 1989. 83. V. N. M. Rao, DuPont: WO (PCT International application) 92/12113, 1992. 84. R. Ohnishi, I. Suzuki, M. Ichikawa, Chem. Lett. 1991, 841. 85. S. P. Von Halasz, Ger. Patent 3 009 760 (Chem. Abstr., 95, 186 620), 1981. 86. T. Saiki, M. Suida, S. Nakano, K. Murakami, US Patent 5 282 379, 1994. 87. C. S. Kellner, V. N. M. Rao, F. J. Weigert, US Patent 5 068 473, 1991. 88. S. Morikawa, S. Samejima, M. Yoshitake, N. Tatematsu, Jpn Patent 3-099 026 (Asahi Glass), 1991. 89. J. Moore, J. O’Kell, EU Patent 508 660 (ICI), 1992. 90. T. A. Ryan, UK Patent 212 6216 A (ICI), 1984. 91. L. Burgess, J. Butcher, T. Ryan, P. P. Clayton, WO (PCT International application) 93/12057 (ICI), 1993.

92. M. J. Nappa, A. C. Sievert, US Patent 5 274 189 (DuPont), 1993. 93. M. J. Nappa, A. C. Sievert, US Patent 5 274 190 (DuPont), 1993. 94. C. G. Krespan, US Patent 5 162 594 (DuPont), 1992 (Chem. Abstr. 117, 69 439). 95. C. G. Krespan, V. N. M. Rao, US Patent 5 171 902 (DuPont), 1992 (Chem, Abstr. 117, 233 426). 96. A. E. Feiring, C. G. Krespan, P. R. Resnick, B. E. Smart, T. A. Treat, R. C. Wheland, US Patent 5 182 342 (DuPont), 1993 (Chem. Abstr. 118, 213 753). 97. A. C. Sievert, C. G. Krespan, F. J. Weigert, US Patent 5 157 171 (DuPont), 1992 (Chem. Abstr. 115, 70 904). 98. R. F. Sweeney, B. R. Sukornick, US Patent 4 192 822, 1980 (Allied Chemical). 99. R. F. Sweeney, US Patent 4 060 469, 1977 (Allied Chemical). 100. J. L. Webster, E. L. McCann, D. W. Bruhnke, J. J. Lerou, L. E. Manzer, W. H. Manogue, P. R. Resnick, S. Trofimenko, US Patent 5 043 491, 1991; US Patent 5 068 472, 1991; US Patent 5 057 643, 1991; US Patent 5 220 083, 1993. 101. D. W. Bruhnke, J. J. Lerou, V. N. M. Rao, W. C. Seidel, F. J. Weigert, US Patent 5 177 273, 1993. 102. M. J. Nappa, V. N. M. Rao, US Patent 5 258 561, 1993. 103. S. Imamura, Catal. Today 1992, 11, 547. 104. Y. Takita, T. Imamura, Y. Mizuhara, Y. Abe, T. Ishihara, Appl. Catal. B: Environmental 1992, 1, 79. 105. K. Hashimoto, H. Habazaki, US Patent 5 220 108 (K. Hashimoto and Yoshida Kogyo), 1993 (Chem. Abstr. 116, 151 333). 106. S. Okazaki, A. Kurosaki, US Patent 5 118 492 (DuPontMitsui), 1992 (Chem. Abstr. 114, 191 662). 107. S. Okazaki, A. Kurosaki, US Patent 5 151 263 (DuPontMitsui), 1992 (Chem. Abstr. 115, 135 502). 108. V. N. M. Rao, S. H. Swearingen, WO PCT International Application 94/13608 (DuPont), 1994. 109. R. E. Fernandez, V. N. M. Rao, US Patent 5 001 287 (DuPont), 1991 (Chem. Abstr. 114, 142 640). 110. G. Engler, U. Gross, D. Prescher, J. Schulze, East German Patent 160 718 (Academy of Sciences), 1984 (Chem. Abstr. 101, 170 689). 111. T. W. Fu, V. N. M. Rao, WO PCT International Application 93/14052 (DuPont), 1993 (Chem. Abstr. 120, 220 857). 112. T. Kuroda, Y. Takamitsu, Jpn Patent 52-122 310 (Onoda Cement Co., Ltd.), 1977 (Chem. Abstr. 88, 38 312). 113. R. K. Jordan, US Patent 4 087 475 (R. K. Jordan), 1978 (Chem. Abstr. 89, 108 083). 114. M. Ajioka, S. Takenaka, H. Itoh, M. Kataita, Y. Kohno, EU Patent 283 198 (Mitsui-Toatsa), 1988 (Chem. Abstr. 109, 213 174).

11.8

Heterogeneous Catalysis in the Troposphere Valentin M. Parmon∗ 11.8.1

Introduction

The Earth’s atmosphere is full of highly dispersed bodies such as liquid and solid aerosols. The aerosol particles ∗

Corresponding author.

11.8.2 The Atmosphere as a Global Catalytic and Photocatalytic Reactor

have a large surface area, and therefore can assist numerous heterogeneous processes which play important roles in the chemistry of the Earth’s atmosphere. For example, they can provide the downstream flux of many atmospheric compounds via sedimentation in the form of the absorbed or adsorbed species, they may facilitate the formation of secondary organic aerosols (SOA) or recombination of free radicals, as well as providing topochemical, thermal catalytic, and photocatalytic chemical reactions in the atmosphere. The relatively low temperatures and partial pressures for most of these reagents, the low catalyst concentrations, and the moderate intensity of the solar light flux do not favor high rates of heterogeneous catalytic and photocatalytic reactions in the atmosphere. However, due to the enormous total volume of the atmosphere, even very slow reactions inside the atmosphere bulk can result in chemical transformations of huge extent for some atmospheric components. As the concentrations of aerosols – and thus their surface area – are much higher in the lowest layers of the atmosphere (i.e., the troposphere) than in the upper layers, the majority of heterogeneous reactions are expected to occur in the troposphere. Due to the low temperatures, the thermal heterogeneous catalytic processes which prevail in the atmosphere seem mostly to be simple hydrolytic reactions (e.g., the hydrolysis of N2 O5 in acidic water droplets to form nitric acid). However, other low-temperature thermal redox heterogeneous process in the atmosphere are now also well known (e.g., the disproportionation of nitrogen oxides). Photocatalysis can provide many more numerous and complicated redox reactions, such as the complete oxidation of various organic substances, or even water splitting over aerosols containing TiO2 . In addition, the acidcatalyzed condensation of some organic compounds over aerosol particles, as well as destructive (photo)adsorption, have also been recognized. The existence of heterogeneous reactions guided by aerosols in the Earth atmosphere was recognized many years ago [1]. However, the actual role of such reactions, both thermal and photochemical, was long underestimated. The role of such catalytic reactions in the atmosphere was reassessed when attempting to understand the phenomenon of ‘‘ozone holes’’ and SOA formation, and today such reactions are considered to make an important contribution to the global chemistry of the atmosphere [2–8]. Until now, few direct experimental studies have been conducted on catalytic reactions with real atmospheric aerosols. Hence, the anticipated role of these reactions in the Earth’s atmosphere is based mainly on estimates from experiments with model catalysts, together with known data that characterize the atmosphere as a form of global catalytic reactor. For example, a drastic acceleration

2435

of some chemical transformations has been observed in the atmosphere after volcanic eruptions [2, 9]. The possible impact on the global chemistry of the Earth’s atmosphere due to photocatalytic reactions involving continental aerosols at background concentrations in the troposphere has also been estimated [3]. The most likely candidates in the role of tropospheric thermal catalysts and photocatalysts, as well as the main types of expected tropospheric catalytic reaction, are discussed in the following sections. 11.8.2

The Atmosphere as a Global Catalytic and Photocatalytic Reactor

The principal components of the Earth’s atmosphere (as vol.%, dry atmospheric air) are dinitrogen (78.09%), dioxygen (20.95%), argon (0.932%) and carbon dioxide (0.03%). The water content varies from 0.1 to 2.8 vol.%. However, some other ‘‘trace’’ components exist which, in spite of their low concentrations, exert strong influence on the atmospheric chemistry [8, 10–13]. The natural content (i.e., average stationary concentrations) of the principal trace components, their average life spans, and rates of supply and removal from the atmosphere are listed in Table 1. The two latter values are in fact equal to each other, and are calculated as the ratio of the stationary concentration of an atmospheric component to its residence time in the atmosphere. The period of vertical stirring of the atmosphere is estimated as 80 days [8, 10, 13, 14]; thus, those trace gases with a residence time less than 80 days are distributed non-uniformly within the atmosphere. Examples of such gases are CO, NOx , NH3 , and SO2 , and some volatile organic compounds such as monoterpenes of biological origin. In the vicinity of their sources, the concentrations of these gases and their removal (supply) rates may be higher than the values shown in Table 1. Atmospheric aerosols may exist in both liquid (droplets of water or water solution) and solid states. Solid aerosols – that is, ensembles of ultrasmall particles that sometimes are embedded into liquid droplets – can be divided roughly into two categories: • primary particles, which consist of dispersed materials lifted from the Earth’s surface (land and ocean) • secondary particles, which consist of materials formed during the course of transformation of gaseous compounds to the liquid or solid phase [7, 14]. Typically, the solid aerosol particles are covered by liquid or frozen water or water (usually acidic) solution layers. References see page 2446

2436

11.8 Heterogeneous Catalysis in the Troposphere

Average natural concentrations of main trace gases in the Earth’s atmosphere, their lifespan τ , and rate of removal (supply) (according to Ref. [10]). The quantum yield, ϕ0.1 , corresponds to the rate of photocatalytic removal (supply) of the these gases to achieve 10% of their total removal (supply) rate [3]

Tab. 1

Gas

Concentration ppb

CO2 CO CH4 CH2 O N2 O NO NO2 NH3 SO2 H2 S CS2 COS (CH3 )2 S H2 H2 O 2 CH3 Cl HCl

3.3 × 105 100 1600 0.1–1 0.3 0.1 0.3 1 0.01–0.1 0.05 0.02 0.5 0.001 550 0.1–10 0.7 0.001

mol

τ

L−1 /

(1 bar, 273 K)

1.5 × 10−5 4.5 × 10−9 7.1 × 10−8 4.5 × 10−(11−12) 1.3 × 10−11 4.5 × 10−12 1.3 × 10−11 4.5 × 10−11 4.5 × 10−(12 – 13) 2.2 × 10−12 8.9 × 10−13 2.2 × 10−11 4.5 × 10−14 2.5 × 10−8 4.5 × 10−(10 – 12) 3.1 × 10−11 4.5 × 10−14

4 years 0.1 year 3.6 years 5–10 days 20–30 years 4 days 4 days 2 days 3–7 days 1 day 40 days 1 year 1 day 6–8 years 1 day 3 years 4 days

Rate of removal/ mol L−1 s−1

Fe2 O3

TiO2

ZnO

1.2 × 10−13 1.4 × 10−15 6.3 × 10−16 (1.0–20) ×10−17 (1.4–2.0) × 10−20 1.3 × 10−17 3.9 × 10−17 2.6 × 10−16 (0.7–17) × 10−18 2.6 × 10−17 2.6 × 10−19 7.0 × 10−19 5.2 × 10−19 (9.9–13) × 10−17 (0.5–52) × 10−16 3.3 × 10−19 1.3 × 10−19

1.7 × 10−2 2.0 × 10−4 9.0 × 10−5 (0.1–2.9) × 10−5 (2.0–2.9) × 10−9 1.9 × 10−6 5.7 × 10−6 3.7 × 10−5 (1–24) × 10−7 3.7 × 10−6 3.7 × 10−8 1.0 × 10−7 7.4 × 10−8 (1.4–1.9) × 10−5 (0.7–7.4) × 10−5 4.8 × 10−8 1.9 × 10−8

1.1 1.3 × 10−2 5.7 × 10−3 (0.9–18) × 10−4 (1.3–1.8) × 10−7 1.2 × 10−4 3.6 × 10−4 2.4 × 10−3 (0.6–16) × 10−5 2.4 × 10−4 2.4 × 10−6 6.5 × 10−6 4.8 × 10−6 (9.2–12) × 10−4 (0.5–48) × 10−3 3.0 × 10−6 1.2 × 10−6

60 0.7 0.32 (0.5–10) × 10−2 (0.7–10) × 10−5 6.5 × 10−3 2.0 × 10−2 0.13 (0.4–8.5) × 10−3 1.3 × 10−2 1.3 × 10−4 3.5 × 10−4 2.6 × 10−4 (4.9–6.5) × 10−2 (0.3–26) × 10−1 1.7 × 10−4 6.5 × 10−5

The average composition of solid aerosol particles (primary and secondary) in normal (background) atmospheric conditions at various altitude is considered usually to be: • below 3 km: 50% (NH4 )2 SO4 , 35% soil (typical soil content is 53% SiO2 , 17% Al2 O3 , 7% Fe2 O3 , 23% other minerals), 15% sea salt (NaCl) • above 3 km: 60% (NH4 )2 SO4 , 40% soil [2, 4, 10, 13–15]. One particular, and important, type of primary aerosol particle is soot particles formed in combustion of either fires or anthropogenic activity [16]. Secondary organic aerosols are of reasonable importance for both urban and rural areas [5]. While dramatic local variations are seen in the composition of aerosols [14, 17, 18], of importance to heterogeneous chemistry is the possibility of high (≈100fold) enrichment of solid aerosols with transition metals such as Zn, Cd, and Pb, over their average (‘‘Clark’’) content in the Earth’s crust [19]. The average concentration of aerosol particles in the troposphere varies from 10−3 g m−3 under arid areas to 10−4 to 10−5 g m−3 under ordinary background conditions [20]. This distribution with respect to altitude is illustrated in Fig. 1a, but within localized area – such as major industrial areas or erupting volcanoes – aerosol concentrations may be dramatically higher. Particles of less than 1 µm diameter generally have atmospheric concentrations in the

Quantum yield (ϕ0.1 )

range from 10 to 10 000 cm−3 ; those exceeding 1 µm diameter typically exhibit concentrations less than 10 cm−3 [14]. Particle size and surface area, as well as the vertical distribution of aerosols, are also important for heterogeneous reactions. The majority of the mass of atmospheric aerosol matter is represented by particles of size 0.1 to 1 µm [14, 15]; that is, their surface area must be ≈10−1 m2 g−1 . The specific surface area A of solid aerosols near the Earth’s surface is considered to be ≈10−7 m2 dm−3 of air under the background conditions, and this may increase by a factor of 100 in urban areas [4]. The vertical distribution of A is shown in Fig. 1b. Due to sedimentation, almost all aerosols are located in the lower layer of the troposphere. For photocatalytic processes, the spectral characteristics and intensity of solar irradiation of various altitudes are important (Fig. 2). It is well known, that solar radiation with wavelengths below 300 nm is almost totally absorbed above the troposphere. Thus, to a first approximation, in the troposphere one should consider only reactions that occur over photocatalysts, which absorb light with λ > 300 nm. Only in the stratosphere and mesosphere may the reactions which occur over photocatalysts that absorb light with λ < 300 nm become important. Thus, the Earth’s atmosphere can be considered as a global heterogeneous catalytic and photocatalytic reactor with the known estimates for

11.8.3 Thermal Heterogeneous Catalytic Processes over Ice Particles, and Liquid and Solid Aerosols

2437

10−6 2

A/(m2 dm−3)

10−7

N/cm−3

101

100

10−8

3

10−9

10−1

10−2

4

1

10−10

1

10−11 1

2

3

4

Altitude/km

(a)

0

5

10

20

30

40

Altitude/km

(b)

(a) Averaged results of concentration, N, measurements at various altitudes for tropospheric particles of radius >2 × 10−7 m in continental air at mid-latitudes (a summary of data from Ref. [15]). (b) Altitude distribution of the specific surface area A of solid atmospheric aerosols (per unity volume of air). Curve numbers: 1, background atmosphere; 2, urban areas; 3, stratospheric volcanic clouds; 4, polar stratospheric clouds. (Fig. 1a adapted from Ref. [4].)

Fig. 1

• the distribution of intensity and spectral characteristics of the light flux, temperature, and pressure with altitude.

Iο(λ)/w · m−2 · µm−1

200

Altitude/km

150

102 100

11.8.3

10−2

100

0

(a)

Hydrolytic Reactions An example of a hydrolytic reaction which proceeds very efficiently in the presence of ice [2] is the transformation of relatively inert atmospheric forms of chlorine (HCl and ClONO2 ) into photochemically active Cl2 : 11.8.3.1

0

100

1000

λ/nm 50

Thermal Heterogeneous Catalytic Processes over Ice Particles, and Liquid and Solid Aerosols

(b)

HCl + ClONO2 −−−→ Cl2 (g) + HNO3 (s) 0

Other examples are the reactions: 0

100

200

300

λ/nm (a) The spectral density of the flux of the virgin incident solar radiation Io (λ) at the upper border of the atmosphere. (b) The altitudes, at which the intensity of the virgin solar radiation at various wavelengths decreases by the factor of 2. (According to Ref. [3].)

Fig. 2

• the average composition of the gas phase • the average composition, concentration, and surface area of solid aerosol particles that can serve as catalysts or photocatalysts

ClONO2 + H2 O −−−→ HOCl + HNO3 (s) HCl + HOCl −−−→ Cl2 (g) + H2 O(s) HCl + N2 O5 −−−→ ClNO2 (g) + HNO3 (s) which are expected to be one of the causes of ozone-hole formation in polar areas [2, 4, 21]. These reactions take place in water droplets and/or water layers that cover dust aerosols. HCl vapor has a strong affinity for liquid water, where it is converted to hydrochloric acid and promotes acid-catalyzed hydrolytic reactions. High concentrations References see page 2446

2438

11.8 Heterogeneous Catalysis in the Troposphere

of absorbed HCl can lead to the melting of ice particles. Experimental evidence suggests that hydrolytic reactions can proceed in the liquid or frozen aqueous phase even at very low (down to ≈150 K) temperatures. Heterogeneous low-temperature hydrolysis and the reaction of BrONO2 and BrO2 , as well as the uptake of ClO radicals with the formation of several highly reactive chlorine oxides such as OClO and ClClO2 on ice surface have also been detected experimentally [21–23]. Recently, it was found that water adsorbed onto iron oxide enhances the low-temperature uptake of CO2 by stimulating its conversion to carbonate and bicarbonate [24]. An important atmospheric heterogeneous reaction is the hydrolysis of N2 O5 , which originates from the gasphase recombination of NO2 and NO3 radicals [2, 4]: N2 O5 + H2 O −−−→ 2HNO3 This reaction can proceed in droplets of water or, more probably, of water solutions of sulfuric acid (which are present in the atmosphere in large quantities) even at temperatures as low as ≈220 K [2]. Volcanic eruptions producing huge clouds of sulfur-containing aerosols are assumed to make a large contribution to the above hydrolytic processes [2, 9]. Redox Reactions Possible low-temperature heterogeneous redox catalytic reactions include the following transformations of NO [2]: 11.8.3.2

−− −− → NO + N2 O5 − ← − 3NO2 −− −− → NO + 2HNO3 − ← − 3NO2 + H2 O These reactions are assumed to proceed in water or sulfuric acid aerosols via acid catalysis. The prehydrolysis of atmospheric SO2 is also considered an important step for the acceleration of SO2 oxidation to SO3 with atmospheric ozone or hydrogen peroxide, due to much higher reactivity of ionized forms of SO2 in water-containing aerosols [2]. Traces of transition metals such as Fe, Cu, Cr, or Mn that are typical for the atmospheric aerosols are expected to promote these oxidation reactions [2, 4, 25, 26]. High catalytic activity is also expected for water-covered particles of elementary carbon or soot which originate from the incineration of organic fuels and wood fires, and are present in the troposphere in large amounts [27]. To date, no evidence has been found for the direct catalytic oxidation of SO2 with atmospheric oxygen, probably due to a very low reactivity of molecular oxygen with respect to SO2 at low temperatures over conventional atmospheric catalysts.

Acid-Catalyzed Reactions of Organic Compounds over Aerosols Inorganic acids such as H2 SO4 and HNO3 are formed in the atmosphere in large amounts. These acids are easily captured by both liquid and solid aerosols and, as a result, the acidic surfaces of atmospheric aerosols can lead to multifold increases in secondary organic aerosol mass and a build-up of stabilized non-volatile organic matter as the particles age [5]. Among such acid-catalyzed reactions over aerosols under discussion are hydration, polymerization, and aldol condensation of volatile aldehydes and carbonyls. Of particular importance for SOA formation is the heterogeneous oligomerization of carbonyls which are produced from volatile monoterpenes of biological origin, following oxidation by atmospheric ozone [5, 7]. 11.8.3.3

11.8.4

Heterogeneous Photocatalytic Reactions in the Troposphere

For the major components and impurities of the Earth’s atmosphere (e.g., N2 , O2 , water, CO2 , methane, methane halides), electronically excited states are formed only upon absorption of light quanta with the energy hν ≥ 5 eV – that is, with a wavelength of λ ≤ 200 nm. Only a small portion of solar light energy is found in this spectral region, which means that most solar energy flux cannot participate in direct photochemical transformations of these atmospheric compounds. On the contrary, photocatalysts can absorb much smaller light quanta, from the visible, near-ultraviolet (UV), and even near-infrared (IR) spectral regions that constitute the main part of the solar energy flux. Upon absorption of such quanta, photocatalysts become chemically active and can participate in numerous chemical reactions, even at low temperatures [1, 3, 6, 27–34, 67]. Dust particles of soil, that contain metal oxides, as well as ions of transition metals captured by water aerosols, may serve as photocatalysts for atmospheric reactions. Reactions over such photocatalytically active particles may occur in all layers of the atmosphere, in contrast to direct photochemical processes, the occurrence of which is restricted to far-UV radiation that exists only in the upper layer of the atmosphere. At least two main mechanisms can be distinguished for the initiation of photocatalytic processes in Nature. ‘‘Real’’ photocatalysis [3, 28, 33, 35] proceeds generally via the photogeneration of either a highly chemically reactive pair – ‘‘a strong chemical oxidant; a strong reducing agent’’ – after the excitation of, for example, a transition ion, or an also chemically reactive electron-hole pair in semiconductor particles (see Fig. 3). In contrast, the so-called ‘‘photoinduced’’ catalysis [35] occurs via

11.8.4 Heterogeneous Photocatalytic Reactions in the Troposphere

Conduction band

Adsorbed water

Ee

B

−

Eg Aoxidized

Aa

Ba

Breduced

hn

+

Forbidden band

A

Valence band

Semiconductor oxide particle Overall reaction: A+B

hn PhC

Aoxidized + Breduced

Basic scheme of the possible photocatalytic action of an atmospheric dust particle with semiconductor properties. The particle is supposed to be covered with a layer of adsorbed atmospheric water. A and B are atmospheric components undergoing, respectively, oxidation and reduction by light-generated holes ⊕ and electrons ; Aa and Ba are forms of A and B that are adsorbed on the particle surface (or absorbed by the water layer). Ee shows the direction of energy change for electrons in the semiconductor particle. Eg is the width of the forbidden band. For simplicity, possible interactions of electrons and holes with water are not shown. Fig. 3

the photogeneration of respectively short-living but highly active catalytic reaction sites which are able to assist thermal transformations of substrates. Species which can serve as these sites include mainly free or bound ions of some transition metals in unusually low oxidation states, which are reactive with respect to dioxygen or to natural substrates even at low temperature. , V+ and some other low-valence Photogenerated Cu+ ads ads ad-cations are known as such superactive, but short living, catalysts [35, 36]. In addition to photocatalytic phenomena, one may expect also an important contribution from another type of light-generated phenomenon, namely photostimulated adsorption processes. The most important solid materials with photocatalytic properties, that are present in the atmosphere, are metal oxides and sulfides. Many of these are semiconductors, and some absorb visible and near-IR light. Of most interest for photocatalytic and photoadsorption processes in the Earth’s troposphere seem to be dust particles containing oxides Fe2 O3 , TiO2 , and ZnO, as well as SiO2 , MgO, and CaO. These oxides are the most plausible candidates for the role of photocatalysts and photosorbents, owing

2439

to their appropriate photochemical properties and rather high concentrations in the atmosphere. When a semiconductor particle absorbs a light quantum with an energy that exceeds its band gap, an electron from the valence band is excited into the conduction band. As a result, this electron can act as a chemical reductant with respect to the molecules adsorbed on the surface of the particle, while the hole in the valence band can play the role of a chemical oxidant (Fig. 3). The width Eg of the forbidden band (i.e., the band gap) of a semiconductor serves as the threshold energy that determines the red (i.e., long-wave) boundary λ0 of the optical absorption spectrum of the material. This parameter is very important for the photocatalytic properties of a semiconductor particle. Data on Eg for some oxides present in the atmosphere are listed in Table 2. Indeed, many semiconductor oxides such as TiO2 , ZnO, and Fe2 O3 are known as efficient photocatalysts for various redox reactions. These photocatalysts are active even under exposure to relatively ‘‘mild’’ light radiation, including quanta of visible light (its spectral range locates at 400–900 nm) (see Table 2). All of the oxides listed in Table 2, except for ‘‘wide band gap’’ pure MgO, SiO2 and Al2 O3 , absorb light with λ > 300 nm; that is, they are able to act as photocatalysts in all layers of the Earth’s atmosphere. For the most common (in Nature) insulator oxides such as SiO2 , Al2 O3 , MgO, and CaO, on exposure to solar light radiation, photoadsorption processes seem to be more typical [33, 37, 38]. The values Eg and λ0 listed in Table 2 refer to pure bulk semiconductor materials. With particles smaller than Band gaps Eg and red boundaries λ0 of the optical absorption of typical atmospheric semiconductor and dielectric oxides [3]

Tab. 2

Oxide

Eg /eVa

λ0 /nm

NiO Cr2 O3 CuO CdO Fe2 O3 TiO2 ZnO MgO SiO2 Al2 O3

0.93 1.4 1.7 2.1 2.2 3.0 3.2 7.2 8.6 9.0

1340 890 735 595 570 420 390 178 145 138

a The value refers to the bulk material and does not take into account

the particle size. References see page 2446

2440

11.8 Heterogeneous Catalysis in the Troposphere

≈10−8 m, these values can be shifted to wider band gaps due to quantum effects, whereas, impurities are known to shift Eg to smaller values. To a first approximation, in the estimates given in Ref. [3], these phenomena that balance each other were not taken into account. 11.8.4.1 Role of Heterogeneous Photocatalytic Processes in the Global Chemistry of the Atmosphere The possible impact on the global chemistry of the atmosphere by abiotic photocatalytic reactions over dust particles has been estimated [3]. In order to estimate, quantitatively, the possible influence of heterogeneous photocatalysis on the composition of the atmosphere, one can compare the rates of the expected photochemical reactions of various atmospheric components over aerosol photocatalysts with the rates of the natural removal of these components from (or supply into) the atmosphere (see Table 1). The rate v of a photocatalytic process on solid semiconductor particles is determined by the spectral density, Ia (λ), of the absorbed light power and the quantum yield, ϕ(λ), of the reaction at a particular wavelength λ: v = Ia (λ) · [hc/λ]−1 · ϕ(λ) · dλ

Here, h is the Planck’s constant and c is the velocity of light in vacuum, so that Ia (λ)·[hc/λ]−1 is the number of quanta with the wavelength λ that are absorbed per second. Ia (λ) was evaluated in Ref. [3] using data on the spectrum and intensity of solar radiation, concentrations of photocatalytically active particles in the atmosphere, and their absorption spectra. As mentioned above, Fe2 O3 , TiO2 and ZnO deserve special attention, as they are present in the atmosphere in rather large quantities and are capable of providing a variety of photocatalytic reactions [1, 3, 6, 27–31, 33, 39, 67]. The spectral density of the absorbed light power (i.e., radiation power, absorbed at a wavelength λ in 1 L) for not very turbid air can be evaluated as: Ia (λ) = I0 (λ) · (1 − e−ε(λ)l ) ≈ I0 (λ) · ε(λ) ·  where I0 (λ) is the spectral density of the flux of the incident solar light (see Fig. 2),  is the total thickness of the oxide layer (in cm), and ε(λ) is the extinction coefficient of the oxide. Ia(λ) was evaluated in Ref. [3] under the assumptions that the average background concentration of dust in the atmosphere is equal to the lower limit of its estimate for the continental area (i.e., 10−5 g · m−3 ), and that at an altitude of less than 3 km 35% of the overall dust quantity is soil particles. It was also

assumed that the average content of chemical elements in soil particles corresponds to that in the Earth’s crust (4.65% Fe, 0.57% Ti, 8.3 × 10−3 % Zn [3]). This gives the following concentrations for atmospheric dust: Fe2 O3 2.3 × 10−7 g m−3 ; TiO2 3.3 × 10−8 g m−3 ; ZnO 3.6 × 10−10 g m−3 . This means that the quantities of these oxides in 1 L of air are equal to those in a layer with an area of 100 cm2 and thickness 4.4 × 10−13 cm for Fe2 O3 , of 8.3 × 10−14 cm for TiO2 , and of 6.3 × 10−16 cm for ZnO. The above figures for the thickness of the hypothetical oxide layers are very small. However, these photochemically active layers can provide a sufficient influence due to the enormous size of the atmosphere and the huge total flux of solar light coming to Earth. The value of ε is approximately constant and equal to 2.6 × 104 cm−1 in the absorption band of TiO2 ; that is, at λ < λ0 = 420 nm, and decreases sharply at longer wavelengths. In rough estimations, the value of ε(λ) for Fe2 O3 and ZnO can also be assumed constant and equal to ≈105 cm−1 in their absorption bands (i.e., at λ < λ0 = 570 nm for Fe2 O3 and λ < λ0 = 390 nm for ZnO; see Table 2). Taking into account that the spectral density of the flux of incident solar light can be approximated by the black-body radiation with a temperature of 5800 K, and that the annual average density of solar light power on the upper border of the Earth’s atmosphere (over the hemisphere exposed to light) is 0.139 W cm−2 , the quantity N of light quanta absorbed by oxides in 1 L of air per second can be calculated. For light with wavelengths between λ0 (red boundary of the oxide absorption band) and 300 nm (the UV boundary for the light that penetrates into the troposphere where the major part of the atmospheric dust is located), the following values of N were obtained: 4.2 × 1011 quanta L−1 s−1 (7.0 × 10−13 mol L−1 s−1 ) for Fe2 O3 ; 6.4 × 109 quanta L−1 s−1 (1.1 × 10−14 mol L−1 s−1 ) for TiO2 ; and 1.2 × 108 quanta L−1 s−1 (2.0 × 10−16 mol L−1 s−1 ) for ZnO [3]. It should be remembered that these values of N refer to the lower limit of at the background concentration of aerosols in the troposphere. In clouds of volcanic dust, or over major industrial regions, the respective figures can be higher by several orders of magnitude (this is easily observed by eye for even a slightly dusty atmosphere). It should be remembered also [19], that the concentration of Zn in air is usually 100-fold higher than the value taken for the estimation in Ref. [3]; this should make N for ZnO also 100-fold larger. At present, it is impossible to forecast a priori quantum yields ϕ of photocatalytic reactions on solid surfaces. For this reason, the estimation of the role of photocatalysis in the global chemistry of the atmosphere was achieved as follows [3]. It was estimated, how large ϕ must be to provide removal or supply of a noticeable part (e.g.,

11.8.4 Heterogeneous Photocatalytic Reactions in the Troposphere

10%) of an atmospheric component via a photocatalytic route. This evaluation was based on the removal (supply) rates provided in Table 1, and has shown at which values of ϕ the photocatalytic formation or decomposition of a particular compound should become important for the composition of the atmosphere. These estimates can be then compared with the values of ϕ that are actually known for the formation or decomposition of the same compound. From this comparison, conclusions can be drawn about whether the known values of ϕ are high enough for particular photocatalytic processes to be important to the global chemistry of the atmosphere. The quantum yield ϕ0.1 that is needed to provide a 10% contribution of the photocatalytic pathway to the observed removal (or formation) rate was estimated [3] using the equation:

ϕ0.1

total rate of removal (or formation) of the substance (mol · L−1 · s−1 ) = 0.1 · number of the light quanta absorbed per second (mol · L−1 · s−1 )

where the rates of removal can be taken from Table 1. The corresponding values of ϕ0.1 are given in Table 1. Experimental Data on Photocatalytic Reactions of Atmospheric Components According to the expected mechanisms of both homogeneous and heterogeneous photocatalysis (see Section 11.8.4.1), mainly redox chemical reactions should be provided. For subsequent discussion of the possible role of photocatalytic reactions in the chemistry of the atmosphere, it is convenient for these to be classified into two groups: (i) reactions of the main atmospheric components that lead to the formation of trace compounds; and (ii) reactions of trace compounds that lead to their removal from the atmosphere. In both of these groups, other trace compounds may also be formed. 11.8.4.2

11.8.4.2.1 Photocatalytic Reactions of the Main AtmoThe spheric Components over Semiconductor Oxides following reactions belong to this group [3, 39]:

• Formation of ammonia and hydrazine from dinitrogen and water: hv

N2 + H2 O−−−→NH3 + O2 TiO2 hv

N2 + H2 O−−−→N2 H4 + O2 TiO2

These processes were observed in many experiments with TiO2 powder in a nitrogen atmosphere in the

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presence of water adsorbed onto the TiO2 surface. The addition of small quantities of α-Fe2 O3 to TiO2 essentially increases the reaction rate. Dinitrogen is also known to be reduced with water to NH3 on Fe2 O3 or vanadium-promoted hydrous ferric oxide. • Oxidation of dinitrogen: hv

N2 + H2 O−−−−−→H2 + HNO2 ZnO/Fe2 O3

• Water decomposition: hv

2H2 O−−−→2H2 + O2 TiO2

Photocatalytic water decomposition on TiO2 was reported in both liquid and gas phases. In the latter case, the water is adsorbed onto the TiO2 surface. • Oxidation of water by dioxygen into hydrogen peroxide: hv

H2 O + O2 −−−→H2 O2 TiO2

This reaction most probably proceeds through the formation of OH radicals on the surface of TiO2 . • Formation of organic compounds from CO2 and H2 O, for example: hv

CO2 + H2 O−−−−−−−−−−−−−−−→HCOOH, TiO2 ,CuOx /TiO2 ,ZnO/TiO2

CH2 O, CH3 OH These processes occur in TiO2 or, better, ZnO−TiO2 and CuOx −TiO2 suspensions in water [33, 35, 40, 41]. ZrO2 and other oxides were also found to be active in the photocatalytic reduction of CO2 . According to the available experimental data, ϕ takes values of 10−3 to 10−4 for the photocatalytic formation of H2 and O2 from H2 O and of organic compounds from CO2 and H2 O over TiO2 ; for CuOx –TiO2 for the CO2 reduction the reported ϕ figures attain even 10% [40]. Thus, these values exceed, or are of the same order of magnitude, as the values of ϕ0.1 for the H2 and CH2 O formation over TiO2 under normal conditions (see Table 1). Thus, photocatalytic reactions on TiO2 may be important as a source of hydrogen and organic compounds in the atmosphere. For the more abundant Fe2 O3 semiconductor, quantum yields of 10−3 to 10−4 would be sufficient to make the photocatalytic pathway important for the regular supply into (removal from) the atmosphere of all gases listed in Table 1, except for CO2 . References see page 2446

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11.8 Heterogeneous Catalysis in the Troposphere

11.8.4.2.2 Photocatalytic Reactions of Trace Components of the Atmosphere over Semiconductor Oxides Photocatalytic reactions of this group are intensively studied for the purposes of wastewater and air purification and photochemical syntheses. Several reviews [1, 27–33, 35, 37, 42–47, 67] have been published on this subject. Here, only few typical examples of such processes are presented.

• Complete oxidation (mineralization) of various hydrocarbons, for example: hv

C2 H6 + O2 −−−→CO2 + H2 O TiO2 hv

C6 H6 + O2 −−−→CO2 + H2 O TiO2

Such reactions may be important for removal from the atmosphere and soil of various hydrocarbons after oil spills. • Complete oxidation (mineralization) of halogenated hydrocarbons, for example: hv

C6 H5 Cl + O2 −−−→CO2 + HCl + H2 O TiO2 hv

CHCl3 + O2 + H2 O−−−→CO2 + HCl + H2 O TiO2

Photocatalytic decomposition of CCl4 , CHCl3 , CFCl3 , CF2 Cl2 , C2 F3 Cl3 and C2 F4 Cl2 in the presence of O2 over particles of ZnO, Fe2 O3 , TiO2 , volcanic ash, chalk, and desert sands has been reported. Such reactions may remove from the atmosphere a certain amount of ozone-depleting halides of methane and ethane. They can also lead to the formation of acid rains containing HCl. • Complete oxidation (mineralization) of oxygencontaining as well as heteroatomic organic compounds, for example alcohols, salicylic acid, phenol, benzoic acid, α-naphthol, herbicides, and even many chemical warfare agents [35, 42, 44, 46]. • Hydrogen production by photocatalytic decomposition of various organic substances associated with complete or partial oxidation of these substances. • Reactions of partial oxidation that are observed on disperse TiO2 , ZnO as well as vanadium oxides [48], for example: C6 H6 + O2 −−−→ C6 H5 OH • H2 S and NOx decomposition: hv

H2 S −−−−→ H2 + S CdS,ZnS

hv

NOx −−−−−−−−−−−−−−−−−→ N2 + O2 metal halides, aluminosillicates

Such reactions may be important for the H2 S and NOx removal from the atmosphere. • SO2 oxidation: hv

SO2 + O2 + H2 O −−−→ H2 SO4 dust

This reaction may be a reason for the formation of acid rain, from SO2 that has been emitted into the atmosphere, since under atmospheric conditions the rate of this photocatalytic reaction can be higher than for the corresponding thermal catalytic reaction (see Section 3). • Photocatalytic destruction of ozone, which is observed on dust of mineral or volcanic origin [9]. The photocatalytic reactions listed above were observed and studied mostly in water suspensions or colloids of semiconductor particles. Taking into account that aerosol particles in the troposphere are usually covered with a layer of adsorbed water, one can suppose that the same reactions might also occur in the atmosphere. Today, much reliable data exist on quantum yields of the above-mentioned photocatalytic reactions on oxide particles, in both water suspensions and gas phase. This allows a quite unambiguous judgment to be made on the ‘‘real’’ role of solid aerosols in the removal of various trace compounds from the atmosphere. For example, the quantum yield 0.1 was reported for isoprene oxidation in the gas phase on solid ZnO upon irradiation with λ > 300 nm [49]. For the CHCl3 oxidation to CO2 and HCl on TiO2 , ϕ was 2 × 10−2 for near-UV light, while on α-Fe2 O3 ϕ was less than 10−3 [50]. The aerobic oxidation of CH3 CH2 OH under illumination of TiO2 with λ < 400 nm at ambient conditions occurs with ϕ ≈ 0.09 in respect to one transferred electron [6, 45]. It should be noted, that experimental conditions in the cited studies differed markedly from those in the real atmosphere; therefore, the values of ϕ in the atmosphere may also differ from the above data. Anyhow, the published values substantially exceed the critical values of ϕ0.1 for many of the photocatalytic reactions on TiO2 (see Table 1). The latter may be regarded as evidence for the actual importance of photocatalysis for the global chemistry of the atmosphere. 11.8.4.2.3 Photocatalysis over Compounds with Wide Band Gaps In addition to semiconductor oxides, some other aerosol species may also contribute to photocatalysis in both the troposphere and the stratosphere. First, are the alkali halide particles which are generated in huge amounts over the surface of the world’s oceans. Alkali halides have been known for almost 80 years to be photocatalysts that assist a large variety of

11.8.4 Heterogeneous Photocatalytic Reactions in the Troposphere

reactions [33, 51], including the uptake of HO2 radicals, the photocatalytic oxidation of H2 , CO, NO2 , and CH4 , and the decomposition of H2 O and CO2 [33, 52–55]. The fundamental absorption bands of alkali halides are at wavelengths below 250 nm; that is, they correspond to the absorption of stratospheric UV. However, alkali halides often contain color centers (generated either by impurities or by rigid UV-light) which may exhibit photoadsorptive and photocatalytic activity at wavelengths up to 800 nm [52]. Similar behavior can be expected of sulfate aerosols, which are the main component of stratospheric aerosols (the so-called Aitken particles). The oxides with wide band gaps, such as γ -Al2 O3 and MgO, are also known to be efficient photocatalysts for the oxidation of ammonia, nitrogen oxides, and hydrocarbons [56–59]. When evaluating the possible role of these most abundant components of continental aerosols in the tropospheric photochemistry, their importance may be due to their photoactivity under irradiation not only in their main absorption bands, but also in the bands of surface chemisorbed species. For example, for dispersed SnO2 , it was found that the surface carbonates formed upon CO2 adsorption sensitize some surface photoreactions to light quanta with energies less than 2 eV (i.e., with λ > 600 nm) [60]. Recently, it was found that carbon and some other carboneous species such as fullerenes are also active in UV photocatalysis. In addition, these can experience a photocatalytic destruction [61, 62]. This suggests a possible contribution to atmospheric photocatalysis by soot particles. 11.8.4.2.4 Photocatalysis by Dissolved Transition Metal Ions A contribution to UV-induced photocatalysis is also expected from the cations of some transition metals dissolved in water droplets, or in the water layer that may cover a solid aerosol. Probable photoreactions in this case would be water-splitting, or redox processes involving atmospheric contaminants that are easily absorbed by the water layer [32, 43, 63]. The primary step in these reactions is expected to be redox transformation of the electronically excited metal compounds. In Nature, among the most abundant transition metals are iron and copper, and hence their ions are widely present in natural waters. These waters are also sometimes polluted with chromium ions originating from technogeneous sources. Recently, Fe(III), Cu(II), as well as Cr(III) and Cr(VI) compounds dissolved in natural waters were recognized as efficient photocatalysts which lead, under aerobic conditions, to the deep oxidation of a variety of both organic and inorganic compounds polluting the environment [25, 32, 43]. The above-mentioned ions

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are present in Nature in the form of either chelates of numerous organic compounds or hydroxides. Under the action of mild UV or even visible light, these ions undergo intramolecular ligand-to-metal charge transfer. The result is the formation of highly reactive organic or OH radicals from the ligands, together with the formation of reduced forms of the metal ions, such as Fe(II) and Cu(I). Such reduced forms of the metals are readily reoxidized by air dioxygen in thermal processes. This closes the photocatalytic cycles that result in the oxidation of different compounds, both organic and inorganic (e.g., compounds of S(IV)). Another feature of photocatalysis involving water-coated solid oxides in the atmosphere is expected to be the generation of hydroxyl and/or peroxide radicals. These species can react rapidly with many water-soluble trace compounds present in the atmosphere, thus providing a pathway for their photocatalytic removal. Photogenerated Catalysis Ultraviolet light is able to generate, on solid surfaces, new states of transition metal ions that can serve as active sites for thermal catalytic reactions. A typical example here is generation of longlived low-valence states of vanadium and copper on the surface of SiO2 [64, 65]. These states catalyze the thermal oxidation of alkenes, even at room temperature. Other examples of the generation with UV light of active catalytic sites on the surface of dispersed oxides are reviewed in Ref. [36]. Unfortunately, to date no data are available for the quantitative estimation of the role of photogenerated catalysis and photocatalytic reactions over halides and oxides with wide band gaps in the global chemistry of the atmosphere. Therefore, in the estimate made above, only photocatalytic reactions over the oxide semiconductors Fe2 O3 , TiO2 and ZnO with relatively narrow band gaps were taken into account. However, new compounds may be added in the future to the list of photocatalysts that are important to the chemistry of the atmosphere. 11.8.4.2.5

11.8.4.2.6 Destructive Photoadsorption by Alkaline Earth An important contribution to cleaning the Oxides atmosphere from halogenated hydrocarbons, including freons, may be expected from MgO- and CaO-containing aerosols. These are present in large amounts in volcanic clouds, as well as in fire or industrial aerosols [4, 19]. Recently, these oxides were found to be efficient photoadsorbents of halogenated compounds, performing stoichiometric substitution of the oxides’ oxygen anions for halogen anions with quantum yields near unity, even under visible or mild UV light [6, 33, 38]. References see page 2446

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11.8 Heterogeneous Catalysis in the Troposphere

It is necessary to note, that the band-to-band absorption region (i.e., absorption with simultaneous formation of a mobile hole in the valence band and a mobile electron in the conductivity band) for the above-mentioned oxidesinsulators is known to be located in the wavelengths region shorter 200 nm (region of the ‘‘vacuum UV’’ radiation) [6, 33, 38]. Nevertheless, air-exposed magnesia reveals a pronounced absorption in the spectrum region of wavelengths 300 to 400 nm, while the same samples of magnesia, after high-temperature oxygen-vacuum pretreatment, do not reveal any absorption in this part of the spectrum. Calcium oxide exposed to ambient atmosphere is known to transmute to calcium hydroxide, and then to calcium hydroxycarbonate, because of the presence of water vapor and CO2 . For long-term air-exposed calcium oxide, an absorption at wavelengths of ≈400 nm is also observed. Following high-temperature pretreatment of calcium oxide, absorption in this spectrum region is also sharply decreased. Thus, absorption in the spectrum region 300–400 nm seems to be connected with the presence of an adsorbed layer or isolated adsorbed molecules of some species. For air-exposed oxides of magnesium and calcium, the absorption of photons at wavelengths shorter than 400 nm, was found to result in the efficient destructive photoadsorption of fluorinated and chlorinated organic compounds (freons). For example, the absolute quantum yield of the photoadsorption of CF3 CHF2 (freon 134a) on air-exposed magnesia attains 0.06 at wavelengths λ = 313 nm. The action spectrum of photoadsorption of the freon is close to the absorption spectrum for the sample of magnesia obtained upon long-term exposure in air. The action spectra of photoadsorption for CHFCl2 (freon 21), CHF2 Cl (freon 22) and CHF3 (freon 23) are similar to the action spectrum of freon 134a photoadsorption, but the quantum yields appear to be lower. Freon 134a photoadsorption on the air-exposed CaO is also observed at wavelengths shorter then 400 nm, as well as its optical absorption. The most important distinctive feature of freones photoadsorption on illumination of the above-mentioned oxides is the quantity of freon which coats the surface. Thus, by illuminating the region of surface absorption, freon coverage can exceed 10% of the monolayer of the surface. In contrast, coverage of the surface of a metal oxide does not exceed 0.1% of the monolayer for the photoadsorption of various ‘‘simple’’ gases at the same illumination [66]. The freones photoadsorption is irreversible and destructive, as neither the freon itself nor its chlorine- or fluorine-containing products of decomposition, release

in the gas phase at the photoadsorbent heating up to temperatures 620 K. The above data allow the following overall reactions for irreversible destructive photoadsorption of hydrogencontaining freones on air-exposed magnesia and calcia to be suggested: RHF + MgO −−−→ MgF2 /MgO + H2 O + CO2 RHF + Ca(OH)2 −−−→ CaF2 /Ca(OH)2 + H2 O + CO2 (here R means a totally or partly halogenated hydrocarbon moieties). According to these reactions, the photoadsorbed freons replace the surface lattice oxygen or surface hydroxyl groups by the fluorine atoms. This explains the high degree of surface coverage by photoadsorbed freons, as well as an absence of fluorine-containing products in the gas phase after freon photodestruction, or after consequent heating of the sorbent to temperatures up to 620 K. A confirmation of such replacement is the absence of any remarkable photoadsorption of compounds not containing fluorine atoms, such as chloroform (CHCl3 ), dichloroethane (CH2 ClCH2 Cl), and trichloroethane (CCl3 CH3 ), to create a water molecule which improves the energetics of the overall reaction. It is of importance that air-exposed magnesia and calcia do not adsorb freones that do not contain hydrogen atoms – CF4 (freon 14) and CF2 ClCFCl2 (freon 113) – both in darkness and under UV exposure. The role of hydrogen atoms in photoadsorption remains to the subject of much debate, but it is possible that these atoms are needed at certain steps of the surface transformation. 11.8.5

Role of Thermal Absorption and Adsorption in Tropospheric Catalysis and Photocatalysis

The important role of absorption into water droplets, for thermal hydrolytic atmospheric reactions, was referred to in Section 11.8.3. Much experimental evidence exists that the absorption of many substances in this case proceeds with conventional values of Henry’s coefficient; that is, the impurities are strongly concentrated in the droplets. Particles of solid atmospheric aerosols may be also covered by a layer of atmospheric water in the liquid or frozen state (see Fig. 3). Many of the compounds listed in Table 1 (e.g., H2 S, SO2 , CH2 O) are readily soluble in water, and may be strongly concentrated in this layer. Thus, their elevated concentrations may serve to enhance photocatalytic processes over aerosols covered with water. In order to be converted, either catalytically or photocatalyticall,y over a solid aerosol particle which is not covered by a water layer, atmospheric components must be adsorbed onto their surface. Experimental data

11.8.6 Conclusions

on such adsorption for real atmospheric aerosols, or their models under natural conditions, are few in number (see, e.g., Refs. [17, 18]). However, as suggested by the estimates made in Ref. [3] with CO2 and H2 S as particular examples, even at very low pressures of a trace component (e.g., 10−10 bar for H2 S) they still may be adsorbed onto a solid aerosol with a sufficiently high coverage, θ. Photocatalysis is assumed to be important for the formation, and also for some properties, of SOAs [68]. The presence of some elements (e.g., Al, Fe, Ti, Mn, V) inside aerosol particles may enhance the processes of heterogeneous oxidations, for example of gaseous hydrocarbons, and their further incorporation into the aerosol mass (e.g., catalytic polymerizations). In contrast, aerosol particles in the atmosphere, containing hydrocarbon vapors, will grow at a higher rate and will be deposited more quicker than others. Hence, the chemical origin or catalytic activity of substances may influence their pathway in the atmosphere. 11.8.6

Conclusions

The above overview suggests that both, thermal catalytic and photocatalytic reactions on solid and liquid aerosol particles, can play an essential role in the global chemistry of the Earth’s atmosphere. Atmospheric photocatalytic reactions are, perhaps, more numerous than thermal catalytic reactions as the latter are more sensitive to the low temperatures of the atmosphere. In contrast to non-catalytic photochemistry, which occurs mostly in the stratosphere and mesosphere under the action of far UV-light, most photocatalytic chemistry is expected to occur in the troposphere under near-UV, visible, and even near-IR light. However, a possible contribution to the chemistry of the stratosphere and mezosphere should not be ignored. This leads to the suggestion that dust in the atmosphere may play an important role, by clearing the atmosphere of harmful compounds, both catalytically and photocatalytically. Thus, large desert areas – which are the main generators of continental dust – may actually play the role of the ‘‘kidneys’’ of the planet. At present, it is not possible to suggest more precise estimates than those made in this chapter, of the actual role of particular photocatalytic reactions in the atmosphere. In order to improve our knowledge in this respect, it will be necessary to continue investigations into the quantitative characteristics of heterogeneous photocatalysis and thermal catalysis over natural aerosols or their models, and under conditions close to those encountered in the atmosphere. The most important characteristic to be measured is the quantum yield

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of photocatalytic reactions of atmospheric components on atmospheric aerosols containing Fe2 O3 , TiO2 and ZnO, as these are the most plausible candidates for the role of photocatalysts due to their appropriate photochemical properties and rather high concentrations in the troposphere. Mathematical modeling of the expected impact of these heterogeneous phenomena over aerosols is also required. Heterogeneous photocatalytic and thermal catalytic processes may seriously influence the atmospheric chemistry in several ways. First, tropospheric photocatalytic processes may accelerate the complete oxidation of sulfur and nitrogen oxides (into H2 SO4 and HNO3 ), as well as the mineralization (via complete oxidation into H2 O, CO2 and HCl) of halogenated organic molecules, thus affecting the strength of acid rains. The catalytic hydrolysis of chlorine and bromine-containing compounds facilitates the evolution of Cl2 and Br2 that may in turn stimulate the subsequent reactions that deplete the ozone layer; alternatively, the photoadsorption of freons should largely protect the ozone layer. The photocatalytic mineralization of halogenated organics and production of dihydrogen and ammonia (or even hydrazine) may also influence the fate of the ozone layer. The photocatalytic oxidation of hydrocarbons may reduce the hazardous consequences of oil spills, while the photocatalytic oxidation of methane and mineralization or destructive photoadsorption of the halogenated organics may reduce concentrations of greenhouse gases in the atmosphere. Photocatalysis also influences the chemical composition of aerosol particles by, for example, accelerating the formation of sulfates and ammonium salts from oxides and salts of alkali metals. An important problem here is the identification of the products actually formed from various chemicals (in particular from complex organic compounds such as pesticides, refrigerants) on solid photocatalysts, under the conditions similar to those in Nature. In this respect, the available data are quite scarce, but the possibility cannot be excluded that, under certain unfavorable conditions in the presence of O2 , these compounds are oxidized not only to CO2 and H2 O, but also to more toxic or environmentally unfriendly compounds. This may be particularly true for compounds that contain such elements as Cl, P, or F. To conclude, it appears that the answer to the question of whether heterogeneous photocatalysis and thermal catalysis are important for global and local chemistry of the atmosphere, seems to be ‘‘yes’’ rather than ‘‘no’’. However, to make this conclusion more unambiguous, further experimental investigations must be conducted. References see page 2446

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11.8 Heterogeneous Catalysis in the Troposphere

References 1. D. R. Schryer (Ed.), Heterogeneous Atmospheric Chemistry, American Geophysical Union, Washington, 1982. 2. J. G. Calvert (Ed.), The Chemistry of the Atmosphere: Its Impact on Global Change. A ‘‘Chemistry for the 21st century’’ Monograph. Blackwell Scientific, Oxford, 1994. 3. K. I. Zamaraev, M. I. Khramov, V. N. Parmon, Catal. Rev.-Sci. Eng. 1994, 36, 617. 4. Yu. M. Gershenson, A. P. Purmal’, Uspekhi Khimii 1990, 59, 1729 (in Russian). 5. M. Jang, N. M. Czocshke, A. L. Northcross, Chem. Phys. Phys. Chem. 2004, 5, 1646. 6. V. N. Parmon, V. S. Zakharenko, Cat. Tech. 2001, 5, 96. 7. W. Schrader, Angew. Chem. Int. Ed. 2005, 44, 1444. 8. R. P. Wayne. Chemistry of Atmospheres: An Introduction to the Chemistry of the Atmospheres of Earth, the Planets, and their Satellites, 3rd Ed., Oxford University Press, New York, 2000. 9. S. Solomon, J. Geophys. Res. 1996, 101, 6713. 10. P. Brimblecombe, Air Composition & Chemistry, Cambridge University Press, Cambridge, 1986. 11. J. Kesselmeier, M. Staudt, J. Atmosph. Chem. 1999, 33, 23. 12. G. Visconti, Fundamentals of Physics and Chemistry of the Atmosphere, Springer, Berlin, 2001. 13. J. E. Andrews, P. Brimblecombe, T. D. Jickells, P. S. Liss, An Introduction to Environmental Chemistry, Blackwell Science, Oxford, 1996. 14. F. Raes, R. Van Dingenen, E. Vignatti, J. Wilson, J.-P. Putaud, J. H. Seinfeld, P. Adams, Atmosph. Environ. 2000, 34, 4215–4240. 15. V. E. Zuev, G. M. Krekov, Optical Models of the Atmosphere, Gidrometeoizdat, Leningrad, 1986 (in Russian). 16. A. Gelencs´er, Carbonaceous Aerosol, Springer, Berlin, 2004. 17. K. P. Kutsenogij (Ed.), Aerosols of Siberia, Nauka, Novosibirsk, 2006. 18. M. Sofiev, G. Petersen, O. Kr¨uger, B. Schneider, M. Hongisto, K. Jylha, Atmosph. Environ. 2001, 35, 2001. 19. S. G. Malakhov, E. P. Makhon’ko, Uspekhi Khimii 1990, 59, 1777 (in Russian). 20. J. Heicklen. Atmospheric Chemistry, Academic Press, New York, 1976. 21. A. Aguzzi, M. J. Rossi, J. Phys. Chem. A 2002, 106, 5891. 22. G. K. Moortgat, A. J. Barnes, G. Le Bras, J. R. Sodeau, LowTemperature Chemistry of the Atmosphere, Springer, Berlin, 1994. 23. J. R. McKeachie, M. F. Appel, U. Kirchner, R. N. Schindler, T. Bender, J. Phys. Chem. B 2004, 108, 16786. 24. J. Baltrusaitis, V. H. Grassian, J. Phys. Chem. B 2005, 109, 12227. 25. D. L. Sedlak, J. Holgn´e, Environ. Sci. Technol. 1994, 28, 1898. 26. N. Coichev, R. van Eldik, New J. Chem. 1994, 18, 123. 27. V. A. Isidorov, Organicheskaya Khimiya Atmosphery (Organic Chemistry of the Atmosphere), Khimiya, St-Petersburg, 1992 (in Russian). 28. K. I. Zamaraev, V. N. Parmon, Catal. Rev.- Sci. Eng. 1980, 22, 261. 29. M. Gr¨atzel (Ed.), Energy Resources Through Photochemistry and Catalysis, Academic Press, New York, 1983. 30. E. Pelizzetti, N. Serpone (Eds.), Photocatalysis. Fundamentals and Application. Wiley, New York, 1990. 31. E. Pelizzetti, M. Sciavello (Eds.), Photochemical Conversion of Solar Energy, Kluwer, Dordrecht, 1991, 660 pp.

32. P. Ciesla, P. Kocot, P. Mytych, Z. Stasicka, J. Mol. Catal. A: Chemical 2004, 224, 17. 33. Yu. A. Artem’ev, V. K. Ryabchuk, Introduction to Heterogeneous Photocatalysis, St. Petersburg University Publisher, St. Petersburg, 1999 (in Russian). 34. H. Nishikawa, S. Kato, T. Ando, J. Mol. Catal. A: Chemical 2005, 236, 145. 35. V. N. Parmon (Ed.), Photocatalysis and Solar Energy Conversion, in Catal. Today 1997, 39, 3. 36. B. N. Shelimov, V. B. Kazanskij, in Photocatalytic Conversion of Solar Energy, K. I. Zamaraev, V. N. Parmon (Eds.), Nauka, Novosibirsk, 1991, pp. 109–137 (in Russian). 37. M. Sciavello (Ed.), Photocatalysis and Environment. Trends and Applications. Kluwer, Dortrecht, 1988. 38. V. N. Parmon, V. S. Zakharenko, Colloids Surfaces A: Physicochem. Eng. Aspects 1999, 151, 367. 39. V. N. Parmon, Colloids Surfaces A: Physicochem. Eng. Aspects 1999, 151, 351. 40. Slamet, H. W. Nasution, E. Purnama, S. Kosela, et al., Catal. Commun. 2005, 6, 313. 41. X.-T. Wang, Sh.-H. Zhong, X.-F. Xiao, J. Mol. Catal. A: Chemical 2005, 229, 87. 42. S. Malato, P. Fernndez, J. Gimnez (Eds.), Catalysis Today, Environmental Application of Photocatalysis, 2005, Vol. 101, Nos. 3–4. 43. K. Stemmler, U. von Gunten, Atmosph. Environ. 2000, 34, 4253. 44. V. N. Parmon (Ed.), Photocatalysis and its Practical Application, in Catal. Today 2000, 58, Nos. 2–3. 45. A. V. Vorontsov, D. V. Kozlov, P. G. Smirniotis, V. N. Parmon, Kinet. Catal. 2005, 46, 203. 46. D. F. Ollis, H. Al-Ekabi (Eds.), Photocatalytic Purification and Treatment of Water and Air, Elsevier, Amsterdam, 1993. 47. Yu. A. Gruzdkov, E. N. Savinov, V. N. Parmon, in Fotokataliticheskoe Preobrazovanie Solnechnoi Energii, V. N. Parmon, K. I. Zamaraev (Eds.), Novosibirsk, Nauka, 1991, p. 138 (in Russian). 48. K. Teramura, Ts. Tanaka, T. Yamamoto, T. Funabiki, J. Mol. Catal. A: Chemical 2001, 165, 229. 49. V. A. Isidorov, E. M. Klokova, P. V. Zgonnik, Vestnik Lenigradskogo Universiteta, ser. 4, 1990, Vyp. 1 (No. 4), 71; Vestnik Lenigradskogo Universiteta, ser. 4, 1990, Vyp. 3 (No. 18), 61 (in Russian). 50. C. Kormann, D. W. Bahnemann, M. R. Hoffmann, J. Photochem. Photobiol., A: Chemistry 1989, 48, 161. 51. A. N. Terenin, Zhurn. Fiz. Khim. 1935, 6, 189. 52. V. K. Ryabchuk, L. L. Basov, Yu. P. Solonitzyn, React. Kinet. Catal. Lett. 1988, 36, 119 53. L. L. Basov, Yu. P. Efimov, Yu. P. Solonitzyn, in Advances in Photonics, vyp. 4, Leningrad State University, Leningrad, 1974, p. 12 (in Russian). 54. D. D. Weis, G. E. Ewing, J. Phys. Chem. A 1999, 103, 4865. 55. R. G. Remorov, Yu. M. Gershenzon, L. T. Molina, J. Phys. Chem. A 2002, 106, 4558. 56. A. V. Alekseev, D. V. Pozdnyakov, A. A. Tsyganenko, V. N. Filimonov, React. Kinet. Catal. Lett. 1976, 5, 9. 57. C. Yun, M. Anpo, Y. Mizokoshi, Chem. Lett. 1980, 7, 788. 58. A. Mansons, J. Chim. Phys. Phys. Chim. Biol. 1987, 84, 569. 59. H. Balard, A. Mansons, J. Chim. Phys. Phys. Chim. Biol. 1987, 84, 907. 60. V. S. Zakharenko, A. E. Cherkashin, React. Kinet. Catal. Lett. 1983, 23, 131

11.9.2 Conversion of Biomass to Biobased Products 61. M. Gevaert, P. V. Kamat, J. Chem. Soc., Chem. Commun. 1992, 1470. 62. L. Stradella, React. Kinet. Catal. Lett. 1993, 51, 299. 63. V. N. Parmon, in Photocatalytic Conversion of Solar Energy, Part 2, K. I. Zamaraev, V. N. Parmon (Eds.), Nauka, Novosibirsk, 1985, pp. 6–106 (in Russian). 64. E. V. Kashuba, L. V. Lyashenko, V. M. Belousov, React. Kinet. Catal. Lett. 1986, 30, 137. 65. V. M. Belousov, E. V. Kashuba, L. V. Lyashenko, React. Kinet. Catal. Lett. 1986, 32, 33. 66. M. D. Driessen, A. L. Goodman, T. M. Miller, G. A. Zaharias, V. H. Grassian, J.Phys. Chem. B 1998, 102, 549. 67. A. I. Kokorin, D. W. Bahnemann (Eds.), Chemical Physics of Nanostructured Semiconductors, VSP, Utrecht, 2003. 68. P. Kutsenogiy, in Macro and Trace Elements, 22nd Workshop, September 24–25, 2004, Vol. 2, Friedrich Schiller University Press, Jena, Germany, 2004, pp. 1436–1441.

11.9

Conversion of Biomass on Solid Catalysts Pierre Gallezot∗ and Alain Kiennemann

11.9.1

Introduction

2447

yield of dry matter per area than conventional cereals or vegetable oils crops. Increasing use should also be made of agricultural wastes and of fast-growing, low-cost lignocellulosic raw materials. Another issue is the high processing costs yielding products that are not cost-competitive with those derived from fuels. Accordingly, extensive R&D efforts in biotechnology, chemistry and engineering will be required to reduce processing costs. Most of the non-food applications of biobased resources fall into four categories: • Traditional uses in timber, paper, fiber, and rubber industries, and fine chemical extraction (flavor and fragrance, dyes, bioactive molecules) • Thermal power generation (biopower) by direct combustion of biomass or after gasification • Production of biobased transportation fuels, for example the production of ethanol by the fermentation of carbohydrates, of biodiesel by transesterification of vegetable oils, of fuel by steam reforming/Fischer–Tropsch (FT) of lignocellulosic materials, and of hydrogen by steam reforming/water gas shift (WGS) of biomass • Production of biobased products by catalytic conversion of carbohydrates, triglycerides, and terpenes. The present chapter deals with the various applications of heterogeneous catalysis for converting biomass either to biobased chemicals and polymers (Section 11.9.2) or to heating and transportation fuels (Section 11.9.3). However, it should be stressed that enzymatic catalysis and homogeneous catalysis play important roles in biomass conversion; indeed, fermentation and hydrolysis are often the first steps employed in biomass processing. Therefore, in many instances, efficient processes combining biotechnologies and heterogeneous catalysis must be developed.

Biosynthesis in plants and trees using sun radiation, atmospheric carbon dioxide, water, and soil nutrients produces huge amounts of biomass estimated up to 200 Gt y−1 , a figure to be compared to 7 Gt y−1 of extracted fossil fuels. Increasing use of biomass for energy, chemicals and material supply is one of the key issues of sustainable development because biobased resources are renewable and CO2 neutral in contrast with fossil fuels. The molecules extracted from biobased resources already contain functional groups, so that the synthesis of chemicals generally requires a lower number of steps than from alkanes. Also, biobased products may have unique properties compared to hydrocarbon-derived products, for instance biodegradability and biocompatibility. In addition, products issued from biomass have a higher added value because of the ‘‘natural’’ or ‘‘bio’’ label. The Biomass Program of the US Department of Energy [1] and the Strategic Research Agenda of SUSCHEM organization in Europe [2], promote the increasing use of renewables for energy and chemical production. However, various hurdles may hamper the development of renewables as feedstock for bioenergy or bioproduct production. Thus, renewables from agricultural crops are primarily processed for food and feed production. To cope with the competition problem, crops dedicated to non-food use should be developed with plant species, giving a better

Strategy for Biomass Conversion to Chemicals Three main sources of renewables – carbohydrates, plant (vegetable) oils, and terpenes – are presently used to produce biobased products via catalytic routes. The first two are issued from crops primarily used for food and feed production, but large amounts (62 × 106 t y−1 of carbohydrates and 14 × 106 t y−1 of plant oils) are also employed in industrial applications. Terpenes (0.4 × 106 t y−1 ) used exclusively in industrial applications are mainly derived from tall oil, a coproduct of the paper pulp industry. In addition, the use of lignocellulosic resources, that can potentially be produced in much larger amounts

∗

References see page 2471

Corresponding author.

11.9.2

Conversion of Biomass to Biobased Products 11.9.2.1

2448

11.9 Conversion of Biomass on Solid Catalysts

than traditional crops, and do not compete with food and feed needs, should be developed by increased research effort for the efficient processing of cellulose and lignin. Three types of strategy can be used to convert biomass into products: • From biomass to products via degraded molecules. In this scheme, biomass is first converted by gasification to synthesis gas, or by pyrolysis to a mixture of small molecules. Synthesis gas may then be converted to hydrocarbons which are subsequently converted to intermediates, using the classical synthesis routes developed for petroleum feedstock. Similarly, small molecules obtained by biomass pyrolysis may, after separation, be converted to valuable chemicals via the existing flow sheets of chemical synthesis. This is not a cost-effective and environmentally sustainable route for chemical production, as highly functionalized molecules from biomass are first degraded to C1 molecules or hydrocarbons which are then subjected to the traditional chemical synthesis steps in order to be functionalized again. • From biomass to products via platform molecules: the biorefinery concept. A biorefinery is a facility that integrates biomass conversion processes to produce fuels and chemicals. According to this scheme, part of the biomass is converted to fuels via pyrolysis and gasification, and the other part is converted by fermentation or chemocatalytic routes to well-identified platform molecules, that can be employed as building blocks in chemical synthesis. The fermentation processes are continuously improved with new, genetically modified bacteria or yeasts [3]. Thus, ethanol produced by fermentation (30 × 106 t y−1 ) is a platform molecule that can be used both as a fuel or fuel precursor and as a building block for chemistry. It is mainly produced by the fermentation of starch derivatives and sucrose, but intensive research is presently under way to identify efficient enzymes to bring cellulose-based ethanol to commercial use. The main platform molecules that have been identified by the US Department of Energy [4] as starting materials to produce chemicals and polymers via catalytic routes are listed in Table 1. In order to convert platform molecules by one or multistep catalytic processes, new routes must be established because the traditional organic synthesis flow sheets developed for hydrocarbons are not adapted. For example, because platform molecules contain several oxygenated groups, both reduction and hydrogenolysis steps would be needed, whereas chemical synthesis from hydrocarbons primarily requires oxidation catalysis. • Direct routes from biomass to products via one-pot reactions. The formulation of commercial products (cosmetics, paper, paints, binders, construction materials) often

Tab. 1

Platform molecules identified by the US DOE [4]

Aspartic acid Glutamic acid Levulinic acid 2-Hydroxypropionic acid 2,5-Furan dicarboxylic acid Glucaric acid

Itaconic acid 1,4-Diacids (succinic, fumaric, malic) 3-Hydroxypropionic acid Glycerol Sorbitol Xylitol/Arabitol

involves a mixture of molecules with the same functionalities (e.g., mixture of diols or polyols). This mixture may well be obtained from biomass (e.g., starch) via one-pot processes, possibly including two or several catalytic steps, avoiding product isolation and thus reducing considerably the processing cost of biomass to chemicals. Catalytic Conversion of Terpenes Heterogeneous catalysis is increasingly used for the conversion of terpenes and derivatives in place of homogeneous processes. The three main starting materials employed in terpene chemistry are α-pinene 1 and β-pinene 2, which are extracted from turpentine oil (350 000 t y−1 ), a coproduct of paper pulp industry, and limonene 3, which is extracted from citrus oil (30 000 t y−1 ). They are used for the synthesis of flavors and fragrances (F&F), although these compounds are often more easily obtained by catalytic routes from hydrocarbons. Thus, while citral synthesis requires five steps from α-pinene, it is obtained with a 95% yield in the BASF process starting from formaldehyde and isobutene [5]. 11.9.2.2

1

2

3

4

p-Cymene 4, a precursor of p-cresol and various F&Fs, was obtained by dehydrogenation of α-pinene at 300 ◦ C in a continuous fixed-bed flow reactor in the presence of 0.5 wt.% Pd/SiO2 [6]. Under similar conditions, but starting from limonene, p-cymene was obtained with a 97% yield and the catalytic activity was stable for 500 h on stream [7]. More interestingly, p-cymene was produced under similar reaction conditions with near-100% yield from a mixture of di-pentene isomers (Sylvapine DP-378), showing that raw materials need not be purified and can still be converted in one step to the desired product [8]. The conversion of α-pinene oxide 5 to campholenic aldehyde 6, an important intermediate in the synthesis of sandalwood-like fragrances, is traditionally carried out in the presence of ZnCl2 or ZnBr2 salts. Alternative routes

11.9.2 Conversion of Biomass to Biobased Products

Tab. 2

Conversion of α-pinene oxide to campholenic aldehyde on acidic solid catalysts

Catalyst

Reaction conditions

Ti-beta zeolite Sulfated alumina Silica-alumina H-US-Y zeolite, HCl-treated

Vapor phase reaction at 90 ◦ C Liquid phase at 0 ◦ C Liquid phase 25 ◦ C Liquid phase, 0 ◦ C

have been described using solid catalysts with Lewis acid sites giving high yield of campholenic aldehydes (Table 2) O O 5

6

The liquid-phase alkoxylation of limonene with C1 −C4 alcohols to 1-methyl-4-[alpha-alkoxy-isopropyl]1-cyclohexene 7 was carried out both in batch and continuous fixed-bed reactors at 60 ◦ C on various acidic catalysts [13]. The best yields were obtained in the batch (85%) or in the continuous reactor (81%), using a beta zeolite with SiO2 : Al2 O3 = 25. ROH H 3C

CH3 OR

3

7

Solid basic catalysts can advantageously be used to replace soluble catalysts in the conversion of terpene derivatives. Thus, the aldol condensation of acetone with citral 8 was achieved in the presence of rehydrated hydrotalcite with a 96% yield to a mixture of cis- and transpseudoinone 9 by operating in liquid phase at 60 ◦ C [14]. O

CHO O +

8

2449

Yield/ %

Reference

94 76, low deactivation 72 80

[9] [10] [11] [12]

Catalytic Conversion of Triglycerides Vegetable oils or triglycerides extracted from the seeds of various plants (rapeseed oil, sunflower oil, soybean oil, palm oil, etc.) are a mixture of fatty acid esters of glycerol with different chain lengths (C12 −C18 ) and C=C bonds. The content of the different fatty acids is different in the various vegetable oils, and can be modified by breeding or genetic modification of crops. There is a serious competition for triglycerides supply between food and feed needs (ca. 100 × 106 t y−1 ), on the one hand, and the growing use for producing biofuels and oleochemicals, on the other hand (ca. 15 × 106 t y−1 ). Oleochemicals produced from triglycerides are mainly employed as surfactants, lubricants, and polymers. 11.9.2.3

11.9.2.3.1 Catalytic Conversion of Triglycerides to Surfactants The scheme given in Fig. 1 summarizes the main catalytic transformations of triglycerides to produce either edible oils and fats or fatty acids, fatty esters, and fatty alcohols used for the production of oleochemicals. While the hydrogenation steps on metal catalysts have long been known and are the subject of continuous improvements in terms of activity, selectivity and stability, the esterification and transesterification reactions catalyzed by acids or bases are still mainly conducted with liquid acids and bases. However, the present trend is to develop ‘‘green’’ processes by replacing soluble catalysts with solid acid and base catalysts. The main surfactants and emulsifiers derived from triglycerides are fatty alcohol sulfates, fatty acid esters, alkyl glucosides, and ethoxylated fatty acid methyl esters. These account for approximately 106 t y−1 , which is a comparatively small amount with respect to surfactants such as alkylbenzene sulfonates produced from petroleum.

9

Solid catalysts are also widely used for selective hydrogenation of terpene derivatives, for example to convert α, β-aldehydes to the corresponding alcohols (cinnamaldehyde to cinnamyl alcohol, citral to geraniol and nerol, etc.). However, this is beyond the scope of the present chapter since, as mentioned previously, the major part of terpene derivatives is not produced from natural terpenes.

Fatty acid esters

ROH

Edible H2 Triglycerides oils & fats Metal

OH− H2O H+

Fig. 1

H2

Metal + ROH H

Fatty acids

H2

Fatty alcohols Metal

Scheme of the main catalytic transformations of triglyceride.

References see page 2471

2450

11.9 Conversion of Biomass on Solid Catalysts Solid acid catalysts for the esterification of glycerol and fatty acids (from Ref. [17])

Tab. 3

T/ ◦ C

Catalyst

Fatty acid

H-Beta (5 wt.%)a H-USY (5 wt.%) Amberlyst 31 (7.2 wt.%) Amberlyst 31 (7.4 wt.%) KA mol sieve (2.2 wt %) Amberlyst 15 (1.57 wt.%) H-USY (1.57 wt.%) Silica gel-SO3 H (1.57 wt.%) HMS-SO3 H (1.57 wt.%) MCM41-SO3 H (1.57 wt.%) MCM41-SO3 H (5 wt.%) end-capped MCM41-SO3 H (5 wt.%) MPMDS

Lauric Oleic Oleic Oleic

100 180 90 90

Lauric Lauric Lauric Lauric Lauric Lauric Lauric

110 110 110 110 110 100 100

a wt.%

Yield/%

Reference

10 5 24 24

20 83 49 77

[18] [19] [20] [21]

11.8 23.5 8 10 24 24 10

44 36 51 52 53 60 60

[22] [22] [21] [22] [22] [23] [24]

of catalyst with respect to the total weight of reaction medium.

HO

HO HO HO

Time/h

HO

HO

O OR

HO

O OH + ROH OH

HO HO

OH

O O OH

HO HO HO

HO HO

O OR

O OR OH

OH

Fig. 2

Scheme of glucose acetalization to glucofuranoside, glucopyranoside, and oligomers.

However, because of their biodegradability and biocompatibility, they have a much higher added value and their production is expanding, though their development may be hampered by the increasing demand for triglycerides used to produce fatty acid esters for biofuels. Fatty acid esters of glycerol are efficient surfactants obtained either by the transesterification of triglycerides with glycerol (glycerolysis), or by the esterification of fatty acids with glycerol. The challenge in both cases is to obtain selectively glycerol monoesters that are non-ionic surfactants with a good hydrophilic/hydrophobic balance employed in food, cosmetics and pharmaceuticals. Glycerolysis reactions have been conducted on basic oxides to replace liquid bases. Blancquart et al. [15] have employed various basic oxides (ZnO, MgO, CeO2 , La2 O3 ) to catalyze methyl stearate glycerolysis. The selectivity to monoglyceride was rather similar to that obtained with molecular catalysts, and the nature of the oxide had little effect. Doping MgO with lithium increased the basicity and conversion, but did not change selectivity. Glycerolysis of rapeseed oil on MgO catalysts gave a 63% yield to monoglyceride [16]. The preparation of glycerol monoester by esterification of fatty acids with glycerol was achieved with acidic solids

to substitute sulfuric acid. The esterification yields of lauric and oleic acids with glycerol on solid acids are listed in Table 3. Fatty acid esters of sugars are also very important biodegradable and biocompatible surfactants which are prepared either by transesterification of methyl ester with sugar on basic catalysts, or by esterification of fatty acids with sugar on acidic catalysts. Liquid acids and bases have been replaced by enzymatic catalysis with lipase giving a higher yield to monoester [25, 26], but solid catalysts have not yet been used extensively. Alkylglucosides are a class of valuable commercial surfactants, particularly for cosmetics applications because of their biocompatibility. They are obtained by the acetalization of carbohydrates with fatty alcohols in the presence of acid catalysts. Glucose acetalization gives a mixture of products with glucofuranoside or glucopyranose rings and various amounts of oligomers (see Fig. 2). Zeolites and MCM-41 have been used as acidic catalysts to achieve the acetalization of glucose with alcohols of different chain lengths [27, 28]. It was shown that shape-selectivity effects decrease the amount of oligomers formed, and that activity and selectivity can be controlled with the Si:Al ratio. In the glucosidation

11.9.2 Conversion of Biomass to Biobased Products

of n-butanol at 120 ◦ C, the furanoside:pyranoside ratio increased with the pore diameter of MCM-41 because of a lower diffusivity of the more bulky pyranoside ring. The kinetics and selectivity of glycosylation between glucose and n-butanol have been investigated in detail over various dealuminated HY zeolites by Chapat et al. [29]. 11.9.2.3.2 Catalytic Conversion of Triglycerides to Lubricants Because almost 50% of lubricants leak and spread out in the environment, there is a need for biodegradable variants, produced from renewables. Fatty acid esters may be suitable lubricants, but their resistance to oxidation and their tribological properties need to be improved. This can be achieved by epoxidation of the fatty acids, followed by alcoholysis of the epoxide (Fig. 3). The epoxidation of fatty acid methyl esters (FAME) is traditionally conducted in strong acidic media, for example, with peracetic acid in sulfuric acid solutions. It has been shown that these reactions can be conducted by an environmentally benign route in the presence of acidic solids. Thus, a mixture of FAME from sunflower oil was epoxided by tert-butylhydroperoxide (TBHP) at 90 ◦ C in the presence of Ti/MCM-41 catalysts, yielding 98% conversion with 85% selectivity to monoepoxy compounds [30]. Ti/SiO2 (aerosil) catalysts showed comparable activity and selectivity, indicating that the presence of the ordered mesoporous framework is not necessary. In the same way, Rios et al. [31] used different Ti/MCM-41 materials with pore diameters ranging from 1.9 to 4.1 nm and amorphous Ti/SiO2 catalysts with different Ti-dispersion to perform methyl oleate epoxidation with TBHP at 70 ◦ C. Selectivities higher than 95% were obtained, irrespective of the structure of the O H3CO H+ O

TBHP H2O2 O

H3CO H+ ROH O

OH

H3CO

supporting material, provided that the titanium was well dispersed. It was shown that no leaching occurs and that the catalysts could be recycled. The alcoholysis with different alcohols of epoxidized FAME was studied on acidic resins of various structure and acid strength [32, 33]. The addition of methanol to epoxidized methyl oleate at 60 ◦ C in the presence of Nafion entrapped in silica (SAC13) or of Amberlyst 15, a sulfonated styrene-divinylbenzene copolymer, resulted in complete conversion with selectivity higher than 98%, but the reaction rate was higher on the more acidic SAC13 catalyst (60 s−1 ) than on Amberlyst 15 (2.4 s−1 ). On the other hand, a too-strong acidity was detrimental to selectivity in the case of branched alcohols such as neopentanol. The viscosity indexes of the final products were improved, for example, from 4.5 mm2 s−1 for methyl oleate to 30.3 mm2 s−1 for hydroxy-neopentoxy methyl stearate obtained by addition of neopentanol on the epoxidized methyl oleate; moreover, their biodegradability was shown to be maintained. Catalytic Conversion of Triglycerides to Polymers Although triglyceride derivatives have long been used in polymer manufacture, heterogeneous catalysis has played only a modest role so far in the production of monomers. Interestingly, the dimerization of fatty acids on montmorillonite clays at 250 ◦ C results in diacids that can be hydrogenated to diols. These diacids and diols can then be used to produce polyesters, polyethers, polyamides, and polyurethanes. Polyethers from dimer diols are valuable for the production of highly hydrophobic polymers stable in acidic media [34]. The epoxidation of triglycerides followed by condensation with alcohols (see Section 11.9.2.3.2) results in the formation of OH groups along the fatty acid chains that can be used as starting materials for the manufacture of thermoplastic polyurethanes. 11.9.2.3.3

11.9.2.3.4 Hydrogenation of Fatty Acids Esters Fatty alcohols are obtained by the hydrogenation of fatty acids or, more easily, of fatty acid esters in the presence of metal catalysts. The reduction is particularly difficult if double bonds need to be preserved. The selective hydrogenation of methyl oleate to the corresponding unsaturated alcohol (Fig. 4) was achieved with Ru-SnB/Al2 O3 catalysts prepared by reduction with NaBH4 of Ru and Sn salts coimpregnated on alumina [35].

Catalytic Conversion of Glycerol Glycerol can be considered as a renewable feedstock, in so far as it is the coproduct of triglyceride saponification, 11.9.2.4

OR

Conversion of fatty acids to lubricants by epoxidation followed by alcoholysis.

2451

Fig. 3

References see page 2471

2452

11.9 Conversion of Biomass on Solid Catalysts

conditions (pH 10–11) an 83% yield to tartronate was obtained at 85% conversion. Abbadi and van Bekkum [39] obtained a 93% selectivity to hydroxypyruvic acid at 95% conversion of glyceric acid on a 5%Bi −5%Pt/C catalyst, without pH regulation. Fordham et al. [40] have studied the preparation of mesoxalic acid by oxidation of sodium tartronate on a PtBi/C catalyst at 60 ◦ C without pH control; the maximum yield was 65% at 80% conversion. A total conversion of tartronic acid was obtained at 80 ◦ C, giving 50% yield of mesoxalic acid with no other products, because all byproducts were totally oxidized into CO2 . More recently, the oxidation of glycerol was conducted in the presence of gold catalysts in basic medium [41, 42]. Glycerol was oxidized with oxygen at 60 ◦ C with a selectivity of 100% at 56% conversion on 1 wt.% Au/activated carbon. Graphite-supported catalysts gave comparable results. The selectivity of glycerol oxidation on gold catalysts on various supports was shown to depend critically upon the size of the gold particles [43].

O OCH3

H2

OH Fig. 4

Hydrogenation of fatty esters to fatty alcohols.

and of the transesterification of vegetable oils to FAME employed as biofuels. As far as production of chemicals is concerned, the main outlet of glycerol is the production of surfactants by esterification with fatty acids (see Section 11.9.2.3.1). Transformation by oxidation and dehydroxylation described in the following sections are also interesting routes to valuable products. Conversion of Glycerol to Oxidized Compounds The selective oxidation of glycerol leads to various valuable oxygenates, as depicted in Fig. 5. Most of these compounds are currently prepared by enzymatic reactions, and have a small market because of their high price. Besson and Gallezot [36] have shown that such products can be obtained by oxidation with air of aqueous solutions of glycerol in the presence of carbon-supported platinum and palladium catalysts. The selectivity can be tuned by promotion of the noble metals with bismuth, or by operating under controlled pH. Garcia et al. [37] found that the oxidation of glycerol at basic pH on palladium and platinum catalysts yielded 70% glycerate. Using bismuth-promoted platinum under similar reaction conditions, the oxidation leads to dihydroxyacetone with a 25% yield. Fordham et al. [38] found that glyceric acid oxidation on a 5%Pt −1.9%Bi/C catalyst yielded 74% hydroxypyruvic acid at 80% conversion at acidic pH (3–4), but on the same catalyst under basic 11.9.2.4.1

11.9.2.4.2 Conversion of Glycerol to Polyglycerols and Derivatives Polyglycerols obtained by the dehydration of glycerol and polyglycerol esters (PGEs) are employed as surfactants, lubricants, cosmetic, and food additives. There are three challenges in the etherification of glycerol: (i) to replace liquid bases by solid catalysts to achieve a greener and more efficient process; (ii) to control the chain length by obtaining a suitable hydrophilic–lipophilic balance in PGEs; and (iii) to avoid the formation of acrolein by internal dehydration. The etherification of glycerol by dehydration in the presence of liquid bases leads to a mixture of dimers, trimers (Fig. 6) and cyclic or branched oligomers. Zeolites have been used to take advantage of shape selectivity effects to minimize oligomer formation, as described in two patents [44, 45] (Table 4). The selectivity

OH HO

OH O

O

b

OH a

OH HO

OH Glycerol

OH

Glyceric acid

O O

c

OH

Tartronic acid

HO

OH

O HO

O OH

Dihydroxyacetone Fig. 5

Oxidation of glycerol to various oxygenates.

OH

Mesoxalic acid

d e

O OH

f

O

O

Hydroxypyruvic acid

11.9.2 Conversion of Biomass to Biobased Products

OH 2 HO

OH OH

HO

2453

OH OH + H2O

O Diglycerol

Glycerol

+ Glycerol

OH HO

OH O

OH O

OH + H2O

Triglycerol

[Polyglycerols] Fig. 6

Dehydration of glycerol.

to monoglycerol was improved over the Cs−ZSM-5 catalyst, but the activity was too low [46]. A fair compromise between activity and selectivity was obtained by Clacens et al. [47] using cesium-impregnated mesoporous MCM41 (Table 4). 11.9.2.4.3 Hydrogenolysis of Glycerol to 1,2- and 1,3Propanediol Glycerol can be selectively dehydroxylated either to 1,2-propanediol (1,2-PDO), a chemical that can advantageously replace ethylene glycol as anti-freezing agent, or to 1,3-propanediol (1,3-PDO), which when copolymerized with terephthalic acid, gives polyesters with unique mechanical properties. 1,3-PDO is currently produced by catalytic routes from ethylene oxide (the Shell route) or acrolein (the Degussa–DuPont route). The microbial production of 1,3-PDO is under development by DuPont–Genencor to produce 1,3-PDO from glucose [48]. In a recent investigation, Chaminand et al. [49] studied the hydrogenolysis of aqueous solutions of glycerol at 180 ◦ C under 8 MPa H2 -pressure in the presence of supported metal catalysts in an attempt to produce, selectively, 1,2- and 1,3-PDO. HO

OH OH

H2 Cat.

HO

OH + HO OH

Different catalysts (Cu, Pd, Rh), supports (ZnO, C, Al2 O3 ), solvents (H2 O, sulfolane, dioxane), and additives (H2 WO4 ) were tested to improve reaction rate and selectivity. The best selectivity (100%) to 1,2-PDO was obtained by hydrogenolysis of an aqueous solution of glycerol in the presence of CuO−ZnO catalysts. To control the selectivity toward 1,3-PDO, the reaction was conducted with rhodium catalysts with tungstic acid added to

Tab. 4

Glycerol dehydration. Selectivity to di- and triglycerol

Catalyst

Diglycerol/ % Triglycerol/ % Reference

A-zeolite β-zeolite Cs-ZSM-5 Cesium-impregnated MCM-41

32 36 94 60

25 18 6 17

[44] [45] [46] [47]

the reaction medium. The best selectivity to 1,3-PDO (1,3-PDO:1,2-PDO = 2) was obtained by operating in sulfolane. The presence of iron dissolved in the reaction medium was also beneficial for the selectivity to 1,3-PDO. A mechanism was proposed to account for the effect of the different reaction parameters. Catalytic Conversion of Carbohydrates Carbohydrates represent the largest amount of renewables employed for the production of biobased products, with two major sources: sucrose (130 × 106 t y−1 ) and starch (40 × 106 t y−1 ), with 50% of the latter being used for industrial purposes. Other polysaccharides, such as inulin, are gaining importance as a source of fructose that can be used as a starting material for the production of 5-hydroxymethylfurfural, from which several diols, dialdehydes and diacids useful for polymer synthesis can be obtained (see Section 11.9.2.5.1). A review of carbohydrates as chemical feedstock has been prepared by van Bekkum and Besemer [50]. 11.9.2.5

11.9.2.5.1

hydrates

Hydrogenation and Hydrogenolysis of CarboGlucose produced from starch or sucrose

References see page 2471

2454

11.9 Conversion of Biomass on Solid Catalysts

H H

OH H

OH

HOH2C OH OH H OH O Gluconic acid OH

HO

OH H2

O HO HO OH

OH

Glucose Fig. 7

CH2OH H

CH2OH

OH

HO

H

H

HO

H

H

OH

H

OH

H

OH

H

OH

CH2OH Sorbital

CH2OH Mannitol

Hydrogenation of glucose to sorbitol and byproducts.

hydrolysis is hydrogenated to sorbitol (ca. 800 000 t y−1 ), a commodity product used in foods and in the pharmaceutical and chemical industries, as well as an additive in many end-products. The main byproducts of the reaction are gluconic acid, formed by the Cannizaro reaction, and mannitol, formed by sorbitol epimerization (Fig. 7). Catalysts allowing a 100% conversion and 99% selectivity are required. These also should be stable after many recycling operations, or for an extended period of time on stream in a continuous reactor. Most of the industrial production is still conducted batchwise on Raney nickel catalysts, promoted with more electropositive metal atoms such as molybdenum and chromium [51]. However, because of the risk of nickel or metallic promoter leaching, these tend to be replaced by more active ruthenium catalysts. Also, because of the large sorbitol production, continuous processes would be preferable. It has been shown by Nicolaus et al. [52] that 1.8% Ru/C catalyst prepared by cationic exchange or anionic adsorption on 0.8-mm Norit extrudates provided up to 99.5% yield of sorbitol for up to 596 h in a tricklebed reactor. No leaching of ruthenium was detected. The activity at low conversion (initial activity) was 1080 mmol h−1 gRu −1 , compared to 50 mmol h−1 gNi −1 on 48% Ni–kieselguhr catalyst. Glucose hydrogenation was also conducted on Ru−Pt/C bimetallic catalysts prepared by the coexchange of Pt and Ru ammino cations. Interestingly, the activity passed through a maximum at 1470 mmol h−1 gRu −1 for the specific atomic composition Ru56 Pt44 . Pt−Ru catalysts were also more selective to sorbitol because the rate of epimerization to mannitol decreased. The contact time with the catalyst can be increased without loss of selectivity, thus allowing operation at total conversion of glucose and at more than 99% selectivity over a large domain of liquid flow rate. The bimetallic catalysts loaded in trickle-bed reactor yielded 2 t day−1 kgRu −1 of sorbitol at 99.5% purity.

Platinum alloying extends the catalyst stability because it may prevent the formation of oligomeric or cracking products, and also prevents any oxidation of ruthenium. The hydrogenation of glucose to sorbitol was achieved on ruthenium catalysts supported on activated carbon cloths (ACC) obtained by carbonization and CO2 activation of woven rayon [53]. A catalyst of 0.9 wt.% Ru−ACC was loaded with ruthenium by cationic exchange or anionic adsorption. Figure 8a is a scanning electron microscopy (SEM) image of the texture of the ACC; Fig. 8b is a transmission electron microscopy (TEM) image, showing the homogeneous distribution of 2-nm ruthenium particles in carbon fibers. The ACC was clamped on a support which fitted along the autoclave walls, thus allowing easy recycling of the catalyst as, unlike powder catalysts, no filtration is required and there is no attrition or leaching. Glucose hydrogenation was achieved with a 99.5% selectivity to sorbitol at 99.7% conversion with an activity of 2.4 mol h−1 gRu −1 comparable to Rucatalyst in powder form. The hydrogenation of glucosone to fructose was achieved on 2.5% Pd−ACC at a rate of 0.5 mol h−1 gPd −1 .

25 nm

500 µm

(a)

(b)

(a) Scanning electron microscopy image of activated carbon cloth; (b) transmission electron microscopy image of Ru dispersion in carbon fibers.

Fig. 8

11.9.2 Conversion of Biomass to Biobased Products

HO OH

CH2OH

OH H+

O

O HO

O

OH

n

Starch

HO HO

O

H2 O 2

OH OH

HO

H

H

OH

H

OH

Glucose

H2

HO

H

H

OH

H

OH

CH2OH

CH2OH

Arabinonic acid Fig. 9

2455

Arabitol

Conversion of glucose to arabitol.

Currently, there is major interest in converting C6 carbohydrates available in plentiful supply from starch or sucrose, into C5 and C4 polyols that are minimally present in biomass, and then to identify many applications in food and non-food products. Thus, glucose can be converted to arabitol by an oxidative decarboxylation of glucose to arabinonic acid, followed by hydrogenation to arabitol (Fig. 9). The main pitfall is to avoid dehydroxylation reactions leading to deoxy-products that are not compatible with the purity specifications required for arabitol. Aqueous solutions (20 wt.%) of arabinonic acid were hydrogenated on Ru-catalysts in a batch reactor [54]. The selectivity was enhanced by adding small amounts of anthraquinone-2-sulfonate (A2S), which decreased the formation of deoxy byproducts. Thus, by adding 260 ppm A2S with respect to arabinonic acid, the selectivity to deoxy-products decreased from 4.2 to 1.6%. A2S acted as permanent surface modifier, since the catalyst was recycled with the same selectivity without further addition of A2S. The highest selectivity to arabitol was 98.9% at 98% conversion, with a reaction rate of 73 mmol h−1 gRu −1 at 80 ◦ C. Deoxyhexitols consisting of C6 diols, triols, and tetrols are well suited to replace polyols derived from petrochemistry for applications in polyester and polyurethane manufacture. Comparatively few investigations on catalytic hydrogenolysis of carbohydrates were performed [55, 56], and most of these were designed to produce C2 −C3 polyols, particularly glycerol, rather than higher molecularweight polyols. Sorbitol was taken as a model molecule to study the hydrogenolysis to C4 −C6 products [57]. To improve the selectivity to deoxyhexitols, the catalysts and reaction temperature were optimized to favor the rupture of C−OH bonds (dehydroxylation reactions) rather than C−C bond rupture. Copper-based catalysts, which have a low activity for the hydrogenolysis of C−C bonds, were employed to treat 20 wt.% aqueous sorbitol solutions in the temperature range 180 to 240 ◦ C. Reactions carried out in the presence of 33% CuO:65% ZnO catalyst at 180 ◦ C under H2 -pressure yielded 73% C4 + polyols and, more specifically, 63% deoxyhexitols. In contrast, operating in the presence of palladium catalysts at 250 ◦ C

under 8 MPa of hydrogen pressure, cyclodehydration reactions of sorbitol and mannitol occurred with formation of cyclic ethers: isosorbide, 2,5-anhydromannitol, 2,5anhydroiditol, and 1,4-anhydrosorbitol [58]. Up to 50% and 90% yield to isosorbide were obtained from sorbitol and mannitol, respectively (Fig. 10). These mixtures of polyols were effectively employed to synthesize alkyd resins and to produce decorative paints with performance comparable to their commercial counterparts. 11.9.2.5.2 Catalytic Oxidation of Carbohydrates Oxidation reactions are widely used for upgrading carbohydrates to a great variety of higher value-added chemicals used in detergents or pharmaceuticals (e.g., vitamin C). Enzymatic and microbial oxidation reactions are widely used for that purpose, but homogeneous catalysis is also common. Thus, oxidation with hypochlorite mediated by TEMPO has been generalized to a large number of carbohydrates [59]. The catalytic system was improved by adding cocatalysts such as laccase enzyme to allow oxidation with oxygen, and by immobilization on a support [60]. Oxidation reactions with H2 O2 , mediated by metal phthalocyanine catalysts, have also proved very efficient, including the oxidation of insoluble substrates such as native starch [61]. This is a very uncommon example of heterogeneous process as a solid substrate was oxidized in the presence of a soluble catalyst.

OH CH2OH OH

OH

CuO-ZnO 180 °C, H2 OH

OH

HO OH OH CH2OH Sorbitol or C6 polyols

Fig. 10

63% yield deoxyhexitols 15% triols, tetrols

HO Pd/Al2O3 250 °C, H2

O O

90% yield isosorbide

OH

Dehydroxylation and cyclodehydration of sorbitol.

References see page 2471

2456

11.9 Conversion of Biomass on Solid Catalysts

Tab. 5

Product distribution in glucose oxidation on bismuth-promoted palladium [65]

Catalysta (run)

PdBi/C (1st ) PdBi/C (2nd ) PdBi/C (3rd ) PdBi/C (4th ) PdBi/C (5th ) Pd/C

Conversion/%b

99.6 99.7 99.8 99.9 99.9 82.6

Yield/mol%

Selectivity/%

1

2

3

4

99.4 98.9 98.5 98.5 99.1 78.1

130 Shell’s XHVI, ExxonMobil’s VISOM

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13.7.2 Mechanism of Skeletal Isomerization of C5+ Alkanes with Acid Catalysts

2809

13.7

13.7.2

Isomerization

Mechanism of Skeletal Isomerization of C5+ Alkanes with Acid Catalysts

Swan Tiong Sie∗

Reaction Cycle in Isomerization of Alkanes Acid-catalyzed skeletal isomerization of alkanes occurs via carbenium (formerly called carbonium) ions as intermediates. It is a part of a chain reaction, that is, a reaction cycle involving chain initiation, carbenium ion rearrangement (the isomerization reaction proper) and chain propagation. Scheme 1 shows the conventional mechanism of acid-catalyzed isomerization of n-pentane. Whereas the notion of a ‘‘free’’ carbenium ion as reaction intermediate is a very useful one to describe the phenomena in acid-catalyzed reactions, carbenium ions may not exist as such in these reactions, but the charged intermediate species may be complexed with the acid catalyst. For example, in the case of an aluminosilicate (zeolite) catalyst having acidic hydroxyl groups, the charged species may be an alkoxy intermediate, rather than a free carbenium ion. For the sake of simplicity, however, we will discuss the mechanism of acid-catalyzed isomerization in terms of ‘‘carbenium ions’’, since in this discussion it is immaterial whether the ion is free or bound to the catalyst. In the chain initiation reactions, carbenium ions are formed from neutral hydrocarbon molecules in contact with the acid catalyst, from alkenes by addition of a proton supplied by the acid catalyst and from alkanes by either a combination of dehydrogenation and proton addition or by abstraction of a hydride ion. The latter ion can be accepted by the acid catalyst, by combining with a proton to form molecular hydrogen. In the chain propagation reaction, a carbenium ion reacts with a neutral feed hydrocarbon molecule whereby transfer of a hydride ion from this molecule to the carbenium ion results in a neutral hydrocarbon product molecule and a new carbenium ion originating from the feed molecule. This new ion can in turn undergo isomerization, thus perpetuating the reaction cycle. 13.7.2.1

13.7.1

Introduction

Probably the most important isomerization reaction in the context of energy-related catalysis is the skeletal isomerization of alkanes and alkenes. Other hydrocarbon isomerization reactions that play a role are the double bond isomerization of alkenes, interconversion of alkylcyclopentanes and cyclohexanes, isomerization of bi- and polycyclic saturated hydrocarbons and isomerization of alkyl aromatics. These other isomerization reactions can occur during acid- or bifunctionally catalyzed conversions, for instance in catalytic reforming, catalytic cracking, hydrocracking, polymerization and alkylation, but in general they are not carried out as dedicated conversion processes in the energy field. Skeletal isomerization of acyclic hydrocarbons, by contrast, is the basis for several important processes in the hydrocarbon processing industry. These include the conversion of light linear alkanes such as n-pentane and n-hexane into their branched isomers to improve the octane quality of the gasoline in which they are present. Conversion of n-butane into isobutane is of importance to increase the availability of the latter hydrocarbon for alkylation or for producing methyl tert-butyl ether (MTBE) via dehydrogenation to isobutene. Skeletal isomerization of higher alkanes is the basis for the conversion of paraffin wax into lubricating base oils with excellent viscometric properties and for improving the cold-flow behavior of paraffinic oils. Skeletal isomerization of normal butenes is another way of obtaining isobutene for MTBE production. In view of the above, this discussion of isomerization focuses primarily on the skeletal isomerization of alkanes. More specifically, we will discuss acid-catalyzed isomerization of alkanes and ignore metal-catalyzed isomerizations, since the acid-catalyzed route is by far the most important one in practice. Extensive reviews of isomerization reactions of pure hydrocarbons in a broader sense have appeared in the literature, for example the monograph by Egloff et al. [1] and the review by Condon [2], covering the literature up to 1958 and dealing mainly with mechanisms, kinetics and thermodynamic aspects. In the present review, more recent concepts and data are discussed and attention is given to the technological aspects of alkane isomerization. ∗

Corresponding author.

The Isomerization Reaction: Rearrangement of the Intermediate Carbenium Ion A simple way to visualize skeletal isomerization is to assume an alkyl shift, such as a methyl shift. Thus, a methyl ion is detached from the carbenium ion chain and reattached at another position in the residual hydrocarbon chain. This simple mechanism is highly unlikely, however, since the methyl ion is a high-energy species so that its detachment from the carbenium ion chain would involve a prohibitively high activation energy (see Table 1). 13.7.2.2

References see page 2828

2810

13.7 Isomerization

Chain initiation: CH3 CH2

CH2

CH2

CH3 + H+ CH3

+

CH

CH2

CH3 + H2

CH2

Carbenium ion rearrangement: CH3

+

CH

CH2

CH2

CH3

CH3

CH

CH2

+

CH2

CH3

CH3

+

C

CH2

CH3

CH3

Chain propagation: CH3

+

C

CH2

CH3 CH3

CH3 + CH3 CH

CH2

CH2

CH2

CH2 CH3 +

CH3 + CH3

CH CH2 CH2 CH3

CH3

Scheme 1

Tab. 1

Reaction cycle in isomerization of n-pentane.

Heats of formation of alkylcarbenium ions [3]

Carbenium ion

Type of ion

Heat of formation/kJ mol−1

Methyl Ethyl n-Propyl n-Butyl Isopropyl sec-Butyl tert-Butyl

Primary Primary Primary Primary Secondary Secondary Tertiary

1080 942 913 883 812 795 733

Another argument against the classical alkyl shift mechanism is that this mechanism would favor the formation of isomers with larger side-groups than methyl groups. Detachment of primary ethyl, n-propyl and n-butyl cations from a long-chain secondary carbenium ion would involve less energy than of a methyl ion, as can be seen in Table 1. Thus, such an alkyl shift mechanism is incapable of rationalizing the very strong preponderance of methyl-branched isomers, as observed in the skeletal isomerization of n-alkanes. The most likely mechanism of skeletal isomerization of the intermediate carbenium ion involves the rearrangement of the classical secondary carbenium ion into a non-classical carbonium ion, namely a protonated dialkylcyclopropane. This mechanism, originally proposed by Condon [4] and Brouwer [5], is depicted in Scheme 2. The transformation of the classical secondary carbenium ion into the protonated dialkylcyclopropane structure will not involve a high energy barrier since

the latter structure can be considered to be a hybrid of resonance structures, as shown in Fig. 1. This resonance will contribute to the stability of the protonated cyclopropane species and will compensate for the inherent strain of the three-membered ring. That resonance is capable of greatly enhancing the stability of this ring is demonstrated by the high stability of the cyclopropenyl cation, which can even display some aromatic character. For example, the triphenylcyclopropenylium cation is so stable that its salts can be isolated [6]. Protonated cyclopropane structures as reaction intermediates have been deduced from the stereochemistry of Wagner–Meerwein-type rearrangements of norbornyl derivatives and evidence for this mechanism has also been obtained by experiments with carbon isotopes [7]. Since these rearrangements occur under mild conditions (below 50 ◦ C with only moderately strong acids), it is unlikely that high energy barriers are involved. Mass spectrometric experiments have provided supporting evidence for the existence of protonated cyclopropane in the gas phase and the ionization potential for formation of this ion proved to be not very much higher than that for a classical secondary ion [8].

H C

R

H +

C

R’

R

H

H

C

+C

C H

H

C H

H

R’ H

H

R

H

H

C+

C

R’

C

H H

H

Some resonance structures of protonated dialkylcyclopropane. In addition to the corner-protonated structures shown, edge- and face-protonated structures are in principle also possible.

Fig. 1

13.7.2 Mechanism of Skeletal Isomerization of C5+ Alkanes with Acid Catalysts

H H H H H H C C C C C H HnH H H Hm

Linear alkane

Hydride abstraction/transfer n ≥ 1 H H H H H H C C C C C H Hn + H H Hm

Classical carbenium ion

H H H H H C C C C H Hn Hm C H+ H H

Non-classical ion

H H H C C Hn C H H

H C +

m≥1

H C H Hm

Classical carbenium ion

H

Hydride transfer H H H C C Hn C H H

H C H

H C H Hm

Isomerized product

2811

corresponding secondary ions, as can be seen from Table 1. Since a difference in activation energy of this magnitude at a temperature of 100 ◦ C corresponds to a difference in reaction rates by a factor of 1014 , the breakage of the one cyclopropane bond that leads to isomerization is so strongly disfavored compared with breakage of the two other bonds that its occurrence can be neglected. The breakage of these other two bonds leads to either the original n-butane molecule or another n-butane molecule in which two carbon atoms have exchanged their position in the chain of carbon atoms (Fig. 2). Using 13 C-labeled n-butane, Brouwer [5] was able to show that this exchange indeed occurred and that the rate of isotope scrambling was comparable to the rate of skeletal isomerization of n-pentane under similar conditions (Fig. 3). Similar evidence for the PCP isomerization mechanism was collected by Chevalier et al. in later experiments with 13 C-labeled n-butane over a Pt/silica–alumina catalyst at 300 ◦ C [11]. Under these conditions, which are more representative for isomerization of alkanes in practical processes, they observed very little skeletal isomerization of n-butane, but scrambling of 13 C occurred. Moreover, in experiments with different catalysts the rate of scrambling proved to correlate very well with the transformation of n-pentane into isopentane, as can be seen in Fig. 4.

H

Scheme 2 Isomerization intermediate.

via

a

protonated

cyclopropane

On the basis of the above arguments, isomerization via a dialkylcyclopropane intermediate seems an acceptable mechanism from an energy point of view. Further evidence for this mechanism is presented below. Supporting Evidence for the Protonated Cyclopropane (PCP) Isomerization Mechanism

13.7.2.3

Reactions with n-Butane Strong evidence for the PCP isomerization mechanism is the finding of Brouwer and Oelderik [9, 10] that n-butane is not isomerized by the superacid HSbF6 at room temperature, whereas n-pentane and n-hexane are rapidly converted to their isomers. This can be understood from Scheme 2. Since for n-butane the value of m is zero, the rupture of the only bond in the cyclopropane ring that would lead to the isobutane structure involves the formation of an energetically unfavorable primary carbenium ion. Breakage of the two other bonds is possible, but will in both cases result in a carbenium ion with a straight chain. The heats of formation of primary carbenium ions are about 100 kJ mol−1 higher than those of the

13.7.2.3.2 Effect of Chain Length on the Relative Rates of Isomerization of n-Alkanes As discussed above, nalkanes larger than n-butane can be isomerized by the PCP mechanism. Since the number of possible cyclopropane structures increases for longer chains, the reactivity for isomerization also increases with chain length. With the simplifying assumption that all cyclopropane structures involved in isomerization form and react with equal chances, it can be deduced that the isomerization reactivity of the n-alkanes with carbon number N should be proportional to N –4. The experimental data of Maslyanskii et al., as reported by Weisser [12], are in line with this prediction, as can be seen from Table 2.

13.7.2.3.1

Location of Branching in Isomerized Alkanes A detailed study of the isomer distribution of the products obtained in the isomerization of n-alkanes at low conversion levels has also provided evidence for the PCP mechanism. According to this mechanism, the formation of the 2-methyl isomer will be less likely than that of the 3-methyl isomer for alkanes with relatively long chains. This has been observed in actual isomerization experiments, for example, by Steijns et al. [13] and by Weitkamp [14]. The abnormally low abundance of the 2-methyl isomer in the product of primary isomerization 13.7.2.3.3

References see page 2828

2812

13.7 Isomerization

H

H C

H H H

H

C

+ C

1

H

C 3

H

C

2

H

H

H

H

4

H

C H

H

H b

C

H C H H C C+ H

H

a

C H+

H

a

H

H

H

C H

b

H H

H H

C

H

+ C

1

2

C H

4

H

H Fig. 2

H

C3

Possible and forbidden rearrangements of the sec-butyl cation.

Observed/equilibrium conversion of butane

0.7

0.6 0.5 0.4 Rate C4 scrambling = 0.25 Rate C5 isomerization

0.3 0.2 0.1 0

0 0.90.8 0.7 0.6

0.5

0.4

0.3

0.2

0.1

Observed/equilibrium conversion of pentane

Conversion of n-butane[1-13 C] to n-butane[2-13 C] compared with isomerization of n-pentane to isopentane. Experiments with a mixture of n-pentane and labeled n-butane over HSbF6 catalyst at 0 ◦ C [5].

Fig. 3

Scrambling of n-Butane [1-13C]/%

30

20

T = 300 °C 0.2% Pt 10

0.5% Pt

Catalyst:Pt/SiO2Al2O3

1.2% Pt

0

0

10

20

30

40

Coversion of n-Pentane/% Fig. 4

Correlation between scrambling of n-butane[1-13 C] and skeletal isomerization of n-pentane [11].

13.7.2 Mechanism of Skeletal Isomerization of C5+ Alkanes with Acid Catalysts Relative rates of isomerization of n-alkanes over a tungsten disulfide catalyst [12]

2813

Alkane

Relative rate

n-Pentane n-Hexane n-Heptane n-Octane

Experimental

PCP theory

1 2.0 3.1 4.2

1 2 3 4

2- Me isomer in methyl isomer fraction/ %

of n-alkanes, especially those of higher carbon number, can be seen in Fig. 5. Figure 6 compares the experimentally observed distributions of monomethyl-branched isomers from isomerization of n-alkanes with distributions predicted from the PCP reaction scheme according to a simplified model as proposed by Weitkamp [14]. In this model, the chances of formation of protonated cyclopropane structures are considered equal for all possible structures, while the likelihood of breaking the bonds in the cyclopropane ring is supposed to be the same for all bonds. It can be seen that there is reasonably good agreement between theory and experiment. However, some systematic deviations can be observed in Figs. 5 and 6: for the hexanes, heptanes and octanes there is an excess of 2-methyl isomers, whereas there is a slight underproduction of these isomers for dodecane and higher alkanes. These systematic deviations have been ascribed to some difference in the rates of the ring-opening modes forming the 3-methyl-2-n-alkyl and 2-methyl-3-n-alkyl cation, related to the relative stability of these cations, as suggested by Martens and Jacobs [15]. The latter authors also suggested that protonated cycloalkanes with larger rings are involved in order to explain the formation of isomers with larger sidegroups than the methyl group. However, compared

80

Estimated for equilibrium and calculated for classical mechanism Experimental (X iso = 1 to 3%)

60

Calculated for PCP mechanism

40 20 0

Abundance / % of methyl isomers

Tab. 2

2– Me

60

Predicted 40

Experiment

20

0

6

8

10

8

10

12

14

Carbon number of feed alkane Formation of 2-methyl isomers in the isomerization of n-alkanes of different chain length [14].

12

14

6

8

10

12

14

Carbon number

(a) 40

4 – Me

5 – Me

6 – Me 7– Me

20

0

8

10

(b)

12

14 10 12 14 12 Carbon number

14 14 15

Distributions of monomethyl branched isomers from isomerization of n-alkanes of different chain length [15].

Fig. 6

with the methyl-branched isomers, the isomers with ethyl, propyl and butyl branches are generally formed in minor amounts only; for example, in the case of n-heptane isomerization only 3% ethylpentane was found among the isomers, the remaining 97% being methylbranched alkanes [14]. Isomerization via a protonated cyclobutane intermediate leading to the ethyl-branched isomer is apparently much less favored than by the PCP route, which may be explained by the lower stability of protonated cyclobutanes. Protonated cyclobutane has been found to be 130 kJ mol−1 less stable than protonated methylcyclopropane [16]. Chain Termination Reactions: Catalyst Deactivation If there were no chain termination reactions, the reaction cycle for isomerization, once initiated, would continue indefinitely, given ample supply of reactants. In practice, 13.7.2.4

6

3 – Me

Fig. 5

References see page 2828

2814

13.7 Isomerization

however, chain termination reactions are responsible for acid consumption or catalyst deactivation. Chain termination occurs when a carbenium ion undergoes a hydride transfer reaction with an alkene molecule, instead of with an alkane molecule as in chain propagation. Hydride transfer between a carbenium ion and an alkene gives rise to an alkane and an unsaturated carbenium ion. The latter ion will be an allylic carbenium ion, which can be considered as a protonated diene that is a much more stable species than the original carbenium ion. In a similar way, a carbenium ion may be involved in a hydride transfer with a conjugated diene to form an alkane and the protonated form of a conjugated triene, and so on. The triene molecule may cyclize to form an aromatic molecule. The highly unsaturated molecules are much more basic and bind the protons much more strongly than simple monoalkenes. They are incapable of participating in the reaction cycle and give rise to a loss of effective acidity or of catalyst activity. Another way in which polyalkenic species can be formed is by cracking, which occurs particularly with carbenium ions having longer chains originating from heavy feed molecules or from oligomerized light alkenes. Cracking gives an alkanoic and an alkenic fragment. If the latter fragment already contained a double bond that was present in the original feed molecule, a diene is formed. Thus, conditions which are inducive to oligomerization/cracking (disproportionation) of alkenes are likely to give rise to catalyst deactivation. Acid-catalyzed Cracking and Isomerization: Effect of Chain Length on Selectivity for Isomerization Since cracking and isomerization are both catalyzed by similar catalysts, it is plausible that cracking may accompany isomerization, thus decreasing isomerization selectivity. This is particularly true for higher alkanes. Acid-catalyzed cracking of alkanes also proceeds with carbenium or carbonium ions as intermediates. The classical, generally accepted mechanism assumes that in a classical secondary carbenium ion the C–C bond in the β-position to the charge center is broken (β-scission). However, from energetic considerations, β-scission is rather unlikely in the case of a secondary ion with a straight chain. Many other facts related to experimentally observed characteristics of acid-catalyzed cracking also strongly argue against this β-scission cracking mechanism [17]. Another mechanism that has been proposed assumes a non-classical carbonium ion, namely the same protonated dialkylcyclopropane species as in isomerization, to be the reaction intermediate in acid-catalyzed cracking [17]. For energetic reasons, namely the avoidance of C–H dissociation in primary carbon atoms, it follows that

cracking may occur if the hydrocarbon has seven or more carbon atoms. The relationship between cracking and isomerization over bifunctional catalysts in the presence of hydrogen (see below) is depicted in Scheme 3. This scheme suggests that hydroisomerization and hydrocracking may occur as parallel reactions, in addition to sequential reactions (cracking of pre-isomerized feed molecules and postisomerization of cracked fragments). Since skeletal isomerization requires a carbon chain of more than four atoms whereas cracking needs at least seven carbon atoms in the chain, pentanes and hexanes are the only members of the homologous series of alkanes that can be easily isomerized, but not easily cracked. Indeed, in practice highly selective isomerization of pentanes and hexanes and their mixtures is readily achievable. For higher alkanes, cracking usually accompanies isomerization, resulting in a lower isomerization selectivity. The clear difference in the behavior of n-pentane and n-hexane versus n-heptane is illustrated by the data in Table 3, obtained over the same catalyst in

H H H H H H C C C C C H HnH H H Hm

n-Alkane n≥1

H H H H H H C C C C C H Hn + H H Hm

Classical carbenium ion

H H H C C Hn

Non-classical ion

13.7.2.5

H H C C H Hm C H+

H

H

m≥1

H H H C C Hn C H H

m≥3

H H C C H H Hm H

Isomerized product

H H H C C Hn C H H

H H H H C H + H C C C H H H H H m-2 H Cracked products

Scheme 3 Isomerization and cracking via a protonated cyclopropane intermediate.

13.7.3 Mechanism of C4 Alkane Isomerization

2815

Isomerization of n-pentane, n-hexane and n-heptane over Pt/H-mordenite [18] (P = 2.4–3.0 MPa, weight hourly space velocity = 1 g g−1 h−1 , H2 : HC = 2.5 : 1)

Tab. 3

Parameter Temperature/ ◦ C Cracked products (100

Isomer in total hexanes / %

Hydrocarbon

60 50 40 2– MP 30

2,2– DMB 3 – MP

20

n – C6 2,3–DMB

10 0

50

100

150

200

250

300

350

400

Temperature / °C Fig. 14 Thermodynamic equilibria of isomeric hexanes in the gas phase. Calculated from thermodynamic data collected in API Project 44 [36].

80 C5 70

90 C4

50 40 30

C5 gas

C5 liquid

60

RON- 0

Iso / iso + normal / %

90

85 C6 liquid 80

0

100

200

300

C6 gas

400

Temperature / °C Thermodynamic equilibria between n-butane and isobutane and between n-pentane and isopentane in the gas phase. Calculated from thermodynamic data collected in API Project 44 [36].

75

Fig. 13

70

0

50

100

150

200

250

300

350

400

Temperature / °C

Processes for Isomerization of Pentanes and Hexanes Industrial processes for the isomerization of C5 and C6 alkanes in light straight-run naphtha can be classified according to the type of catalyst used: 13.7.5.2

(i) processes using HCl/AlCl3 as a monofunctional catalyst (ii) processes using a noble metal on chlorided alumina as a bifunctional catalyst (iii) processes using a noble metal on amorphous silica–alumina as a bifunctional catalyst (iv) processes using a noble metal on an acid form of a zeolite as a bifunctional catalyst.

Fig. 15

Octane numbers of equilibrium mixtures of pentanes and

hexanes.

The most important processes in these categories are briefly discussed below. 13.7.5.2.1 Processes Using Aluminum Chloride/ A proHydrochloric Acid as Monofunctional Catalyst cess belonging to this category of older processes is the light naphtha isomerization process of Standard Oil Co. of Indiana [37]. The reaction is carried out in fixed beds of References see page 2828

13.7 Isomerization

a solid, supported catalyst. The feedstock used is purified to remove contaminants, and HCl is added to it as catalyst promoter. Because of catalyst deactivation, several reactors are used, which allows periodic regeneration of a reactor without interrupting the overall operation. HCl is removed from the reactor effluent by distillation or stripping and is recycled. Use of a liquid instead of a solid catalyst allows easier maintenance of catalyst activity since the deactivated catalyst can be removed and replenished during operation. This concept was applied in the Isomate process of the Standard Oil Co. of Indiana [38]. Purified naphtha, with dissolved HCl as catalyst promoter, is fed to the bottom of a reactor containing a pool of molten catalyst. Catalyst activity is maintained by periodically injecting a slurry of

Recycle HCI

Settler

C1– C4

Water

Hydrogen

HCI stripper

AlCl3 slurry

Fig. 16

Makeup HCI

Reactor

HCI absorber

Drier

Feed

Deisopentanizer

Gas to fuel

fresh aluminum chloride in naphtha. Effluent from the top of the reactor is cooled and flashed for separation of entrained catalyst. The liquid product is stripped from hydrogen chloride, which is recycled. A flow scheme of the Isomate process is shown in Fig. 16. The operating temperature in an isomerization process with liquid catalyst can be lowered by employing AlCl3 in combination with SbCl3 in the form of a lowmelting eutectic mixture. This catalyst system was used in Shell’s liquid-phase isomerization process [39–41], a flow scheme of which is shown in Fig. 17. The cold light naphtha feed is dried and contacted countercurrently with a stream of catalyst from the reactor to extract active components from the latter stream, leaving a waste stream of spent catalyst. After addition of recycled HCl, the naphtha

Water wash

Stabilizer

2820

Isomate product

Caustic wash

Spent catalyst

Caustic

Flow scheme of the Isomate process for light naphtha isomerization [38].

Drying

Catalyst scrubbing

Reaction

Catalyst Gas recovery scrubbing

HCL stripping

Soda treating

Recycle gas Product Vent gas Water

Spent catalyst Feed Fig. 17

Caustic Catalyst recycle

Flow scheme of the Shell liquid phase isomerization process [39].

13.7.5 Isomerization of Pentanes and Hexanes for Octane Enhancement

stream enters the reactor, where isomerization takes place at a temperature between 60 and 100 ◦ C. The reactor effluent containing dissolved catalyst is passed to a column where catalyst is recovered as a bottoms stream which is recycled to the reactor. Fresh catalyst is added to the latter stream for activity maintenance. The isomerized naphtha stream is freed from HCl by stripping and HCl is recycled. Typical data obtained in isomerization of a mixed C4 –C6 feed with Shell’s liquid-phase isomerization process are listed in Table 6. The data show the relatively high proportion of isoalkanes, notably isopentane in the pentanes fraction and 2,2-dimethylbutane in the hexanes fraction, which is consistent with the relatively low operating temperature (see Figs. 13 and 14). The data also show the relatively low proportion of isobutane in the product, illustrating the difficulty of isomerizing n-butane under conditions where pentanes and hexanes are readily isomerized to concentrations close to equilibrium (see Section 13.7.2). 13.7.5.2.2 Processes Using Bifunctional Catalysts Consisting of Pt Supported on Chlorinated Alumina The stable operation possible with bifunctional catalysts operating in the presence of hydrogen have rendered alkane isomerization processes using monofunctional catalysts with their complications arising from catalyst deactivation practically obsolete. Among the earlier processes for isomerization of pentanes and hexanes with platinum supported on chlorinated alumina is the BP isomerization process of British Petroleum [42–46]. A flow scheme of this process is shown in Fig. 18. Isomerization of mixed C4 –C6 feeda by the Shell liquid-phase isomerization process [41]

Tab. 6

Component

C3 and lighter Iso-C4 n-C4 Iso-C5 n-C5 2,2-Dimethylbutane 2,3-Dimethylbutane 2-Methylpentane 3-Methylpentane n-C6 Others

Concentration in isomerizate/wt.%

Relative proportion in carbon number fraction/%b

0.7 3.8 10.3 49.6 13.5 5.0 1.4 4.0 2.2 1.4 8.1

– 27.9 (72.3) 72.1 (27.7) 78.7 (81.0) 21.3 (19.0) 35.7 (41.3) 10.0 (9.9) 28.6 (27.0) 15.7 (14.0) 10.0 (7.8) –

a Straight-run feed contained 0.7 wt.% propane, 13 wt.% butanes, 62 wt.% pentanes, 14 wt.% hexanes and 10 wt.% C5 and C6 cycloalkanes. b Values in parentheses are equilibrium data.

2821

The light naphtha feed is purified to remove water and sulfur compounds and freed from aromatics by prehydrogenation over a nickel catalyst prior to passage over the isomerization catalyst consisting of platinum on an alumina support, which is activated by chlorinating with HCl (added as such or as an organic chloride which generates hydrogen chloride in situ). The process is carried out between 100 and 160 ◦ C at a pressure of about 2 MPa and at a hydrogen : hydrocarbon molar ratio of about 2. Hydrogen is recovered from the effluent and recycled. Notwithstanding the use of predried feed, some chloride is stripped from the catalyst so that continuous replenishment of chloride is necessary. Hydrogen chloride stripped from the catalyst is recovered from the product and recycled via an absorber system. In a later version, the BP process was simplified considerably since it was found that the level of chloride as catalyst activator can be lowered considerably to the point where HCl recovery facilities can be eliminated. It was also recognized that prehydrogenation over a separate catalyst is not needed, since the platinum on the isomerization catalyst in the inlet part of the reactor can perform the same function [47]. This simplified version of the BP process has much in common with the Penex process to be discussed below. The Penex process of Universal Oil Products (UOP) is among the oldest processes based on the use of platinum on chlorided alumina as catalyst and is very widely used [48–54]. A flow scheme of the Penex process is shown in Fig. 19. Desulfurized C5 , C6 or C5 –C6 feed is passed through a molecular sieves drier, combined with recycle hydrogen and passed through the isomerization reactor or two isomerization reactors in series. The catalyst consists of platinum on alumina which, as in the BP process, is activated with chloride. Hydrogen is separated from the reactor effluent, combined with make-up hydrogen to compensate for hydrogen losses (consumed in aromatics saturation, in hydrocracking and solution in product) and recycled. The liquid product is stabilized, and the stabilizer overhead vapors containing HCl (from decomposition of an organic chloride compound added to the feed) are scrubbed with caustic solution to remove the acid. Continuous addition of chloride activator to the feed ensures maintenance of catalyst activity. A life of up to 4 years has been claimed for the I-8 catalyst [54]. The operating temperatures and pressures are typically 120–170 ◦ C and 2–7 MPa, at a hydrogen : hydrocarbon molar ratio of 1 : 2. Due to the relatively low isomerization temperature, a product with a relatively high degree of branching can be obtained and an octane number References see page 2828

2822

13.7 Isomerization

Desulfurization Isomerization

Dearomatization Furnace Ni catalyst

Activator recovery

Recycle hydrogen

Cat reformer off-gas

Recycle hydrogen

H2S / H2O removal Raw feedstock H2 + H2S

Activator

Product

Make-up hydrogen

Fig. 18

Flow scheme of the BP isomerization process [42].

Make-up hydrogen Dryer

Recycle gas Reactors

Gas to fuel Scrubber

Separator Stabilizer

Fresh/spent caustic

Dryer

C5 /C6 charge

Fig. 19

Isomerate

Flow scheme of the Penex process of UOP [54].

of about 84–85 (RON-0) is attainable in once-through isomerization of a C5 plus C6 feedstock. A later version of the Penex process applies hydrogen in once-through operation. In this HOT (hydrogen once-through) Penex process, several major items of plant equipment can be omitted, namely the recycle compressor, the product separator and associated heat exchangers, resulting in equipment cost savings of about 15% [31].

13.7.5.2.3 Processes Using a Noble Metal on Amorphous Silica–Alumina as Bifunctional Catalyst The use of silica–alumina as a solid catalyst eliminates the drawback of continuous chloride addition, removal of HCl from effluent streams and the precautions needed in the design and operation to avoid corrosion problems. An example of this class of processes is the Pentafining process of Atlantic Refining [55–57]. However, due to the

2823

13.7.5 Isomerization of Pentanes and Hexanes for Octane Enhancement

relatively low acid strength of amorphous silica–alumina, a relatively high operating temperature is required (between 424 and 480 ◦ C) at an operating pressure in the range 2–5 MPa [55]. Because of unfavorable thermodynamic equilibrium at these high temperatures (see Figs. 13 and 14), the octane levels attainable are relatively low (see Fig. 15). For this reason, C5 and C6 isomerization processes using amorphous silica–alumina have become obsolete and have been replaced by processes using chlorinated alumina (see above) or processes applying zeolites (see below).

RON -3 ml TEL /USG

95 ∆T = 5 °C 94

93 10 ppm H2O in gas

100

Processes Using a Noble Metal on Zeolite as Bifunctional Catalyst The most prominent example in this category is the Hysomer process, which is offered for license by Shell and by UOP (formerly by Union Carbide) [60–62]. The zeolite used as acidic component, H-mordenite, is a much stronger acid than amorphous silica–alumina and allows operation at temperatures in the range 240–280 ◦ C. With platinum as the hydrogenactivating component on the catalyst, the process operates under hydrogen at a total pressure between 0.8 and 3 MPa. An important advantage of zeolitic catalysts over chlorinated alumina-based catalysts is their much greater robustness. In particular, water in the feed can be tolerated up to fairly high concentrations, obviating the need for expensive drying facilities. Continuous addition of a chloride activator, removal of HCl from effluent streams and precautions against chloride corrosion are unnecessary. The Hysomer catalyst can also tolerate a fair amount of sulfur in the feedstock, rendering hydrotreatment of the feedstock superfluous in most cases. The tolerance of the Pt/H-mordenite catalyst of the Hysomer process for water and sulfur is shown in Figs. 20 and 21. A flow scheme of the Hysomer process is shown in Fig. 22, which illustrates the much greater simplicity of this process compared with processes using aluminum chloride. The ruggedness of the Hysomer catalyst is also demonstrated by its long life: catalyst charges have been used for up to 7 years in commercial operation. A catalyst deactivated by operational mishaps can in most cases be regenerated by a simple carbon burn-off. Table 7 lists some representative data on the isomerization of a C5 plus C6 feed with the Hysomer process. Due to the higher operating temperature, the octane levels of the products are somewhat lower than for the Penex process (RON-0 81–82 versus 84–85). The Hysomer process, which saw its first commercial application in 1970, is now employed in a large number of plants. More recently, a rather similar process based on a

200 ppm H2O in gas

92 98

13.7.5.2.4

96

94

92

C5+ yield / wt.%

RON - 3 ml TEL / USG

Fig. 20 Effect of water on the isomerization of light naphtha over Pt/H-mordenite.

95

94 No S added 14 wt. ppm added to feed

93

100

98

96

94

92

C5+ yield / wt.% Fig. 21 Effect of sulfur on the isomerization of light naphtha over Pt/H-mordenite.

modified zeolite impregnated with platinum (Procatalyse IS-632) has been developed by the Institut Fran¸cais du P´etrole (IFP) [63]. Another recent development is that of the Hysopar catalyst, which reportedly is a zeolite catalyst containing noble metals and which is applied in a newly developed process, the CKS (Cepsa–Kellogg–S¨udChemie) Isom process [64]. Combining C5 and C6 Alkane Isomerization with Physical Separations Due to the thermodynamic equilibrium limitations discussed earlier, complete conversion of n-alkanes into isoalkanes is not achieved in once-through operation. To 13.7.5.3

References see page 2828

2824

13.7 Isomerization

Reactor Recycle

C5 /C6 feed

H2 Product

Fig. 22

Flow scheme of the Shell Hysomer process.

Typical compositions of feed and products of the Hysomer and TIP processes

Tab. 7

Component

Feed

Hysomer product

TIP product

C4 and lighter/wt.% Iso-C5 /wt.% n-C5 /wt.% 2,2-Dimethylbutane/wt.% 2,3-Dimethylbutane/wt.% 2- and 3-Methylpentane/wt.% n-C6 /wt.% Cyclo-C5 and -C6 /wt.% RON-0 C5+ yield/wt.%

0.7 29.3 44.6 0.6 1.8 13.9 6.7 2.4 73.2 –

1.8 49.6 25.1 5.0 2.2 11.3 2.9 2.1 82.1 97.5

2.8 72.0 2.0 5.5 2.5 13.4 0.1 1.8 90.7 96.8

obtain complete isomerization, the isomerization process can be combined with a physical separation process which allows the isolation of residual n-alkanes from the isomerization product for recycling and conversion to extinction. Since isoalkanes boil at lower temperatures than their straight-chain isomers, this separation can be effected by distillation. In the case of pentane isomerization, n-pentane can be isolated from the isomerate as bottom product in a deisopentanizer column. Similarly, for hexane isomerization a deisohexanizer column can be installed. A disadvantage of separation by distillation is that the proportion of residual n-alkanes in the isomerized product is generally rather low, whereas the capacity of the fractionator is determined by the bulk of isoalkanes to be distilled off as overhead product. In the case of isomerization of mixed C5 and C6 feed, complete isomerization becomes quite involved.

Particularly in the last case, iso/normal separation by selective molecular sieve adsorption becomes of interest. Zeolite 5A selectively adsorbs n-alkanes since isoalkanes have a too large molecular diameter to enter the intracrystalline pores. Since the required capacity of the adsorption unit is mainly determined by the amount of n-alkanes to be adsorbed, molecular sieve adsorption is advantageous at relatively low residual contents of n-alkanes. Examples of iso/normal separation using zeolite 5A in a pressure-swing adsorption process are the BP pressure swing process of British Petroleum [65] and the Isosiv process of Union Carbide [66–68]. An alternative molecular sieve iso/normal separation process is the Molex process of UOP [69–71]. In contrast to the BP and Isosiv processes, which operate at relatively high temperatures in the gas phase, the Molex process operates at a relatively low temperature in the liquid phase. Desorption is effected by using a light n-alkane as desorption aid, and this desorber is recovered from the extract phase by distillation. A special feature of the Molex process is the use of a single cylindrical adsorbent vessel provided with multiple inlet/outlet ports connected to a multiport rotary valve. Thus, with a stationary adsorbent column, a continuous chromatographic separation is simulated. The BP isomerization process has been combined with the BP pressure swing n-alkane adsorption process to achieve complete isomerization of n-alkanes and a commercial plant was built in the mid-1970s [72]. The Penex–Molex combination was commercialized in 1990 [31]. A special advantage of a zeolite-based isomerization process such as Hysomer is that within the window of feasible operating conditions (temperature, pressure and hydrogen flow-rate), these can be chosen so as to fit the requirements of an Isosiv process mode with desorption by purging with hydrogen instead of by pressure reduction. Thus, rather than by applying two separate processes, isomerization and separation may be closely integrated in a single process with significant savings. This concept was realized in the total isomerization package (TIP) process [73–78]. Figure 23 shows a flow scheme of the TIP process, in which the Hysomer reactor is situated upstream of the Isosiv adsorbers (Hysomer lead option). This is the preferred configuration if the feed is rich in n-alkanes. However, if the feed is rich in isoalkanes and cyclic hydrocarbons, the reverse sequence will be more advantageous (Isosiv lead option). These alternative configurations are shown in Fig. 24. The application of molecular sieve separation and recycle of n-alkanes not only leads to a higher octane number of the product, but also decreases the effect of

13.7.5 Isomerization of Pentanes and Hexanes for Octane Enhancement

2825

Hysomer Isosiv

Feed

Product H2

Fig. 23

Flow scheme of the TIP process (Hysomer lead option).

a: Hysomer lead

92

C4− Isosiv

Stabilizer

Hysomer Feed

With n - alkanes recycle 90

Isoalkanes product b: Isosiv lead

RON - 0

88

n -alkanes recycle

86 Once through 84

Isosiv

Hysomer

Feed

Stabilizer

C4−

82 Pt / Cl /Alumina Pt / H-MOR 80 50

iso- + n -alkanes recycle

Fig. 24

Alternative configurations of the TIP process.

isomerization temperature on product quality. Therefore, the disadvantage of the zeolite-based isomerization processes with their higher operating temperature than the processes using chlorided alumina largely disappears in recycle operation, as can be inferred from Fig. 25. Table 7 compares the product compositions obtained in isomerization of light naphtha with the Hysomer process in the once-through mode and with the TIP process. The TIP process was commercialized in 1975 and since that time an appreciable number of plants have been built. A more recently proposed extension of the TIP process is its integration with UOP’s SafeCat technology. This SafeCat isomerization scheme is shown in Fig. 26. Instead of installing a stand-alone hydrotreater to remove sulfur from the feed, the hydrotreater reactor is incorporated in the same gas circuit as the Hysomer and Isosiv units. Because of the similar pressures in hydrotreating, isomerization and adsorption, a simple line-up is possible with a single gas loop and a minimum of heating and cooling equipment.

100

150

200

250

300

350

Operating temperature / °C Fig. 25 Effect of temperature on attainable octane numbers in isomerization without and with recycle of n-alkanes. Feed: 60% pentanes, 30% hexanes, 10% cyclics [63].

Another process that combines isomerization and physical separation of n- and isoalkanes is the more recently developed Ipsorb Isom process of IFP. The isomerization reactor product is separated in a molecular sieve separation section of the vapor phase, pressure swing adsorption type (Ipsorb). A first unit was put into service around 1996 [79, 80]. Probably the closest integration of isomerization and normal/iso separation of alkanes is the application of reactive distillation or a membrane reactor. Conceptual schemes for membrane catalysis in alkanes isomerization are shown in Fig. 27 for the production of either isoalkanes or n-alkanes [81]. Recently, some experimental proof of this concept has been obtained by Gora and Jansen in a laboratory reactor containing a silicalite-1 membrane and a chlorided Pt/alumina catalyst [82]. However, the development of an industrial process on References see page 2828

2826

13.7 Isomerization

Hydrogen recycle

HT

ISOM

ADS

H2 makeup

Feed

DES

Gas

HT = Hydrotreating ADS = Adsorption DES = Desorption ISOM = Isomerization

Safecat

Isomerate

Fig. 26

SafeCat isomerization scheme [31].

Isoalkanes H2

Selective membrane Isomerization catalyst

n

n

n

n

n

n

n

n H2 i + n Feed

i + n Feed H2

n + H2 i+n

(a)

Fig. 27

Normal alkanes

(b)

Conceptual schemes for membrane catalysis. (a) For production of n-alkanes; (b) for production of isoalkanes.

this basis will be far from easy and still represents a great challenge for the future. 13.7.6

Processes for Isomerization of Higher Alkanes

For reasons explained earlier, selective isomerization of alkanes higher than hexanes is difficult since cracking generally accompanies isomerization. For this reason,

isomerization of heptanes with the objective of octane enhancement has found little application, since cracking to gaseous hydrocarbons represents a loss of product value. Another reason is that the octane values of the main monobranched isomers are still relatively low, namely RON-0 values of 45 and 65 for 2- and 3-methylhexane, respectively. The isomerization of higher n-alkanes that may be solid at room temperature is a way to lower the pour

13.7.7 Processes for Conversion of n-Butane to Isobutane

point. Thus, lubricating oils can be produced from waxy oil fractions by isomerization in the presence of hydrogen using bifunctional catalysts. In this application, some hydrocracking can be accepted, since the cracked fragments that still boil in the desired range do not lead to loss of product yield but merely to some reduction in average molecular weight. Moreover, smaller fragments boiling below the main product range can still be valuable by-products as long as they are liquid. A process for converting slack wax to lubricating base oil of very high viscosity index (>145) by hydrocracking/isomerization has been developed by Shell and is applied commercially as part of a complex to produce lubricating base oils by catalytic hydroprocessing of waxy distillates and deasphalted oils [83]. A more recent process in the field of wax isomerization for lube oil production is the Isodewaxing process of Chevron [84]. The Isodewaxing catalyst typically contains a hydrogenation component on an intermediate pore silicoaluminophosphate molecular sieve (SAPO). The intermediate pore size suppresses the formation of highly branched structures, which is favorable since monobranched structures are more desirable components as they combine a relatively low pour point with a high viscosity index. A Pt/SAPO-11 catalyst was found capable of isomerizing linear alkanes of long chain length with a higher selectivity than can be achieved with a catalyst consisting of platinum on amorphous alumina. This can be seen from Table 8, which shows a comparison of these two catalysts in the isomerization of n-hexadecane [85]. Isomerization of heavier alkanes also plays a role in the conversion of the product of the Fischer–Tropsch synthesis step of the Shell middle distillate synthesis (SMDS) process. This Fischer–Tropsch product, which consists mainly of long-chain n-alkanes, is hydrocracked to produce hydrocarbons boiling in the kerosene and gasoil ranges. In this hydrocracking step, isomerization of n-alkanes which accompanies the cracking is essential

Isomerization of n-hexadecane at 6.9 MPa, 3.1 weight hourly space velocity, H2 : HC = 30 and 96% conversion [85]

Tab. 8

Parameter

Temperature/◦ C Isomerization selectivity/wt.% n-C16 in C16 product/wt.% Methyl-C15 in C16 product/wt.% Dimethyl-C14 in C16 product/wt.% Other C16 in C16 product/wt.% Pour point/◦ C

Pt/SAPO-11 catalyst

Pt/silica– alumina catalyst

340 85 4.7 53.3 29.8 12.2 −51

360 64 6.0 21.6 37.8 34.6 −28

2827

to meet the requirements on cold flow properties of the products [86–89]. A similar application of hydroisomerization of a waxy Fischer–Tropsch product is the conversion of this product to a liquid that can be easily shipped by pipeline or in a conventional crude tanker, as part of Exxon’s advanced gas conversion technology [90, 91]. 13.7.7

Processes for Conversion of n-Butane to Isobutane

As mentioned before, n-butane is hardly isomerized under the conditions of C5 and C6 isomerization, but may be converted to isobutane under more severe conditions by the alkylation–isomerization–dealkylation mechanism, using the same type of catalysts. As can be seen in Fig. 13, the conversion of nbutane into isobutane is limited by the thermodynamic equilibrium, even more so than the isomerization of higher alkanes. Therefore, isomerization of n-butane is generally coupled with iso/normal separation in a deisobutanizer or isobutane is allowed to react selectively in the isomerized mixture, for example by reaction with alkenes in the alkylation process. Processes for the manufacture of isobutane from n-butane either use AlCl3 activated with HCl as a monofunctional catalyst or Pt supported on chlorided alumina as a bifunctional catalyst. Monofunctional catalysts based on AlCl3 /HCl have been applied in older processes developed before and shortly after World War II. An example of this type of process is the catalytic isomerization process of Phillips Petroleum [92, 93]. This uses a solid catalyst consisting of AlCl3 sublimed on to bauxite. Dried n-butane is vaporized, superheated and mixed with recycled HCl vapor and passed over the catalyst in a number of fixed-bed reactors. HCl is removed from the condensed product by stripping and recycled. A similar catalyst is used in the butane vaporphase isomerization process of Shell [94, 95]. This process operates at temperatures of 100–140 ◦ C and pressures in the range 1–2 MPa. The liquid-phase isomerization process of Shell, discussed earlier in the context of C5 and C6 isomerization, has also been used for n-butane isomerization in a number of commercial units [39–41, 95]. Butane isomerization is typically carried out at 90–110 ◦ C and about 2–3 MPa. As in the case of C5 and C6 isomerization, monofunctional catalysts suffer from deactivation in n-butane processing and catalyst is consumed in the process. Catalyst deactivation is minimized by using bifunctional catalysts operating under hydrogen pressure. Among the References see page 2828

13.7 Isomerization

butane isomerization processes using bifunctional catalysts, the most prominent example is the Butamer process of UOP [96, 97]. This process, which was actually a predecessor of the Penex process, uses a similar platinum on chlorided alumina catalyst, which can be operated for a long period in the presence of hydrogen. Another process which uses platinum on chlorided alumina is the BP C4 isomerization process. Typical conditions are a pressure between 1.5 and 3 MPa and a temperature in the range 150–210 ◦ C [98].

15

Million tons per year

2828

10

5

13.7.8

Role of Alkane Isomerization in the Hydrocarbon Processing Industry

1970

75

80

85

90

95

2000

Year

Isomerization of light alkanes has been applied commercially for many years in the hydrocarbon processing (petroleum refining) industry. Particularly during World War II, demand for alkylate as component for aviation gasoline caused a rapid increase in butane isomerization: from the first commercial unit coming on stream in 1941 to nearly 40 units by the end of the War. The replacement of high-octane gasoline by kerosene as the main aviation fuel has stopped the growth of butane isomerization, although the process still retains an important place in hydrocarbon processing. In the 1990s, interest in isobutane manufacture revived because of the rapid growth in demand for methyl tert-butyl ether (MTBE) as an octane-boosting component in gasoline, spurred by the advent of unleaded gasoline and the specified presence of oxygenated compounds in so-called reformulated or green gasoline. Because the availability of isobutene for MTBE manufacture from the traditional source (catalytic cracker gases) is limited, production of isobutene by isomerizing n-butane followed by dehydrogenation gained interest. An alternative possibility is the skeletal isomerization of n-butenes present in catalytic cracker gases and a suitable zeolitic catalyst for this conversion has been found [99]. Isomerization of C5 and C6 alkanes commenced towards the end of World War II, to provide additional blending stock for aviation gasoline. The advent of catalytic reforming over platinum-type catalysts has reduced the need for isomerization of light naphtha for octane enhancement in the production of gasoline for cars. In the 1980s, however, the drive to remove lead additives from gasoline led to a rapid growth in the need for C5 /C6 isomerization capacity. This is illustrated by Fig. 28, showing the increase in Hysomer and TIP capacity in that period. With the increasing global demand for gasoline and environmental concerns (spreading to developing countries as well) that led to a drive for a reduction in

Fig. 28 Growth of capacity of installed Hysomer units, either as standalone unit or integrated in TIP.

aromatics in gasoline and of MTBE, isomerization of light naphtha remains of interest in recent times. The total capacity of light alkane isomerization has been estimated at more than a million barrels (160 000 m3 ) per day in 1993. This capacity is about equally split between butane isomerization, light naphtha isomerization over Pt/Cl/alumina and light naphtha isomerization over Pt/zeolite [31] and implies a large number of plants. UOP, the leading licensor in C5 /C6 isomerization technology, was reported to have commissioned 188 C5 /C6 isomerization units as of the second quarter of 2002 [100]. Isomerization of higher alkanes in the wax range is of growing importance as the moderately branched alkanes produced are excellent components for lubricating oils. Hydrocracking/hydroisomerization of synthetic heavy wax produced by a Fischer–Tropsch process provides a possibility for the manufacture of high-quality fuels and other hydrocarbon products from sources other than petroleum, such as natural gas and coal. This option may well become more important in the future when demands on the quality (cleanliness) of fuels increase while petroleum resources diminish. References 1. G. Egloff, G. Hulla, V. I. Komarewsky, Isomerization of Pure Hydrocarbons, Reinhold, New York, 1942, pp. 21–50. 2. F. E. Condon, in Catalysis, P. H. Emmett (Ed.), Vol. 6, Reinhold, New York, 1958, Ch. 2, pp. 43–178. 3. J. L. Franklin, in Carbonium Ions, G. A. Olah, P von R. Schleyer (Eds.), Vol. 1, Interscience, New York, 1968, p. 8. 4. F. E. Condon, in Catalysis, P. H. Emmett (Ed.), Vol. 6, Reinhold, New York, 1958, p. 121. 5. D. M. Brouwer, Recl. Trav. Chim. Pays Bas 1968, 87, 1435.

References 6. C. D. Nenitzescu, in Carbonium Ions, G. A. Olah, P von R. Schleyer (Eds.), Vol. 1, Interscience, New York, 1968, pp. 19–20. 7. J. D. Roberts, C. C. Lee, W. H. Saunders, J. Am. Chem. Soc. 1954, 76, 4501. 8. P. N. Rylander, S. Meyerson, J. Am. Chem. Soc. 1956, 78, 5799. 9. D. M. Brouwer, J. M. Oelderik, Presented at the Division of Petroleum Chemistry, American Chemical Society, Meeting, San Francisco, CA, April 1968. 10. D. M. Brouwer, J. M. Oelderik, Recl. Trav. Chim. Pays Bas 1968, 87, 721. 11. F. Chevalier, M. Guisnet, R. Maurel, in Proceedings of the 6th International Congress on Catalysis, G. C. Bond, P. B. Wells, F. C. Thompkins (Eds.), Vol. 1, Chemical Society, London, 1977, p. 478. 12. O. Weisser, Int. Chem. Eng. 1963, 3, 408. 13. M. Steijns, G. Froment, P. Jacobs, J. Uytterhoeven, J. Weitkamp, Ind. Eng. Chem. Prod. Res. Dev. 1981, 20, 654. 14. J. Weitkamp, Ind. Eng. Chem. Prod. Res. Dev. 1982, 21, 550. 15. J. A. Martens, P. A. Jacobs, J. Catal. 1990, 124, 357. 16. A. Fiaux, D. L. Smith, J. H. Futrell, Int. J. Mass Spectrom. Ion Phys. 1977, 25, 281. 17. S. T. Sie, Ind. Eng. Chem. Res. 1992, 31, 1881. 18. S. T. Sie, Ind. Eng. Chem. Res. 1993, 32, 403. 19. J. M. Oelderik, J. C. Platteeuw, in Proceedings of the 3rd International Congress on Catalysis, W. M. H. Sachtler, G. C. A. Schuit, P. Zwietering (Eds.), North-Holland, Amsterdam, 1965, p. 736. 20. H. W. Kouwenhoven, in Molecular Sieves, W. M. Meier, J. B. Uytterhoeven (Eds.), Advances in Chemistry Series, Vol. 121, American Chemical Society, Washington, DC, 1973, p. 529. 21. J. Dwyer, Chem. Ind. (London) 1984, April 2, 258. 22. P. B. Koradia, J. R. Kiovski, M. Y. Asim, J. Catal. 1980, 66, 290. 23. S. van Donk, A. Broersma, O. L. J. Gijzeman, J. A. van Bokhoven, J. H. Bitter, K. P. de Jong, J. Catal. 2001, 204, 272. 24. M. Tromp, J. A. van Bokhoven, M. T. Garriga Oostenbrink, J. H. Bitter, K. P. de Jong, D. C. Koningsberger, J. Catal. 2000, 190, 209. 25. A. J. Koster, U. Ziese, A. J. Verkleij, A. H. Janssen, K. P. de Jong, J. Phys. Chem. B 2000, 104, 9368. 26. P. A. Lawrance, A. A. Rawlings, in Proceedings of the 7th World Petroleum Congress, Vol. 4, Elsevier Applied Science, Barking, 1967, p. 135. 27. G. F. Asselin, H. S. Bloch, G. R. Donaldson, V. Haensel, E. L. Pollitzer, Presented at the Symposium on Advances in Gasoline Technology, Division of Petroleum Chemistry, American Chemical Society, Meeting, New York, August 1972. 28. W. C. van Zijll Langhout, in Proceedings of the 9th World Petroleum Congress, Vol. 5, Elsevier Applied Science, Barking, 1975, p. 197. 29. International Petroleum Encyclopedia, Penwell, Tulsa, OK, 1978, p. 368. 30. M. E. Reno, R. S. Haizmann, B. H. Johnson, P. P. Piotrowski, A. S. Zarchy, Hydrocarbon Int. 1990/1991, 73. 31. P. J. Kuchar, J. C. Bricker, M. E. Reno, R. S. Haizmann, Fuel Process. Technol. 1993, 35, 183. 32. G. Martino, in Proceedings of the 12th International Congress on Catalysis, A. Corma, F. V. Melo, S. Mendioroz, J. L. G. Fierro

33. 34.

35.

36.

37. 38. 39. 40.

41. 42.

43. 44. 45. 46. 47.

48. 49. 50.

51.

52. 53. 54. 55. 56. 57. 58. 59. 60. 61.

2829

(Eds.), Studies in Surface Science and Catalysis, Vol. 130, Part A, Elsevier, Amsterdam, 2000, p. 83. B. Domergue, L. Watripont, World Refiner 2000, 10(4), 26. O. Clause, L. Mank, G. Martino, J. P. Franck, in Proceedings of the 15th World Petroleum Congress, Vol. 2, Wiley, Chichester, 1998, p. 695. P. J. Kuchar, R. D. Gillespie, C. D. Christopher, W. C. Martin, M. J. Cleveland, P. J. Bullen, Int. J. Hydrocarbon Eng. 1999, 4(3), 50. American Petroleum Institute, Project 44, in Selected Values of Physical and Thermodynamic Properties of Hydrocarbons and Related Compounds, F. D. Rossini, K. S. Pitzer, R. L. Arnett, R. M. Braun, G. C. Pimentel (Eds.), Carnegie Press, Pittsburgh, PA, 1953. Anon., Hydrocarbon Process. Pet. Refiner 1964, 43(9), 177. H. G. Krane, E. W. Kane, Pet. Refiner 1957, 36(5), 177. H. D. Evans, E. B. Fountain, W. S. Reveal, W. E. Ross, Hydrocarbon Process. Pet. Refiner 1961, 40(9), 171. W. M. J. Ruedisulj, H. D. Evans, E. B. Fountain, in Proceedings of the 6th World Petroleum Congress, Verein f¨ur die F¨orderung des 6. Welt-Erd¨ol-Kongresses, Hamburg, 1963, Sect. III, Paper 9, p. 1. W. M. J. Ruedisulj, H. D. Evans, E. B. Fountain, Hydrocarbon Process. Pet. Refiner 1963, 42(7), 125. T. C. O’May, in Proceedings of the 6th World Petroleum Congress, Verein f¨ur die F¨orderung des 6. Welt-Erd¨olKongresses, Hamburg, 1963, Section III, paper 4, p. 17. B. W. Burbidge, J. K. R. Rolfe, Hydrocarbon Process. Pet. Refiner 1966, 45(8), 168. A. H. Richardson, J. H. D. Hooper, Presented at the NPRA Meeting, San Francisco, CA, March 1971. A. H. Richardson, M. F. Olive, Presented at the AIChE Meeting, Houston, TX, March 1971. A. H. Richardson, J. H. D. Hooper, Petrochem. Int. 1971, 11(12), 52. B. W. Burbidge, P. Watson, A. H. Richardson, G. J. Wanless, Presented at the NPRA Meeting, San Antonio, TX, March 1975. D. H. Belden, Petr. Refiner 1956, 35(10), 149. L. E. Dean, H. R. Harris, D. H. Belden, V. Haensel, Oil Gas J. 1958, 56(9), 54. D. H. Belden, V. Haensel, W. C. Starnes, R. C. Zabor, Presented at the American Petroleum Institute Meeting, Philadelphia, PA, 1957. H. S. Bloch, G. R. Donaldson, V. Haensel, Presented at the Symposium on Isomerization and Related Processes, Division of Petroleum Chemistry, American Chemical Society, Meeting, Boston, MA, April 1959. R. A. Erickson, G. F. Asselin, Chem. Eng. Prog. 1965, 61(3), 53. G. Bour, C. P. Schwoerer, G. F. Asselin, Oil Gas J. 1978, 76 (October 26), 57. R. J. Schmidt, J. A. Weiszmann, J. A. Johnson, Oil Gas J. 1985, 83 (May 27), 80. G. R. Worrell, Oil Gas J. 1956, 54(46), 183. G. R. Worrell, Pet. Refiner 1956, 35(4), 138. G. R. Worrell, World Pet. 1956, 27(8), 74. H. R. Grane, J. K. Ozawa, G. R. Worrell, Pet. Refiner 1957, 36(5), 172. Anon., Process Handbook Hydrocarbon Process. 1964, 43(9), 180. H. W. Kouwenhoven, W. C. van Zijll Langhout, Chem. Eng. Prog. 1971, 67(4), 65. Anon., Oil Gas J. 1971, 69 (March 8), 44.

2830

13.8 Alkylation of Isobutane with Light Alkenes on Solid Catalysts

62. Anon., Oil Gas J. 1971, 69 (August 16), 62. 63. A. Hennico, J.-P. Cariou, Hydrocarbon Int. 1990/1991, 68. 64. F. F. Mike, M. F. Gilbert, E. K¨ohler, Int. J. Hydrocarbon Eng. 1998, 3(8), 42. 65. J. Grebbell, G. J. Wanless, Presented at the NPRA Meeting, San Antonio, TX, March 1975. 66. M. F. Symoniak, A. C. Frost, Oil Gas J. 1971, 69 (March 15), 76. 67. J. H. Olive, M. F. Symoniak, Oil Gas J. 1972, 70 (June 26), 68. 68. G. J. Griesmer, W. F. Avery, M. N. Y. Lee, Hydrocarbon Process. 1965, 44(6), 147. 69. D. B. Carson, D. B. Broughton, Pet. Refiner 1959, 38(4), 130. 70. D. B. Broughton, A. G. Lickus, Pet. Refiner 1961, 40(5), 173. 71. M. J. Sterba, Hydrocarbon Process. 1965, 44(6), 151. 72. C. W. Cartwright, R. J. Stock, Oil Gas J. 1978, 76 (September 18), 141. 73. M. F. Symoniak, R. A. Reber, R. M. Victory, Hydrocarbon Process. 1972, 51(9), 122. 74. J. J. Collins, M. F. Symoniak, W. C. J. Quick, Presented at the NPRA Meeting, San Antonio, TX, March 1978. 75. W. C. J. Quik, J. J. Collins, Presented at the DGMK Meeting, Nuremberg, October 1972. 76. W. C. J. Quik, J. J. Collins, Erd¨ol, Kohle-Erdgas-Petrochem. 1972, 25, 706. 77. M. F. Symoniak, R. A. Reber, R. M. Victory, Hydrocarbon Process. 1973, 52(5), 101. 78. M. H. Hainsselin, M. F. Symoniak, G. R. Cann, Presented at the NPRA Meeting, San Antonio, TX, March 1975. 79. Anon., Hydrocarbon Process. 1996, 75(11), 138. 80. Anon., Hydrocarbon Process. 2002, 81(11), 130. 81. S. T. Sie, in Advanced Zeolite Science and Applications, J. C. Jansen, M. St¨ocker, H. G. Karge, J. Weitkamp (Eds.), Studies in Surface Science and Catalysis, Vol. 85, Elsevier, Amsterdam, 1994, p. 625. 82. L. Gora, J. C. Jansen, Oil Gas Eur. Mag. 2005, 1, 11. 83. S. Bull, A. Marmin, in Proceedings of the 10th World Petroleum Congress, Vol. 4, Heyden, London, 1980, p. 221. 84. S. J. Miller, M. A. Shippey, G. M. Masada, Presented at the NPRA National Fuels and Lubricants Meeting, Houston, TX, November 1992. 85. S. J. Miller, in Zeolites and Related Microporous Materials: State of the Art 1994, J. Weitkamp, H. G. Karge, H. Pfeifer, W. H¨olderich (Eds.), Studies in Surface Science and Catalysis, Vol. 84, Part C, Elsevier, Amsterdam, 1994, p. 2319. 86. M. J. van der Burgt, J. van Klinken, S. T. Sie, Presented at the 5th Synfuels Worldwide Symposium, Washington, DC, November 1985. 87. M. J. van der Burgt, J. van Klinken, S. T. Sie, Erd¨ol Erdgas Kohle 1986, 102, 351. 88. M. J. van der Burgt, C. J. van Leeuwen, J. J. Del’Amico, S. T. Sie, in Methane Conversion, D. M. Bibby, C. D. Chang, R. F. Howe, S. Yurchak (Eds.), Studies in Surface Science and Catalysis, Vol. 36, Elsevier, Amsterdam, 1988, p. 473. 89. J. Eilers, S. A. Posthuma, S. T. Sie, Catal. Lett. 1990, 7, 253. 90. C. C. Lahn, R. F. Bauman, B. Eisenberg, J. M. Hochman, Presented at the European Applied Conference on Natural Gas, Eurogas 92, Trondheim, Norway, June 1992. 91. B. Eisenberg, L. L. Ansell, R. A. Fiato, R. F. Bauman, Presented at the 73rd Annual Gas Processors Association Convention, New Orleans, LA, March 1994. 92. Anon., Process Handbook, Hydrocarbon Process. 1964, 43(9), 174.

93. G. H. Unzelman, C. J. Wolf, in Petroleum Processing Handbook, W. F. Bland, R. L. Davidson (Eds.), McGraw-Hill, New York, 1976, pp. 3–51. 94. H. A. Cheney, C. L. Raymond, Trans. Am. Inst. Chem. Eng. 1946, 42, 595. 95. Anon., Pet. Refiner 1952, 31(9), 168. 96. H. W. Grote, Oil Gas J. 1958, 56(13), 73. 97. Anon., Oil Gas J. 1971, 69 (August 16), 66. 98. Anon., Hydrocarbon Process, 1982, 61(9), 170. 99. H. H. Mooiweer, K. P. de Jong, B. Kraushaar-Czarnetzki, W. H. J. Stork, B. C. H. Krutzen, in Zeolites and Related Microporous Materials: State of the Art 1994, J. Weitkamp, H. G. Karge, H. Pfeifer, W. H¨olderich (Eds.), Studies in Surface Science and Catalysis, Vol. 84, Part C, Elsevier, Amsterdam, 1994, p. 2327. 100. Anon., Hydrocarbon Process. 2002, 81(11), 132.

13.8

Alkylation of Isobutane with Light Alkenes on Solid Catalysts Yvonne Traa and Jens Weitkamp∗

13.8.1

Introduction

The alkylation of isobutane with light alkenes to produce higher hydrocarbons in the gasoline range has been known for a long time. Light alkenes such as propene, butenes and pentenes can be used as alkylating agents. In commercial practice, however, mostly butenes are applied and this is why we shall often refer to isobutane/butene alkylation in this chapter. In industry, the reaction is catalyzed by liquid HF or H2 SO4 . Therefore, the alkylation of isobutane with alkenes is a classic example of old liquid acid-catalyzed processes where environmentally benign solid catalysts have been sought for a long time. However, until recently, this search had not been successful, and considerable research efforts are still being undertaken. The importance of the topic is documented by recent review articles focusing on new research in this field [1–4]. The literature up to about 1995 was reviewed in the first edition of this Handbook [5]. The current review will focus on the main lines of development using solid catalysts since then. 13.8.2

Stoichiometry of the Alkylation of Isobutane with Light Alkenes

In chemistry, alkylation is the generic term for a broad variety of reactions which have in common that an alkyl ∗

Corresponding author.

13.8.3 Role of Isobutane/Butene Alkylation in Petroleum Refining

group of arbitrary (but usually well defined) carbon number is introduced into an arbitrary substrate by means of an alkylating agent, typically an alkene, an alcohol or an alkyl halide. In petroleum refining, the term alkylation more specifically refers to a process located downstream of the fluid catalytic cracking unit (FCCU) (see Chapter 13.5) and meant to convert part of the C4 hydrocarbons formed as by-products in the FCCU into the so-called alkylate, a most valuable component in the gasoline pool of the refinery. In a strongly oversimplified manner, the stoichiometry of the alkylation of isobutane with 2-butene can be represented by Eq. (1): +

→

(1)

However, isobutane/alkene alkylation differs fundamentally from numerous other alkylation reactions, in that it is inadequate to describe the stoichiometry by this simple equation. For example, under typical alkylation conditions, the conversion of isobutane with 2-butene (cis- or trans-isomer) does not lead to high yields of 2,2,3trimethylpentane, either in HF or in H2 SO4 . Rather, alkylate is a complex mixture of isoalkanes, which is not only composed of all possible trimethylpentanes (2,2,3-, 2,2,4-, 2,3,3- and 2,3,4-trimethylpentane), certain dimethylhexanes and, even though in small concentrations, the methylheptanes, but also of isoalkanes with less (isopentane, isohexanes and isoheptanes) and more (isononanes, etc.) than eight carbon atoms (‘‘odd carbon numbers’’). Examples of product distributions obtained from a real C4 cut in an industrial plant are listed in Tables 1 and 2. Such product distributions unambiguously indicate that, in spite of the mild reaction conditions applied, intensive skeletal rearrangements and even some carbon–carbon bond ruptures must have taken place. Indeed, the chemistry of isobutane/alkene alkylation in hydrofluoric acid and sulfuric acid is extremely complex, and numerous acid-catalyzed types of reaction have been claimed to be involved [8–10], such as hydrogen transfer, destructive alkylation, disproportionation, conjunct polymerization, ester formation, formation of acid sludges and numerous others. In this context, the systematic and thorough studies by Albright and coworkers deserve particular attention [11–23]. Albright’s work focuses on the H2 SO4 -catalyzed alkylation, and a good summary of his mechanistic views is presented in Ref. [20]. Hutson and Hays discussed in detail the peculiarities and various steps involved in the HF-catalyzed reaction [24]. Specific features of the mechanism of isobutane/alkene alkylation on solid catalysts are discussed in Section 13.8.5.

2831

Alkylation of isobutane with butenes: typical carbon number distributions in wt.% obtained from a technical C4 cut, after Refs. [6, 7]

Tab. 1

Catalyst HF Light ends: i-C5 H12 i-C6 H14 i-C7 H16 -------------------------

H2 SO4

5 4 4 --------------

8 7 6 ---

i-C8 H18 74 64 -----------------------------------------Heavy ends: i-C9 H20 and higher

13

15

13.8.3

Role of Isobutane/Butene Alkylation in Petroleum Refining

In modern petroleum refining, the alkylation of isobutane with butenes is often combined with the manufacture of methyl or ethyl tert-butyl ether (MTBE, ETBE), which are other desirable gasoline components (see Chapter 13.10), as shown schematically in Fig. 1. The C4 cut from the FCCU may contain some small amounts of butadiene, which would have harmful effects both in the etherification and in the alkylation. One method to remove butadiene is its partial hydrogenation to nbutenes. In the etherification, isobutene reacts selectively and quantitatively with methanol or ethanol on an organic ion-exchange resin catalyst (see Chapter 13.10). The remaining C4 hydrocarbons are fed to the alkylation unit together with a relatively large stream of recycled isobutane. In the acid-catalyzed alkylation, the butenes react selectively and completely with isobutane, the product being alkylate, whereas n-butane is inert under the alkylation conditions. The excess isobutane is separated from the reactor effluent and recycled. In industrial practice, two acids have found application as catalysts for isobutane/butene alkylation, viz. hydrofluoric acid and concentrated sulfuric acid. Figure 2 shows that the production capacity for alkylate has increased considerably within the last 10 years (for the conversion of barrels per calendar day into tons per year, the typical density of alkylate from an HF process was used, i.e. 0.697 g cm−3 ). By January 1, 2006, the worldwide production capacity for alkylate amounted to 39.1 × 106 and 33.6 × 106 t a−1 for the HF and the H2 SO4 processes, respectively, which adds up to 72.7 × 106 t a−1 [25]. In addition, other alkylation units existed with a combined References see page 2851

2832

13.8 Alkylation of Isobutane with Light Alkenes on Solid Catalysts

Tab. 2 Alkylation of isobutane with butenes: typical isomer distributions in the C5 H12 to C8 H18 fractions in mol% produced from a technical C4 cut, after Refs. [6, 7]

Catalyst HF

H2 SO4

n-Pentane 2-Methylbutane 2,2-Dimethylpropane

0 100 0

3 97 0

n-Hexane 2-Methylpentane 3-Methylpentane 2,2-Dimethylbutane 2,3-Dimethylbutane

0 25 11 0 64

0 17 9 0 74

n-Heptane 2-Methylhexane 3-Methylhexane 3-Ethylpentanea 2,2-Dimethylpentane 2,3-Dimethylpentane 2,4-Dimethylpentane 3,3-Dimethylpentane 2,2,3-Trimethylbutanea

0 6 3 0 4 35 52 0 0

0 3 2 0 4 34 57 0 0

n-Octane 0 0 -----------------------------------------2-Methylheptane  3-Methylheptane 3-Ethylhexane 4-Methylheptanea --------------------------

0 0 ----------------

2,2-Dimethylhexane 2,3-Dimethylhexane 2,4-Dimethylhexane 2,5-Dimethylhexane 3,3-Dimethylhexane 3,4-Dimethylhexane 3-Ethyl-2-methylpentanea 3-Ethyl-3-methylpentane --------------------------

0 7 6 5 0 0.8 0 0 ------------

0 6 5 8 0 0.4 0 0 ----

2 53 12 14 -------------

2 39 18 21 ----

2,2,3-Trimethylpentane 2,2,4-Trimethylpentane 2,3,3-Trimethylpentane 2,3,4-Trimethylpentane ------------------------2,2,3,3-Tetramethylbutane a Probably

0.1

0.1

0.3

0.4

0

as a blending component in modern gasoline [26, 27]. In fact, alkylation gasoline is relatively rich in hydrogen and, hence, environmentally friendly, and it exhibits excellent combustion properties in spark ignition engines, including high research octane numbers (RON) around 95 and motor octane numbers (MON) which are only slightly lower, viz. in the vicinity of 93 [28, 29]. Figure 3 shows that, while North America continues to be the realm of isobutane/butene alkylation, the process does exist in refineries all over the world. The commercially most important alkylation processes working with HF as catalyst are the UOP and the Phillips processes, and the relevant processes using H2 SO4 are the so-called effluent refrigeration and cascade autorefrigeration processes. Typical process conditions are listed in Table 3. A good description of these commercial liquid-phase alkylation processes has been given by Corma and Mart´ınez [29]. The existing alkylation technology, while producing high-quality, environmentally benign gasoline components, suffers from a number of disadvantages and drawbacks. In particular, the sulfuric acid processes tend to suffer from an unusually high catalyst consumption, which may reach 70–100 kg per ton of alkylate produced. The spent sulfuric acid contains tarry hydrocarbons (ca. 7 wt.%) and water. It has been estimated that replacement of sulfuric acid in alkylation plants accounts for about 89% of the total volume of refinery catalysts [30]. Spent sulfuric acid can be regenerated, typically by burning at ca. 1000 ◦ C, which converts the acid into SO2 . Clean SO2 is oxidized on a V2 O5 catalyst to SO3 , which is subsequently absorbed in H2 SO4 . The cost of the thus recovered acid has been estimated to be two to three times that of sulfuric acid available on the market [30]. Newer process options are said to reduce acid consumption by 10–20% by processing propene, butenes and pentenes (amylenes) in different reactors [31] and by at least 50% by using a novel contactor [32].

Typical process conditions for isobutane/butene alkylation, after Ref. [28]

Tab. 3

0

present in very small amounts, but not detected.

capacity of 4.5 × 106 t a−1 , for which the nature of the catalyst has not been disclosed, and one further unit with an alkylate capacity of 0.3 × 106 t a−1 was under construction in 2006 [25]. This huge process capacity totaling almost 80 × 106 t a−1 reflects the desirability of alkylate

Temperature/ ◦ C Pressure/bar Residence time/min V˙ acid /V˙ total hydrocarbons a V˙ isobutane /V˙ alkenes a Acid consumption per mass of alkylate produced/kg t−1 a Liquid

HF processes

H2 SO4 processes

25–40 8–20 5–20 1–2 10–20 0.4–1

4–10 4–6 20–30 1–2 8–12 70–100

flow-rates (volume) including recycled isobutane.

2833

13.8.3 Role of Isobutane/Butene Alkylation in Petroleum Refining

C4 from FCCU

H2

i -Butane recycle Methanol or ethanol

Feed pretreatment

n -Butane i -Butane

n -Butane i -Butane

Etherification

n -Butenes i -Butene

n -Butane

n -Butenes

Alkylation Alkylate

MTBE or ETBE

Typical position of an alkylation unit in a modern refinery. FCCU: fluid catalytic cracking unit.

30

2.5 2.0

60

1.5 HF Processes

40

1.0

20

0.5

H2SO4 Processes 1996

1998

2000 2002 Year

2004

2006

H2SO4 Processes HF Processes

20

28

10

4

0 0

52 53

Total

80

Capacity / 106 t a−1

Capacity / 106 t a−1

100

Capacity / 106 bpcd

Fig. 1

0.0

7

19 15

North South Europe America America

17 0

Asia Pacific

6

Africa

Geographical distribution of processes for the alkylation of isobutane with butenes, as of January 1, 2006 (North America includes Mexico). The circled numbers above the bars indicate the number of alkylation units in the respective region. Data from Ref. [25].

Fig. 3

Worldwide alkylate capacity (bpcd: barrels per calendar day). Data from Ref. [25]. The fact that the sum of the capacities of the HF and the H2 SO4 processes does not exactly equal the total alkylate capacity is due to insufficient information on the kind of liquid catalyst used in some alkylation units.

Fig. 2

In the hydrofluoric acid processes, the catalyst consumption is considerably lower (below 1 kg t−1 ); however, there has been much concern about the toxicity of HF combined with its volatility and corrosiveness. To mitigate these potential problems and reduce the risks, a package of safety measures has been introduced in recent years by refiners who operate HF alkylation units. These safety measures include high-flux water sprays, a reduced acid inventory, a rapid acid dump system, a diminution of flange connections, on-site HF regeneration and additives reducing the volatility and aerosol-forming tendency of HF [1, 27, 33–36]. Thus, HF units can be considered as much safer nowadays. Together with the increasing demand for alkylate, this increased safety may tend to favor HF alkylation over H2 SO4 alkylation, since the HF units are not only easier to expand, but also have lower maintenance and operating costs than H2 SO4 units [27].

The ultimate goal is, of course, to replace the existing alkylation technology by a completely new process which relies on a non-toxic, non-corrosive, easy-to-handle and environmentally friendly catalyst. In this context, a new alkylation process using a composite ionic liquid catalyst has recently been retrofitted into an existing 65 000 t a−1 H2 SO4 alkylation unit in China [37, 38]. Thus, higher yields of alkylate with higher research octane numbers can be produced than in the H2 SO4 unit, but in a safer, environmentally friendly and virtually non-corrosive environment [38]. However, more attention has been paid to solid catalysts, not only true solid acids such as zeolites or heteropoly acids, but also more traditional liquid and gaseous acids artificially ‘‘solidified’’ on suitable carriers. In fact, a broad chemical variety of solid or solidified acids has been tried out as catalysts in isobutane/alkene References see page 2851

2834

13.8 Alkylation of Isobutane with Light Alkenes on Solid Catalysts

alkylation in the 1980s and early 1990s. The pertinent older literature has been reviewed in detail in Ref. [5]. The present chapter focuses on the more recent literature on solid acid-catalyzed alkylation and on developments with regard to the regeneration of deactivated alkylation catalysts and the design of the process. Currently, two solid acid-catalyzed processes are offered for licensing at an industrial scale [1, 32], but it appears that up to 2005 only one agreement had been signed [32]. However, it has also been argued that political issues such as the current European discussion and possible legislation on the use and transportation of chemicals could have a positive impact on isobutane alkylation over solid catalysts (see Chapter 13.1, Section 13.1.3.2.2D). Thus, isobutane/alkene alkylation on solid acids is, to the best of our knowledge, not yet operated on an industrial scale. A detailed description of the processes designed so far is given in Section 13.8.10. However, a variant of the above-described (direct) alkylation of isobutane, the so-called indirect alkylation or isooctane technology, has been licensed recently, and this is why this process is briefly mentioned here [26, 31, 32, 39–43]. The indirect alkylation can be described as dimerization of isobutene followed by hydrogenation of the isooctene products to isooctane according to Eq. (2): 4

+

+H2

2

(2) In the presence of linear alkenes, the reaction is more complex because linear butenes can react with isobutene to form codimers with different molecular structures [43]. Thus, the product of the indirect alkylation is very similar to the product of isobutane/alkene alkylation, C4 from steamcracker or FCCU

and the two processes are competing with each other. The advantages of the indirect alkylation are (i) the high flexibility with respect to the feed: the C4 cut from steam crackers or from the FCCU can be used as a feedstream [40] (cf. the simplified process flow diagram in Fig. 4); (ii) the possibility to revamp idle MTBE units into indirect alkylation units in regions where ethers such as MTBE are phased out; (iii) the product flexibility: isooctenes can directly be used as gasoline blending components or, with restriction of the gasoline alkene content, after hydrogenation to the isooctanes. In addition, isooctenes and isooctanes have higher research octane numbers than alkylate produced by direct isobutane/alkene alkylation [31, 40, 43]. All these advantages have led to the rapid commercialization of the isobutene dimerization technology: nine plants have been licensed, of which five were already in operation in 2005 [32, 41]. All subsequent sections of this chapter will be devoted exclusively to the direct alkylation of isobutane with butenes on solid catalysts. 13.8.4

Salient Features of Isobutane/Alkene Alkylation on Solid Catalysts

If isobutane and an alkene are reacted over certain solid acids, mixtures of higher hydrocarbons with all features of alkylates produced conventionally in HF or H2 SO4 can indeed be obtained. However, the salient feature of solid acids in isobutane/alkene alkylation is the rapid deterioration of their catalytic properties while on stream, which seems to be the major obstacle for a breakthrough of these materials in industry. A typical time-on-stream

H2

Solvent recovery

n -Butane i -Butane n -Butenes

Feed pretreatment

n -Butane i -Butane n -Butenes i -Butene

Dimerization

H2

Olefin saturation

Fig. 4

Light products

Alkylate

Process flow diagram of the dimerization of isobutene followed by hydrogenation of the olefinic products (‘‘indirect alkylation’’).

13.8.4 Salient Features of Isobutane/Alkene Alkylation on Solid Catalysts

Butene conversion / %

100 80 Catalyst: Ce-Y

60 40 20 0

0

100

200

300

of large amounts of octenes. At this point, the question arises as to how long the product is considered to be an alkylate. Arbitrarily, we have set this limit at a content of alkanes in the C8 product fraction of 90 mol%. In the example shown in Fig. 6, the alkylation stage then ends after 31 min. Typical product distributions in the alkylate formed on the Ce-Y zeolite catalyst at various times on stream during the initial alkylation stage are given in Tables 4 and 5. A comparison with the data in Tables 1 and 2 unambiguously shows that the alkylate formed on the solid catalyst is composed of the same hydrocarbons as industrial alkylates produced either in HF or H2 SO4 . It is furthermore evident from Table 4 that pronounced changes occur in the alkylate with time on stream: as the reaction proceeds, more and more octanes and heavy ends are produced at the expense of the light ends. The only reasonable way to account for the occurrence of light and heavy ends, i.e. products with odd carbon numbers, is to invoke the formation of C8 , C12 , C16 and perhaps even higher molecular weight intermediates by butene oligomerization from the C4 feed hydrocarbons followed by cracking steps. A closer look at the product distributions reveals that all possible isoalkanes with at least one tertiary carbon atom are formed, whereas alkanes lacking a tertiary carbon atom, such as n-alkanes, 2,2-dimethylpropane, 2,2-dimethylhexane or 2,2,3,3-tetramethylbutane, are completely absent (see Table 5). Even within a given carbon number fraction, pronounced changes occur with time on stream with respect to the distribution of individual isomers. For example, in the C7 H16 fraction, the content of 2,3-dimethylpentane increases at the

. . . nOctanes / (nOctanes + nOctenes) / %

behavior is shown in Fig. 5 [44]. The experiment was conducted with a Ce-Y zeolite with a high concentration of acid sites generated by exchange of 98% of the original sodium cations by trivalent cerium. It should be noted that only the conversion of butene into hydrocarbons with five or more carbon atoms was considered, i.e. the doublebond isomerization into cis- and trans-2-butene was not included. In fact, when n-butenes appeared in the product after ca. 30 min, they were always recovered as a mixture with a composition close to the equilibrium composition, which is approximately 5 mol% 1-butene, 30 mol% cis-2butene and 65 mol% trans-2-butene. Initially, butene is completely converted. However, this complete conversion lasts only a limited time, in this particular case ca. 30 min, whereupon the butene conversion drops rapidly and finally stabilizes at a value far below 100% (in the example in Fig. 5 at 30–35%). Perhaps much more serious than the decline in activity is the concomitant loss of selectivity. Initially, during the stage of complete butene conversion, a perfect alkylate is formed which is composed of isoalkanes with 5 to ca. 10 carbon atoms exclusively. However, approximately at the time when the butene conversion begins to drop, alkenes appear in the C+ 5 product spectrum in increasing amounts, and the carbon number fractions representing multiples of the carbon number of the feed alkene, i.e. C8 and C12 , begin to grow considerably. In other words, there is a complete shift away from isobutane/butene alkylation towards butene oligomerization. A typical result is shown in Fig. 6 for isobutane/1-butene conversion on the same highly exchanged Ce-Y zeolite. The composition of the C8 product fraction is plotted versus time on stream. Initially, octanes are formed exclusively, after ca. 30 min, more and more octenes appear. In the dotted region of the curve, the accuracy is relatively low, because it is difficult to determine small amounts of octanes in the presence

2835

100 Alkylate

90 mol% Alkanes

80 60

Catalyst: Ce-Y

40 Oligomerizate

20 0

0

20

40

60

80

100

Time on stream / min

400

Time on stream / min Conversion of a liquid isobutane/1-butene mixture on a Ce-Y zeolite, after Ref. [44]. Composition of the C8 fraction (reaction conditions as in Fig. 5).

Fig. 6

Conversion of a liquid isobutane/1-butene mixture on a Ce-Y zeolite in a fixed-bed reactor, after Ref. [44]. Temperature: 80 ◦ C; liquid feed rate at room temperature: 7.5 cm3 h−1 ; n˙ isobutane /˙n1−butene = 11; pressure: 31 bar; catalyst mass: 1.42 g.

Fig. 5

References see page 2851

2836

13.8 Alkylation of Isobutane with Light Alkenes on Solid Catalysts

Alkylation of isobutane with n-butenes on a Ce-Y zeolite catalyst at 80 ◦ C: typical carbon number distributions in wt.% as a function of time on stream, after Refs. [45, 46] Tab. 4

Alkylation of isobutane with n-butenes on a Ce-Y zeolite catalyst at 80 ◦ C: typical isomer distributions in the C5 H12 to C8 H18 fractions in mol% as a function of time on stream, after Refs. [45, 46]

Tab. 5

Time on stream/min Time on stream/min 1 Light ends: i-C5 H12 i-C6 H14 i-C7 H16 --------------

5

22.6 14.3 11.9 9.6 9.0 8.3 -----------------

15

8.7 7.0 7.6 --------

30

6.1 5.0 5.3 ---

i-C8 H18 47.0 49.5 50.6 53.8 -----------------------------------------Heavy ends: i-C9 H20 i-C10 H22 to i-C12 H26

6.4 3.1

7.0 11.4

7.5 18.6

8.1 21.7

expense of 2,4-dimethylpentane. Most pronounced in the C8 H18 fraction is the enhanced formation of 3,4dimethylhexane with time on stream, which can also be explained by butene dimerization [see Eq. (3)] starting from a butyl cation formed by addition of a proton to n-butene. The isooctyl cation shown can be desorbed as 3,4-dimethylhexene by leaving a proton behind or as 3,4dimethylhexane by uptake of a hydride ion from a hydride ion donor. +

+

→

+

(3)

Table 5 moreover reveals that the overall alkylate quality drops with time on stream: the total content of the most desired trimethylpentanes in the C8 H18 fraction diminishes from 77.8 to 53.9%. At this point, it should be mentioned that all these pronounced trends in the carbon number distribution (Table 4) and isomer distributions (Table 5) remain hidden if the product analysis is done on an integral, time-averaged sample collected over an extended time on stream. It is obvious that such analytical data are questionable, or even meaningless, if the distribution of the products formed on the catalyst changes significantly during the time of sampling. In addition, as Tables 4 and 5 show, alkylate is a complex mixture of isomeric hydrocarbons extending over a large range of carbon numbers (C4 to ca. C10 ) and boiling points. Taking representative samples of such mixtures without partial loss of light or heavy components, the storage of these product samples and the complete separation of all components during the gas chromatographic analysis require experimental skill and carefully optimized procedures. Differential, instantaneous sampling combined

1

5

15

30

n-Pentane 2-Methylbutane 2,2-Dimethylpropane

0 100.0 0

0 100.0 0

0 100.0 0

0 100.0 0

n-Hexane 2-Methylpentane 3-Methylpentane 2,2-Dimethylbutane 2,3-Dimethylbutane

0 14.7 22.3 0 63.0

0 17.5 22.2 0 60.3

0 19.4 22.4 0 58.2

0 17.3 26.9 0 55.8

n-Heptane 2-Methylhexane 3-Methylhexane 3-Ethylpentane 2,2-Dimethylpentane 2,3-Dimethylpentane 2,4-Dimethylpentane 3,3-Dimethylpentane 2,2,3-Trimethylbutane

0 5.8 6.8 1.0 0 29.0 51.7 0 5.8

0 7.0 8.9 1.2 0 35.4 43.6 0 3.9

0 7.7 10.8 1.2 0 38.9 38.3 0 3.1

0 5.9 10.9 2.2 0 41.9 34.7 0 4.4

n-Octane 0 0 0 0 -----------------------------------------2-Methylheptane  3-Methylheptane 3-Ethylhexane 4-Methylheptane --------------------

0.1 0.1 0.2 0.1 ----------------------

2,2-Dimethylhexane 2,3-Dimethylhexane 2,4-Dimethylhexane 2,5-Dimethylhexane 3,3-Dimethylhexane 3,4-Dimethylhexane 3-Ethyl-2-methylpentane 3-Ethyl-3-methylpentane --------------------

0 0 0 4.8 8.4 13.1 6.2 6.4 7.4 3.0 3.1 3.8 0 0 0 6.9 11.3 13.5 0.6 1.0 1.3 0 0 0 -------------------

0 11.7 4.9 2.3 0 24.3 1.8 0 ---

2,2,3-Trimethylpentane 2,2,4-Trimethylpentane 2,3,3-Trimethylpentane 2,3,4-Trimethylpentane --------------------

4.4 22.3 27.9 23.2 ------

2.8 12.2 19.6 19.3 ---

2,2,3,3-Tetramethylbutane

0.2

0.3

0.3

0.2

0.4

0.7

0.8

0.8

0

3.6 18.3 24.9 21.9 -----0

3.2 17.7 20.7 18.0 ------0

0

with high-resolution gas chromatography is therefore mandatory if one aims at an in-depth understanding of the most complex chemistry of isobutane/alkene alkylation on solid acids. A more detailed description of suitable experimental techniques may be found in Ref. [5].

13.8.4 Salient Features of Isobutane/Alkene Alkylation on Solid Catalysts

The principal problem that has so far prevented solid acid catalysts from being used in isobutane/alkene alkylation on an industrial scale is obvious from Figs. 5 and 6: after an unacceptably short time on stream, the catalysts lose their activity with a concomitant shift of the selectivity towards oligomerization of the alkene. The species that are responsible for this serious deterioration of the catalyst properties are primarily formed from the alkene, rather than from isobutane. Consequently, the alkene concentration at any location inside the reactor should be kept as low as possible. One way to achieve this is to use a high n˙ isobutane /n˙ alkene ratio in the feed. This measure, however, has its economic limitations, because any excess isobutane has to be separated from the product downstream of the alkylation reactor and recycled (cf. Fig. 1), and the cost of this separation increases with the amount of unconverted isobutane (in industrial practice with liquid acids, n˙ isobutane /n˙ alkene ratios up to ca. 10 or 20 seem to be tolerable). Another important factor that influences the local alkene concentration is the type of reactor: in any reactor with no or little back-mixing [plug-flow reactor (PFR)], in particular in a fixed-bed reactor, relatively high alkene concentrations will inevitably exist at the entrance; conversely, in an ideally back-mixed reactor [continuous stirred tank reactor (CSTR)], as approximated by a continuous stirred slurry reactor, the alkene concentration is very low at any location inside the reactor. It is for this reason that today, most groups involved in research on isobutane/alkene alkylation are using CSTRs [47–49]. Another option, described by Ketikidis et al. [50], is to use a slurry transport reactor with addition of the alkene feed at various locations along the reactor axis. The results presented in Figs. 5 and 6 were obtained using a continuously operated fixed-bed plug-flow reactor. In a continuous stirred slurry reactor, under otherwise comparable conditions, the duration of the initial alkylation stage with complete alkene conversion is considerably extended, say from ca. 30 min to times of the order of several or ten hours [47, 48]. However, apart from this extension of the time scale, the principle features of the reaction remain unchanged. With regard to the catalyst, it has already been mentioned that a broad variety of acid solids were shown to possess alkylation activity. Examples are acid forms of zeolites, liquid Brønsted or Lewis acids coated onto solid (often porous) supports, perfluorinated polymers containing sulfonic acid groups, sulfated zirconia and titania, etc. [5]. In recent years, some other systems have been added to this list, e.g. WO3 /ZrO2 [51], immobilized ionic liquids [52], Nafion/silica composites [53–56], amorphous microporous molecular sieves [57] and chlorinated alumina modified with alkali metal cations [58]. It appears, nevertheless, that zeolite catalysts continue to be

2837

favored by those groups who conducted recent in-depth studies on isobutane/butene alkylation [47, 48, 59–62]. Medium-pore zeolites such as H-ZSM-5 or H-ZSM-11 were shown to be unsuitable (see Section 13.8.5). Among the large-pore zeolites, Brønsted acid forms of zeolites X and Y (FAU) appear to be a good choice [47, 48, 63–66], inter alia because of their low price and their high ionexchange capacity, permitting a high concentration of catalytically active Brønsted acid sites. In a typical experiment performed with an isobutane/2butene feed and an La-X zeolite catalyst in a continuous stirred slurry reactor, Lercher and coworkers [47] achieved alkylate yields (mass flow of product hydrocarbons with five and more carbon atoms out of the reactor divided by the mass flow of cis-2-butene into the reactor) slightly above 2. This is the value expected for the idealized alkylation stoichiometry [Eq. (4)]. i-C4 H10 + C4 H8 −−−→ i-C8 H18

(4)

With the characteristic alkene conversion of 100% and an initial isooctane selectivity of 85 wt.% [47], the catalyst cannot be improved much further with regard to activity and selectivity. The principal hurdle which remains to be overcome before the solid acid-catalyzed isobutane/butene alkylation can be realized industrially is the loss of activity and selectivity after a too short time on stream. Absolute times on stream, however, are not meaningful if the performance of various catalysts or the performance of one and the same catalyst under different reaction conditions are to be compared. Rather, it is appropriate (see Ref. [5], Section 3.14.4.3 for details) to convert the absolute time on stream into the dimensionless catalyst age, which is defined as catalyst age ≡ t × m ˙ butene,fed /mcatalyst

(5)

where t is the time on stream, m ˙ butene,fed is the mass flow of the butene used in the feed mixture at the reactor entrance and mcatalyst is the mass of catalyst. The numerator in Eq. (5) represents the mass of butene fed cumulatively to the reactor from the beginning of the experiment until the time t. With the above-mentioned La-X zeolite catalyst, the butene conversion was 100% until a catalyst age of about 2. Subsequently, the conversion dropped sharply and so did the trimethylpentane production [47]. The precise value of the catalyst age at which these undesired events occur is, of course, dependent on many parameters. However, for a good alkylation catalyst and with a well back-mixed reactor, it will be of the same order of magnitude. Therefore, the main lines of today’s research, which are References see page 2851

2838

13.8 Alkylation of Isobutane with Light Alkenes on Solid Catalysts

also major topics of the rest of this chapter, are (i) to establish the reasons for and eventually eliminate catalyst deactivation or (ii) to accept the fast catalyst deactivation, but to design processes with alkylation and intermittent catalyst regeneration which overcome this problem. There has recently been limited progress with the first topic (see Section 13.8.6). By contrast, fresh ideas in the work on the second topic can be perceived (see Section 13.8.8). Prior to discussing the progress in these two areas, the common views on the mechanism of isobutane/butene alkylation are discussed in the next section. 13.8.5

Mechanistic Considerations

For an advanced discussion of the mechanistic views on isobutane/alkene alkylation, reference can be made to the publications of Ipatieff and Schmerling [8, 10], Kennedy [9], Buiter et al. [67], Albright [20, 68, 69], Hutson and Hays [24], Kazansky and Vasina [70, 71], Daage and Fajula [72, 73], Corma and Mart´ınez [29], Guisnet and Gnep [74], Feller et al. [47], Paukshtis et al. [75], Hommeltoft et al. [76, 77] and to previous reports from our group (see Ref. [5] for a review). The following discussion will focus on the alkylation mechanism on solid acid catalysts. It is to be understood, however, that the striking similarities in the composition of alkylates produced in liquid acids (Tables 1 and 2) and on the surface of solid acids (Tables 4 and 5) lead one to invoke the same principal chemistry and the same reactive intermediates in both cases. There is indeed general consensus that isobutane/alkene alkylation proceeds via carbocations, both in liquid and on the surface of solid acids. Recent work [78–80] indicates that most hydrocarbons are adsorbed at acid sites of a solid catalyst in the form of alkoxy species, i.e. with a considerable contribution of homopolar bonding. While adsorbed, these alkoxy species can undergo a variety of chemical reactions, of which true carbocations are transition states. It is therefore justified to discuss hydrocarbon transformations on solid acid catalysts in terms of carbocation chemistry known to a large extent and in considerable detail from investigations in liquid superacids [79, 80]. A classification of carbocations into classical, tricoordinated carbenium ions and non-classical, pentacoordinated carbonium ions and the reactions by which these are formed and which they typically undergo can be found in, e.g., Chapter 5.4. Carbocation chemistry not only governs the mechanism of isobutane/alkene alkylation, but also those of other important refinery processes, such as catalytic cracking (see Chapter 13.5), hydrocracking (Chapter 13.6) and isomerization of light naphtha (Chapter 13.7). Most research groups believe that Brønsted acid centers on the solid surface are active sites in isobutane/alkene

alkylation [81, 82], e.g. a bridging OH group near a framework aluminum atom in a zeolite (H+ Zeol.− ). It has also been claimed [47, 62, 63] that, in active alkylation catalysts, the acid strength must be beyond a certain threshold value. According to Kob et al. [83], this threshold may be around 130 kJ mol−1 , expressed as the neutralization enthalpy in the calorimetric titration with pyridine. It should be mentioned, however, that it has also been reported that the butene conversion depends on the Lewis acid site strength only and that there is no correlation between the Brønsted acid sites and the alkylation activity [75]. Carbocation reactions, which play an important role in the complex mechanism of isobutane/alkene alkylation, include: (I) Protonation of the alkene, whereby an alkylcarbenium ion is formed, e.g. + −− −− → n-C4 H8 + H+ − ← − CH3 −CH −CH2 −CH3

(6)

(II) Alkylation, i.e. addition of an alkene to an alkylcarbenium ion, e.g. + −−− −− → i-C4 H9 + + n-C4 H8 ← − i-C8 H17

(7)

In the alkylation step, a new carbon–carbon σ -bond is formed by addition of the double bond in the alkene to the positively charged carbon atom in the alkylcarbenium ion. The reverse step (which governs the mechanisms of catalytic cracking and hydrocracking; see Chapters 13.5 and 13.6) is called β-scission. β-Scissions are usually classified [84–86] into type A, B and C β-scissions (Fig. 7), depending on the nature of the carbenium ions involved. Type A β-scissions proceed easily, whereas type B and even more so type C β-scissions are more demanding. It is very likely that, at the low temperatures of isobutane/alkene alkylation, only type A β-scissions can proceed. (III) Intramolecular hydride shift, i.e. migration of a hydride ion from an adjacent carbon atom to the positively charged carbon atom in an alkylcarbenium ion, e.g. 3,4,4-trimethylpentyl-(2) cation −−− −− → ← − 2,2,3-trimethylpentyl-(3) cation

(8)

(IV) Rearrangement of the carbon skeleton in a chemisorbed alkylcarbenium ion, e.g. a shift of a methyl group along the main chain in an isooctyl cation, such as 2,2,3-trimethylpentyl-(3) cation −−− −− → ← − 2,3,3-trimethylpentyl-(2) cation

(9)

13.8.5 Mechanistic Considerations

Fig. 7

n≥6

Type C:

R1

+

n≥7

Type B1:

R1

+

n≥7

Type B2:

R1

+

n≥8

Type A:

R1

+

sec.

sec.

R1

R2

sec.

+

R1

tert.

sec. R1

R2

tert. R2

R2

+

tert.

R2

2839

tert. R1

R2

+

+

R2

R2

Classification of β-scissions of alkylcarbenium ions, after Refs. [84–86]; n is the carbon number.

The reaction shown in Eq. (9) is a so-called type A rearrangement, i.e. the number of branchings in the cation remains constant, as opposed to type B rearrangements in which the number of branchings increases or decreases [87–89]. Type B rearrangements proceed much more slowly and after a completely different mechanism, i.e. via protonated cyclopropanes [87, 88] (see also Chapter 13.7), than type A rearrangements. It is likely that, under the mild reaction conditions of isobutane/alkene alkylation, only type A rearrangements can occur. (V) Intermolecular hydride transfer, typically from an isoalkane to an alkylcarbenium ion, e.g. + −−−→ i-C4 H10 + i-C8 H+ 17 ←−−− i-C4 H9 + i-C8 H18

(10) Hydride transfer is the mechanism for interconversion of alkanes and alkylcarbenium ions chemisorbed at the acidic catalyst sites. It is relatively rapid if both the reactant and the product carbenium ion are tertiary, but it is much less favored if a chemisorbed tertiary carbenium ion is to react with an alkane which can only give a secondary or a primary carbenium ion [90, 91]. Hydride transfer is perhaps the least well understood step in the mechanism of isobutane/alkene alkylation on solid catalysts, even though its role is of utmost importance and intermolecular hydride transfer may well be rate-controlling for the overall reaction (see below). In its simplest version, the mechanism of isobutane/nbutene alkylation is usually described in terms of the cycles depicted in Figs. 8 and 9, where (meant just as an example) the negatively charged framework of a zeolite Zeol.− is supposed to act as the counterion of the carbocations. The two schemes hold for any nbutene, since double bond shift between 1-butene and

2-butene and even more so cis-/trans-isomerization in 2butene, are usually very rapid reactions which occur under alkylation conditions even at moderately strong acidic sites. An equilibrium or near-equilibrium mixture of the three n-butenes will therefore be available for alkylation, regardless of which individual n-butene isomer is used in the feed. In such equilibrium mixtures, 2-butene strongly prevails (>90 mol%). The tert-butyl cation left behind, e.g. from an intermolecular hydride transfer [Eq. (10)], is alkylated by 2-butene (Fig. 8), whereby the secondary 3,4,4-trimethylpentyl-(2) cation forms. An intramolecular hydride shift [Eq. (8)] leads to the tertiary 2,2,3trimethylpentyl-(3) cation, which, upon intermolecular hydride transfer from an isobutane molecule [Eq. (10)] is desorbed as 2,2,3-trimethylpentane, and this closes the cycle. Alternatively, the tert-butyl cation can be alkylated by 1-butene (Fig. 9) which results in the 5,5dimethylhexyl-(3) cation. An intramolecular hydride shift [Eq. (8)] followed by a methyl shift [Eq. (9)] gives the 2,3dimethylhexyl-(2) cation and, ultimately, through an intermolecular hydride transfer [Eq. (10)], 2,3-dimethylhexane. In complete analogy with these pathways of isobutane/ n-butene alkylation, the simplest mechanism of isobutane/isobutene alkylation is shown in Fig. 10. This reaction is also referred to as self-alkylation, with isobutene being generated by hydride transfer from isobutane to n-butene [47]. The pathways shown in Figs. 8 to 10 nicely account for the formation of several highly branched isooctanes from isobutane and C4 alkenes on acid catalysts. On the other hand, they obviously oversimplify the true chemistry in that they fail to explain (in addition to other findings) (i) the formation of light and heavy ends, i.e. of isoalkanes with odd carbon numbers References see page 2851

2840

13.8 Alkylation of Isobutane with Light Alkenes on Solid Catalysts

III +

+

Zeol.−

Isobutane

II

V

+

2-Butene

2,2,3-Trimethylpentane

Simplest mechanism for the formation of a trimethylpentane isomer in the isobutane/n-butene alkylation on an acid catalyst, e.g. a zeolite (H+ Zeol.− ). The Roman numerals refer to the carbocation reactions explained in the text.

Fig. 8

+

IV

III

+ Zeol.−

+

Isobutane V

II + 1-Butene

2,3-Dimethylhexane

Simplest mechanism for the formation of a dimethylhexane isomer in the isobutane/n-butene alkylation on an acid catalyst, e.g. a zeolite (H+ Zeol.− ). The Roman numerals refer to the carbocation reactions explained in the text.

Fig. 9

(ii) the formation of small amounts of methylheptanes in the isobutane/n-butene alkylation (iii) the formation of both dimethylhexanes and small amouts of methylheptanes in the isobutane/isobutene alkylation (iv) alkylate yields (m ˙ product hydrocarbons/ m ˙ butene fed ) other than 2.04 (v) the formation of trimethylpentanes other than 2,2,3trimethylpentane in the isobutane/n-butene alkylation (cf. Fig. 8), the formation of dimethylhexanes other than 2,3-dimethylhexane in the isobutane/nbutene alkylation (cf. Fig. 9) and the formation of trimethylpentanes other than 2,2,4-trimethylpentane in the isobutane/isobutene alkylation (cf. Fig. 10). To account for the obvious occurrence of all isomers ‘‘forbidden’’ according to item (v), extensive skeletal type A rearrangements at the level of trimethylpentyl cations (Figs. 8 and 10) or dimethylhexyl cations (Fig. 9) are usually invoked. Another option is to assume the formation of alkyl-substituted cyclic propenyl cations

at the level of C8 H+ 17 [75]. However, even then the experimental finding remains unexplained why, for instance, surprisingly little 2,2,3-trimethylpentane often occurs in the product of isobutane/n-butene alkylation (cf. Table 5). If skeletal isomerization started from the carbocation with this skeleton, one would conversely expect that 2,2,3-trimethylpentane is formed with some preference. In the fresh state, the solid acid seems to be a good catalyst for intermolecular hydride transfer [step (v)]. Later in the alkylation stage, this step begins to proceed sluggishly, and the alkylcarbenium ions are no longer desorbed as isoalkanes, but rather as isoalkenes by leaving behind something like a free Brønsted site. The mechanism is then better described by a scheme as in Fig. 11: Two 2-butene molecules give a 3,4dimethylhexene, and isooctenes with this skeleton are indeed predominantly found during the oligomerization stage. Interestingly, there seems to be an intermediate stage slightly before the end of the alkylation stage, where

13.8.5 Mechanistic Considerations

2841

+ Isobutane II

V

Zeol.−

+

Isobutene

2,2,4-Trimethylpentane

Simplest mechanism for the formation of a trimethylpentane isomer in the isobutane/isobutene alkylation on an acid catalyst, e.g. a zeolite (H+ Zeol.− ). The Roman numerals refer to the carbocation reactions explained in the text.

Fig. 10

+ III 2-Butene

II + −

+

Zeol.

I

I

2-Butene H+ 3,4-Dimethylhexene

Mechanism for the dimerization of 2-butene on an acid catalyst, e.g. a zeolite (H+ Zeol.− ). The Roman numerals refer to the carbocation reactions explained in the text.

Fig. 11

the carbon skeleton of 3,4-dimethylhexane is already preferred, but the desorption via hydride transfer is still working (Table 5, at 30 min time on stream). Coming back to the reactions of alkylcarbenium ions during the initial alkylation stage, the simple mechanism outlined in Figs. 8 to 10 must obviously be supplemented by additional steps which account for the ubiquitous occurrence of isoalkanes with odd carbon numbers: During the formation of all products, the carbon number of which is not a multiple of 4, a C−C bond cleavage must occur. Alkylation of a butyl cation (formed either by abstraction of a hydride ion from isobutane or by addition of a proton to butene) with a butene gives an octyl cation and this may be further alkylated by butene to a dodecyl cation and so on. In fact, we believe that, at typical alkylation temperatures, a distribution of C4n H+ 8n+1 cations builds up inside the zeolite pores, and there is no argument for claiming that the integer n is limited to low values of, say, 4. However, to account

for the formation of the odd carbon numbers observed experimentally, it is sufficient to limit the size of the carbenium ions at the surface to C16 H+ 33 , i.e. n = 4. Type + A β-scissions of highly branched C12 H+ 25 and C16 H33 cations directly lead to all odd carbon numbers observed experimentally, if one allows for some mild skeletal type A rearrangements while the carbenium ions are adsorbed at the acid sites. With the classification of β-scissions and skeletal rearrangements of alkylcarbenium ions, even the product selectivity can be explained. Whenever the type A rearrangements lead to an α,α,γ -arrangement of three branchings in the carbon skeleton, type A β-scission can readily take place. The formation of all isoalkanes with 5, 6, 7 and 8 hydrocarbons found experimentally in an alkylate can be accounted for if one assumes type A βscissions starting from C12 H+ 25 cations. Interestingly, it is References see page 2851

2842

13.8 Alkylation of Isobutane with Light Alkenes on Solid Catalysts

even possible to explain the low concentration of 2,2,3trimethylpentane in the isooctane fraction: To arrive at the 2,2,3-trimethylpentyl cation, β-scission has to start from a C12 H+ 25 cation with one ethyl branching. It is known from isomerization studies with long-chain alkanes [88, 89] that such isomers with an ethyl side-chain form slowly and at relatively low concentrations only. The formation of isoalkanes with 9, 10 or 11 carbon atoms can be readily explained if one invokes type A β-scissions starting from C16 H+ 33 cations. All these considerations have been explained in more detail in Ref. [5]. The use of medium-pore zeolite catalysts in isobutane/1butene alkylation supports this mechanistic view: zeolites H-ZSM-5 and H-ZSM-11 are inactive at temperatures up to about 100 ◦ C, but above ca. 150 ◦ C, the typical time-onstream behavior shown by large-pore zeolites at ca. 80 ◦ C is observed [92]. Due to the significantly higher temperatures and the spatial constraints inside the pores, the alkylate produced differs markedly from that obtained in zeolite Y: (i) much more light ends (C5 H12 to C7 H16 ) are formed; (ii) large amounts of monomethyl isomers are produced in all carbon number fractions, even n-alkanes occur in the alkylate; (iii) considerable amounts of cyclic hydrocarbons are formed; (iv) trimethylpentanes are absent at any alkylation temperature. While all these product features render the alkylate produced on H-ZSM-5 or H-ZSM-11 zeolites commercially unattractive, they are fully consistent with and furnish completely independent support for the important role of type A β-scissions and the mechanistic pathway sketched above for largepore zeolites. The highly branched precursor carbenium ions required for type A β-scissions are too bulky to be formed in zeolites ZSM-5 or ZSM-11. Hence, alkylation inside these zeolites – although not principally different from the low-temperature alkylation in faujasites – has to proceed via energetically more demanding carbenium ion reactions. These require elevated temperatures and lead to alkylate components with a much lower degree of branching. Also consistent with the mechanistic picture outlined above were kinetic studies under experimental conditions at which catalyst deactivation was reasonably slow and finite ( mordenite(6.5 × 7) > ZSM-5(5.3 × 5.6) [16]. The largepore zeolites such as Omega and HY also undergo rapid deactivation during these processes as a result of the ease of formation of the bulky polynuclear aromatic coke precursors leading to pore blockage, and mordenite is unsuitable for this reaction as a result of its onedimensional structure which leads to rapid deactivation due to pore-mouth blockage. 1 H NMR studies have shown that, on average, the number of carbon atoms between branches for ZSM-5 is four as opposed to two for amorphous silica–alumina and one for solid phosphoric acid. At higher carbon numbers, these values increase for ZSM-5 and decrease for the other two. This property of 13.9.4.1

13.9.4 Catalysts Used in Alkene Oligomerization

ZSM-5 makes it particularly suitable for the production of good-quality distillate fuels. The product spectrum usually consists mainly of alkenes with about 10% of these being 1-alkenes. The hydrogen transfer activity of the zeolite is undesirable since it results either in rapid and early chain termination or in the formation of highly unsaturated molecules, which lead to the production of aromatics and coke species. The hydrogen transfer ability increases stepwise in orders of magnitude in the sequence: silicalite-1 (Si/Al > 1000) < ZSM-5 (Si/Al ca. 40) < SiO2 −Al2 O3 < mordenite < zeolite Y [16]. At high reaction temperatures (>350 ◦ C) cracking and aromatization reactions become more prominent, resulting in the formation of benzene, toluene and xylene. It has been shown that in the absence of free oxygen no aromatics were formed at temperatures as high as 450 ◦ C, suggesting that HZSM-5 can activate oxygen, which is then able to abstract hydrogen from aliphatics with the consequent formation of aromatics and water [17]. Table 1 shows a typical example of the product spectrum obtained in the oligomerization of propene over various acid catalysts including ZSM-5 (Si/Al = 35) and mordenite at 5 MPa and 220 ◦ C. ZSM-5 has excellent lifetime characteristics. At low temperatures good diesel and at high temperatures good gasoline yields are obtained. Steaming of the ZSM-5 results in the partial removal of framework aluminum and when carried out under optimum conditions may result in greater catalyst lifetimes. There is a strong inverse relationship

2859

between the zeolite crystal size and the catalyst utilization value (grams liquid product produced/gram catalyst) due to the increased pore diffusional resistances in larger crystals [18]. As referred to above, inertizing the external sites of the crystallites results in the formation of more linear products by reducing the influence of the nonshape-selective external surface activity [19]. Other zeolites have been investigated for their alkene oligomerization activity, and some of the results from these studies are shown in Table 1. In most instances their performance is usually determined by their pore geometry. The most common observation is that largepore zeolites such as HY allow the formation of bulky polynuclear aromatics, which lead to deactivation by coke formation, whereas one-dimensional zeolites such as mordenite rapidly deactivate due to pore-mouth blockage. Medium-pore molecular sieves such as SAPO-11, -31 and -34 have been reported to be alkene oligomerization catalysts, but their catalyst utilization value is significantly less than that of ZSM-5 [20]. None of these catalysts had lifetimes or selectivities that were in any way comparable to those of ZSM-5. Suppressing hydride transfer reactions can lead to improved activity, stability and selectivity in alkene oligomerization. A report on the performance of HZSM57 has indicated such an effect, and this is ascribed to its lobate-pore structure, which favors crosswise arrangement of alkenes, thus reducing hydride transfer [21]. In References see page 2863

Tab. 1 Typical product distributions (mass%) obtained from propene oligomerization over various acid catalysts at high conversions (>80%). For propene oligomer fractions refer to the carbon numbers: dimer (5–7), trimer (8–10), tetramer (11–13), pentamer (14–16) and hexamer+ (17+). For butene oligomers these correspond to multiples of four [18, 28–34]

Catalyst

ZSM-5 Mordenite Mordenite Solid phosphoric acid Solid phosphoric acid SMMa Ni–SMMa SiO2 –Al2 O3 Ni–SiO2 –Al2 O3 Amberlyst 15 Aluminum tungstophosphoric acid a SMM

Feed

Reaction T/ ◦ C

Reaction P/ MPa

Mass% Dimer

Trimer

Tetramer

Pentamer

Hexamer

C3 H6 C3 H6 C4 H8 C3 H6

220 300 250 200

5 5 5 3

18 28 48 9

30 31 40 65

27 23 7 16

14 10 5 10

11 8 – –

C4 H8

200

3

80

14

6

–

–

C3 H6 C3 H6 C3 H6 C3 H6 C3 H6 C3 H6

150 150 200 80 130 230

5 5 3 3 5 5

9 8 16 60 55 12

25 30 35 22 30 44

21 26 27 13 10 25

21 15 20 5 5 14

24 21 2 – – 5

= synthetic mica–montmorillonite.

2860

13.9 Oligomerization

the case of the oligomerization of alkenes over ZSM-22 [7], the reaction appears to occur predominantly at or near the outer surface of the zeolite crystallite. The higher degree of linearity compared with ZSM-5 is thought to be due to the presence of active sites located at the pore mouths which drive the shape-selective behavior of the zeolite. Mesoporous Catalysts Although the major focus of attention over the past 20 years has been on ZSM-5, recently there has been significant interest in other classes of catalysts. Eni Technologie [13] have announced the application of an amorphous mesoporous silica–alumina catalyst (MSA) as an alternative to the use of supported phosphoric acid. This catalyst is reported to possess good catalytic activity for the oligomerization of light alkenes to gasoline and jet fuel [22–26]. The gasoline boiling range fraction has RON values between 97 and 102 with low benzene contents and the kerosene boiling range has low smoke and freezing point values. In another study of the oligomerization of butene over a mesoporous aluminosilicate catalyst at 250 ◦ C and pressures in the range of 1.5–2 MPa, it was shown that catalysts with pore openings of approximately 3 nm can exhibit high selectivity and good stability with time for the production of branched dimers [27]. It was suggested that the behavior of these catalysts is related to the moderate strength and high dispersion of the acid sites in the mesoporous structure. 13.9.4.2

Other Acid Catalysts Alkene oligomerization has been evaluated using many other types of solid acid catalysts. Amorphous silica–alumina is active for the oligomerization of alkenes and, at about 200 ◦ C and 3 MPa, is particularly selective for the production of trimers and tetramers, the major dimer product formed during propene oligomerization being 3-methylpentene (Table 1). The product of propene oligomerization is mainly in the gasoline fraction. The ability to give high yields at relatively low reaction temperature reduces the possibility of coke formation and thus the catalyst has excellent lifetime characteristics with no indication of deactivation occurring after many hours onstream [29]. Macroreticular sulfonic acid resins, such as Amberlyst-15, produce excellent yields in the oligomerization of alkenes, although the catalyst is thermally unstable at temperatures higher than about 130 ◦ C [34]. 13.9.4.3

oligomerization uses homogeneous catalysis in which ethene is converted almost with 100% selectivity to 1-butene and n-butenes and then into octenes over a nickel complex [30]. Supported nickel catalysts are usually prepared by impregnation or ion exchange with supports such as zeolites, silica, alumina or silica–alumina. These catalysts are particularly excellent for alkene dimerization. The molecular weights of the products increase with increasing acid strength of the support phase. When the acid sites are stronger, a broad range of hydrocarbons are formed by oligomerization, cracking, isomerization and hydrogenation. The nature of the active site is not entirely clear. It is proposed that it is either a coordinatively unsaturated nickel(I) ion or metal hydride complex which is formed during reaction. This protonates the alkene, resulting in a linear alkylmetallo complex or an isoalkylmetallo complex. In the case of propene, further chain growth onto this complex results in the formation of predominantly methylpentene products. The acidity of the support will mainly affect the extent of secondary cracking reactions. Propene oligomerization leads mainly to the formation of dimers (Table 1), whereas ethene produces higher oligomers via an Eley–Rideal mechanism. The use of catalysts in which nickel is supported on SiO2 –Al2 O3 results in very high conversions at temperatures as low as 80 ◦ C, as opposed to the reaction on silica–alumina, which requires a temperature of at least 180 ◦ C [29]. Clay-type structures such as 2 : 1 layer aluminosilicates, e.g. synthetic mica–montmorillonite (SMM), in which aluminum is substituted for silicon in the tetrahedral layer, produce high yields in the oligomerization of propene at temperatures as low as 150 ◦ C [37]. Various metals such as nickel, cobalt and zinc may be incorporated into the catalyst by ion exchange or by incorporation into the octahedral layer of the framework [31]. Reduction of this framework nickel results in increased Brønsted acidity due to the formation of protons arising from the reduction of the divalent nickel [38]. Cobalt oxide-oncarbon catalysts produce high yields of dimers when ethene is the feed [39]. At 5 MPa and about 220 ◦ C, heteropo1y acids, such as aluminum tungstophosphoric acid, convert alkenes such as propene and butene into middle distillates with no aromatics formation [32]. Results of typical studies are shown in Table 1. 13.9.5

Catalysts Incorporating Metals Soluble nickel complexes are a most important class of molecular alkene oligomerization catalysts converting, for example, ethene into highly linear l-alkenes with even-numbered oligomers in the C4 –C20 range [35, 36]. One of the most important industrial applications of 13.9.4.4

Industrial Applications of Alkene Oligomerization Zeolite-Based Processes There has been limited application of the use of ZSM-5 for alkene oligomerization. Mobil’s MOGD (Mobil Olefin to Gasoline and Distillate) process is proposed 13.9.5.1

13.9.5 Industrial Applications of Alkene Oligomerization

to be able to produce, after hydrogenation, mainly low-branched alkanes and scarcely any aromatics [40]. The high selectivity to alkenes in the diesel mode of operation is due to the restricted transition shape selectivity, which both favors alkene formation and inhibits the formation of typical cyclic coke precursors. These factors, together with a high reaction pressure (typically 5 MPa) and moderate reaction temperature (200–220 ◦ C), are conducive to the formation of diesel fractions. In a typical proposed commercial application, ZSM-5 has an Si/Al ratio of about 35 and is extruded with 35 wt.% alumina. In Mossel Bay, South Africa, S¨ud-Chemie’s COD Process is being used for the conversion of olefins to diesel and gasoline in a ca. 500 000 tons per year plant [41]. Typical operating conditions are temperatures of 200–250 ◦ C, pressures of 6 MPa and a fresh feed WHSV of 0.5 h−1 with a gasoline recycle of 0.3 wt./wt. The feed composition is typically 82% alkenes, 15% alkanes, 1.5% aromatics and 1.8% oxygenates, and the liquid fuel yields, based on alkenes, are 97%. When operated in distillate mode, the yields of distillate and gasoline were 78 and 19%, respectively, with alkene conversions per pass varying from 99% for propene to 53% for C7 alkenes. In gasoline mode the process is capable of producing product with a research octane number (RON) of ca. 92 and motor octane number (MON) of ca. 80 and in diesel mode the product after hydrogenation can have a cetane number of ca. 56 with tertiary. The smallest tertiary alcohol, tertbutanol (2-methyl-2-propanol), is unreactive under typical etherification conditions [59]. Among primary alcohols, the etherification rate increases with increasing carbon number within the C1 –C4 alcohols on ion-exchange resins [65]. The equilibrium constants are also affected: in the etherification of isobutene with various alcohols it has been shown that the reaction equilibrium constant decreases with increasing size of the alcohol [27, 60]. An increase in the molecular size of the alkene sharply decreases the reaction rate and equilibrium constant of the etherification reaction [66]. Polyfunctional Ethers Polyols are also reactive in etherification with isoalkenes. Reactions of 1,2-ethanediol (ethylene glycol) and 1,3-propanediol (propylene glycol) and also 1,2,3propanetriol (glycerol) have been studied mainly with isobutene [67–72]. When multiple hydroxyl groups are available in the alcohol, various products are formed depending on the extent of the etherification and product isomers also exist. The product distribution can be controlled by adjusting the process conditions such as temperature and feed composition [73, 74]. The glycol ethers have been studied for solvent applications in paints, inks and coatings and for trapping agents [67]. tert-Butyl ethers of glycerol have mainly been studied for fuel applications. The octane numbers for the ether product mixture are 112–128 (blending research octane number, BRON) and 91–99 (blending motor octane number, BMON) [75], and these values indicate that in this sense they are suitable to be used as gasoline components. A study of tert-butyl ethers of glycerol as biodiesel components has also been performed [76]. 13.10.4.2

13.10.5

Catalysts Acidic Ion-Exchange Resins The etherification reaction is an acid-catalyzed reaction. The catalysts used in the commercial etherification processes are strongly acidic cation-exchange resins which are sulfonated organic polymers. The most important catalyst support is a copolymer of styrene and divinylbenzene (DVB), the structure of which is presented schematically in Fig. 5. 13.10.5.1

H2 C CH

H2 C CH

CH SO3− H+

SO3− H+

H2 C

SO3− H+ C H

Fig. 5

Acidic ion-exchange resin.

The amount of DVB determines directly the degree of crosslinking and the rigidity of the structure. That is, the more DVB in the catalyst, the more rigid is the resin and its structure. Resins with a DVB content less than 8 wt.% are of the gel type without permanent porosity. These types of resins function only in the presence of polar components that swell the resin structure. Resins with a DVB content of 12 wt.% or more have permanent macroporosity. Still, these materials also have a microporous gel phase consisting of gel-type microspheres [77]. The DVB content specifies the specific surface area and the pore size distribution of the macroporous resin [78]. The ionexchange resins used as catalysts in the etherification usually have a macroporous structure. Macroporous ion-exchange resins consist of uniformsized gel-type microspheres and pores that are formed between them [79]. Two types of mass transfer are present as the reactant molecules diffuse through the pores to the surfaces of the microspheres and also inside the microspheres to the polymer matrix. The structure of the resin is presented in Fig. 6. The polymeric matrix is insoluble in water and organic solvents. The diameter of the polymer beads lies in the range 0.3–1.3 mm. The distribution of bead sizes can be influenced by the polymerization parameters. Abrams and Millar have reviewed the origin and development of macroporous ion-exchange resins [80]. The catalytic activity of the support is obtained by treating the polymer matrix with sulfuric acid. If all the benzene rings are monosulfonated, an acid capacity in the order of 5 mmol g−1 cat should be obtained. The sulfonic acid sites are situated in the easily accessible macropores, but also inside the gel-type microspheres [81]. In some studies it has been concluded that the sulfonic acid groups within the microparticles are catalytically more active than those on the surface of the microparticles [82, 83]. The sulfonic acid groups are linked by a network of hydrogen bonds [84]. Polar components break into the strong hydrogen-bonded network and cause the protons of the acid groups to be dissociated. This type of network of acid groups solvated by the hydrogen-bonded alcohol is catalytically much less effective than the original

13.10.5 Catalysts

Macroporous bead Fig. 6

Microporous bead

Structure of an ion-exchange resin.

are shown. The properties are those reported by the producers. Over the years, few comparisons of commercial ionexchange resins in MTBE synthesis have been published. Panneman and Beenackers [91] tested several commercial macroporous resins. Unfortunately, they used a feed containing 2 mol% 2-methylpropene and 98 mol% methanol, which is very far from the composition of the real MTBE feed. Parra et al. [92] compared the performances of 12 ion-exchange resins in the liquid-phase synthesis of MTBE (Fig. 7). The results show that the resins with high acid capacities were the most active. In fact, this parameter had the greatest influence on the catalytic performance, whereas other properties such as the specific surface area and the average pore diameter had hardly any effect.

network of acid groups themselves, because the proton donor–acceptor strength of the network is reduced [84]. The polarity of the reaction medium influences the swelling of ion-exchange resins and thus the number of accessible sites [85, 86]. The strongly acidic cation-exchange resins are thermally stable up to temperatures of 390–400 K. Stability up to 420 K has been reported for some resins [87]. Above these temperatures, a loss of catalytic activity is observed, due to the hydrolysis of active sulfonic acid groups. Thermal stability generally decreases as crosslinking increases [88]. The major producers of ion-exchange resins are Sybron Chemicals [89] (Lewatit resins), Dow Chemical [90] (Dowex resins), Purolite [88] (Purolite resins) and Rohm and Haas [87] (Amberlyst resins). In Table 4 [87–90], the main properties of the commercial resins recommended by the producers for etherification

References see page 2879

r /mol (h meq)−1

0.025 0.02 0.015 0.01 0.005

9 17

24 Pu

ro

lit e

C T-

5 5/ 17

T-

ro

lit

e

C

ro Pu Pu

MTBE formation rates with different catalysts [92].

28

17

1 17 lit e

C T-

9 16 Pu

ro

lit e

C T-

5 16 Pu

ro

lit e

C T-

1 lit e

Pu

ro

lit e ro Pu

C T-

15

2 C T-

M

-3

35 ex ow D

Am

be

rly

st

15

1 Am

be

rly

st

50 -1 C

O ye r

Ba

Ba

ye r

K2

63

1

0

Fig. 7

2871

2872 Tab. 4

13.10 Etherification Physical properties of some commercial ion-exchange resins [87–90]

Catalyst

Amberlyst 15 W Amberlyst 16 W Amberlyst 35 W Dowex M-31 Lewatit K2621 Lewatit K2624 Lewatit K2431 Purolite CT-175

DVB content/wt.%

Ion-exchange capacity/(meq H+ L−1 )

20

1.7 1.7 1.9 1.85 1.4 1.4 1.2 1.8

In addition to the ion-exchange resins described above, some modifications have been tested in ether synthesis. Kunz et al. [93] developed a polymer–ceramic composite catalyst to replace the conventional resin beads in the catalytic distillation for MTBE production. A crosslinked styrene–divinylbenzene polymer was incorporated into a macroporous ceramic support. After the sulfonation, a mechanically stable acid catalyst was obtained. This catalyst could be prepared in various shapes such as Raschig rings, honeycombs and corrugated packings. Another possibility is to add the sulfonic acid groups on the surface of polymer fibers. The Smopex-101 catalyst developed by Smoptech is prepared by grafting polyethene fibers with styrene and subsequently sulfonating the grafted fiber with chlorosulfonic acid [57]. To the best of our knowledge, these catalysts are not used in commercial etherification units. Zeolites and Modified Zeolites In addition to ion-exchange resins, several zeolites have been studied as possible catalyst candidates for etherification reactions [7]. Zeolites such as H-ZSM-5, H-Y, H-Beta, partially alkali metal-exchanged H-Y and dealuminated zeolites have been tested [94–97]. The main advantage of the zeolite catalysts is their well-known better thermal stability compared with ion-exchange resins. However, at the temperatures of the etherification processes the resins have been proven to be stable enough (see Table 4). Improved ether selectivities have been reported with zeolites but the activities have remained at a lower level than with the resins [94, 95]. The reaction tests have typically been carried out in the gas phase, which means that the total ether productivity does not reach that obtained with resins in the liquid phase. Comparisons of the catalysts have been carried out at temperatures between 333 and 353 K [95] and between 303 and 393 K [94]. Even in the gas phase, the resin Amberlyst 15 gave the best performance in MTBE synthesis. In reactive distillation studies of ETBE synthesis, zeolites have been reported to show activities comparable to or higher than 13.10.5.2

Particle size range/mm 0.355–1.18 0.355–1.18 0.355–1.18 0.40–0.65 0.4–1.3 0.4–1.3 0.5–1.3 0.4–1.2

Tmax /K

Pore diameter/nm

Specific surface area/m2 g−1

393 403 423 – 393 403 403 418

30 25 30 22 65 65 40 65

53 30 50 30 40 40 25 30

those of ion-exchange resin catalysts. Zeolites with high aluminum content were more selective, and the highest selectivity (99%) was obtained for zeolite Beta (Si/Al = 8) at 363 ◦ C [98]. Even though some improvements have been seen in using zeolites in etherification synthesis, it is not expected that zeolites will replace resins as catalysts in commercial etherification processes. Other Acidic Catalysts Over the years, some other solid and molecular catalysts have been tested in laboratory-scale etherification. These include clays [99–101], heteropolyacids [102–105], Ti–silicalite-1 [106], fluorinated zeolitebased catalysts [107], sulfided zirconia [108–111] and p-toluenesulfonic acid [61]. The catalysts have been used in both gas- and liquid-phase syntheses. So far these catalysts have remained in the research stage. 13.10.5.3

Catalyst Deactivation The acid sites of the cation-exchange resins are easily poisoned by basic compounds which can be present in the feeds or be carried to the reactor through disturbances in preceding process units. Both sodium and ammonium ions are strongly bound into the acid groups, and these cause shrinkage of the catalyst. The activity of the deactivated resins was observed to be independent of the nature of the deactivating cation, but to be very strongly dependent on the amount of unchanged hydrogen ions left on the catalyst [15]. The ether selectivity remained at the same level even when the hydrogen ion content was reduced to less than half of the original value. Nitrogen-containing compounds such as ethanenitrile (acetonitrile) react on the sulfonic acid resin forming ethanamide (acetamide). This compound has also been detected on the used MTBE catalyst. Further, it has been shown that there is a direct relationship between the ethanamide yield and the loss of catalyst activity [112]. Important catalyst poisons are dienes, which are dealt with in more detail later. As indicated in Section 13.10.3.2, 13.10.5.4

13.10.6 Raw Materials and Products

they react under the etherification reaction conditions forming gum, which can block the pores of the catalyst. Also, any excessive water in the process feed streams will produce alcohols as side products, viz. tertiary and secondary butyl alcohols (TBA, 2-methyl-2-propanol and SBA, 2-butanol) from the C4 hydrocarbons and tertiary and secondary pentyl alcohols (2-methyl-2-butanol and 2-pentanol) from C5 hydrocarbons. Water is more strongly bound on the resins than the alcohols and thus reduces the catalytic activity [113]. If there are any dissolved metal ions (e.g. Fe2+ /Fe3+ ) in the feed streams, they will replace the active H+ from the catalyst and thus deactivate it. The metal ions deactivate the catalyst permanently. In principle, the resin catalysts deactivated by exchanged ions can be regenerated by treating the resin with concentrated sulfuric acid. However, the amounts needed are typically so huge that in practice the regeneration of the catalysts is not carried out. 13.10.6

Raw Materials and Products

The etherification processes use two basic raw material streams: • the hydrocarbon stream containing reactive isobutene (for MTBE and ETBE) or isopentenes (for TAME and TAEE) and • the alcohol stream, either methanol (for MTBE and TAME) or ethanol (for ETBE and TAEE). The product streams are the ether product and the hydrocarbon stream. The hydrocarbon stream contains inert hydrocarbons (such as alkanes), unconverted isoalkenes and non-reactive alkenes. Additionally, a waste water stream containing dissolved alcohols may be created if the alcohol is removed from the other products using water washing and no alcohol recovery unit is installed. Tab. 5

2-Methylpropene 2-Methylpropane n-Butane 1-Butene 2-Butenes

bC 4 cC 4

Typical Raw Materials Feed streams are normally obtained in a refinery from a steam cracker (SC) or a fluid catalytic cracking (FCC) unit. The isobutene content of the C4 stream from a catalytic cracking unit can be increased in a skeletal isomerization unit [114–117]. Alternatively, so-called field butanes via isomerization and dehydrogenation processes create a feed stream for MTBE and ETBE production. The field butane stream is first pretreated to remove heavier compounds, and then n-butane is isomerized to isobutane and finally dehydrogenated to a mixture of isobutene and isobutane [115]. Isobutene can also be obtained as a very concentrated stream by dehydrating 2-methyl-2-propanol (tert-butanol, TBA), which is obtained as a by-product from 1,2epoxypropane (propylene oxide) production [115, 118]. The composition ranges of the C4 sources are shown in Table 5 [114, 117]. Corresponding composition ranges of C5 sources can be found in Refs. [114, 117]. Methanol and ethanol are normally obtained from outside the ether production plant (see Chapter 13.13 for methanol synthesis). Ethanol is most commonly produced by fermentation from agricultural raw materials containing sugars, but small amounts are produced synthetically from ethene [116, 119]. 13.10.6.1

Feed Stream Treatment Methanol and ethanol must have a low water content to prevent side product formation. Depending on the alcohol production processes, some minor amounts of other impurities exist. The feed specifications for the alcohols depend on the process configuration [117]. Basic nitrogen compounds tend to neutralize the acidic sites of the resin catalyst, hence they must be removed upstream of the etherification reactor. 13.10.6.2

References see page 2879

Compositions of hydrocarbon feeds for the MTBE process [114, 117]

Component

a Fluid

2873

FCCa /wt.%

C4 raffinateb /wt.%/vol.%

Isomerization and dehydrogenation of field butane/wt.%

Skeletal isomerizationc /wt.%

15 35 11 13 26

45/44–49 2/2–3 5/6–8 28/24–28 20/19–21

48 52 1 1 0

17 6 44 10 23

catalytic cracking of gas oil. fraction from steam cracking (SC) that has been treated by butadiene extraction and selective hydrogenation. fraction from SC treated with skeletal isomerization.

2874

13.10 Etherification

Some of the steam-cracked C4 streams contain such high concentrations of 1,3-butadiene that these streams are directed to butadiene recovery units. Ethanenitrile (acetonitrile) is a commonly used solvent in these units and often some traces of it might be present in the streams leaving the butadiene recovery units. These streams must be treated by water washing and an adsorption bed to remove the ethanenitrile [120]. Propanenitrile that exists in the heavier C5 streams does not wash off so easily with water, and it is often found to be uneconomical to remove it from the TAME and TAEE feed streams [120]. Fortunately, the conversion of propanenitrile per reactor pass is low compared with that of ethanenitrile, and therefore its deactivating effect is not so significant provided that propanenitrile is not allowed to accumulate in the process. This is avoided if the alcohol is not recycled. Both ethanenitrile and propanenitrile form azeotropes with methanol and are accumulated in methanol recycle streams [115]. Dienes must be removed by selective hydrogenation to prolong the resin catalyst lifetime. Dienes form alkenic ethers that are colored and they can also polymerize during storage [115, 120]. Also, the members of the family of vinylacetylenes (e.g. but-3-en-1-yne, 3-methylbut-3en-1-yne and pent-3-en-1-yne) are harmful. Selective hydrogenation typically uses palladium-based catalysts. It is possible to combine the selective hydrogenation with the etherification process so that no gas–liquid separation is required before the etherification. The key feature of the selective hydrogenation is that as little alkenes are lost to alkanes as possible. One way to protect the catalyst is to use a guard bed that adsorbs some impurities. This guard bed is normally smaller than the actual reactors and operates at lower temperatures in order to prevent any noticeable amount of reactions on it. The most reactive alkadienes and alkenynes have a possibility of polymerizing on this bed. To prevent oligomerization of the alkenes the guard bed reactor must also have the alcohol on its feed. Basic nitrogen compounds neutralize the sulfonic acid groups of the guard bed, leading to a longer lifetime of the catalyst in the etherification reactors [120]. In addition to ion-exchange resins, some adsorption beds such as activated alumina or activated carbon are used to remove impurities. Ether Products MTBE is produced either in the refinery or in an onpurpose plant. The MTBE from an on-purpose plant can be sold as a pure chemical-grade product, but it is most commonly used as a gasoline component. On the other hand, if MTBE is made in a refinery complex it is not produced as a pure component, and the product 13.10.6.3

is used as a blending component in the gasoline pool. There are some limits on the amount of alcohols allowed in the ether products, because the blending Reid vapor pressures of methanol and ethanol are higher than those of the corresponding ethers (see Table 1) [121]. If some TBA is formed in the process, it is left in the ether product, as TBA is also a good gasoline component [115]. Hydrocarbon Products The C4 hydrocarbon stream leaving the MTBE or ETBE unit normally contains inert n-butane and isobutane in addition to unreacted isobutene and the linear butenes. This stream has to be treated according to its usage. It can be used in alkylation where isobutane reacts with the butenes to form alkylate (see Chapter 13.8). The most commonly used alkylation process is based on hydrogen fluoride. This process does not tolerate large amounts of oxygenates, i.e. oxygen-containing molecules such as alcohols and ethers. That is why in many etherification processes the oxygenate specification is met by using a water wash. The hydrocarbon stream is often processed in an oxygenate removal column, prior to directing it to the alkylation unit. In this oxygenate removal column the alcohols and light ethers (such as dimethyl ether as a reaction by-product) are obtained on the column top as these compounds form minimum boiling point azeotropes with the hydrocarbons (see the process scheme in Fig. 8). When the etherification process is based on the use of field butane (see Section 13.10.6.1), the main constituents of the C4 stream leaving are isobutane and unreacted isobutene. Often this stream is recycled back to the dehydrogenation stage [115]. 13.10.6.4

13.10.7

Processes

The etherification reaction using an acidic ion-exchange catalyst is very selective when the temperature control is done properly and the methanol to isobutene molar ratio is near the stoichiometric value (1.1–1.2). A slight alcohol excess is used to prevent dimerization and oligomerization of isobutene. Operating Conditions Practically all of the etherification processes in industrial use are carried out in the liquid phase and in the temperature range from room temperature to 390 K. The pressure in the process is kept just above the bubble point pressure of the reacting mixture (typically 0.7–2.0 MPa). Typically the liquid hourly space velocity (LHSV, i.e. total liquid volumetric inlet flow at 298 K divided by the catalyst volume) is from 4 to 6 h−1 in the MTBE process [117]. 13.10.7.1

13.10.7 Processes

Under the MTBE reaction conditions, compounds other than isobutene and methanol are practically inert, which means that the C3 and C5 compounds, left in the feed stream, do not react. As the etherification reaction is exothermic, the temperature must be controlled by proper means. However, for dilute isobutene (up to 20%) feeds such as the C4 fraction from an FCC unit, the temperature rise is not very high. Three types of reactors are normally used, namely fixed-bed reactors, isothermal tubular reactors and the fixed-bed reactors operated in boiling-point conditions. The choice of the reactor type depends on the isobutene content of the feed stream and on the usage of recycle streams [115]. 13.10.7.2

Processes for MTBE Production

13.10.7.2.1 Conventional MTBE Units The thermodynamic equilibrium limitation of the etherification, as discussed in Section 13.10.3.3, has guided the MTBE process design (Fig. 8). These conventional MTBE units have two or three reactors in series. The temperature of the first reactor is the highest to obtain a high reaction rate. Then the temperature is dropped in the following reactors to obtain higher total yields with good selectivity [115–117]. The catalyst lifetime is normally several years. Temperature control is obtained by using a reactor where the catalyst is packed in multiple tubes and the cooling medium flows around these tubes. When the reactors are of the fixed-bed type, heat exchangers between the reactors are used to adjust the inlet temperature of each reactor. As Fig. 8 shows, it is even possible to lead a recycle stream from the outlet of the first reactor to its inlet through a heat exchanger. This also helps in controlling the reactor temperature. One option is to utilize the fixed-bed reactors in the boiling-point mode. This means that there should be a separate gas outlet from the reactor and liquid level control. When the boiling-point mode is used, pressure control determines the operating temperature. The ion-exchange resin beads are small in size (typically from 0.3 to 1.3 mm in diameter, see Table 4), requiring a specific catalyst support structure [115]. The conventional processes described above achieve up to 96% isobutene conversion with very high selectivity.

Reactive Distillation in MTBE Production To overcome the thermodynamic equilibrium limitation, reactive distillation was developed for ether production during the 1980s [115]. Reactive distillation (also called catalytic distillation) has been treated in detail by Miracca et al. in Chapter 10.6 of this Handbook and by Sundmacher and Kienle [122]. A good introduction to 13.10.7.2.2

2875

reactive distillation for ether production is also given by J¨arvelin [115] and Kolah et al. [123]. When reactive distillation is applied for MTBE production, there are normally one to three reactors before the reactive distillation column. The process scheme is almost identical with the conventional MTBE process shown in Fig. 8. Especially the methanol recovery and oxygenate removal parts of the processes are very similar. The greatest difference is that part of the catalyst is placed inside the MTBE separation column, making this column a reactive distillation column. The catalyst can be packed into the reactive distillation column in the liquid phase on woven bags or belts, integrated into structured packing or in forms of catalytically active random packings. The investment cost of reactive distillation is higher than that of the conventional MTBE process, and higher isobutene conversion can be obtained. With reactive distillation in the MTBE process it is possible to obtain isobutene conversions over 99%. Processes for ETBE Production ETBE can be produced in a very similar way as described above for the MTBE production [124]. The conversion of isobutene is about 10% lower than in the MTBE processes. The ethanol used in the process must be free from water, otherwise TBA is formed. This is important for such process configurations where water wash is used to extract ethanol from the hydrocarbons. The recovered ethanol must then be made anhydrous, which cannot be achieved by ordinary distillation due to the azeotrope formation between water and ethanol. Therefore, the recycled ethanol must be purified using, for example, membrane units or adsorption-based separation methods [115]. 13.10.7.3

Conversion of MTBE Plants for ETBE Production The bio-directive and support for renewable energy sources, especially in Europe [6], have led to the situation where many of the MTBE plants have been converted into ETBE plants. Methanol is replaced with ethanol, which is produced by fermentation of some agricultural sugarcontaining products or products that can be hydrolyzed to sugars. Tsai et al. [125] listed the following requirements for necessary changes when converting an MTBE unit into an ETBE unit: 13.10.7.4

• If the MTBE unit has only one reactor, another reactor should be added in order to have reasonably high conversion (92%) for isobutene. • A water separation unit should be added to the alcohol recycle stream so that excess TBA formation is prevented. References see page 2879

2876

13.10 Etherification

Oxygenates and C4

Methanol recycle Methanol C4 feed

Water wash column

MTBE column

Water C4s to alkylation

Reactor 1 Reactor 2 MTBE

Fig. 8

Methanol recovery column

Oxygenates removal column

Conventional MTBE process with alcohol recovery and oxygenates removal system.

Wells and Buckland [126] gave the following points to consider when converting an MTBE unit into an ETBE unit: • The alcohol mass flow into the process should be increased as the molar mass of ethanol is higher than that of methanol. This means that a larger pump might be needed for the alcohol feed. The piping and flow measuring equipment should also be checked. • The temperature on the column bottoms should be increased as the normal boiling point of ETBE is higher than that of MTBE. This means that the column reboiler duties should be checked. • Because the molar mass of ethanol is higher than that of methanol, the mass flow of the ether will be increased provided that the hydrocarbon feed rate is kept constant. Therefore, the cooling capacity of the product coolers must also be checked. As discussed earlier, the bioethanol may contain small amounts of components that are harmful in the ETBE process. This means that the ethanol requires suitable purification prior to feeding it to the ETBE process. If ethanol is obtained as a dilute aqueous solution, the distillation and dehydration system costs should be evaluated. If the ethanol usage is small, it might be more economical to buy anhydrous ethanol. Processes for TAME Production TAME can be produced from a C5 alkene-containing stream. The conventional process is very similar to that shown for MTBE production in Fig. 8. The hydrocarbon stream is often obtained from the FCC unit as an FCC light gasoline fraction. Methanol is mixed with the hydrocarbons coming from the selective hydrogenation and led into the reactor train consisting typically two or more reactors. After the last 13.10.7.5

reactor, the effluent is distilled to separate the TAME product as the column bottoms. The distillate from this column is treated in a similar way as in the MTBE process (Fig. 8) methanol recovery and oxygenate removal unit. The hydrocarbon feed to the process can contain also reactive C6 and C7 isoalkenes that convert into heavier ethers with the methanol. These improve the total ether yield of the process. Reactive distillation can also be used in the TAME process to improve the alkene conversion. Unfortunately, the catalyst deactivation is a much greater problem in the TAME process than in the MTBE process. Catalyst exchange in a reactive distillation column is not a trivial task, and this makes reactive distillation unattractive for the TAME process. On the other hand, reactors and distillation can be combined in a more sophisticated way, utilizing so-called side reactors connected to the distillation column. The side reactors can also be called reactive pump-arounds or the whole system as reaction with distillation. The side-reactor concept for the TAME process is similar to that shown in Fig. 9 [127]. This process, called NExTAME , was developed by Neste Oil [128]. The main advantage of this system is the reduced number of distillation columns as no methanol recovery is required. Methanol forms an azeotrope with the C4 components present. This azeotrope has a maximum concentration near the column top. A side stream is taken out from the distillation column from this location, and this methanol–C4 azeotrope is recycled back to the reactor feed. Thus the two columns from the methanol recovery are eliminated. The ether is obtained as a bottom product with unreacted hydrocarbons. TAEE can in principle be produced in a TAME process when methanol is replaced with ethanol. However, it can be expected that some process changes will be required.

13.10.8 Diesel Ethers

FCC light gasoline and C4 fraction Methanol

C4 components and methanol as a sidestream from column 2 C3, C4 and methanol C and C components 5

C3 stream to fuel

and methanol as a sidestream from column 1

Reactor 1

Fig. 9

6

2877

Reactor 2

Ethers and C4s to unreacted C5 + alkylation hydrocarbons Distillation 2 Distillation 1

NExETHERS process [128, 130] shown for methanol. Ethanol can be used instead of methanol.

Multipurpose Processes for Production of Ethers Based on the market situation for different ethers and on the availability and price of methanol and ethanol, it might be advantageous to change the etherification process rapidly from MTBE to ETBE or vice versa. As an alternative for such a process change, G´omez et al. [129] reported the simultaneous production of MTBE and ETBE using both methanol and ethanol as the process alcohol feed. They suggested that with catalyst particles of larger diameter methanol and ethanol react with similar rates. One of the most versatile process configurations is NExETHERS by Neste Oil [128, 130]. This can be used with a C4 hydrocarbon feed with methanol or ethanol to obtain MTBE or ETBE. With the use of the FCC light gasoline and methanol a TAME process can be obtained and, in addition, a TAEE configuration is possible. The principle is shown in Fig. 9. The process consists of two adiabatic fixed-bed reactors with acidic ion-exchange resin catalysts and two distillation columns. Different side streams are taken from the distillation columns and recycled into suitable locations. Alcohol recovery columns are not required. In ETBE production this process reaches 95% conversion for isobutene, and the purity of the ETBE is around 99.5%. 13.10.7.6

properties as a diesel fuel. It has a high cetane number (55–60) and does not contain any sulfur or metal compounds. Furthermore, due to the absence of carbon–carbon bonds, DME does not form any soot or other particulate emissions and has a potential to reduce NOx emissions [131]. In addition, it is not corrosive to any metal and not harmful to the human body. Because DME can be produced independently of crude oil fractions for example from natural gas and biomass, it is currently referred to as the ‘‘fuel of the 21st century’’ [132]. The drawback of DME is its low normal boiling point (248 K), but it can be easily liquefied under pressure close to 6 bar (1 bar = 105 Pa). It has to be handled and stored as LPG with the same safety precautions. With an energy density of about half that of diesel oil, it needs large on-board tanks for an equivalent driving range. The properties of DME are given in Table 6 [133].

Tab. 6

Properties of DME [133]

Property

DME

13.10.8

Diesel Ethers

Diesel ethers are formed by dehydration of an alcohol, especially linear alcohols. The simplest of these is DME (dimethyl ether, methoxymethane), which is considered in more detail. Another candidate for diesel usage is DNPE (di-n-pentyl ether, pentoxypentane). These have good cetane numbers and burn cleanly. 13.10.8.1

Dimethyl Ether

13.10.8.1.1 Properties of Dimethyl Ether In the early 1990s, it was discovered that dimethyl ether has attractive

Formula Molar mass/g mol−1 Normal boiling point/K Specific density, gas vs. air Liquid density at 298 K/kg m−3 Cetane number Vapor pressure at 298 K/kPa Lower heating value/kJ kg−1 Lower explosion limit in air/vol.% Lower flammability limit in air/vol.% Auto ignition temperature/K Critical pressure/kPa Critical temperature/K

References see page 2879

C2 H6 O 46.07 248.25 1.59 668 55–60 510 28430 3.4 1.7 508 5269 400

2878

13.10 Etherification

DME is formed in the dehydration reaction of two methanol molecules releasing one water molecule: −  2CH3 OH −  −− − − CH3 −O−CH3 + H2 O

(2)

This reaction is slightly exothermic, r H 0 (gas phase) = −24 kJ mol−1 , and limited by thermodynamic equilibrium only at higher temperatures. At 473 K the equilibrium methanol conversion with the pure methanol feed is 92% [134]. If methanol or the synthesis gas used for the methanol synthesis is produced from biomass, DME can be considered as biodimethyl ether [6]. Processes for Dimethyl Ether Production Traditionally, linear ethers have been produced by catalytic dehydration of alcohols or by reaction of alkyl halides with alkoxides. The dehydration can be carried out either in the liquid phase using acids, particularly sulfuric acid, or metal and non-metal halides as catalysts or in the gas phase over solid acidic catalysts such as alumina or modified alumina. However, the main source of these ethers has been as by-products of the production of the corresponding alcohols by hydration of alkenes over acid catalysts [135]. Until about 1975, DME was obtained as a by-product in the high-pressure production of methanol [136]. However, by 1980 almost all of these processes had been replaced by low-pressure processes in which only very small amounts of DME are formed. As a consequence, specific processes have been developed for DME production. DME can be produced in two ways: the so-called indirect route is a two-step process in which methanol is first synthesized and then separately converted to DME; the direct route involves the conversion of synthesis gas to DME in one step [132]. Methanol synthesis is strongly limited by thermodynamic equilibrium, as discussed in detail in Chapter 13.13. Even though it is evident that methanol is the 13.10.8.1.2

intermediate also in the direct synthesis, its immediate conversion to DME leads to higher once-through yields. The direct DME synthesis is highly exothermic and can proceed via two reactions: −−  3CO + 3H2  −− − − CH3 −O−CH3 + CO2

(3)

−  2CO + 4H2 −  −− − − CH3 −O–CH3 + H2 O

(4)

Reaction (3) is in fact a combination of the following steps: −  2CO + 4H2 −  −− − − 2CH3 OH −−  2CH3 OH  −− − − CH3 −O−CH3 + H2 O −  CO + H2 O −  −− − − CO2 + H2

Purge gas to fuel Methanol from distillation

Boiler feed water Steam generator First reactor Synthesis gas as make-up gas Second reactor

Fig. 10

Haldor Topsøe DME process [139].

(6) (7)

The difference between reactions (3) and (4) is that, instead of CO2 , the extra oxygen atoms come out as water, which means that no water gas shift reaction [Eq. (7)] takes place. The catalysts for the direct synthesis must be bifunctional combining metallic sites characteristic for methanol synthesis and acidic sites characteristic for dehydration. The two functions can be attained either by mechanically mixing the Cu/Zn methanol catalyst and an acidic catalyst such as modified alumina or a zeolite or by supporting the metallic sites of the methanol catalyst on an acidic support [137, 138]. Processes developed for DME production include those from Haldor Topsøe (Denmark) [139] and JFE (Japan) [140]. The former process is presented in Fig. 10. The process has two reactors. The first reactor is a cooled reactor where the highly exothermic methanol synthesis takes place. The less exothermic DME formation is carried out in the second fixed-bed reactor operated adiabatically. The advantage of this process is that both reactors can be operated separately under their optimal conditions. In the JFE process, the synthesis gas is fed to a slurry

Compressor

Steam

(5)

Flash drum

DME, methanol and water to distillation

References

2879

reactor in which a finely powdered catalyst mixed to an inert high boiling point oil converts it directly to DME. The overall reaction (3) is very exothermic, and the temperature can be controlled through intensive mixing of the liquid phase. The key factor in both processes is the proprietary bifunctional catalysts. The purity required for the DME product has a strong influence on the production costs. Clear savings can be reached if fuel-grade DME, that is, DME containing some amounts of methanol, water and other oxygenates, is produced. Fuel-grade DME can also be used for power generation. In 2004, it was announced that the first large-scale DME plant based on the technology developed by Haldor Topsøe for the dehydration of methanol will be constructed at Bandar Assaluyeh, Iran, with a capacity of 800 000 t a−1 . The implementation of this project will increase worldwide DME production by more than 200% [141].

Some surplus isobutene from previous MTBE production can be utilized, for example, by dimerization followed by hydrogenation to produce isooctane, the so-called synthetic alkylate [146, 147]. However, methanol is much cheaper than ethanol and that may still keep MTBE production volumes high for a long time, for example in Asia. The use of tetraethyllead in gasoline to improve fuel octane number is coming to the end in many Asian, South American and African countries and octane enhancers have to be used to keep the gasoline quality at the required level. If methanol usage in MTBE and TAME declines, then methanol might find a new application in DME production, as DME seems to be a very promising new diesel fuel. There are several advantages in the use of ion-exchange resins as catalysts in etherification processes. At present no competing alternative is in sight.

Di-n-Pentyl Ether Longer chain linear ethers such as DNPE (di-n-pentyl ether, pentoxypentane) are interesting components for diesel fuels. DNPE is formed in a dehydration reaction between two 1-pentanol molecules:

References

13.10.8.2

−  2C5 H11 OH −  −− − − C5 H11 −O−C5 H11 + H2 O

(8)

DNPE has a normal boiling point of 460 K and a cetane number of 109 [142]. With respect to its viscosity and density, it behaves as a light diesel fuel. Ion-exchange resins have been suggested as catalysts for the reaction to replace sulfuric acid and some kinetic studies have been performed [143]. The use of resins is restricted by their low thermal stability. Resins start to deactivate at temperatures above 420 K and, to achieve moderate conversions in the reaction, the temperature should be higher. New resins with better thermal stability have been developed. It has been suggested that gel resins are more active and selective than the microporous ones in the formation of DNPE [144]. 13.10.9

The Future

Due to the legislative restrictions concerning the use of MTBE in gasoline, especially in the USA [4], the production of MTBE is declining. The environmental issues and different legislative actions and directives [6] to use bio-component-based fuels are especially increasing the use of ethanol as such or in the form of ETBE or even TAEE [145, 146]. This means that many MTBE units have already been converted to ETBE units, and this trend is expected to continue. Similarly, TAME processes might be converted into TAEE units in the near future.

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120. I. Noopila, Master’s Thesis, Helsinki University of Technology, Espoo, 2000. 121. M. A. Ali, H. Hamid, in Handbook of MTBE and Other Gasoline Oxygenates, H. Hamid, M. A. Ali (Eds.), Marcel Dekker, New York, 2004, pp. 37–64. 122. K. Sundmacher, A. Kienle (Eds.), Reactive Distillation: Status and Future Directions, Wiley-VCH, Weinheim, 2003, 308 pp. 123. A. H. Kolah, L. K. Rihko-Struckmann, K. Sundmacher, in Handbook of MTBE and Other Gasoline Oxygenates, H. Hamid, M. A. Ali (Eds.), Marcel Dekker, New York, 2004, pp. 159–202. 124. C. R. Marston, Fuel Reformulation 1994, 4(4), 42. 125. M. J. Tsai, R. Kolodziej, D. Ching, Hydrocarbon Process. 2002, 81(2), 73. 126. W. J. Wells, M. D. Buckland, Presented at the 1996 World Conference of Refinery Processing and Reformulated Fuels, 1996. 127. J. Ignatius, H. J¨arvelin, P. Lindqvist, Hydrocarbon Process. 1995, 74(2), 51. 128. Neste Jacobs, NExETHERS , Next Generation Etherification Technologies, http://www.nestejacobs.com/, 12 July 2006. 129. C. G´omez, F. Cunill, M. Iborra, F. Izquierdo, J. Tejero, Ind. Eng. Chem. Res. 1997, 36, 4756. 130. M. Koskinen, E. Tamminen, H. J¨arvelin, Presented at the NPRA Annual Meeting, San Antonio, TX, 1997. 131. A. M. Rouhi, Chem. Eng. News, 1995, May 29, 37. 132. E. F. Sousa-Aguiar, L. G. Appel. C. Mota, Catal. Today 2005, 101, 3. 133. S. C. Sorenson, J. Eng. Gas Turbines Power 2001, 123, 652. 134. HSC Chemistry for Windows, Chemical Reaction and Equilibrium Software, Version 5.1, Outokumpu Research, Pori, Finland, 2002. 135. W. Heitmann, in Ullmann’s Encyclopedia of Industrial Chemistry, http://www.mrw.interscience.wiley.com/ueic/articles/ a10023 /toc.html, 20 June 2005. 136. M. M¨uller, U. H¨ubsch, in Ullmann’s Encyclopedia of Industrial Chemistry, http://www.mrw.interscience.wiley.com/ueic/ articles/a08541 /frame.html, 20 June 2005. 137. J. Hu, Y. Wang, C. Cao, D. C. Elliot, D. J. Stevens, J. F. White, Ind. Eng. Chem. Res. 2005, 44, 1722. 138. A. Y. Rozovskii, Kinet. Catal. 2002, 44, 360. 139. G. R. Jones, H. Holm-Larsen, D. Romani, R. A. Sills, Presented at the Petrotech–2001 Conference, New Delhi, January 2001; http://www.topsoe.com/site.nsf/all/BBNN– 5PFHXH?OpenDocument, 10 November 2005. 140. Y. Ohno, Presented at the Japan DME Forum, 2003; http://www.aboutdme.org/, 5 November 2005. 141. Contract License and Basic Engineering of the First Largescale Plant for Production of DME, http://www.topsoe.com/, 16 June 2004. 142. M. Marchionna, R. Patrini, F. Giavazzi, G. C. Pecci, in 212th National Meeting, American Chemical Society, Washington, DC, 1996, p. 585. 143. F. Cunill, J. Tejero, C. Fite, M. Iborra, J. F. Izquierdo, Ind. Eng. Chem. Res. 2005, 44, 318. 144. J. Tejero, F. Cunill, M. Iborra, J. F. Izquierdo, C. Fite, J. Mol. Catal. A 2002, 182–183, 541. 145. B. T. Eskew, C. L. Geisler, in Handbook of MTBE and Other Gasoline Oxygenates, H. Hamid, M. A. Ali (Eds.), Marcel Dekker, New York, 2004, pp. 353–373. 146. G. Parkinson, Chem. Eng. 2005, 112(April), 21. 147. P. Hunszinger, H. J¨arvelin, M. Nurminen, R. Birkhoff, Hydrocarbon Process. 2003, 82(9), 57.

2882 Steam Reforming1 Jens Richard Rostrup-Nielsen∗

13.11.1

Introduction Reactions Steam reforming is the reaction between steam and hydrocarbons to give a mixture of hydrogen, carbon monoxide, carbon dioxide, methane and unconverted steam [1–4]. Steam may be replaced by carbon dioxide as reactant. The reforming reactions are accompanied by the water gas shift reaction. In addition, there may be the potential for carbon-forming reactions on the catalyst and thermal cracking of higher hydrocarbons. The reactions are listed in Table 1. The term steam reforming is also used for the reaction between steam and alcohols (methanol and ethanol) [5, 6] and liquid-phase reaction with carbohydrates [7] or bio-oils [8]. The term steam reforming should not be confused with catalytic reforming, used for the conversion of paraffinic hydrocarbons to high octane hydrocarbons such as isoalkanes and aromatics. A better term may be ‘‘oxygenolysis’’ [1], as the reaction involves the breakage of C−H and C−C bonds by means of oxygen-containing species. An alternative reaction route is partial oxidation by oxygen with or without the presence of steam, as practiced in non-catalytic processes (Texaco, Shell), catalytic partial oxidation (CPO) [9] and the hybrid, autothermal reforming process (ATR) [10]. This chapter deals with gas-phase steam reforming of hydrocarbons. 13.11.1.1

Thermodynamics The reforming reactions are catalyzed by Group VIII metals and the product gas will normally be close to the equilibrium of Eqs. (1) and (2) (Table 1). Industrial practice applies the term ‘‘approach to equilibrium’’ expressed by a temperature difference: 13.11.1.2


T (approach) = T (QR ) − Texit

The reforming reaction involving two stable molecules such as methane and water is strongly endothermic and it leads to the formation of more molecules. This means that Eq. (1) is favored by high temperature and low pressure, as illustrated in Fig. 1. Stoichiometric reforming (H2 O/CH4 = 1, CO2 /CH4 = 1) is rarely feasible [11], as most reformers operate at the supply pressure of the product gas to save energy for compression of the larger gas volume. At high pressure, the methane conversion is low, as shown in Fig. 1. Reaction (4) in Table 1 is the reverse Fischer–Tropsch synthesis, but the conversion of higher hydrocarbons is irreversible at the normal reforming temperatures. The higher hydrocarbons react on the metal surface to give C1 components or stay as carbonaceous deposits. At temperatures above 600–650 ◦ C, the catalytic reaction may be accompanied by thermal cracking. The negative heat of reaction of the reforming reaction and the high exit temperatures at typical process conditions mean that heat must be supplied to the process typically in a fired tubular reactor, the tubular reformer. However, depending on reaction conditions, the overall heat of reaction may become positive, for instance when converting higher hydrocarbons at low temperature. The heat of reaction of the methanation reaction [reverse 100% 1 bar abs 80% 60% 40% Methane conversion / %

13.11

0% 20 bar abs 80% 60% 40%

1 A list of abbreviations/acronyms used in the text is provided at the end of the chapter. ∗ Corresponding author.

S/C = 5.0 S/C = 2.5 S/C = 1.0

20%

(1)

where T (QR ) is the equilibrium temperature corresponding to the reaction quotient, QR , in the reforming reaction. The shift reaction can be assumed to be at equilibrium.

S/C = 5.0 S/C = 2.5 S/C = 1.0

20%

0% 400

500

600

700

800

900

1000

Reforming equilibrium temperature / °C

Steam reforming and methane conversion [11] (reproduced with permission from Elsevier).

Fig. 1

13.11.1 Introduction

Tab. 1

2883

Reforming reactions −
H◦298 /kJ mol−1

Steam reforming

LnKp = a + b/T a a

1. CH4 + H2 O = CO + 3H2 2. CO + H2 O = CO2 + H2 3. CH4 + CO2 = 2CO + 2H2
4. Cn Hm + nH2 O = nCO + n + Ethane n-Heptane 5. CH4 = C + 2H2 6. 2CO = C + CO2 7. CO + H2 = C + H2 O

m 2



−206 41 −247 H2 :

b −27106 4160 −31266

30446 −3798 32244

−347 −1107 −75 172 131

60.42 21053 12.69b −21.09 −17.29

−45256 −141717 −10779b 20486 16326

state: 25 ◦ C (298.15 K) 1 bar, linear regression 500–900 ◦ C. carbon.

a Standard b Whisker

estimated from enthalpy tables (Table 2) when the product gas composition is known from equilibrium calculations.

Enthalpy of formation (standard conditions 298 K, 1 bar) (linear regression 500–900 ◦ C)

Tab. 2

a H2 O H2 CO CO2 CH4 C2 H6 n-C7 H16 aH

b

−256555 −9435 −121704 −413856 −106592 −135527 −356368

40.87 30.17 32.93 53.93 70.58 120.3 375.6

in kJ kmol−1 , T in K.

reaction (1), Table 1] and the water gas shift reaction may become larger than that of the reforming of the higher hydrocarbons. The overall heat requirement can be Tab. 3

Applications The steam reforming process is a key technology for the manufacture of synthesis gas and hydrogen. By choosing the operating parameters (H2 O/C, CO2 /C, P , Texit ), a variety of gas compositions can be manufactured. Typical industrial examples are listed in Table 3. Synthesis gas for ammonia synthesis (H2 /N2 = 3) is manufactured by adding nitrogen with air in a second reforming step in which oxygen reacts with non-converted methane from the primary reformer, as illustrated in Fig. 2. Carbon monoxide is converted in high-temperature shift (360–420 ◦ C) and low-temperature shift (230 ◦ C) reactions and, after removal of CO2 by washing and remaining CO by methanation, the syngas is fed to the ammonia 13.11.1.3

H = a + bT a

Compound

References see page 2903

Typical process conditions for steam reforming

Process

Feed

General Ammonia synthesis Methanol synthesis Hydrogen

Natural gasa Natural gasa Natural gasa

Specific examples Methanol CO-rich gas CO-rich gas Pre-reformer GTL Pre-reformer SOFC

Natural gas Naphtha Natural gasb Natural gas Natural gas/diesel

a Natural b SPARG.

gas or naphtha.

H2 O/C/mol atom−1

2.8–3.5 1.5–2.0 1.7–4.0 2.5 1.5 0.9 0.6 0.8

CO2 /C/mol atom−1

0 0 0 0.2 0.4 0.54 0

P/bar absolute

Texit / ◦ C

30–45 30–45 25–40

750–820 700–750 850–920

25 22 8 31 1.4

920 950 900 396 450

2884

13.11 Steam Reforming

Ammonia

Hydrogen

Hydrogen for fuel cell

GTL synthesis

Process schemes using steam reforming. NG, natural gas; HDS, desulfurization, normally hydrodesulfurization (CoMo/ZnO); TSR, tubular steam reforming (Ni); SR, secondary reforming (Ni); WGS, water gas shift, high temperature (Fe3 O4 /Cr2 O3 ), low temperature (Cu/ZnO/Al2 O3 ); Meth, methanation (Ni); PSA, pressure swing adsorption; PR, pre-reforming (adiabatic) (Ni); GTL, gas-to-liquid; ATR, autothermal reforming (Ni); FT, Fischer–Tropsch synthesis; HC, hydrocracking.

Fig. 2

synthesis loop [12]. The process involves eight catalytic steps, including desulfurization of the feed. Ammonia plants are built at capacities of typically 2000 t day−1 . For the manufacture of hydrogen, the reforming process is followed by water gas shift carried out over copper catalysts (210–230 ◦ C) to ensure complete conversion of carbon monoxide into carbon dioxide and hydrogen [13]. From thermodynamics (Fig. 1), the optimum conversion for hydrogen would be favored by low pressure, high exit temperature and low steam to carbon ratio. The industrial supply pressure is typically 20 bar (1 bar = 105 Pa) and it turns out to be cheaper to operate with low steam to carbon ratio when using pressure swing adsorption (PSA) for cleaning the hydrogen (up to 99.999% H2 ). The PSA offgas is then used as fuel for the reformer as shown in Fig. 2. Steam reforming of hydrocarbons results in a H2 /CO ratio close to 3 [Table 1, reaction (1)]. By replacing steam with CO2 , the ratio can be reduced to one [Table 1, reaction (3)]. The required properties of the syngas are different for the different syntheses [11] as illustrated in Table 4. Both CO and CO2 are reactants in the methanol synthesis and the feedgas should have a module M = (H2 −CO2 )/(CO + CO2 ) close to 2. In contrast, CO is the only reactant for the low-temperature Fischer–Tropsch synthesis for liquid fuels and an H2 /CO ratio close to 2 is optimum. None of these ratios can

be achieved by steam reforming, although the addition of CO2 when available may be helpful (Table 1). If 25% of steam is replaced by CO2 , the overall stoichiometry corresponds to that of methanol [11]. This is rarely the situation. As a result, large-scale methanol plants and plants converting natural gas into synthetic fuels [gas-toliquid (GTL) plants] use autothermal reforming coming closer to the optimum ratios [11]. These large plants may convert natural gas flows of the size of order 250 000 Nm3 h−1 , corresponding to 25 000 barrels per day of synthetic fuels in each process line. At the other end of the scale, the steam reforming process is applied as hydrogen generator for fuel cells with a consumption of 0.25 N m3 h−1 of natural gas to produce 1 kW in a PEM fuel cell. PEM fuel cells are sensitive to CO, which must be removed either by methanation or preferential oxidation (PROX) over noble metal catalysts [14]. Phosphoric acid fuel cells (PAFC) operating at ca. 200 ◦ C are less sensitive to CO, whereas CO can be used as fuel in high-temperature fuel cells such as solid oxide fuel cells (SOFC) and molten carbonate fuel cells (MCFC). In high-temperature fuel cells (SOFC, MCFC), the steam reforming process may take place in the anode chamber. The coupling with the heat produced by the electrochemical process results in high electrical efficiency [5, 15, 16].

13.11.1 Introduction Tab. 4

Optimum syngas compositions [11]

Process Ammonia Methanol DME High-temperature Fischer–Tropsch Low-temperature Fischer–Tropsch Acetic acid Higher alcohols Industrial hydrogen Hydrogen for PEMFC Reducing gas (iron ore)

Optimum composition H2 /N2 = 3 H2 − CO2 M= =2 CO + CO2 H2 − CO2 =2 M= CO + CO2 H2 − CO2 M= =2 CO + CO2 H2 /CO ≈ 2 CO H2 /CO = 1 99.99 H2 95%) by radiation from the furnace gas and in a side wall-fired furnace also from the furnace walls, the remainder being convection. About 50% of the heat input is transferred through the tubes for the reforming reactions and for heating the gas to the exit temperature. For a typical ammonia plant reformer, about 60 and 40% of the transferred heat is consumed by the reaction and by heating the gas, respectively. Typical values of the overall heat transfer coefficient (U ) are 300–500 W m−2 K−1 [2]. The thermal efficiency of the tubular reformer and the waste heat recovery section is about 80% (related to fired heat input) and 95% (reformer feed plus fuel) as 13.11.2.1

H2 -plant based on steam reforming of natural gas. The tubular reformer is shown in the background with the PSA unit to the right (capacity: 39 000 Nm3 H2 equiv. h−1 ). Fig. 3

most of the heat that is recovered from the flue gas. The waste heat is used for steam production and for preheating of the reformer feed, combustion air, etc. The same is true for the heat contained in the hot product gas leaving the reformer. Whereas the steam can easily be used in ammonia and methanol plants for compression of synthesis gas, etc., surplus steam is often a problem for hydrogen plants. Optimum design aims at minimizing the steam export. One solution is to operate at a low steam to carbon ratio (Table 3), thus minimizing the flow through the reformer [13]. Processes based on steam reforming show high thermal efficiencies, as illustrated in Table 5. The tubular reformer, however, is an expensive piece of equipment. This has been the driving force behind strong efforts to reduce the size of the reformer by improving the heat transfer and hence reducing the number of tubes. Tubular reformers today are designed for operation at average heat fluxes exceeding 100 000 kcal m−2 h−1 (0.12 MW m−2 ), almost double what was industrial practice 20 years ago [13]. The critical parameter is not the average heat flux, but the temperature difference across the tube wall. Development of new high-alloy steels has allowed the use of tube wall temperatures well above 1000 ◦ C. The tube life is determined by the creep phenomenon, which is highly dependent on temperature; a 10 ◦ C higher tube wall temperature may result in halving of tube life [2]. The size of the reformer can also be decreased by increasing the inlet temperature. This involves the risk of thermal cracking of higher hydrocarbons in the preheater, which then may operate like a steam cracker for ethene. The conversion to alkenes and the risk of coke formation depend on temperature and residence time and may be expressed by the kinetic severity function [2, 31]. By installing an adiabatic low-temperature prereformer [11, 32], as shown in Fig. 5, it is possible to eliminate this problem and to increase the preheat temperature to about 650 ◦ C. In the pre-reformer, all higher hydrocarbons are converted in the temperature range 350–550 ◦ C and methane reforming and shift reactions are brought into equilibrium [33]. From a catalytic point of view, the steam reforming may appear straightforward as gas compositions and heat balances are determined by simple thermodynamics, but, as illustrated, it is a complex coupling of catalysis, heat transfer and mechanical design. It is not sufficient to consider an overall balance [2]: Transferred heat = reaction heat + sensible heat dT 4 U (TW − TCM ) =
H rv + us ρg cp dt dZ

(2)

13.11.2 The Reforming Process

Waste heat section

Fig. 4

2887

Reformer furnace

Reformer furnace. Waste heat section (only every second tube is shown. Number of tubes: 66).

Tab. 5

Efficiencies of processes based on steam reforming [30]

Process

Ammonia Methanol Hydrogen Synfuels

Yearly production/106 t a−1

Energy consumption/GJ t−1

120 30 20 17d

29 28 141b 67

Thermal efficiency (LHV)/% Practical

Ideal

65 72 84c 60

89.2 84.2 91.8 78

CO2 /t t−1

1.6a 0.28 9.0 1.18

a Including

CO2 converted into urea. GJ per 10 000 Nm3 H2 . c No steam export. d Excluding 8 × 106 t a−1 under construction. b 12.6

Flow diagram of process with tubular reformer with pre-reformer [32] (reproduced with permission from Elsevier).

Fig. 5

It is necessary to make a detailed local analysis to determine the tube wall temperature and to assess the risk

of carbon formation (see Section 13.11.4.3). This requires a two-dimensional reactor model considering the radial temperature concentration gradients [2, 34–37]. Tubular reformers are built today for capacities up to 300 000 Nm3 h−1 of hydrogen or synthesis gas [38]. The economy of scale is almost linear, with total costs related to the number of tubes and related hairpins, etc. This means that the partial oxidation processes become more economical at very large capacities because the economy of scale of the oxygen plant is more favorable [39], as illustrated in Fig. 6. As an example, steam reforming is the cheapest route for methanol plants at capacities below ca. 2500 t day−1 , whereas autothermal reforming (ATR) is cheapest at capacities of ca. 7000 t day−1 . For mid-size capacities, a hybrid two-step reforming (Fig. 6) becomes the optimum choice. With improvements of the tubular reformer the limit moves upwards. References see page 2903

2888

13.11 Steam Reforming

H2O/CH4

Air H2O/CH4

Log costs

Tubular reforming

O2

O2-plant Syngas Air H2O/CH4 O2

Log capacity

Syngas

Fig. 6

Impact of scale of operation [39] (reproduced with permission from Elsevier).

Convective Reformers It is possible to increase the amount of heat transferred to the process gas from about 50% of the supplied heat in the tubular reformer to about 80% (Fig. 7) when utilizing the heat of the hot product gas as heat input for the process by heat exchange [38, 40]. This is achieved in a convective reformer (Fig. 8), in which both the flue gas and the hot product gas are cooled by heat exchange with the process as flowing through the catalyst bed. This results in a more compact piece of equipment. However, in all heat exchange reformers, the heat exchange is by convection and this generally leads to lower heat fluxes than in reformers with radiant heat transfer. Hence the fired reformer remains the most economical solution for large-scale steam reforming. A major application of convective reformers is for the manufacture of hydrogen, as smaller amounts of waste heat mean that the steam export can be eliminated. They are also well suited for fuel cell plants because of their compactness and their capability to respond to quick load changes [14]. Another application of convective reforming is associated with recovery of the process gas heat in ammonia and methanol plants [34, 41] and lately also in ATR-based synthesis gas units for GTL plants [42].

Fig. 7

Reforming Rates and New Reformer Concepts Most industrial catalysts have a high activity for the reforming reaction provided the feed is sulfur free. As shown in Fig. 9, nickel and ruthenium catalysts are able to convert methane even at 300 ◦ C [43]. The conversion is low, however, because of thermodynamics.

It can be shown that there is a huge surplus of activity in a tubular reformer [11]. High conversion can be achieved at space velocities (SVs) of the same order of magnitude as reported for CPO [6] [SVC1 ≈ 106 vol.

13.11.2.2

13.11.2.3

Feed

Fuel

100

50

100

20

(a)

Feed

Fuel (b)

Energy conversion in steam reforming. (a) Radiant box reformer; (b) convective reformer. The non-transferred heat (fuel) is recovered in a waste heat section not shown.

2889

13.11.2 The Reforming Process

Flue gas outlet

Process gas inlet

Process gas outlet

Reformer tubes

Burner air

Burner fuel

Haldor Topsøe convection reformer [38] (reproduced with permission from Palladian Publishers).

Fig. 8

20 Ni - catalyst Ru - catalyst Conversion/%

15

10

ium

ibr

il qu

reformer is soon close to equilibrium, with the axial temperature gradient being the main driving force for further conversion, as illustrated in Fig. 10 [2, 45]. The catalyst tubes are filled with large catalyst pellets to minimize pressure drop and improve heat transfer; however, this results in low catalyst utilization (effectiveness factor) [2, 41]. The low catalyst utilization creates a potential for new reformer concepts. One attempt is to utilize the high reforming activity by carrying out the reaction at low temperature. This would allow the use of cheaper construction materials and the application of heat at lower temperature. The thermodynamic constraint may be circumvented by combining the reforming reaction with extraction of hydrogen through a palladium membrane [13, 46–48], thus pushing the reforming equilibrium to higher conversion. As shown in Fig. 11, this results in hydrogen at low pressure and CO2 at high pressure [13]. The last is good for CO2 sequestration and a low supply pressure of hydrogen may be acceptable for fuel cells. Still, there is a need to develop stable, thin Pd membranes (thickness of a few microns) and reliable manifold systems. In principle, a fuel cell with internal reforming operates as a membrane reformer. Another attempt is to utilize the surplus activity by catalyzing the heat transfer surfaces (catalyzed hardware) and by leaving the tubular design. Plate-type reformers and multi-channel reformers are being developed for compact units for small-scale operation [49–52], for instance, for use in cars. Some designs involve a reforming catalyst on one side of the metal wall and a combustion catalyst on the other side [49, 53]. This means that the heat needed for the reforming reaction is provided by conduction, but still the heat transfer over

E

5

Temperature

Rate

250

300

350

Temperature /°C

Steam reforming at low temperatures. Conditions: H2 O/CH4 = 4, H2 O/H2 = 10, P = 1 bar [43] (reproduced with permission from Sekiyu Gakkaishi). Fig. 9

CH4 (vol. cat.)−1 h−1 ]. However, the tubular design imposes constraints on the utilization of the catalyst [SV = 2000–4000 vol. CH4 (vol. cat.)−1 h−1 ]. The tube diameter and the total tube surface area are fixed from mechanical design criteria and the constraint from maximum allowable temperature difference over the tube wall. This leaves a large catalyst volume as the dependent parameter [2, 44]. As a result the process gas in a tubular

Parameter indicated

0 200

CH4 CH4 equilibrium

η effectiveness factor 0 Inlet

Tube length

100% Outlet

Fig. 10 Typical profiles in a tubular steam reformer [11] (reproduced with permission from Elsevier). References see page 2903

2890

Fig. 11

13.11 Steam Reforming

Membrane reforming. Process scheme [13] (reproduced with permission from Elsevier).

the gas film determines the heating of the process gas. The multi-channel reformer designs are compact, with estimated hourly productivities in the order of 1000 Nm3 H2 (m3 reactor)−1 [49, 50]. The multi-channel reformers may result in fast heating of the feed stream within milliseconds, hence with a possibility of avoiding coke from thermal cracking in preheaters and over deactivated catalyst [54].

Activity of various metals: steam reforming of methane [15]. 500 ◦ C, H2 O/CH4 = 4, H2 O/H2 = 10, P = 1 bar absolute, support MgAl2 O4 (12 m2 g−1 )

Tab. 6

Catalyst

Metal/wt%

TOF/molecules s−1

E/kJ mol−1

9.1 0.4 0.5 0.5

0.7 3.5 11.9 1.3

26 23 29 –

Ni Ru Rh Pt

13.11.3

Catalysts Catalyst Activity Group VIII metals are active for the steam reforming reactions [2, 11]. Nickel is the preferred choice. Cobalt and iron are also active, but cobalt will be oxidized under normal reforming conditions and iron requires a strongly reducing atmosphere, as for instance in shaft furnaces for direct iron ore reduction. All these metals are subject to oxidation above a certain H2 O/H2 ratio, being lower than that predicted by bulk thermodynamics and depending on the specific catalyst (metal particle size, support, etc.) [2]. Catalysts should be compared in terms of intrinsic activity referred to unit metal surface area, i.e. turnover frequencies (TOF). There are a few pitfalls when measuring the intrinsic activity for steam reforming in differential gradientless reactors [11]. Even with small catalyst particles, the high heat of reaction may cause temperature gradients over the gas film. This makes measurements at high temperatures (above 550 ◦ C) difficult. The high space velocities required to obtain partial conversion (SV 105 –106 vol. vol.−1 h−1 ) means that there is no back diffusion of hydrogen (or hydrogen from the interior of the small catalyst particles). As a result, the H2 O/H2 ratio at the inlet easily exceeds the critical value for catalyst oxidation. Hence reliable measurements should include hydrogen in the feedstream. Several methods are used to measure the nickel surface area [11], including chemisorption of hydrogen or hydrogen sulfide, X-ray diffraction and electron

25

13.11.3.1

Forward CH4 turnover rate / mol (g atom surface metal)−1 s−1

Pt catalyst Ir catalyst

20

15

Rh catalyst

10

5 Ru catalyst 0

0

0.2 0.4 0.6 Metal fraction dispersion

0.8

Activity of noble metal catalysts: 650 ◦ C, 1 bar absolute, H2 O/CH4 [55] (reproduced with permission from the authors and ACS Publications). Fig. 12

microscopy. These methods can also be applied for the other Group VIII metals. Rhodium and ruthenium are significantly more active than nickel, with the rest of the noble metals having activities in between. Examples of comparisons are shown in Table 6 [1, 15]. In contrast, Wei and Iglesia [55] found Pt to be the most active metal (Fig. 12).

13.11.3 Catalysts Reforming experiments at atmospheric pressure: various hydrocarbons, 500 ◦ C, 1 bar absolute, catalyst Ni/MgO [1]

Tab. 7

Feed

Specific initial activity (500 ◦ C)a × 103

Apparent activation energy, Ea /kJ mol−1

mol m−2 h−1

g-atom m−2 h−1

61 120 138

61 240 552

CH4 C2 H6 n-C4 H10

110 76 78

a Rates

calculated at the same partial pressure of hydrocarbon (H2 O/Cn Hm = 8).

Alkali has a pronounced negative effect on the activity of Group VIII metals [1, 11, 15], hence it is important to check for trace amounts of alkali in the catalysts before comparing activities. The activity for steam reforming of methane on nickel catalysts correlates linearly with that for steam reforming of ethane and of naphtha [1]. Methane is the least reactive of the alkanes [1], as illustrated in Table 7. The sequence of reforming activities also correlates with that of activities for hydrogenolysis and methanation, as shown in Table 8 [1], but with a much stronger impact of alkali on the latter. This may be explained by enhanced adsorption of CO in the presence of alkali leading to self-poisoning [56]. The impact of alkali on steam reforming and hydrogenolysis may be explained by electrostatic considerations as the adsorption of potassium on nickel creates a dipole in the same direction as that created by adsorption of the hydrocarbon. This leads to repulsion and an increase in the activation energy [57]. The situation is the opposite for the activation of nitrogen in ammonia synthesis for which alkali promotes the rate. Apparently, alkali has no impact on the activity of decomposition of ammonia. Tab. 8

Non-Metal Catalysts Non-catalytic steam reforming requires high temperatures. Methane cracks above 1000 ◦ C into radicals, leading to the formation of ethene, acetylene and coke [58]. The radicals may react with steam radicals, but temperatures above 1500 ◦ C are necessary for significant conversion [59]. One approach to improve rates is the use of plasma technology [60], the key issue being the power consumption. The thermal cracking of higher alkanes becomes significant above 650 ◦ C [31, 58], with the formation of alkenes, aromatics and coke. This is applied in steam crackers, where steam is added as a diluent and for minimizing coke formation. When applying an alkaline catalyst for gasification of coke and coke precursors, it has been possible to make syngas. A process was tested by Kellogg [61] in the mid-1960s in which coal was dissolved in molten potassium carbonate and was reacted with steam to produce synthesis gas at high temperatures and pressures. Another example is the co-production of syngas and light alkenes which was possible from heavy gas oil and naphtha over a potassium-promoted zirconia support under conditions close to those applied in industrial steam crackers [58, 62]. Alkaline catalysts have also been studied for steam reforming of methane. However, rates over a calcium aluminate catalyst at 850 ◦ C were still a fraction of what is obtained on nickel at 500 ◦ C [63]. The same is true for attempts at steam reforming on SOFC electrodes. CeO2 has an activity almost two orders of magnitude less than nickel [64]. Another approach has been to use molybdenum carbide and tungsten carbide catalyst for steam and CO2 reforming [65, 66]. Again, rates are significantly lower than those obtained on Group VIII metals [66], as shown 13.11.3.2

References see page 2903

Relative specific activities at atmosphere pressure [1, 2]

Catalyst

A: Ni/MgO B: Ni/MgO C: Ni/MgO D: Ni/MgAl2 O4 E: Ni/MgAl2 O4 F: Ni/Al2 O3 G: Ni/Al2 O3 H: Ni/Al2 O3 I: Ni/C

Alkali/wt.%

Na 0.07 K 0.5 K 1.6

K 5.8

Reforming of C2 H6 500 ◦ C

CH4 500 ◦ C

1.0 2.4 0.03 3.0 0.02 0.7 2.5 0.09 0.04

1.0 2.0 0.09 1.4 0.02

Hydrogenolysis of C2 H6 300 ◦ C

1.0 4.2 0.08

8.7 0.04

2891

0.01

Methanation of CO 250 ◦ C

1.0 9.1 0.002

2.7 20.6 0.001 0.4

Decomposition of NH3 500 ◦ C

1.0 0.9 0.9 3.2 1.5 0.4 0.5 0.5 0.5

2892

13.11 Steam Reforming

by Xu and Froment [72] was based on a classical Langmuir–Hinshelwood approach and resulted in the following expressions:

Log (activity/ mol g−1 s−1)

−4.0 −4.2 −4.4 −4.6

Ru/MgAl2O4 Ru/Al2O3 (Claridge et al.) (Sehested et al.)

a: CH4 + H2 O = CO + 3H2 : k1 PCH4 PH2 O r1 = (1 − β) PH2 2 Z 2

−4.8 −5.0 −5.2 −5.4

Mo2C (Sehested et al.) Mo2C (Claridge et al.)

CO + H2 O = CO2 + H2 : k2 PCO PH2 O r2 = (1 − β) PH2 Z 2

b:

−5.6 −5.8 −6.0 0.80

1.00

1.20

1.40

1.60

1.80

1000 / T / K−1

(3)

(4)

c: CH4 + 2H2 O = CO2 + 4H2 : r3 =

Reforming with carbide catalysts [66] (reproduced with permission from Elsevier). Fig. 13

k3 PCH4 PH2 O 2 PH2 3.5 Z 2

(1 − β)

(5)

where

Kinetics Most kinetic studies of the steam reforming process deal with initial rates, although the reaction in the tubular reformer proceeds close to equilibrium (Fig. 10) with the temperature gradient as the driving force. Reactions close to equilibrium become first order in any parameter describing the distance from equilibrium [45]. The first kinetic studies were strongly influenced by diffusion restrictions and it was not before Temkin’s group [68] studied the reaction on nickel foil in the 1960s that more precise information was achieved [68, 69]. Recent work by Wei and Iglesia [70] gives strong evidence that activation of methane is the rate-determining step. Isotope exchange studies on Ni/MgO catalysts at 600 ◦ C with CH4 /CD4 showed no formation of mixed isotopes, which would have been the case if the methane adsorption step was at equilibrium. Furthermore, Wei and Iglesia found no impact on the rate of other reactants. Although attempts have been made to establish kinetics for steam reforming on basis of a microkinetic approach [71], most work is based on empirical kinetics, which has been sufficient to develop highly sophisticated models for tubular reformers. The comprehensive work 13.11.3.3

Z = 1 + Ka,CO PCO + Ka,H2 PH2 + Ka,CH4 PCH4 + Ka,H2 O

PH2 O PH2

Because the three reactions are not independent, it is necessary to combine the three rate equations into two: one for conversion of methane and one for production of CO2 : rCH4 = r1 + r3 rCO2 = r2 + r3

(6)

These two expressions include five temperaturedependent constants listed in Table 9. They show a small negative reaction order of the overall pressure, in agreement with the data in Fig. 14 [11]. There is general agreement that the reaction is first order with respect to methane with activation energies 100 Activity / mol g−1 h−1

in Fig. 13. Molybdenum carbide will hardly be stable in a plug flow reactor [66]. It might be argued that low catalytic activity should be sufficient for the tubular reformer. This, however, would be at the expense of a high tube wall temperature, as illustrated by Eq. (1). Low-activity catalysts might have potential for internal reforming in high-temperature fuel cells, where the balance between the rate of reforming and the rate of the electrochemical reaction is critical [15, 67].

10 bar 10

4 bar 30 bar

1

0.1 1.1

1.15

1.2

1.25 1.3 1000 /T / K−1

1.35

1.4

Steam reforming of CH4 in a plug flow reactor with H2 O/CH4 = 4, H2 O/H2 = 10 and 0.2 g of catalyst (with particle diameters in the range 0.16–0.3 mm); space velocity = 1.7 × 106 vol. total feedgas (vol. cat. bed)−1 h−1 [11] (reproduced with permission from Elsevier). Fig. 14

13.11.3 Catalysts

was not observed in the studies of Wei and Iglesia [55, 70], probably due to the higher temperatures applied. CO2 reforming of propane was studied by Olafsen et al. [78]. Under industrial conditions, reaction rates for steam reforming are strongly influenced by mass and heat transport restrictions. This is reflected mainly by a temperature gradient over the gas film surrounding the pellet and concentration gradients inside the pellet. The effectiveness factor is typically below 5% (Fig. 10) [2]. Even in a continuous stirred tank reactor (CSTR), significant temperature and concentration gradients are present, as shown in Fig. 15 [79].

Reforming kinetics: parameters for Eqs. (3)–(5) [72]; Ni/MgAl2 O4 catalyst, Ni surface area 3 m2 g−1

Tab. 9

Value 4.225 × 1015 exp(−240.1/RT) 1.995 × 106 exp(−67.1/RT) 1.020 × 1015 exp(−243.9/RT) 8.23 × 10−5 exp(70.65/RT) 6.12 × 10−9 exp(82.90/RT) 6.65 × 10−4 exp(38.28/RT) 1.77 × 105 exp(−88.68/RT)

a Rate constants are in units of mol g−1 h−1 ; activation energies

and K adsorption constants are in kJ mol−1 .

in the range 100–120 kJ mol−1 [1, 68, 70, 73]. This is not reflected by Eq. (3). Studies of steam reforming of ethane (1 bar, 500 ◦ C) [1] showed that the reaction order with respect to steam varies with composition of the catalyst with the presence of alkali resulting in large negative reaction orders. The retarding effect of water decreases with temperature. This is another inconsistency with expressions (3)–(5) including a negative heat of adsorption of steam (Table 9). Recent work by Wei and Iglesia [55, 70] confirmed the first-order kinetics with respect to methane, but with no impact on the rates for other reactants. This was achieved on a number of Group VIII metals (Ni, Pt, Rh). This is in contrast to an overall pressure dependence that is slightly negative reported in other studies [2, 11, 72]. The discrepancy might be related to the fact that the studies of Wei and Iglesia were carried out at high temperatures (580–730 ◦ C). The steam reforming of higher hydrocarbons shows reaction orders with respect to hydrocarbons of Al2 O3 > TiO2 > SiO2 . 13.13.4

Catalyst Deactivation Sintering Methanol synthesis catalysts undergo relatively fast deactivation even in the absence of poisons. More than one-third of the activity is lost during the first 1000 h of operation [18, 185], as seen in Fig. 5. Despite this fact, which often determines the economic lifetime of an industrial catalyst charge, relatively little has been published on the subject. Thermal sintering of dispersed crystallites should not, according to Tammann’s rule: 13.13.4.1

Tmobility > 0.5Tbulk melting point occur for copper with a melting point of 1358 K at industrially used operating temperatures of 480–580 K. 4

TF

3

2

1

0 0

200

400

600

800

Days on-stream Relative activity versus time on-stream: (•) boiling water reactor; (◦) quench reactor, first bed.

2929

XRD analyses of spent methanol synthesis catalysts do, however, reveal a growth in the copper crystallite sizes, from typically 7 to more than 20 nm. Sintering can, however, also occur by release of atomic or molecular species from crystallites above the H¨uttig temperature, one-third of the melting temperature, Tm , according to Flynn and Wanke [186]. This leads to the following equations describing the deactivation: −

dS = kS n dt

(7) 1

S(t) = S0 [1 + k(n − 1)t] n−1

(8)

where S is the surface area, implicitly assumed to be proportional to activity, t is time and k and n are constants; n would typically range from 2 to 16. Industrial catalysts are therefore normally not operated above 573 K [187, 188]. Irreversible deactivation was observed when Cu/ZnO was operated in CO/H2 gases without CO2 or H2 O [16, 189], which has been interpreted as reduction of Cu+ from the ZnO matrix. Other explanations could be evaporation of Zn or formation of brass (Cun Zn). The latter has been observed in low-temperature shift catalysts above 532 K. Rapid formation of brass has been observed in methanol synthesis catalysts using H2 /CO mixtures above 570 K, leading to rapid deactivation [190]. The beneficial effect of adding alumina (or chromia) and ZnO to the catalysts has been explained by rather crude models invoking a mechanical spacing effect, which prevents sintering [84, 191, 192]. In very CO2 rich synthesis gases [leading also to high water contents after reactions (1) and (3)], accelerated aging can also be observed, perhaps related to failure of the alumina phase to stabilize the Cu/ZnO constituent of the catalyst. This fact has been used to devise accelerated aging tests by subjecting the catalysts to CO2 -rich gases at high temperature, thus generating substantial amounts of water [193]. That a high CO2 content in itself does not necessarily induce rapid aging, but that it is rather the resulting water which is responsible, is indicated by the results reported [194, 195] from stability tests with a commercial Cu/ZnO/Al2 O3 catalyst in slurry phase. Although the feed gas contains 13% CO2 , the simultaneous presence of 51% CO will result in a low water partial pressure, because water will be removed by the water gas shift reaction. Additional information on deactivation models dealing with temperature and gas composition effects is available [23, 196]. Dynamic modeling and genetic algorithms have been used to study optimum operating strategies to counteract aging [197–199].

Fig. 5

References see page 2943

2930

13.13 Methanol Synthesis

Sulfur Poisoning In plants based on natural gas as feedstock, sulfur poisoning of the methanol synthesis catalyst rarely represents a problem, because the feedstock has to be carefully desulfurized in order to protect the reforming catalysts, and sulfur in the form of H2 S would normally be condensed out with the surplus water after the reforming step. Lubrication oil from the make-up gas compressor and the recirculator can, however, contain sulfur compounds and, in plants based on gasification of coal or heavy oil, sulfur can be a problem. The industrial synthesis catalyst does contain ZnO, which provides some self-guarding properties to the catalyst because sulfur, most likely as hydrogen sulfide, will be absorbed by the ZnO according to the equilibrium 13.13.4.2

−−− −− → ZnO + H2 S ← − ZnS + H2 O

(9)

It should be noted that in the hydrogen-rich methanol synthesis gas, H2 S will be the most stable sulfur species and the content of other compounds such as COS, CS2 , SO2 and thiols will be minute if the gas is equilibrated. According to the temperature and water content of the gas, this equilibrium defines a residual H2 S content, which is still able to block the active sites considered to be metallic copper or a copper–zinc surface alloy (see Section 13.13.6). It is well known that metal surfaces can be poisoned by sulfur, notably H2 S [200, 201]. A sulfur for copper isotherm has been derived [202]. Synthesis catalysts containing several weight percent of sulfur can retain a significant fraction of their activity [19, 203]. In a study of the influence of different sulfur compounds in an internal recycle reactor [204], hydrogen sulfide and thiophene behaved similarly, probably because thiophene was transformed to H2 S over the catalyst, whereas COS had hardly any effect in the very CO-rich synthesis gas used. It is believed by others, however, that compounds such as COS [205, 206] (which does not directly react with ZnO) attacks the Cu surface and results in rapid poisoning in a similar manner to the poisoning action of thiophene and other organic sulfur compounds [207]. Other Poisons Chlorine is not often present in the feed gas to methanol reactors, but it is a very severe poison, because it induces accelerated sintering of both Cu and ZnO. The mechanism has been studied especially in connection with copper-based, low-temperature shift catalysts and is believed to be related to the formation of volatile copper and zinc chlorides [46, 208, 209]. Carbonyls of iron such as Fe(CO)5 and Ni(CO)4 have also been studied as catalyst poisons [210]. These carbonyls can be present in the make-up gases from 13.13.4.3

gasification plants or be generated within the synthesis loop itself from the steel in heat exchangers or in the reactor itself. The carbonyls are catalytically decomposed to free metal over the methanol synthesis catalyst. Fe is a strong poison, even below 1000 wt. ppm [211–213]. Part of the poisoning by Fe can be explained by the Fischer–Tropsch activity of Fe, which will ultimately cover the catalyst with high-boiling waxes (C20 –C50 aliphatic alkanes). Tests of spent catalysts used in industry containing significant amounts (0.4 wt.%) of Ni have shown only a very modest deactivating effect [213], in contrast to what is claimed elsewhere [214]. Other studies have also shown that adding Ni during the precipitation of the catalyst gives a fully active methanol catalyst in which Ni is found as Ni3 ZnCu0.7 [193, 215]. In plants using a coal gasifier for producing synthesis gas, significant amounts of arsenic have been found on deactivated catalysts [207, 216]. Studies with controlled addition of contaminants to a synthesis gas gave the following ranking of poisons: C4 H4 S = AsH3 > CH3 Cl > CH3 SCN > CS2 > COS > PH3 > CH3 F. Industrial examples of poisoning with fluorine and phosphorus have not been found. A surprising finding was a nonpoisonous action of N-containing compounds such as HCN, CH3 NH2 and CH3 CN. Others [217] have found a modest reversible poisoning of methanol catalysts by NH3 and a stronger reversible poisoning by acetic acid and other organic acids. 13.13.5

By-Product Formation

Although abundant literature exists on the intentional formation of higher alcohols by modified low-temperature methanol synthesis catalysts [21], studies of selectivity in proper methanol synthesis are relatively scarce. Modern copper-based methanol catalysts are very selective. In fact, selectivities above 99.9% are not uncommon. This is truly remarkable, because all of the by-products normally found: • higher alcohols, predominantly ethanol, butanols and propanols • esters, notably methyl formate and methyl acetate • ethers, especially dimethyl ether • ketones, mainly acetone and methyl ethyl ketone • hydrocarbons such as normal paraffins • minute amounts of acids and aldehyde are thermodynamically much more favored products than methanol; formaldehyde and formic acid are exceptions. Higher alcohol formation is favored by the presence of alkali and all commercial catalysts contain minor residual amounts. The distribution of the higher alcohols

13.13.6 Reaction Mechanism(s) and Active Sites

Tab. 6

high temperatures and CO partial pressures) unless the catalyst contains iron or nickel from the manufacturing step or by deposition from carbonyls as reported [222].

Adapted from Ref. [185]

Parameter

Tinlet /K Toutlet /K CO:H2 ratio CO:CO2 ratio Ethanol/wt. ppm Propanols/wt. ppm Butanols/wt. ppm Acetone/wt. ppm Methyl ethyl ketone/wt. ppm

Gas type CO-rich

CO2 -rich

470 568 0.33 13.132 2840 921 651 48 83

470 569 0.16 0.80 287 166 110 90–95%) of the make-up gas. This recycle is dictated by the unfavorable gas-phase thermodynamics at the temperature required by today’s catalysts. Recycling requires investment costs for the recirculator and higher cost for methanol condensers, etc., and the energy consumption also increases. In order to overcome these equilibrium limitations, several solutions have been suggested. Using fine alumina powder as an adsorbent for methanol in a gas–solid–solid, trickle flow reactor system eliminates the 13.13.8.12

References see page 2943

2942

13.13 Methanol Synthesis

equilibrium limitations by removing the methanol from the gas phase [385]. The methanol was desorbed from the alumina by lowering the pressure. The mechanical problems connected with handling the solids and the cyclic operation make the process too expensive and complicated. Another approach is to use a liquid that selectively absorbs methanol. High-boiling tetraethylene glycol dimethyl ether (TEGDME) serves this purpose [386, 387]. Unfortunately, the presence of TEGDME in the catalyst pores also lowers the efficiency factor of the catalyst. TEGDME can also be used to remove methanol between reactors in series, as for example a series of boiling water reactors. The methanol and water are flashed off from the solvent, which is returned to the absorbers. More than 97% CO conversion can be achieved by using four reactors in series [388, 389]. The use of a supercritical solvent has been investigated using 2-butanol [390]. In addition to the solvent effect as a supercritical fluid, the 2-butanol was found to have a catalytic effect similarly to other studies using an alcohol co-solvent and formate as intermediate. Thermodynamic calculations with the Soave–Redlich–Kwong equation of state were performed using different alkanes, identifying hexane and heptane as the most suitable [391]. Experimental studies in a three-phase slurry reactor have confirmed that the gas-phase equilibrium limitations could be broken using n-hexane as a supercritical solvent, albeit the −1 space velocities were rather low ( Windows

Cages = Windows

Cages = Windows

8-ring window

10-ring window

12-ring window

SAPO-18 (AEI)

Fig. 7

Schematic illustration of the channel–cage relations in relevant topologies.

Conversion or selectivity / %

13.14.4 Reaction Mechanisms

2955

100 Total olefins

80 60

Conversion

Ethene 40 Propene 20 Propane 0

Time Conversion and selectivities as a function of time on stream over SAPO-34.

Wt.% in product stream

Fig. 8

60 SAPO-34

50

ZSM-5

40 30 20 10 0 C1-C3 Paraffins

Ethene

Propene

C4 + & others

Comparison of product yields over H-ZSM-5 and SAPO-34 [5].

Fig. 9

systems, catalyst optimization and modifications have been reviewed previously [16]. Less work is published in this area for SAPO-34, but Wilson and Barger [25] have published results on the effect of crystallite size and Si content for SAPO-34. It is reported that a reduction in Si content decreases byproduct propane formation and increases catalyst life. Reduction in crystallite size has the same effect. A number of SAPO-type catalysts have been studied as potential MTO catalysts, including MeAPSOs, ElAPSOs, MeAPOs and also Ni-modified SAPO-34. A comparison of the catalytic properties of these different materials is difficult since the results obtained depend on many factors such as morphology, composition and test conditions. A detailed analysis of this work is outside the scope of this chapter. It should only be stated that, although several improvements have been claimed [16], none has resulted in any commercial significance. 13.14.4

Reaction Mechanisms

Immediately after the publication of the original papers revealing the methanol conversion reaction, a number

of laboratories started work on the reaction. Because the reaction was wholly unexpected, the majority of papers had elucidation of the reaction mechanism as their main objective. Simultaneously with the formation of hydrocarbons, there is also the formation of dimethyl ether (DME). This reaction is actually much faster than the hydrocarbon formation, so under most circumstances the reactant system is essentially an equilibrium mixture of methanol, DME and water. The question has therefore been raised of whether the reaction takes place from methanol or DME or both. A large number of mechanistic pathways were proposed during the early years. They can be broadly categorized into four groups, each having several subgroups that will not be discussed here. Detailed surveys can be found in earlier reviews [13–16]. Carbene−Carbenoid Mechanisms This mechanism was first proposed by Venuto and Landis [26] to explain the formation of small amounts of olefins (alkenes) that were obtained when methanol was reacted over NaX zeolite at 260 ◦ C. They suggested the sequence 13.14.4.1

H−CH2 −OH −−−→ H2 O+ : CH2 followed by n : CH2 −−−→ (CH2 )n Chang and Silvestri [12] pointed out that a direct reaction between the carbene fragments was unlikely because they were likely to be present in only very low concentrations. They therefore refined the concept to involve two methanol molecules, where a transient formation of carbene leads to insertion in the C−O bond in methanol, whereby a C−C bond is formed. References see page 2964

2956

13.14 Methanol-to-Hydrocarbons

Carbocationic Mechanisms Ono and Mori [27] studied the reaction at low temperatures (220–260 ◦ C) and found that the reaction displayed characteristics of an autocatalytic reaction. IR studies gave indications that methoxy groups could be formed within the zeolite structure. They speculated that the methoxy group could split off a methyl cation that could attack a methanol/DME molecule and form a C−C bond. They pointed out that after formation of the first C−C bond the reaction might proceed by a chain propagation process. 13.14.4.2

Trimethyloxonium Intermediacy Van den Berg et al. [28] postulated that on an acid catalyst formation of trimethyloxonium (TMO) ions might take place via a reaction between a protonated DME molecule and a DME molecule, leading to a TMO and a methanol molecule. The TMO might then transfer a proton to the zeolite framework and the ylide group which would then be formed might attack the neighboring methyl group, i.e. a Stevens-type rearrangement. 13.14.4.3

Radical Mechanisms Clarke et al. [29] utilized electron spin resonance spectrometry and detected some formation of free radicals. On this background, it was suggested that the hydrocarbonforming reaction might involve free radicals. 13.14.4.4

The Hydrocarbon Pool The initial work on the MTH mechanism, referred to above, focused on how two or more C1 entities (e.g. methanol, DME, TMO ions) could react so that C−C bonds are formed. Later, Dahl and Kolboe [30–32] performed co-feeding experiments of 13 CH3 OH and 12 C ethanol or propanol (which were rapidly converted to the corresponding alkenes) over SAPO-34 and observed that only a small portion of the products contained 12 C. They further observed that the 12 C present was fully scrambled among the products. This result led Dahl and Kolboe to suggest that the MTH reaction proceeds mainly by a mechanism where a pool of adsorbed hydrocarbons is all the time adding methanol and splitting off ethene, propene and possibly even higher homologues, as illustrated schematically in Fig. 10. Today, the importance of direct C−C bond formation is considered to be minor and the hydrocarbon pool mechanism is gaining general acceptance [33–35]. The hydrocarbon pool was initially not further specified, but during the last few years it has become clear that methylbenzenes play central roles in the hydrocarbon pool mechanism and are essential parts of the catalytic cycle. Already in 1983, Mole and coworkers [36, 37] observed that toluene acts as a ‘‘cocatalyst’’ for the 13.14.4.5

C2H4 C3H6

−n H2O

n CH3OH

(CH2)n

Saturated Aromatic

C4H8 Fig. 10

The carbon pool reaction mechanism.

MTH reaction over H-ZSM-5. In 2000, Mikkelsen et al. [38] observed isotopic scrambling in the olefinic products when cofeeding [12 C]toluene and [13 C]methanol over H-ZSM-5. Arstad and Kolboe [39, 40] studied the organic material trapped inside SAPO-34 after switching from a [12 C]methanol to a [13 C]methanol feed. They observed isotopic scrambling in all polymethylbenzenes, especially the higher homologues. Further, they analyzed the olefinic products after the 12 C/13 C switch and observed that the isotopic distribution of the olefinic products changed only gradually, strongly pointing to polymethylbenzenes as being the active hydrocarbon pool in the SAPO-34 catalyst. Sassi et al. [41] showed that polymethylbenzenes fed over H-Beta zeolite are active for olefin formation. An interesting feature when considering the cited contributions is that similar conclusions have been drawn for very different zeotype systems. The role of polymethylbenzenes as the major hydrocarbon pool species appears to be independent of the zeotype catalyst chosen. The exact nature of the hydrocarbon pool species may, however, depend on the catalyst type and reaction conditions, and polymethylnaphthalenes and cyclopentenyl species have also been shown to function as hydrocarbon pool species [18]. 13.14.4.5.1 Product Formation Most detailed studies of the reactivity of polymethylbenzenes have been carried out over H-Beta zeolite, not least because the 12membered ring pore structure allows direct feeding of polymethylbenzenes. In H-Beta zeolite, it has been shown that hexamethylbenzene has, compared with lower polymethylbenzenes, by far the highest reactivity for product formation [41]. It has further been shown that the amount of HexaMB retained inside the pores during the MTH reaction decreases dramatically when the methanol feed is stopped and the zeolite is flushed with inert gas [42]. The decrease is accompanied by continued production of gas-phase products. Under methylating conditions, HexaMB may take up a CH+ 3 ion and form the heptamethylbenzenium ion (HeptaMB+ ). This was demonstrated already in 1958 by Doering et al. [43], who studied the

2957

13.14.4 Reaction Mechanisms

+

− H+

Zeolite−

+

CH3OH

− H+

Zeolite−

H-Zeolite

H-Zeolite CH3OH

CH3OH − H+

+

+

Zeolite−

Fig. 11

Zeolite−

The exocyclic methylation reaction.

Friedel–Crafts methylation of benzene. They also observed an equilibrium between HeptaMB+ and its corresponding base, 1,2,3,3,4,5-hexamethyl-6-methylene1,4-cyclohexadiene (HMMC). In 2002, Song et al. [44] observed the HeptaMB+ ion inside H-Beta zeolite by NMR spectroscopy. Slightly later, Bjørgen et al. [45] isolated HeptaMB+ trapped inside H-Beta during the methanol + benzene reaction, as HMMC, and identified it ex situ by high-resolution NMR spectroscopy. Recently, protonation of polymethylbenzenes in H-Beta zeolite has been used to probe the acid strength of the zeolite, as measured by IR and UV/visible spectroscopy. It was shown that TetraMB and higher homologues inside the micropores of H-Beta zeolite are protonated, whereas TriMB is not [46, 47]. The acid strength of the zeolite is a key parameter in discussing possible reaction paths whereby olefins may be formed from the HeptaMB+ ion. Two main reaction paths have been proposed, the exocyclic methylation route and the ‘‘paring’’ reaction (Figs. 11 and 12, respectively). In the exocyclic methylation route, first proposed by Mole and coworkers [36, 48], to explain the co-catalytic effect of toluene, and later refined by Haw et al. [19], HeptaMB+ is first deprotonated to give HMMC. The exocyclic double bond then undergoes reaction with an incoming methanol molecule, resulting in an ethyl group on the benzene ring, which is subsequently dealkylated as ethene. Deprotonation followed by a new methylation leads to propene formation. Theoretical calculations, considering reactions in the gas phase, indicated that the rate-determining step in the exocyclic methylation route is deprotonation of the HeptaMB+ ion and that the essential intermediate in the MTH cycle is gem-dimethyl groups on the aromatic ring [49]. Evidence was found that the propene-to-ethene ratio could be increased by increasing the number of methyl groups on the aromatic ring.

CH3OH

+

CH3OH

−H+

+

−H+ + +

Fig. 12

+

The paring reaction.

The ‘‘paring’’ reaction mechanism, introduced by Sullivan et al. in 1961 [50] to explain isobutane formation from HexaMB, suggests alkyl side-chain formation by ring contraction/expansion, as shown in Fig. 12. The paring reaction mechanism will lead predominantly to propene and isobutene formation. It will further lead to a carbon atom interchange between the ring and the methyl substituents. In a recent study by Bjørgen et al. [45], methylbenzenes consisting of a [12 C]benzene ring and [13 C]methyl groups were obtained by reacting [13 C]methanol and [12 C]benzene over zeolite H-Beta. The isotopic distribution of the gas-phase products strongly suggested a paring reaction-type mechanism to dominate in the olefin formation over this zeolite. However, it might be possible to favor the exocyclic methylation route and thereby adjust the reaction selectivity by using a catalyst with sites having a lower acid strength. References see page 2964

2958

13.14 Methanol-to-Hydrocarbons

13.14.4.5.2 Parallel Reactions In the 1980s, Dessau and LaPierre [51, 52] suggested that the MTH reaction proceeds via sequential methylation of light olefins to higher olefins, which are easily cracked into lower olefins that are again methylated. Aromatics were supposed to be formed by secondary alkene interconversion reactions and otherwise not take part in the reactions. Later, Svelle and coworkers [53, 54] studied the methylation of [12 C]ethene, -propene and -butene with 13 C-labeled methanol at low contact times over an H-ZSM-5 catalyst. It was observed that the methylation reaction is first order in the olefin and zero order in methanol, for all olefins studied, and that the methylation rate increases in the order ethene, propene, butene. A major finding was that alkene interconversion reactions are severely suppressed by the presence of methanol. The isotopomer distribution in the products further confirmed that methylation, oligomerization, polymethylbenzene (hydrocarbon pool) formation and cracking reactions all take place in parallel in the catalyst, again underlining the extremely complicated reaction pattern that constitutes the MTH reaction.

Coke Formation Coking is a major challenge in MTH processes. When the topology of the zeolite permits, methylbenzenes may be gradually converted to polycyclic aromatics and, finally, to coke [42, 45, 55, 56]. The mechanism of coke formation is not yet fully understood. Sassi et al. [41] observed the formation of tetrahydrodimethylnaphthalene during co-reaction of methanol and various methylbenzenes over H-Beta zeolite and speculated that the second aromatic ring is formed by coupling of two isopropyl substituents on the first benzene ring. Bjørgen et al. [42] reported that the isotopomer distribution observed after testing H-Beta zeolite under MTH-like conditions was very similar for the lowest naphthalene derivative dihydrotrimethylnaphthalene and for the HeptaMB+ ion. The authors suggested that the HeptaMB+ ion may be an important intermediate not only for gaseous product formation, but also for coke formation, in this zeolite. Irrespective of the mechanism, formation of polyaromatics requires ample space in zeolitic cages or channels. The rate of coke formation varies among the various zeotype catalysts and the influences of catalyst topology, acid strength and acid site density have all been studied. Yuen et al. [23] performed comparative tests of catalysts with CHA topology with varying acid strength (SAPO-34 and SSZ-13, respectively) and acid site density. They observed that an intermediate acid site density (10% Al vs. 18% Al and 3.3% Al, respectively, in SSZ13) was advantageous for the stability of an SSZ-13 catalyst. In comparison, a SAPO-34 catalyst with the same 13.14.4.6

structure, but lower acid strength and 10% Si, was far more stable than the corresponding chabazite sample. A comparison of product selectivity showed that chabazite produced more coke deposits and more paraffins than SAPO-34. The authors suggested a relationship between acid strength and hydride transfer ability. 13.14.5

Reaction Kinetics

The MTO reactions are very fast, with an overall first-order rate constant of roughly 250 m3 gas m−3 catalyst s−1 for SAPO-34 at 450 ◦ C [57, 58]. Evidence for intracrystalline diffusion limitations has been found at crystal sizes >2.5 µm [59, 60] and for intraparticle diffusion limitations for particles >1 mm [57]. This may be the reason why the reported activation energies vary over a broad range from 35 to 90 kJ mol−1 [58, 61–63]. The reaction rates are generally reported to decline as carbon is accumulated in the zeolitic structure. Bos and Tromp [57] tested several deactivation models and found that an exponential model gave the best fit to experimental data obtained on SAPO-34, whereas Chen et al. [61] reported that a linear dependence gave the best representation of the experimental data. Later, Chen [58] reported a maximum in the conversion rate for methanol and DME at approximately 5 wt.% carbon on SAPO-34, in accord with the carbon pool model. Gayubo et al. [62] developed a kinetic model for simulation of the MTO reaction over SAPO-18 that accounted for the initiation period and the maximum in olefin production. Several kinetic models have been developed in order to describe conversion and product selectivity as a function of activity, coke content and gas-phase composition. The reaction network is complex and most researchers have applied lumped kinetics to describe the product distribution. The reaction of methanol to give DME has been found to be fast compared with the MTO reaction, and these two compounds are normally treated as a single kinetic species. For H-ZSM-5 and other zeolites with medium-sized pores, a consecutive reaction mechanism has been applied to describe the product versus residence time pattern reported for these types of catalysts (cf. Fig. 6). The development of kinetic models for HZSM-5 type catalyst systems has been reviewed by Chang [1]. The importance of secondary reactions is less pronounced for eight-membered ring structures such as SAPO-34, and a parallel mechanism approach in accordance with the carbon pool model has found wide acceptance. Bos and Tromp [57] developed a reaction scheme with 12 reactions involving six product lumps plus coke. Some of the reactions were of first order, others of second order. The effect of coke on selectivity is

13.14.7 Process Descriptions

described by introducing different deactivation rates for individual rate constants. Their reactor model suggests that the circulating fast fluidized-bed reactor and a turbulent fluidized-bed reactor are the most promising reactor systems for MTO and that a certain carbon content on the catalyst is required to obtain maximum yield of light olefins. Chen et al. [61] and Gayubo et al. [63] applied a similar approach and introduced activation energies and competing adsorption to describe the effect of temperature and water on selectivity and deactivation. Gayubo et al. found that the reaction network could be reduced to four parallel reactions without losing accuracy even at conversions up to almost 100%. The parallel reaction network has been found to describe adequately the effect of carbon content, temperature, residence time and H2 O-to-methanol ratio on selectivity and conversion. By their nature, the lumped kinetic models cannot predict the detailed product composition. As early as 1983, Mihail et al. [64] described the MTO reaction mechanism by 27 reactions and developed kinetic constants for each. Later, Park and Froment [65–67] developed detailed kinetic models at the elementary step level applying the single event kinetics approach for ZSM-5 and SAPO34 [68]. Eight rival mechanisms were considered. After discrimination between these, the surface-bound oxonium methylide mechanism proposed by Hutchings and Hunter [69] was retained. It should be noted that the carbon pool mechanism with a methylated aromatic ring as the reaction center was not among the mechanisms considered. It is debated in the literature whether product selectivity in SAPO-34 is controlled by kinetic, steric or diffusion effects. Chen et al. [59] studied the effect of crystal size on the conversion of methanol and dimethyl ether. They concluded that the selectivity was not controlled by diffusion and explained the increased ethene-to-propene ratio with increasing coke content by a transition state selectivity, suggesting that the formation of longer olefinic chains requires a larger intermediate, which is sterically hindered at high coke contents. Barger [70] plotted the ethene-to-propene ratio observed over SAPO-34 at different temperatures and pressures versus the thermodynamic equilibrium ratio of these products at the same temperature and found a linear relationship. Based on these results, Barger suggested that ethene and propene are in equilibrium in the cages of SAPO-34, but that ethene diffuses more easily out of the cages, leading to higher ethene selectivity than predicted from thermodynamic equilibrium. Dahl et al. [60] studied the effect of crystal size on the conversion of ethene and propene. They suggested that the product selectivity in SAPO-34 is controlled by diffusion limitations. Song and Haw [71] modified the cages in SAPO-34 by introducing PH3 that was subsequently converted into phosphate groups. They

2959

observed an improved ethene selectivity, which can be explained both by the transition-state selectivity suggested by Chen et al. [59] and by stronger diffusion limitations as suggested by Barger [70] and Dahl et al. [60]. 13.14.6

Reactor Technology

The characteristics of the different types of MTO catalysts discussed above are important for the choice of reactor technology and process design. A fluidized-bed reactor and regenerator system is for several reasons ideally suited for an MTO process based on SAPO-34. First, the product distribution in the MTO reaction is dependent on catalyst coke content, as illustrated in Fig. 8. In a system with fluidized-bed reactor and regenerator steady-state yields are obtained. Second, the MTO reaction and regeneration are exothermic reactions and fluidized-bed reactors offer excellent temperature control. Third, fluidizedbed reactors are well-proven industrial technology (see Chapter 10.2) and have the potential for very high singletrain capacities. Different types of fluidized-bed reactors can be used for the methanol conversion. A dense-phase reactor with continuous or periodic removal of catalyst through a slipstream to a regenerator is one alternative. By its nature, however, dense-phase fluidized-bed reactors operate at fairly low superficial gas velocities in the order of less than 1 m s−1 and low space velocities. Hence it is advantageous to minimize catalyst inventory requirements by employing a higher level of fluidization that in commercial practice can range from a circulating fast fluidized-bed reactor to a dilute-phase riser-type reactor. Broad commercial experience exists in the commercial design and utilization of circulating fast fluidized-bed or turbulent fluidized-bed reactor systems in FCC applications. Because of its flexibility, a fast fluidizedbed reactor offers significant operating advantages while reducing the catalyst inventory. Lurgi’s MTP process is based on a medium-pore-sized catalyst (MFI type) and, as discussed above, this catalyst deactivates more slowly than a catalyst based on SAPO-34. This led to the choice of a system with several fixed-bed reactors with interstage direct cooling for temperature control. 13.14.7

Process Descriptions UOP/Hydro MTO Process A simplified flow scheme for the UOP/Hydro MTO process is shown in Fig. 13. Methanol is fed to a reactor 13.14.7.1

References see page 2964

2960

13.14 Methanol-to-Hydrocarbons

where the conversion of methanol and DME to light olefins proceeds to completion (>99% conversion) at a very short residence time. The reactor operates in the vapor phase at temperatures between 350 and 600 ◦ C and pressures between 1 and 6 bar. After the oxygenate recovery section, the effluent is further processed in the fractionation and purification section to separate the key products from the byproduct components. This section is very similar to the corresponding section in a conventional steam cracker. Ethene and propene are produced as polymer-grade products and sent to storage. In the UOP/Hydro MTO process the ethene content in the C2 fraction and the propene content in the C3 fraction are about 97%, hence for chemical-grade ethene or propene production the splitter columns are not required. The UOP/Hydro MTO process design includes recovery of oxygenates that are produced in minor quantities for recycle conversion to olefins. The remaining trace oxygenated impurities are removed from the product streams using proven technology that has been in commercial operation for many years, including UOP’s Oxygenate Removal Unit process. The UOP/Hydro MTO process offers a wide range of flexibility for altering the relative amounts of ethene and propene products by adjusting the operating severity in the reactor. The process can be designed for an ethene/propene ratio between 0.75 and 1.5 [2, 5, 72] and a single-train capacity greater than 106 t a−1 of light olefins. The SAPO-34-based catalyst has been produced on a commercial scale. A typical mass balance for a unit producing 500 000 t a−1 of ethene and propene each is shown in Quench Reactor Regenerator tower

Caustic wash

Table 1. Approximately 3 t of methanol are required per ton of light olefins or approximately 8700 t day−1 . This represents a carbon-based yield of almost 80%. Integration of MTO with Olefin Cracking The mass balance in Table 1 shows that a significant amount of C4+ is produced in the process. UOP and Total Petrochemicals have jointly developed an Olefin Cracking Process (OCP) that upgrades C4+ to ethene and propene. By integrating MTO with OCP, combined ethene and propene yields close to 90% and propene-to-ethene ratios >2 can be obtained [5]. This also leads to a reduction in the specific methanol consumption per ton of light olefins from approximately 3 to 2.7. 13.14.7.2

Lurgi MTP Process Figure 14 shows a simplified process flow diagram for Lurgi’s MTP process [7, 73]. Methanol is sent to an adiabatic DME pre-reactor where methanol is converted 13.14.7.3

Tab. 1

Material balance for the UOP/Hydro MTO process Feed/103 t a−1

Methanol Ethene Propene Butenes C5+ hydrocarbons Fuel gas Other (COx , water, coke) Totals

De-C2

De-C1

C2 splitter

Products/103 t a−1

3016 500 500 167 70 47 1732 3016

3016

C3 De-C3 splitter

De-C4

Tail gas Ethene

Regeneration gas

Propene

Dryer

DME recovery

Mixed C4

C2H2 Reactor C5+

Air

Water

Propane Ethane

Methanol

Fig. 13

Simplified process flow scheme of the UOP/Hydro MTO process.

13.14.8 Development and Commercial Status of the MTH Processes

Methanol

MTP reactors (2 operating + 1 reg.)

DME prereactor

2961

Optional: Fuel

Ethene

Propylene Product conditioning LPG

Gasoline Product fractionation

Olefin recycle Water recycle Process water

Fig. 14

Simplified process flow scheme for Lurgi’s MTP process [7]. Figure provided by courtesy of Lurgi AG.

to DME and water. Nearly thermodynamic equilibrium is achieved. The methanol–water–DME stream is routed to the MTP reactor together with steam and recycled olefins. Methanol and DME are more than 99% converted with propene as the predominant hydrocarbon product and the balance consisting mostly of gasoline-range hydrocarbons. Process conditions in the five or six catalyst beds per reactor are chosen so as to obtain similar reaction conditions and maximum overall propene yield. Introducing small streams of fresh feed between the beds controls the conditions. Two reactors are operating in parallel while the third one is in the regeneration or stand-by mode. Regeneration is necessary after about 500–600 h of cycle time when the active catalyst centers become blocked by coke formed in the side reactions. By using diluted air, the regeneration is performed at mildest possible conditions, nearly at operating temperature, thus avoiding thermal stress on the catalyst. The product mixture from the reactors is cooled, and the product gas, organic liquid and water are separated. The product gas is compressed and traces of water, CO2 and DME are removed. The cleaned gas is then further processed yielding chemical-grade propene with a typical purity of more than 97%. Per pass yields of ethene plus propene are about 50% with a ZSM-5-type catalyst, several olefin-containing streams are sent back to the main synthesis loop for their recycle conversion leading to an additional propene production. In this respect, the Lurgi MTP process combines the C4+ olefin cracking with its MTP operation. To avoid accumulation of inert materials in the loop, a small purge is required for light-ends and the C4 /C5 cut. Gasoline is produced as a by-product. Water is recycled to steam generation for the process. A typical mass balance for the process is given in Table 2.

According to Lurgi [73], the basic process design data were derived from extensive pilot plant work and also from a demonstration unit with continuous operation with real methanol feedstock. This work was done in cooperation with Statoil at their methanol plant at Tjeldbergodden, Norway. ExxonMobil MTO Technology ExxonMobil has not published much information on their MTO technology, aside from patent applications. From the information in these applications it can be concluded that the process is based on SAPO-34 as active component in the MTO catalyst and that fluidized-bed reactor technology is used. This is confirmed in Ref. [4]. 13.14.7.4

13.14.8

Development and Commercial Status of the MTH Processes

At the time of writing this chapter, several commercial projects have been announced for the UOP/Hydro and Lurgi MTP process and there are also confidential projects being studied. It is considered outside the scope of this review to discuss these potential commercial projects, and emphasis in this section is given to a brief summary of the development status of the different MTH processes. Mobil MTG Mobil’s MTG process was commercialized in New Zealand in 1985 [1]. The plant was later closed for economic reasons. No other commercial plans for this technology have been announced. Mobil also operated 13.14.8.1

References see page 2964

2962 Tab. 2

13.14 Methanol-to-Hydrocarbons Material balance for Lurgi’s MTP process Feed/103 t a−1

Methanol Ethene Propene LPG Gasoline Fuel gas Other (COx , water, coke) Totals

Products/103 t a−1

1670

1670

– 471 41 185 34 939 1670

a fluidized-bed plant in Wesseling, Germany, with the purpose of demonstrating MTO technology based on a ZSM-5 catalyst. UOP/Hydro MTO Process An MTO demonstration plant has been operated at Hydro’s research facilities in Porsgrunn, Norway, since 1995. The plant is a fully integrated reactor and regenerator unit. The original design capacity of the unit was 0.5 t day−1 of methanol, but with process and catalyst optimization the plant capacity has been increased to 1 t day−1 . One of the major goals of the plant was to demonstrate the stability of the process with a commercially produced MTO catalyst and crude methanol from a commercial plant. Excellent product stability was demonstrated [2, 72]. Product quality is a concern for any new olefin technology. Work in the MTO demonstration plant in Norway has shown that no unique or unusual contaminants are found in the MTO effluent and the effluent only requires conventional processing to produce polymer-grade ethene and propene products. Regarding oxygenates, a direct side-by-side comparison of oxygenates present in the effluent from an MTO unit and from an LPG cracker has confirmed that the same oxygenates are present in both streams. The levels of acetylenic and diolefinic compounds in the effluent from an MTO unit are significantly lower than those from a cracker, simplifying the product recovery section. The UOP/Hydro MTO process is offered for license. 13.14.8.2

Lurgi MTP Process Lurgi has carried out a demonstration project for the MTP process. A demonstration plant was constructed at Tjelbergodden, Norway, in 2001 using feedstock from Statoil’s methanol plant. The demonstration unit was started up in January 2002, and in September 2003 the unit had completed the 8000-h life-cycle test. According to Lurgi [7, 73], it was demonstrated that the catalyst meets the commercial target of 8000 h on-stream. Deactivation 13.14.8.3

rates of methanol conversion decreased with operation time, propene selectivity and yields were in the expected range and polymer-grade propene was produced. The catalyst is produced on a commercial scale. The process is offered for license. ExxonMobil According to Ref. [4], ExxonMobil has a large-scale MTO pilot unit at the Baytown, USA, research center. This is a complete demonstration, from methanol feed to the polymerization of ethene and propene. 13.14.8.4

13.14.9

GTO Project Scenarios

Several commercial MTO and MTP projects are at various stages of development at this time. MTO (MTP) projects are often closely tied to the utilization of remote natural gas generally located far from large olefin or olefin derivative markets. Therefore, the gas must be converted into an easily transportable product to be cost-effective. It follows that MTO (MTP) projects generally fall into two scenarios, viz. integrated gas-to-polymer (GTP) plants located close to the supply of remote gas or ‘‘segregated’’ gas-to-methanol and methanol-to-olefins or methanol-topolymer plants. For the integrated projects the objective is to produce polymer pellets because these are commonly transported solids and can be used domestically or exported. For the segregated projects the objective is to convert the gas to methanol and ship the methanol to an MTO (MTP) plant located in proximity to the olefin or olefin derivative markets. Either approach can be economical, so the choice depends on other factors, such as: • • • • •

locations of target markets for products location of remote gas source project ownership structure project finance structure and incentives transportation and delivery costs.

13.14.10

MTO and GTO Economics

MTO and MTP economics are not discussed in much detail in this chapter since economics are very sitedependent and the main focus here is catalysis. However, presentations by UOP/Hydro [5, 6] and Lurgi [7] clearly show that GTO concepts based on lower cost natural gas (less than US$2 GJ−1 or ca. US$2 mmbtu−1 ) are very competitive with conventional steam crackers for olefin production. In order to illustrate the competitiveness of an MTO plant with a naphtha cracker, Fig. 15 was developed

13.14.12 Conclusions

US$ t−1 0

100

200

300

400

500

100 % 90 % 80 % 70 %

ROI

60 % 50 % 40 % 30 % 20 % 10 %

2963

These economics are, of course, generic and simplified. Furthermore, it should be emphasized that methanol is an intermediate in the GTO chain and that gas cost could be used instead of methanol feedstock cost in such an analysis, since the methanol production cost always will be directly linked to the gas price in a GTO chain. New large-scale plants based on remote gas can produce methanol at the costs indicated in Fig. 15 with gas prices below approximately US$2 GJ−1 (ca. US$2 mmbtu−1 ). Therefore, the crude oil prices of US$290–440 t−1 (US$40–60 bbl−1 ) experienced in 2004–06 have led to a very significant commercial interest in GTO, MTO and MTP.

0% 0

10

20

30

40

50

60

70

CO2 Aspects

Crude oil benchmark / US$ bbl−1 MTO @ 110 US$ t−1 MTO @ 130 US$ t−1 MTO @ 150 US$ t−1

Naphtha cracker MTO @ 70 US$ t−1 MTO @ 90 US$ t−1

13.14.11

Fig. 15 Light olefins production. MTO vs. naphtha cracking economics (ROI, return on investments).

by Andersen et al. [6]. The figure shows return on investments (ROI) for a naphtha cracker as a function of crude oil price. This curve is based on cracker investment, operating costs and historical relationships between the prices of crude oil, naphtha, olefins and by-products. The figure also shows ROI for an MTO unit with different methanol feedstock costs. These curves are based on the same prices for olefins and by-products as used for the naphtha cracker curve. The results illustrate how different methanol feedstock costs give different breakeven economics with the naphtha cracker. With crude oil prices above US$145–160 t−1 (US$20–22 bbl−1 ), the MTO plant is competitive based on a methanol feedstock price in the order of US$90–100 t−1 . If future crude oil prices average in the US$220–360 t−1 (US$30–50 bbl−1 ) range, the MTO process becomes very attractive.

CO2 emissions are an important issue. Compared with a naphtha cracker, an MTO unit emits approximately 1 t of CO2 less per ton of light olefins produced (see Fig. 16) [5]. However, if the methanol production is included, the overall CO2 emissions are only slightly lower for GTO. 13.14.12

Conclusions

New technologies for conversion of methanol to hydrocarbons are on the verge of being commercialized, in particular for the production of light olefins such as ethene and propene through the UOP/Hydro MTO process and the Lurgi MTP process. The main drivers for the emergence of these new technologies are the search for alternative, lower cost feedstocks for olefin production, monetization of remote natural gas and developments in methanol technology. Both technologies are based on innovative catalyst systems. The strong industrial interest in MTH technology has stimulated academia to undertake extensive fundamental research in MTH chemistry, References see page 2964

t CO2 / t (C2 + C3)

1.4 1.2 1

MTO

0.8

MeOH

0.6 0.4 0.2 0 Steam cracker

Fig. 16

CO2 emissions for olefin production technologies.

Propane de hydrogenation

GTO

2964

13.14 Methanol-to-Hydrocarbons

in particular reaction mechanisms. This work has led to the development of the hydrocarbon pool reaction mechanism that is now gaining more and more acceptance. It is expected that the importance of MTH technologies will increase in the future, as natural gas and coal become more and more important petrochemical feedstocks. References 1. C. D. Chang, in Handbook of Heterogeneous Catalysis, 1st Ed., G. Ertl, H. Kn¨ozinger, J. Weitkamp (Eds.), Vol. 4, VCH, Weinheim, 1997, p. 1894. 2. B. V. Vora, T. L. Marker, P. T. Barger, H. R. Nilsen, S. Kvisle, T. Fuglerud, in Natural Gas Conversion IV, M. de Pontes, R. L. Espinoza, C. P. Nicolaides, J. H. Scholz, M. S. Scurrel (Eds.), Studies in Surface Science and Catalysis, Vol. 107, Elsevier, Amsterdam, 1997, p. 87. 3. M. Rothaemel, H.-D. Holtmann, in Proceedings of the 13th Ethylene Producers’ Conference, American Institute of Chemical Engineers, New York, 2001, p. 37. 4. ExxonMobil, 2004 Technology Webcast, August 31, 2004. 5. J. Q. Chen, A. Bozzano, B. Glover, T. Fuglerud, S. Kvisle, Catal. Today 2005, 106, 103. 6. J. M. Andersen, Presented at CMAI World Methanol Conference, Miami, FL, December 12–14, 2005. 7. W. Liebner, Lurgi, Presented at Propylene Trade and Derivatives Markets, Singapore, October 24–25, 2005. 8. R. Gist, Presented at CMAI World Methanol Conference, Miami, FL, December 12–14, 2005. 9. B. V. Vora, P. Pujado, in Encyclopedia of Chemical Processing and Design, S. Lee (Ed.), Marcel Dekker, New York, 2005, p. 379. 10. D. McCaskill, Presented at the CMAI World Methanol Conference, Miami, FL, December 12–14, 2005. 11. S. L. Meisel, J. P. McCullough, C. H. Lechthaler, P. B. Weisz, Chemtech 1976, 6, 86. 12. C. D. Chang, A. J. Silvestri, J. Catal. 1977, 47, 249. 13. C. D. Chang, Catal. Rev.-Sci. Eng. 1983, 25, 1. 14. C. D. Chang, Catal. Rev.-Sci. Eng. 1984, 26, 323. 15. G. F. Froment, W. J. H. Dehertog, A. J. Marchi, Catalysis 1992, 9, 1. 16. M. St¨ocker, Microporous Mesoporous Mater. 1999, 29, 3. 17. C. D. Chang, in Shape-Selective Catalysis: Chemicals Synthesis and Hydrocarbon Processing, C. Song, J. M. Garc´es, Y. Sugi (Eds.), ACS Symposium Series, Vol. 738, American Chemical Society, Washington, DC, 2000, p. 96. 18. J. F. Haw, Phys. Chem. Chem. Phys. 2002, 4, 5431. 19. J. F. Haw, W. Song, D. M. Marcus, J. B. Nicholas, Acc. Chem. Res. 2003, 36, 317. 20. S. W. Kaiser, US Patent 4 499 327, assigned to Union Carbide, 1985. 21. S. W. Kaiser, Arabian J. Sci. Eng. 1985, 361. 22. F. A. Wunder, E. I. Leupold, Angew. Chem. 1980, 92, 125. 23. L.-T. Yuen, S. I. Zones, T. V. Harris, E. J. Gallegos, A. Auroux, Microporous Mater. 1994, 2, 105. 24. R. Wendelbo, D. Akporiaye, A. Andersen, I. M. Dahl, H. B. Mostad, Appl. Catal. A: General 1996, 142, 197. 25. S. Wilson, P. Barger, Microporous Mesoporous Mater. 1999, 29, 117.

26. P. B. Venuto, P. S. L. Landis, in Advances in Catalysis, D. D. Eley, H. Pines, P. B. Weisz (Eds.), Vol. 18, Academic Press, London, 1968, p. 259. 27. Y. Ono, T. Mori, J. Chem. Soc., Faraday Trans. 1 1981, 77, 2209. 28. J. P. van den Berg, J. P. Wolthuizen, J. H. C. van Hooff, in Proceedings of the 5th International Conference on Zeolites, L. V. C. Rees (Ed.), Heyden, London, 1980, p. 649. 29. J. K. A. Clarke, R. Darcy, B. F. Hegarty, E. O’Donoghue, V. Amir-Ebrahim, J. J. Rooney, J. Chem. Soc., Chem. Commun. 1986, 425. 30. I. M. Dahl, S. Kolboe, Catal. Lett. 1993, 20, 329. 31. I. M. Dahl, S. Kolboe, J. Catal. 1994, 149, 458. 32. I. M. Dahl, S. Kolboe, J. Catal. 1996, 161, 304. 33. J. F. Haw, Phys. Chem. Chem. Phys. 2002, 4, 5431. 34. M. Seiler, U. Schenk, M. Hunger, Catal. Lett. 1999, 62, 1999. 35. M. Hunger, M. Seiler, A. Buchholz, Catal. Lett. 2001, 74, 61. 36. T. Mole, G. Bett, D. Seddon, J. Catal. 1983, 84, 435. 37. T. Mole, in Methane Conversion, D. M. Bibby, C. D. Chang, R. F. Howe, S. Yurchak (Eds.), Studies in Surface Science and Catalysis, Vol. 36, Elsevier, Amsterdam, 1988, p. 145. 38. Ø. Mikkelsen, P. O. Rønning, S. Kolboe, Microporous Mesoporous Mater. 2000, 40, 95. 39. B. Arstad, S. Kolboe, Catal. Lett. 2001, 71, 209. 40. B. Arstad, S. Kolboe, J. Am. Chem. Soc. 2001, 123, 8137. 41. A. Sassi, M. A. Wildman, H. J. Ahn, P. Prasad, J. B. Nicholas, J. F. Haw, J. Phys. Chem. B 2002, 106, 2294. 42. M. Bjørgen, U. Olsbye, S. Kolboe, J. Catal. 2003, 215, 30. 43. W. von E. Doering, M. Saunders, H. G. Boyton, H. W. Earhart, E. F. Wadley, W. R. Edwards, G. Laber, Tetrahedron 1958, 4, 178. 44. W. Song, J. B. Nicholas, A. Sassi, J. F. Haw, Catal. Lett. 2002, 81, 49. 45. M. Bjørgen, U. Olsbye, D. Petersen, S. Kolboe, J. Catal. 2004, 221, 1. 46. M. Bjørgen, F. Bonino, S. Kolboe, K. P. Lillerud, A. Zecchina, S. Bordiga, J. Am. Chem. Soc. 2003, 125, 15863. 47. M. Bjørgen, F. Bonino, B. Arstad, S. Kolboe, K. P. Lillerud, A. Zecchina, S. Bordiga, ChemPhysChem 2005, 6, 232. 48. T. Mole, J. A Whiteside, D. J. Seddon, J. Catal. 1983, 82, 261. 49. B. Arstad, J. B. Nicholas, J. F. Haw, J. Am. Chem. Soc. 2004, 126, 2991. 50. R. F. Sullivan, C. J. Egan, G. E. Langlois, R. P. Sieg, J. Am. Chem. Soc. 1961, 83, 1156. 51. R. M. Dessau, R. B. LaPierre, J. Catal. 1982, 78, 136. 52. R. M. Dessau, J. Catal. 1986, 99, 111. 53. S. Svelle, P. O. Rønning, S. Kolboe, J. Catal. 2004, 224, 115. 54. S. Svelle, P. O. Rønning, U. Olsbye, S. Kolboe, J. Catal. 2005, 234, 385. 55. H. Fu, W. Song, J. F. Haw, Catal. Lett. 2001, 76, 89. 56. D. M. Marcus, W. Song, L. L. Ng, J. F. Haw, Langmuir 2002, 18, 8386. 57. A. N. R. Bos, P. J. J. Tromp, Ind. Eng. Chem. Res. 1995, 34, 3808. 58. D. Chen, PhD Thesis, Technical University of Trondheim, 1998. 59. D. Chen, K. Moljord, T. Fuglerud, A. Holmen, Microporous Mesoporous Mater. 1999, 29, 191. 60. I. M. Dahl, R. Wendelboe, A. Andersen, D. Akporiaye, H. Mostad, T. Fuglerud, Microporous Mesoporous Mater. 1999, 29, 159. 61. D. Chen, H. P. Rebo, A. Grønvold, K. Moljord, A. Holmen, Presented at the 6th World Congress of Chemical Engineering, Melbourne, 23–27 September, 2001.

13.15.2 FT Synthesis as a Source of Fuels and Chemicals and its Viability 62. A. G. Gayubo, A. T. Aguayo, A. Alonso, A. Atutxa, J. Bilbao, Catal. Today 2005, 106, 112. 63. A. G. Gayubo, A. T. Aguayo, A. E. Sanchez del Campo, A. M. Tarrio, J. Bilbao, Ind. Eng. Chem. Res. 2000, 39, 292. 64. R. Mihail, S. Straja, G. Maria, G. Musca, G. Pop, Ind. Eng. Chem. Process Des. Dev. 1983, 22, 532. 65. T.-Y. Park, G. F. Froment, Ind. Eng. Chem. Res. 2001, 40, 4172. 66. T.-Y. Park, G. F. Froment, Ind. Eng. Chem. Res. 2001, 40, 4187. 67. T.-Y. Park, G. F. Froment, Ind. Eng. Chem. Res. 2004, 43, 682. 68. S. M. Alwahabi, G. F. Froment, Ind. Eng. Chem. Res. 2004, 43, 5098. 69. G. J. Hutchings, R. Hunter, Catal.Today 1990, 6, 279. 70. P. Barger, in Zeolites for Cleaner Technologies, M. Guisnet, J.-P. Gilson (Eds.), Catalysis Science Series 3, Imperial College Press, Danvers, 2002, pp. 239–260. 71. W. Song, J. F. Haw, Angew. Chem. Int. Ed. 2003, 42, 892. 72. S. Kvisle, H. Reier Nilsen, T. Fuglerud, A. Grønvold, B. V. Vora, P. R. Pujado, P. T. Barger, J. M. Andersen, Erd¨ol Erdgas Kohle 2002, 118, 361. 73. Lurgi A. G, personal communication to S. Kvisle, January 2006.

2965

13.15.2

FT Synthesis as a Source of Fuels and Chemicals and its Viability

The Fischer–Tropsch synthesis is the catalytic hydrogenation of CO to give a range of products, which can be used for the production of high-quality diesel fuel, gasoline and linear chemicals such as 1-alkenes, alkanes and oxygenated hydrocarbons. In 1902, Sabatier and Senderens [1] reported that when syngas was passed over a nickel catalyst, methane was produced. The German firm BASF found in 1913 that when operating at higher pressures, some oil was produced [2]. In 1923, two German chemists, Fischer and Tropsch, published their famous work using alkalized iron catalysts [3]. In 1933, Ruhrchemie constructed the first FT plant using a nickel catalyst at atmospheric pressure. By 1938 there were nine operating FT plants in Germany using cobaltbased catalysts at about 0.1 MPa. Their total capacity was about 0.66 million tons per year. FT plants were also briefly operated in Japan, Manchuria and France. After the Second World War, all these plants were shut down as they were uneconomic. A fuller early history of the FT synthesis can be found elsewhere [4–6] (in Ref. [6], see Chapter 1). The development of the FT process after 1945 is discussed in Sections 13.15.3 and 13.15.4.

World-wide, liquid fuels and the bulk of all chemicals are mainly produced from crude oil. Table 1 shows an approximate comparison of the known carbon sources relative to oil. At the current world consumption rate, crude oil supplies may last for about 50 years. Nevertheless, for both political and economic reasons, alternative energy sources could be utilized to a greater extent in the near future. For instance, natural gas is more widely distributed over the world than are the crude oil deposits, and so the supply is less sensitive to local political and other upheavals. Oil is being produced from the tar sand deposits in Alberta, Canada. The extraction of oil from shale was demonstrated at a plant in Colorado, which operated for several years before being shut down for economic reasons. The oils from tar sand or shale are highly aromatic and so are not suitable for the production of either linear alkenes or of high-quality diesel fuel. Both of these products are readily produced via the FT synthesis. Due to the higher efficiency of diesel engines relative to gasoline-powered engines, the demand for diesel fuel has increased and is projected to continue to do so (see Section 13.15.10 for other advantages of FT fuels). The syngas required for the FT synthesis is produced by the gasification of either coal or methane. This is a highly endothermic process, and the required energy is provided by oxygen combustion of part of the carbon source. As methane has four or more times the number of hydrogen atoms per carbon atom, the use of methane is more efficient. Typically in the case of methane, about 20% of the carbon is converted to CO2 whereas for coal it is about 50%. As syngas production can account for up to 70% of the capital cost of the total overall process, the efficiency of syngas production is clearly important. To be economically viable, natural gas or coal must be available at low cost. In the case of natural gas, this applies to stranded gas (gas that is far from the market, thus making transport uneconomical) and associated gas (gas that is co-produced with crude oil and which is currently simply being flared). In the long term, when considering the huge amount of coal reserves compared with the other carbon reserves (see Table 1), an increase in the use of coal may be inevitable. The FT process has to compete directly with crude oil and so the price of crude is of vital importance. The variation in the price of crude oil is depicted in Fig. 1. A few years ago it was estimated that the FT process would be economically viable when the price of crude was

∗

References see page 2992

13.15

The Fischer–Tropsch (FT) Synthesis Processes Mark E. Dry∗

13.15.1

Brief History

Corresponding author.

13.15 The Fischer–Tropsch (FT) Synthesis Processes World relative to oil

Tab. 1

Source Crude oil Tar sands Shale oil Natural gas Coal

reserves

of

‘‘carbon’’

Reserves, oil equivalent 1.0 0.7 1.2 1.5 26

US$ / ton

400

50

300

40

200

30 20

100

10

US$ / barrel

2966

0 0 1940 1960 1980 2000 2020 Year

Variation of the price of crude oil. (The density of crude oil was taken as 0.858 t m−3 .) The right hand ordinate is in US$ per barrel. Note that in the mid-1980s the price was over US$30 per barrel, in the late 1990s it was about US$10, by 2004 it had risen to over US$50 and during 2005 it peaked at US$70 per barrel.

Fig. 1

above about US$20 per barrel [7]. The history of the three Sasol FT plants in South Africa shows the importance of the oil price. In the 1940s there was a world-wide perception that due to the limited amount of the known crude oil reserves the price would rise with time as demand increased. Based on this, and also for political reasons, it was decided to build the Sasol One plant. Before the completion of the plant, however, huge new deposits of crude oil were discovered in the Middle East and so the price of crude did not rise as anticipated. In the 1970s, the price of crude did start to rise sharply, and it was decided to build two more much larger plants which came on-line in the early 1980s when the price of crude was over $30 per barrel. In 2004 the price of crude rose to over $50 per barrel and during 2005 it peaked at about $70 per barrel. As it is not foreseen that it will drop below about US$40 per barrel in the future, the construction of FT plants world-wide could be viable, given low-cost methane or coal. 13.15.3

Laboratory and Commercial FT Reactors

The FT reaction is highly exothermic: the heat released per ‘‘CH2 ’’ is about 145 kJ mol−1 , which is much higher than for typical reactions in oil refining operations. Because of this, it is important that there is efficient transfer of heat

from the catalyst particles to the surrounding medium. Should this not be the case, the particle temperature would increase, resulting in an increase in the selectivity of the undesired methane (see Section 13.15.7.3). In addition, the catalyst would lose activity due to sintering and fouling, as discussed in Section 13.15.5.4. To ensure effective heat transfer in a two-phase system, an important requirement is that the gas linear velocity is high. In laboratory reactors this if often not the case, which can lead to erroneous selectivity and kinetic results since the actual catalyst temperature could be higher than the measured gas temperature. If the gas velocity is increased to improve heat transfer, the conversion levels will be low and so the partial pressures of the reactants and products will be very different from those in high-conversion units. This could mean that neither the overall selectivity nor the kinetics of commercial reactors can be predicted from lowconversion laboratory reactor runs. For laboratory studies it is recommended that Berty or slurry reactors are used. In Berty reactors, the conversion levels can be varied by altering the fresh feed flow-rate (the high internal recycle ensures a high gas velocity through the bed). When a considerable amount of wax is produced, the reactors are three-phase systems. Because the liquid wax is both inside and outside the catalyst particles, overheating is less likely and so either fixed or slurry bed laboratory reactors can be used. The low-temperature FT (LTFT) produces mainly waxes and the high- temperature FT (HTFT) produces gasoline and alkenes. The reactor types used in commercial FT operation are depicted in Figs. 2 and 3. In the multitubular reactor (Fig. 2) the catalyst particles are packed in narrow tubes (diameters about 3–5 cm). This ensures high gas velocities at the gas feed flows used and thus high heat transfer rates. The narrower the tubes, the higher are the heat transfer rates. Since cobalt-based catalysts are more active than iron-based catalysts (see Section 13.15.6), the tube diameters for cobalt catalysts need to be narrower. Multitubular or slurry bed (Fig. 3c) reactors are used at ‘‘low’’ temperatures (190–250 ◦ C) when the production of wax is the main objective. Tubular reactors are simpler to operate and, since the liquid wax simply trickles down the catalyst bed, there is no need for a wax–catalyst separation unit. The latter is essential for slurry bed reactors as the small catalyst particles used are suspended in the liquid wax medium. In the event of a breakthrough of some catalyst poison such as H2 S, only the catalyst at the bed entrance of the fixed-bed reactors will be deactivated, whereas in a slurry reactor all of the catalyst will be affected (see Section 13.15.5.4). Slurry bed reactors, however, have several advantages: (a) lower construction costs, (b) lower compression costs due to lower pressure drops over the reactor, (c) lower catalyst loadings required for the same

13.15.3 Laboratory and Commercial FT Reactors

Gas inlet

Steam heater

Steam collector

Steam outlet Feed water inlet

Inner shell Tube bundle

Gas outlet

Wax outlet

Fig. 2

Multitubular reactor for LTFT operation.

production because of the higher activity of the smaller particles, (d) more isothermal beds and (e) longer reactor runs because on-line removal of used catalyst and the addition of fresh catalyst can be applied. Multitubular

References see page 2992

Product gases

Hopper

Cyclones Fludised bed

Steam

Slurry bed Boiler feed water Wax

Steam

Boiler feed water

Standpipe Slide valve Gas in (a)

reactors are not used at temperatures above about 280 ◦ C as carbon deposition, when using iron catalysts, would result in reactor plugging. Two-phase fluidized bed reactors (Fig. 3a and b) are used with iron catalysts at about 340 ◦ C when the objective is the production of alkenes and gasoline. Conditions resulting in the production of wax should be avoided, as the presence of liquid wax on the outside of the catalyst particles will result in defluidization of the catalyst bed. Because of the turbulent nature of fluidized-bed reactors at the gas velocities used, the rate of heat exchange from the catalyst to the heat exchanger tubes is very high, resulting in near isothermal catalyst beds. The combination of high rates of heat exchange and high temperatures means that high conversions and high syngas throughputs can be achieved. The fixed fluidized bed (FFB) reactor is preferred over the circulating fluidized bed (CFB) for the following reasons: an FFB reactor is cheaper to construct; because the reaction section is wider, more cooling coils can be installed, which increases its syngas conversion capacity; and the compression costs are lower because of the lower gas velocities. More details about the various commercial reactors, their development and their modes of operation can be found elsewhere [6, 8] (in Ref. [6], see Chapter 2). In commercial operations, the exit gas is cooled to condense out the water and the heavier hydrocarbon fractions. It is common practice to recycle a portion of the remaining gas back to the reactor entrance. This results in higher overall conversion of the syngas. Recycling also increases the gas velocity through the reactor and hence the heat exchange rate. Recycling also lowers the partial pressure of water in the reactor,

Products

Out

Catalyst

2967

(b)

Gas distributor

Gas distributor

Total feed

Syngas in (c)

Fluidized bed FT reactors: (a) and (b) are for HTFT operation. (a) High gas velocity CFB (circulating fluidized bed) unit; (b) FFB (fixed fluidized bed, i.e. an ebullating bed) unit; (c) three-phase slurry bed unit for LTFT operation. In the HTFT reactors all the products exit the reactors as gases whereas in the LTFT reactors the wax produced leaves as a liquid. Note that the units are not drawn to the same scale: the CFB unit is much taller than the other two types.

Fig. 3

2968

13.15 The Fischer–Tropsch (FT) Synthesis Processes

which can have a negative effect on the kinetics (see Section 13.15.6). An important advantage of using fluidized-bed reactors at high temperatures, where the rates of catalyst deactivation inevitably are higher, is that deactivated catalyst can be removed and fresh catalyst added on-line, which means that long continuous runs can be achieved. 13.15.4

Commercial FT Plants 13.15.4.1

Past and Present Plants (Post-1945)

Brownsville, TX, USA During the 1950s Carthage-Hydrocol constructed and operated a plant at Brownsville, TX. Syngas was produced from natural gas. The capacity of the plant was about 360 × 103 t a−1 . The FT reactor was a two-phase fluidized bed reactor developed by Hydrocarbon Research [9]. After a few years, the plant was shut down for economic reasons. 13.15.4.1.1

Rentech, Pueblo, CO, USA Fuelco operated a small plant (designed by Rentech) at Pueblo, CO. The syngas was produced by reforming landfill methane. The FT synthesis was performed in a slurry phase reactor using an iron-based catalyst [10, 11]. The capacity was about 7500 t a−1 . The plant was dismantled in 1996. 13.15.4.1.2

13.15.4.1.3 Sasol, South Africa [6–8] The first Sasol plant came on stream in 1955. Syngas was produced from coal. Two types of FT reactors were used: the HTFT CFB reactors, developed by Kellogg, produced gasoline and linear 1-alkenes; the LTFT multitubular Arge reactors, developed by Lurgi/Ruhrchemie, produced linear alkanes and waxes. In later years the CFB reactors were replaced by a single LTFT slurry phase reactor, which has the same capacity as the original five Arge reactors. Currently syngas is being produced by reforming natural gas. The much larger second and third Sasol HTFT plants came on-stream in 1980 and 1982, respectively. Coal gasification provides the bulk of the syngas, the balance being produced by catalytic autothermal reforming of the methane from the FT tailgas. Figure 4 depicts the plant layout. The original CFB reactors have all been replaced with the more efficient FFB reactors. The capacity of the Sasol plants is about 7500 × 103 t a−1 . A detailed description of the Sasol plants and of the various FT reactors is given elsewhere [6, 8] (in Ref. [6], see Chapter 5). 13.15.4.1.4 PetroSA, South Africa In 1991, the PetroSA (Mossgas) plant came on stream. Syngas is produced by catalytic reforming of offshore natural gas. The HTFT reactors are CFB units. After recovering the desired products, the remaining tailgas is recycled to the autothermal reformer. The total fuel production is about 1020 × 103 t a−1 [6, 8].

Syngas

HTFT SAS reactors

O2 Autothermal reformers

H2

Steam

CH4 Water

Fuel gas

Separation C4

Oxygenates Acids

C 2H 4 C 3H 6

C5-C10 Diesel

Recovery Olig. Extraction LPG Water to bioworks

Separation purification

Hydrofine

I-alkenes (C5 to C8)

Gasoline

Pt reform

Diesel

Gasoline

General layout of the Sasol FT plant at Secunda, South Africa. The linear C5 –C8 1-alkenes are utilized in the production of linear 1-alcohols and also as copolymers in the production of polyethylene.

Fig. 4

13.15.4 Commercial FT Plants

2969

Natural gas

POX SGP

O2

LTFT (Multitube)

Reform

Steam

Tail gas Separation H-crack H-lsom.

Oils Hydrofine

Naphtha

Diesel

Wax Naphtha Diesel / Kerosene

Kerosene

General layout of the Shell plant at Bintulu, Malaysia. POX SGP is the partial oxidation Shell gasification process. H-crack and H-isom are hydrocracking and hydroisomerization operations.

Fig. 5

13.15.4.1.5 Shell, Bintulu, Malaysia [12, 13] In 1993, the Shell SMDS (Shell middle distillate synthesis) plant came on-line at Bintulu, Malaysia. The layout of the plant is shown in Fig. 5. Offshore natural gas is converted to syngas by non-catalytic partial oxidation. Since the H2 /CO ratio of the product gas is 1.7, which is lower than that required for the FT synthesis, the additional required hydrogen is produced by catalytic reforming of the FT tailgas. The FT units are large multitubular reactors loaded with a cobalt-based catalyst. The production capacity of the plant is about 500 × 103 t a−1 . 13.15.4.2

FT Plants Being Constructed or Planned

13.15.4.2.1 Qatar [14] The natural gas from Qatar’s North Field is estimated to contain 9% of the world’s reserves. An FT complex came on line in 2007. The planned capacity is about one million tons of product per year. The technologies that are utilized are Haldor Topsøe for gas reforming, Sasol for slurry phase FT synthesis, using a cobalt-based catalyst, and ChevronTexaco for hydroprocessing. The main product will be diesel fuel (see Section 13.15.8.2). 13.15.4.2.2 Escravos, Nigeria [14] Nigeria intends to eliminate the flaring of natural gas which is associated with the recovery of crude oil. An FT complex at Escravos in the Niger Delta is being constructed. The capacity and the technologies are the same as those at Qatar. 13.15.4.2.3 Exxon [15–17] Exxon has developed their AGC-21 process (Advanced Gas Conversion for the

21st century) for converting natural gas to liquid fuels. The process involves methane reforming, slurry phase FT synthesis with a cobalt-based catalyst and hydroisomerization and hydrocracking of the waxes. A commercial plant has been designed to operate in Qatar. 13.15.4.2.4 Syntroleum [14, 18, 19] The oxygen plant accounts for about 55% of the investment for the production of syngas from methane. To lower these costs, Syntroleum have operated a demonstration unit in which air instead of oxygen is fed to a low-pressure autothermal reformer. The problem, however, is that the syngas then contains about 50% nitrogen and hence the partial pressures of the reactants in the lowpressure downstream FT reactors are very low and so the reactor capacities too are low (see Section 13.15.6). The tailgas, after condensation of the liquid FT products, will contain about 85% nitrogen and so cannot be recycled to the reformers or be used as fuel gas. Unless it is demonstrated otherwise, the process does not appear to be competitive. 13.15.4.3 Other Activities in and Proposals for FT Operations [6, 14] British Petroleum, ConocoPhillips, Energy International and Statoil have run pilot units converting methane to syngas followed by FT synthesis. Iran, Australia, Indonesia and Oman are considering similar FT routes to fuels. China is investigating the possibility of constructing coal-based FT plants. References see page 2992

2970

13.15 The Fischer–Tropsch (FT) Synthesis Processes

13.15.5

Catalysts for the Fischer–Tropsch Synthesis

The German plants operated with cobalt-based catalysts. Various iron-based catalysts were tested at Schwarzheide [5] but the first utilization of iron catalysts was at the Sasol and Brownsville plants. Only Ni, Co, Fe and Ru have sufficient FT activity for commercial utilization. Iron catalysts can be prepared from either steelworks mill scale or from suitable scrap iron. Relative to these sources of iron, the prices of the other metals are shown in Table 2. Ruthenium is the most active but, because of its low availability and consequent very high price, it is ruled out for commercial application. Nickel is very active but since it also is very active for hydrogenation the selectivity to undesired methane is high and that of alkenes is low. When used at lower temperatures and high pressures for the production of wax, volatile nickel carbonyls are formed and so the catalyst is slowly lost from the reactor. This leaves cobalt and iron as the only viable catalysts. To date, all the FT plants in South Africa use iron catalysts while the Shell plant in Malaysia uses cobalt. Because cobalt is much more active than iron all the plants being built for the production of diesel fuel (see Section 13.15.8) will probably use cobalt. For the production of gasoline

Approximate relative cost of metals active for the FT synthesis

Tab. 2

Fea Ni Co Ru a Fe

Tab. 3

a%

as scrap metal.

13.15.5.1

13.15.5.1.1 Catalysts for Wax Production A Preparation The catalyst used at Sasol, except for some changes, is similar to that originally developed by Ruhrchemie [20, 21]. Scrap iron together with copper metal is dissolved in nitric acid and the oxides are precipitated by the addition of sodium carbonate solution. The type of oxide, the porosity and specific surface area depend on various factors such as the concentrations of the solutions, the precipitating temperature and the final pH [6, 21]. The hydrated ferric oxide is washed, re-slurried with water and the appropriate amount of potassium waterglass added to yield the required amount of silica and potassium. The filter cake can either be extruded for use in fixed-bed reactors or re-slurried and spray-dried for use in slurry phase reactors. The BET

Unreduced

Reduced in H2

Pore volume/ cm3 g−1

Specific surface area/m2 g−1

Area in pores >4.5 nm/ m2 g−1

Pore volume/ cm3 g−1

Specific surface area/ m2 g−1

Area in pores >4.5 nm/ m2 g−1

Reduction/ %a

0.37 0.47 0.74 0.71 0.75 NAb

275 345 375 390 370 405

41 59 90 94 96 NA

0.22 0.43 0.48 0.61 0.65 NA

35 190 250 270 265 280

35 68 80 84 85 NA

100 80 46 58 57 NA

of total Fe present in metallic state after a fixed time at a fixed temperature. not available.

b NA,

Iron-Based Catalysts

Influence of SiO2 on precipitated Fe2 O3 [6, 21]

g SiO2 /100 g Fe

0 8 19 25 29 50

1 250 1000 48 000

and of the high-value linear 1-alkenes, iron catalysts operating in the HTFT mode remain the best option (see Section 13.15.7). A catalyst not only needs to have a high activity and good selectivity, but also other requirements such as particle size, porosity and strength must be met. Under commercial conditions, the rate of pore diffusion is a factor for larger sized particles. To minimize this, either the particle size must be decreased or the average pore size must be increased. Decreasing the particle size in fixed beds increases the differential pressure over the bed and so increases compression costs. If the particles disintegrate due to a too low strength, the situation is much more severe. A compromise must be made between porosity and strength. In two-phase fluidized beds, if the particle sizes are too small, catalyst is lost from the reactor despite the use of efficient cyclones (see Section 13.15.5.4).

13.15.5 Catalysts for the Fischer–Tropsch Synthesis

Tab. 4

2971

Changes in surface area and pore structure of precipitated iron catalysts [21]

State

Unreduced Partially reduced After some FT synthesis (all wax extracted)

a Catalysts

Catalysta

Total pore volume/cm3 g−1

Total surface area/m2 g−1

Surface area in pores >4.5 nm/%

Volume in pores >16.0 nm/%

A B A B A

0.39 0.67 0.30 0.46 0.22

355 340 195 150 51

5 25 20 45 100

0 19 1 22 10

B

0.38

47

100

44

A and B are differently prepared catalysts with different initial pore volumes.

surface area and the porosity of the catalyst increase as the silica content is increased (see Table 3). Using carbonates as the precipitating agents results in higher catalyst porosities than when hydroxides are used. More details of the preparation methods and their effect on the characteristics of the catalyst are reported elsewhere [6, 21] (in Ref. [6], see Chapter 7). As the Fe/Si ratio is typically >4, the silica does not act as a classical support but rather as a binder, to improve strength, and as a spacer, to minimize sintering. A typical catalyst contains 25 g SiO2 , 5 g Cu and 5 g K2 O per 100 g Fe. B Reduction and Conditioning The ease of hydrogen reduction decreases with increasing silica content. This could be due to the formation of iron oxide–silica complexes. The reason for adding copper is to facilitate the reduction of the iron oxide at lower temperatures, typically 220 ◦ C [21]. Increasing the reduction temperature of a Cu-free catalyst increases the rate of reduction, but it is then found that the FT activity of the catalyst is inferior. Full reduction is not necessary as further reduction and conversion to Fe5 C2 occurs in the FT reactors. Some pre-reduction is necessary, however, to shorten the time required for the catalyst to attain full FT activity. Over the whole procedure, from unreduced to fully operational carbided catalysts, the BET surface area decreases while the pore volume and average pore size increase [6, 21] (see Table 4). Pre-reduction can also be carried out with COcontaining gases. Davis et al. [22] reported that, compared with activation with hydrogen-rich syngas, activation with CO or CO-rich syngas resulted in more active catalysts. Bukur and coworkers [23–25] reported that a catalyst reduced with H2 at 250 ◦ C had a higher activity but a lower wax selectivity than the same catalyst reduced with CO at 280 ◦ C.

Catalysts for the Production of Gasoline and Chemicals

13.15.5.1.2

A Preparation and Characterization As the catalysts are used in high-velocity FFB/CFB reactors, they need to be robust. Iron oxide is fused at about 1500 ◦ C with the chemical promoter K2 O and structural promoters such as MgO or Al2 O3 . The melt is poured into ingots and rapidly cooled. The main phase present is magnetite (some wustite is also present). The ingots are then crushed to the particle size distribution required for effective fluidization. During the solidification the structural promoter cations enter into solid solution in the magnetite phase [21]. These promoter cations are therefore atomically dispersed in the magnetite phase, and on subsequent hydrogen reduction small aggregates of Al2 O3 or MgO are precipitated in between the reduced iron crystallites. Being non-reducible, high melting point oxides, they act as spacers between the iron metal crystals. This inhibits sintering and hence results in higher surface areas. The relative effectiveness of various oxides as structural promoters on the BET surface areas of the reduced catalysts is shown in Fig. 6. More details are presented elsewhere [6, 21, 26] (in Ref. [6], see Chapter 7). The basicity of the iron catalyst determines its activity and, in particular, the selectivity (see Section 13.15.7). The amount of CO2 chemisorbed on the reduced catalyst appears to reflect the basicity of the catalyst. The amount increases with increasing alkali level and also the chemisorption on catalysts promoted with K2 O is higher than when the weaker base Li2 O is used. The presence of silica lowers the chemisorption of CO2 on alkali metal promoted catalysts [21]. These findings are in keeping with the observations that lithium is an inferior chemical promoter to potassium and also that silica lowers the heavy hydrocarbon selectivity. It References see page 2992

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13.15 The Fischer–Tropsch (FT) Synthesis Processes

Area / m2 g−1 unreduced sample

20

15 Ti + Mg

Al

Ti

Cr Mg

10

Mn 5

Ca

Be

Si 0

0

1

2

3

4

5

6

Promoter concentration Al Ti Ti + Mg

Cr Mg Ca

Mn Si Be

BET surface areas of fused catalysts after complete reduction with hydrogen. Note that the areas are given as m2 g−1 unreduced catalyst. The concentrations of the various promoters are atoms promoter added per 100 atoms of iron. Fig. 6

has been found that chemisorbed nitrogen lowers the amount of CO2 that can subsequently be chemisorbed, which indicates that surface nitrides lower the basicity. This is in agreement with the observation that nitriding iron catalysts results in a lower heavy hydrocarbon selectivity [27]. CO2 chemisorption data nevertheless need to be interpreted with care. For instance, promotion with CaO increases the CO2 chemisorption but it has little effect on the FT selectivity. Neither the K+ nor Si4+ ions enter into solid solution with the magnetite and so, if some silica is present in the iron oxide used, small occlusions of alkali metal silicates are present as separate phases in the fused catalyst. Microscopic investigations of the milled catalyst showed that, whereas the larger particles still contain alkali metal silicate occlusions, the finer particles consist of a mixture of separate alkali metal silicate and magnetite particles [21]. Hence the distribution of alkali in the milled fused catalyst is heterogeneous. During reduction and FT synthesis, however, the alkali does to some extent spread over the catalyst surface [6]. B Reduction and Conditioning Unlike in the unreduced precipitated iron catalysts, the surface area of fused iron oxides is virtually zero. Pre-reduction is therefore essential to develop the area required for activity. Reduction is carried out with H2 at 350–450 ◦ C. The presence of structural promoters retards the reduction rate.

Precipitated catalysts are readily activated for FT synthesis by mixtures of H2 and CO. However, the reduction of fused iron oxide is markedly retarded when CO is present in the H2 [6, 21]. In any event above 300 ◦ C the Boudouard reaction producing elemental carbon proceeds rapidly and so the use of CO-containing gas for reducing fused catalysts is undesirable. The water vapor produced markedly depresses the rate of reduction [6, 21]. The rate is insignificant at an H2 O/H2 molar ratio of only about 20% of the thermodynamic equilibrium value. To ensure a high reduction rate, it is therefore necessary to maintain a low water vapor pressure by using a high hydrogen linear velocity. The presence of water vapor also enhances the rate of sintering of the metallic iron crystals and so lowers the final surface area. In general, there appears to be an inverse relation between the initial rate of reduction and the surface area of the fully reduced catalyst [6, 21]. The presence of structural promoters retards the rate but yields higher final surface areas. When a reduced catalyst is re-oxidized with water vapor at 400 ◦ C and then re-reduced, the rate of reduction is higher and the surface area is lower. Thus, when re-oxidized at 400 ◦ C, the structural promoters are not re-incorporated into the magnetite lattice. When reduced iron is exposed to synthesis gas, the iron is rapidly converted to H¨agg carbide. This highly exothermic reaction can result in sintering, fouling and excessive carbon deposition. The normal procedure is to start up the reactors pressurized with hydrogen, operating in the recycle mode and gradually introducing synthesis gas. C Carbon Deposition During the FT Synthesis Above about 300 ◦ C, elemental carbon, as distinct from coke, is deposited on iron catalysts at a constant rate. The main source is the Boudouard reaction:

2CO −−−→ C + CO2 The activation energy of this reaction is higher than that of the FT reaction, and the rate of the former therefore increases more rapidly than that of the FT reaction with an increase in temperature. This agrees with the findings that catalysts such as Co, Ru or Fe operating below about 240 ◦ C do not deposit elemental carbon during normal FT reactions. When elemental carbon is deposited the density of the particles is lowered and also catalyst fines are produced. The indirect effect of this on the FT performance of the catalysts is discussed in Section 13.15.5.4. The influence of various promoters on the rate of carbon deposition has been described [6, 21, 28] (in Ref. [6], see Chapter 7). It was found, for instance, that promotion with alumina markedly increased the rate, whereas calcium

13.15.5 Catalysts for the Fischer–Tropsch Synthesis

correlated with the value of PCO /PH2 2 inside the reactor and not simply with the PCO /PH2 ratio [6, 21, 30]. A speculative derivation of the carbon deposition factor was presented elsewhere [30]. The partial pressure of CO2 had no effect on the rate. An important feature of the factor PCO /PH2 2 is that, if (at a fixed temperature) the production capacity of an HTFT reactor is increased, the rate of carbon deposition actually decreases despite the higher hydrocarbon production rate (see Table 6). (The production capacity can be increased, for instance, by increasing the pressure and simultaneously increasing the feed gas flow, thereby maintaining the same residence time.) Another advantage of operating at higher pressures

oxide lowered the rate. However, when the rates were calculated on a unit Fe metal surface area basis, the rates of the various cases were not significantly different. It has been claimed that promotion with chromium oxide decreases the rate of CO decomposition over iron catalysts [29]. Alkali promotion increases the intrinsic rate of carbon deposition, with K2 O having a greater effect than the weaker base Na2 O [27, 28]. The presence of silica lowers the basicity of the alkali and in keeping with this the rate of carbon deposition is lower. The rates of carbon deposition during the FT reaction at a fixed temperature, but at different pressures, fresh feed gas compositions and recycle ratios, were determined. As shown in Table 5, the rate of carbon deposition Tab. 5

H2

CO

CO2

32 24 16.4 17.0 15.6 12.5 13.4 9.2 11.5 8.6 8.5 8.1 8.8

6.4 5.1 4.0 3.6 3.4 2.6 2.7 1.6 2.8 1.8 1.7 1.8 2.3

6.4 4.6 4.1 1.3 3.7 2.5 1.1 0.6 3.4 0.7 1.5 1.7 1.6

ag

Tab. 6

Seta

A

a Sets bg

PCO /PH2

2 × PCO /PH 2 100/bar−1

Relative carbon deposition ratea

0.20 0.21 0.24 0.21 0.22 0.21 0.20 0.17 0.24 0.21 0.20 0.22 0.26

0.6 0.9 1.5 1.2 1.4 1.7 1.5 1.9 2.1 2.4 2.4 2.7 3.0

1.2 1.4 1.6 1.9 2.5 2.5 3.0 3.4 3.4 3.4 4.5 5.0 5.0

C per 100 g Fe per unit time.

Influence of pressure on carbon deposition [21] Total pressure/ bar

9.4 12.9 14.6 18.1 21.5 20.8 30.8 60.8 75.8

B

References see page 2992

Reactor entry gas composition and carbon deposition rate [21]

Entrance partial pressure/bar

PH2 /PCO

(CO + CO2 ) converted/ mol h−1

Relative carbon deposition rateb

4.3 5.2 5.3 5.8 5.9 5.0 4.8 4.7 5.0

49 68 80 94 120 106 157 320 377

5.9 3.3 2.1 1.6 1.1 4.5 2.5 1.4 1.2

Feed characteristics 2 × PCO /PH 2 100/bar−1

7.7 3.7 3.3 2.3 1.7 2.3 1.7 0.9 0.6

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A and B are for two different HTFT Fe catalysts. C per 100 g Fe per unit time.

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13.15 The Fischer–Tropsch (FT) Synthesis Processes

is that a synthesis gas with a higher PCO /PH2 ratio can be fed without the formation of excessive carbon deposition. Higher operating pressures would, however, increase the capital and operating costs of the FT reactors. 13.15.5.1.3 Alternative Iron-Based Catalysts Instead of fusing, iron oxides can be sintered together with the desired promoters. In general, the FT performances of these catalysts are similar to those of their fused counterparts [27]. Alternatively, the iron oxide powders can be bound together by compounds such as aluminum nitrate or silicates. These preparations are less effective than those prepared by fusion or sintering [27]. Possibly, the binders chemically react with the alkali resulting in less basic and hence less effective catalysts. A common method of catalyst preparation is to impregnate the metal precursor onto high surface area supports such as alumina or silica. In the case of iron, however, the essential alkali promoter will also adsorb on or chemically react with the support. The result is that the effectiveness of the alkali with respect to the iron component could be markedly lowered. The activities and the wax selectivities of impregnated catalysts are inferior to those prepared by precipitation methods. The use of chemically inert supports such as wide-pore charcoals could possibly be a better choice.

Cobalt-Based Catalysts The original German cobalt catalyst was prepared by coprecipitating the nitrates of cobalt and thorium with a basic solution in the presence of kieselguhr. A typical mass ratio was 100 Co:18 ThO2 :100 kieselguhr [27]. It was found that the addition of 2% copper enhanced the rate of reduction [31]. However, copper increased the rate of activity decline [32] and so was not added. The high cobalt content would today make such a catalyst very expensive. Because of the high cost of cobalt, it is important to minimize the amount used, but nevertheless have a high cobalt metal surface area. This is achieved by supporting the cobalt on stable high surface area oxides. A common technique is impregnation of the support with a cobalt salt solution, drying, calcining (to convert the salt to the oxide) and finally reduction with hydrogen. In order to obtain a high Co dispersion, each preparation step needs to be optimized. 13.15.5.2

13.15.5.2.1 Cobalt Salts Used Using titania as a support, various ways of adding the cobalt and also using different cobalt salts were studied [33]. It was found that the catalysts prepared from cobalt oxalate by ‘‘spreading’’ (heating the mechanically mixed components to 250 ◦ C) yielded the most active catalyst. Incipient wetness impregnation with cobalt(III) acetylacetonate produced

a more active catalyst than the commonly used nitrate. The effect of using different Co precursors for catalyst preparation by incipient wetness impregnation of alumina was investigated [34]. When loading 2.5% Co, using ammonium cobalt citrate, very small Co oxide particles were formed. On thermal treatment under reducing conditions the particles reacted with the alumina to form aluminates, which were inactive. The catalysts prepared from Co nitrate formed larger oxide particles. This catalyst did reduce and was active for the FT reaction. When the Co loading was increased to 5%, using the citrate salt, the oxide particles were larger and they were reducible. The use of Co nitrate and two Co carbonyls as precursors were compared for the preparation of Co/silica catalysts [35]. The dispersion of Co on the reduced catalysts decreased in the order Co2 (CO)8 > Co4 (CO)12 Co nitrate, and consequently the carbonyl-derived catalysts had a higher initial FT activity. The effect of impregnating alumina with different Co salts was investigated [36]. The acetate yielded smaller crystals than the nitrate, and the initial FT activity was higher but it declined faster than in the nitrate case. Acetylacetonate yielded a poor catalyst, apparently due to pore blockage by undecomposed acetylacetonate deposits. 13.15.5.2.2 Supports for Cobalt Catalysts Comparing alumina and silica as supports, temperature-programmed reduction (TPR) studies showed that after identical hydrogen reduction procedures, the silica-supported material contained little or no unreduced cobalt, whereas the cobalt on alumina was only about 35% reduced (the balance of the Co being present as some form of ‘‘cobalt aluminate’’) [36]. Despite this, the FT activity of the Co/alumina was higher than that of the Co/silica catalyst. Transmission electron microscopy (TEM) and chemisorption studies of the reduced catalysts showed that the Co/silica catalyst contained smaller Co metal crystals. Under FT conditions, the small Co crystals are apparently oxidized (see Section 13.15.5.3), and this could possibly account for the lower activity. Goodwin and coworkers [37] investigated a series of catalysts supported on TiO2 , SiO2 and Al2 O3 . The influence of various promoters such as Ru, Re and La or Zr oxides was investigated. They concluded that the Al2 O3 supported ruthenium-promoted catalyst had the best FT performance whereas the TiO2 catalysts, due to their lower surface areas, were inferior. Ru or Re increased the activity of the Al2 O3 - or TiO2 -supported Co, and ZrO2 did the same for the Co/SiO2 catalysts. Bartholomew and coworkers [38, 39] compared the performances of unsupported Co and Co supported on alumina, silica, titania, carbon and magnesia. The tests were carried out at atmospheric pressure and at

13.15.5 Catalysts for the Fischer–Tropsch Synthesis

low conversion levels, i.e. far removed from practical FT conditions. The order of decreasing activity for catalysts containing 3% Co was Co/titania > Co/silica > Co/alumina > Co/carbon > Co/magnesia. This order is different from that found by others under more realistic FT conditions [36, 37] and with catalysts containing higher Co loadings. Shell’s commercial plant at Bintulu uses a cobalt-based catalyst. Exxon and Sasol have operated demonstrationscale slurry bed FT units with cobalt catalysts. Conoco, Statoil and Syntroleun have also developed cobalt-based catalysts. The methods of preparation of these commercial catalysts have not been fully revealed but the patents taken out could give some indication of the preparations. A list of patents is presented in Chapter 7 in Ref. [6]. Shell appears to have concentrated on silica as a support, Sasol and Statoil on alumina and Exxon on titania. A review of many of the patents appears elsewhere [37]. Sasol found that during their slurry impregnation procedure, some of the alumina is dissolved and reprecipitated. This results in inferior bonding of the cobalt with the support, which leads to some physical loss of cobalt during the FT process. The problem was solved by precoating the alumina support with silica [40–42]. Sasol claims that loading the cobalt salt on the support by the slurry phase impregnation method results in a more active catalyst than when the incipient wetness method is used. (In the former method an amount of the aqueous metal solution in excess of the support pore volume is added, and the slurry is then dried under sub-atmospheric pressure.) The slurry phase method was found to result in a more uniform distribution of the cobalt. The FT performance of Co supported on kieselguhr, silica, alumina, bentonite, Y-zeolite, mordenite and ZSM-5 has been compared by Bessell [43]. Whereas the low-acidity supports such as silica and alumina produced linear products, the acidic supports produced more branched products and at higher temperatures also produced aromatics. The isomerization and aromatization are secondary, acid-promoted reactions of the FT olefins. This is then equivalent to a combination of the FT and the Mobil olefins-to-gasoline (MOG) process. (Note that with iron-based catalysts this approach is unlikely to be successful as alkali promotion is essential and the alkali would neutralize the required acid sites on the zeolite support.) Calleja et al. [44] studied the FT performance of Co/HZSM-5 prepared by incipient wetness impregnation of the zeolite. Promotion with thorium, being basic, decreased the acidity of the zeolite and so less aromatics were formed and consequently more of the heavier hydrocarbons emerged from the reactor because of the lowered level of secondary reactions. They also compared the performance of this catalyst with that of a physical mixture of a Co FT catalyst and HZSM-5. As expected from

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the effect of the basic thorium, they found that for the mixture the amount of aromatics formed was somewhat higher but the amount of heavier hydrocarbons formed was significantly lower. Reduction and Promoters Cobalt catalysts are pre-reduced with hydrogen before being loaded into the FT reactors. Reduction is carried out from 250 to 400 ◦ C, at high linear gas velocities to minimize the vapor pressure of the product water, which enhances sintering [6]. TPR studies show that reduction of Co3 O4 starts at about 230 ◦ C. It has been claimed that the activity of a cobalt catalyst could be increased by reducing in hydrogen, then re-oxidizing the catalyst followed by rereduction [45]. Exxon found that the addition of low levels of Ru enhanced not only the initial reduction rate but also the in situ regeneration of the cobalt [46]. Goodwin and coworkers [47] confirmed the positive effect of Ru promotion and reported a three-fold increase in FT activity. Sasol investigated the effects of various promoters and found that Pt was as effective as Ru whereas Re resulted in inferior performance [48]. Others found that, while Re promotion enhanced the rate of the initial reduction, the catalyst activity decline rate was higher [49, 50]. La and Zr promotion of Co catalysts has also been extensively studied by Goodwin and coworkers [37, 51–53]. 13.15.5.2.3

13.15.5.3 Catalysts Based on Metals Other than Iron or Cobalt Ru is more active than Ni, Co or Fe. It also is very versatile in that at high temperatures the only product is methane, whereas at low temperatures it produces large amounts of waxes in the low molecular mass polyethylene range [54, 55]. The high price of Ru, however, makes its commercial application very unlikely. Of the other Group VIII noble metals, only Rh and Os have some moderate FT activity [54]. The products obtained on Rh catalysts contain large fractions of oxygenated molecules. Mo as a catalyst has been investigated, the concept being that the purification of sulfur-containing syngas would not be required [56]. Even though the Mo catalyst was active in the presence of H2 S, it was nevertheless more active in its absence. Other studies found that, to achieve reasonable conversions, temperatures around 400 ◦ C were required, but then the methane selectivity was about 90%. Promotion with alkali lowered it to about 50% [21].

Deactivation of FT Catalysts The factors involved in the decline of activity with time on-stream are the following: (1) fouling of the catalyst 13.15.5.4

References see page 2992

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13.15 The Fischer–Tropsch (FT) Synthesis Processes

surface by coke deposits; (2) poisons in the feedgas; (3) hydrothermal sintering; (4) oxidation of the active phase to the inactive oxide; and (5) Boudouard carbon deposition.

A

Poisoning by Sulfur Compounds Sulfur compounds in the syngas rapidly deactivate Co and Fe catalysts. Even at levels as low as 0.03 mg S m−3 poisoning is observed [21]. The extent of poisoning depends on the type of reactor. For fluidized-bed reactors all of the catalyst would be poisoned, whereas for a long fixed-bed reactor the bulk of the sulfur would be absorbed by the upper sections of the bed, leaving the lower sections relatively unscathed. For an Fe catalyst operating at about 230 ◦ C in a fixed bed, it was found that organic sulfur resulted in a higher rate of activity decline than H2 S [21]. This appears to indicate that H2 S reacts more strongly with iron and thus would be more completely absorbed in the upper layers of the catalyst bed. The organic S would penetrate deeper into the bed before reacting and so would poison a larger portion of the catalyst. In the case of short catalyst beds, as used in laboratory reactors, the opposite effect would be observed, namely that H2 S is a more powerful poison. Hydrogen re-reduction of a sulfur-poisoned catalyst does not reactivate the catalyst. 13.15.5.4.1

Deactivation of Iron Catalysts A Low-Temperature Fischer–Tropsch (LTFT) Operation The surface area of an LTFT catalyst which has been operating for only a few days is about 200 m2 g−1 . The XRD pattern of the catalyst is broad and indistinct, confirming the presence of very small iron carbide crystals. The area of the catalyst, after it has lost about 20% of its original activity, is about 50 m2 g−1 , and the XRD pattern is sharper [21]. These changes indicate that crystal growth had occurred with time on stream. In confirmation, it was found that, if the amount of SiO2 , which acts as a spacer, added to the catalyst is lowered, the rate of activity decline is faster. When water vapor was deliberately added to the feed gas, thereby resulting in a higher water partial pressure throughout the reactor, the rate of activity decline was increased. This could have been due to both enhanced hydrothermal sintering and to oxidation. Using the kinetics presented in Section 13.15.6, it is estimated that if half of the active Fe surface sites are lost then the conversion will drop by a factor of 0.3. Several separate FT runs were carried out in pilot-plant tubular reactors, and the catalysts were then carefully unloaded to obtain samples over the reactor length. (The reactor tubes were of the same length and diameter as those in the multi-tube commercial reactors.) Each sample was then analyzed and the FT performance determined 13.15.5.4.2

Activity

B

C

Entry

Exit Bed position

LTFT multitubular catalyst deactivation profiles. Time on-stream increases from A to B to C.

Fig. 7

under fixed conditions in laboratory units [57]. The results of the FT tests are illustrated in Fig. 7. As can be seen, for young catalysts there was little difference in the activity of the catalyst in the different sections of the bed. With time, however, the activity of the catalyst near the entry declined markedly, ending up completely inactive. The catalyst samples near the exit also declined in activity but the catalyst samples from the middle sections declined the least. The BET surface areas (after thorough wax extraction) of the samples declined progressively from the entry to the exit of the reactor whereas the porosity of the samples increased. X-ray line broadening studies confirmed that the average crystallite sizes increased from the entry to the exit of the reactor. The percentage oxidation of the iron to Fe3 O4 also increased in the same direction. Chemical analyses showed a high S content in the top section of the bed, the S content decreasing with bed depth. It thus appears that sulfur poisoning is the main cause of loss of activity of the front sections of the catalyst bed. Note that the BET surface areas of the samples near the reactor entrance were the highest of all the samples and so sintering was not the main factor there. As the FT reaction progresses down the reactor, the partial pressure of water increases and thus the likelihood of hydrothermal sintering increases. This matches the observation of decreasing surface areas. In addition to the increasing H2 O concentrations, those of H2 and CO also decrease, which means that the system becomes more oxidizing down the tube, as was in fact observed from the analysis of the catalyst samples. The increased sintering and degree of oxidation explain the progressive loss in FT activity in the lower sections of the catalyst bed. In a slurry bed reactor, the catalyst is fluidized and hence back-mixing of the solids occurs freely in the bed. This

13.15.5 Catalysts for the Fischer–Tropsch Synthesis

means that on average all the catalyst particles are exposed to any poisons and to the various gas compositions along the bed height and so all the catalyst particles should deactivate at the same rate. B High-Temperature Fischer–Tropsch (HTFT) Operation Despite operating at 340 ◦ C, the catalyst contains significant amounts of ‘‘waxes’’. In contrast to the LTFT process, these products are highly aromatic, a portion of which is pyridine-insoluble and can only be removed by hydrocracking at higher temperatures, yielding oils and gases [21]. (Note that hydrogen treatment of carbidic or Boudouard carbon does not produce oil.) It is therefore likely that for HTFT fouling ‘‘coke’’ contributes to activity decline. The presence of liquid inside the catalyst pores will also lower the reaction rate by lowering diffusion rates. This would not in itself result in activity decline with time on stream unless there is a continuous build-up of heavier, more viscous oils in the pores. At HTFT temperatures the Boudouard reaction (2CO → CO2 + C) results in a continuous increase in the elemental (‘‘free’’) carbon content of the catalyst. After several weeks there is, on an atomic basis, more carbon present than iron. Since the decline in activity is relatively small, it appears unlikely that carbon deposition directly effects the decline. There is, however, a strong indirect negative influence. Because of the vigorous movement of the catalyst in the high gas velocity fluidized beds, scouring of the carbon-rich carbide outer layers (see below) of the catalyst particles and also break-up of the particles take place. Hence more fine carbide-rich particles are generated. Since these small particles have low densities, because of the high free carbon content, they pass through the cyclones and so catalyst is lost from the reactor. The loss of carbide-rich fines is probably one of the reasons why the oxide content of the remaining catalyst increases with time on stream [21]. A freshly carbided catalyst has the same BET surface area as the reduced catalyst but with time there is a continuous increase in area. This is due to the deposition of high surface area ‘‘free’’ carbon. Because of this, the BET measurements cannot provide evidence of sintering and oxidation. Analysis of aged catalyst shows that the smaller particles have higher levels of carbide and of ‘‘free’’ carbon and lower levels of oxide than the larger particles. Scanning electron microscopy/energy-dispersive X-ray (SEM/EDX) examination of polished sections revealed that the small particles consisted of crystals of iron carbide embedded in a matrix of carbon, whereas the larger particles had cores of magnetite surrounded by the carbide–carbon mixture [57]. The formation of the oxidized cores of

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the larger particles is due to the fact that, as the gas moves towards the center of the particle, it becomes more oxidizing. (Due to the FT reaction, the water concentration increases and the H2 and CO concentrations decrease.) Overall, the reasons for the observed activity decline with time on stream are possibly due to three factors: first, the loss of the small, loose alkali metal silicate promoter particles, and also of the small carbide-rich particles from the reactors; second, the poisoning of the alkali metal/Fe sites by sulfur compounds in the feed gas; and third, the fouling by ‘‘coke’’ of the more active alkali metal/Fe centers. The latter concept is supported by the observation that H2 re-reduction at 350 ◦ C of used catalysts largely restores both the activity and the selectivity [21]. (H2 treatment at 350 ◦ C does not reduce iron sulfide, i.e. S poisoning cannot be reversed.) The observation that the regeneration of the activity by rereduction is as marked for a catalyst containing 5% oxide as for one containing 60% oxide also indicates that catalyst oxidation during synthesis is probably not the main cause of activity decline [21]. 13.15.5.4.3 Deactivation of Cobalt Catalysts Because of the high price of Co, it is essential to minimize the rate of FT activity decline. Initially the decline is relatively rapid, after which it is slower [58, 59]. As the FT temperatures are mostly below 230 ◦ C, the deposition of ‘‘coke’’ or ‘‘free’’ carbon is unlikely, as evidenced by the absence of aromatics and the low levels of CO2 , respectively. A possible reason for the initial ‘‘rapid’’ decline is the build-up of long-chain waxes in the pores, which retards diffusion. In practice, there must be a flow of liquid wax from within the catalyst particles to the outside surface and so after a steady state has been reached it is unlikely that further build-up will occur. Another probable reason for the initial loss of activity is the rapid sintering of the smaller Co crystallites. It has been shown that ruthenium-promoted Co can be regenerated by H2 at reaction temperatures, whereas the unpromoted catalyst cannot [46]. Ru may result in a more strongly hydrogenating surface, thereby lowering the surface coverage by heavy waxes by hydrocracking them [60]. An important factor in the deactivation of cobalt appears to be the water vapor produced. Three related aspects may be involved, viz. sintering of the cobalt metal particles, surface or total oxidation of the smaller cobalt particles and the formation of inactive cobalt–support compounds. With a Co/silica catalyst it was observed that with time on-stream there were a loss of activity References see page 2992

2978

13.15 The Fischer–Tropsch (FT) Synthesis Processes

and an increase in the amount of silicates formed [61]. Niemel¨a and Krause [59] reported that the higher the Co dispersion on SiO2 , the higher are the deactivation rates. Thermodynamically, bulk phase oxidation of cobalt metal by water, under FT conditions, is not possible. A portion of the Co atoms on the surface, however, may be oxidized. A popular concept of the FT mechanism is the dissociation of chemisorbed CO to adsorbed C and O atoms. The latter is equivalent to a surface oxide which is subsequently hydrogenated to water. Whatever the detail, the continuous cyclic oxidation and reduction of the surface cobalt atoms can bring about sintering, which results in a loss of activity. The smaller the Co crystallites, the higher is the degree of unsaturation of the exposed atoms and so probably the higher the coverage of the surface by inert ‘‘oxide’’. Crystallites below a critical size, estimated at 5 µm [62], could be fully oxidized and hence rendered completely inactive. Since the reaction of Co with the support to form silicates or aluminates apparently requires cobalt to be in the oxidized state, the presence of water vapor should enhance this reaction. These compounds are not reducible under FT conditions and so this results in permanent activity loss. It has been found that for a Co/Al2 O3 catalyst the addition of water to the syngas results in surface oxidation of the cobalt and in permanent loss in activity [50, 63]. XPS and gravimetric analyses indicated that only a small fraction of the bulk Co metal was oxidized but significant oxidation of the surface Co atoms occurred and/or highly dispersed Co reacted with the support. In situ M¨ossbauer studies showed that oxidation of cobalt metal supported on alumina does occur under FT conditions [64]. From XANES studies, Davis and coworkers [65] concluded that, on noble metal-promoted Co/alumina, the small Co metal clusters oxidized in the presence of water and that during the initial deactivation stage cobalt cluster growth took place, resulting in a lower Co metal surface area. On Co/silica catalysts, the FT product water or water deliberately added to the feed gas was found to increase the FT activity at low partial pressures of water but to decrease the activity at high water partial pressures [66]. Under high water partial pressure conditions, the activity loss was permanent due to the formation of irreducible Co silicates [67]. It has been reported that over a Co/titania catalyst, operating at low conversion, the addition of water also increased the FT activity [62]. Others have found that at low conversions water addition had little effect but at higher conversion it resulted in a decrease in activity [68]. The effect of water, at low levels, on the activity of the different Co catalysts is confusing but at higher levels (i.e. at high conversions) water does apparently result in catalyst deactivation.

13.15.6

Kinetics of the Fischer–Tropsch Synthesis

In the past there have been discrepancies regarding the relative activities of Ru, Ni, Co and Fe catalysts for the FT reaction (see Chapter 7 in Ref. [6]). This may be due to the fact that in general the observed rates will depend not only on the intrinsic rates but also on diffusion rates in and out of the porous catalyst particles. The latter will in turn depend on the pore sizes, the particle sizes and whether or not liquid wax is present in the pores. As far as commercially viable catalysts are concerned, cobalt catalysts are more active than iron catalysts. Iglesia and coworkers [69] reported that the FT turnover rates on cobalt catalysts are about three times higher than those on iron catalysts. Iron-Based Catalysts For all iron catalysts, alkali promotion is essential (see Section 13.15.7.3.2B). For fused catalysts at high temperatures (where the wax selectivity is low), adding alkali results in higher conversions but further alkali additions have little effect (see Table 12). At lower temperatures the wax selectivity is higher, and then it is found that increasing the alkali content results in a lower activity [21]. For precipitated catalysts operating in the high wax selectivity mode, the activity also decreases when the alkali content is raised above a certain level. At higher alkali levels the wax formed within the catalyst pores becomes progressively longer chained and thus more viscous. The diffusion rates, and consequently the overall reaction rates, are lowered. When an LTFT fixed bed is periodically treated with a liquid solvent, which extracts the wax from the catalyst, there is an immediate large increase in activity, which then rapidly declines to its previous level as the catalyst particles are again filled with wax [21]. This clearly indicates that diffusion restrictions due to the presence of wax in the catalyst pores lower the overall reaction rate. In keeping with this, it was found that decreasing the catalyst particle size resulted in higher activities [6, 21]. In the LTFT slurry phase reactor, the catalyst particles are much smaller, and so diffusion effects within the particles are absent. Permanent fouling by ultra-high molecular mass waxes is unlikely in systems where liquid wax is being produced, as they would be continuously flushed out. It was also found that nitriding iron catalysts resulted in lower wax selectivities and in higher activities [6, 21]. The activation energies (E) of diffusion are normally lower than those of chemical reactions. For precipitated catalysts, which produce high yields of wax, E = 55–62 kJ mol−1 in the range 200–250 ◦ C [70]. For fused catalysts, which produce much less wax, E = 71 kJ mol−1 [71]. The above values 13.15.6.1

13.15.6 Kinetics of the Fischer–Tropsch Synthesis

confirm the presence of diffusion restrictions. In practice, however, only the partial pressures of the reactants and products in the gas phase outside the catalyst particles are measured, and these are used to derive the overall kinetics. Often kinetic studies are carried out in high-flow, i.e. low-conversion, single-pass reactors. This ignores the possibility that some of the products, which would be more abundant at higher conversion levels, could have an influence on the kinetics. To avoid these pitfalls, it is recommended that laboratory studies be carried out in high gas velocity, high-recycle ‘‘stirred tank’’-type units such as Berty reactors. Over the years, various kinetic equations have been proposed for the FT reaction over iron-based catalysts. A few of these equations are as follows: r = aPH0.62 PCO 0.4 − br 0.5 PH0.52 O [72]

(1)

r = aPH2 PCO /(PCO + bPH2 O ) [27, 70]

(2)

r=

aPCO PH2 2 /(PCO PH2

+ bPH2 O ) [73]

r = aPCO PH2 /(PCO + bPH2 O + cPCO2 ) [74]

(3) (4)

where r is the FT reaction rate and a, b and c are the relevant constants. A large number of experiments were carried out in pilotplant fixed- and fluidized-bed reactors using commercial iron-based catalysts [21]. The variables included different catalyst bed heights, total pressures (from 0.8 to 7.6 MPa), recycle ratios and feed gas compositions. The effect of varying the individual partial pressures of CO, CO2 , H2 and H2 O were also investigated. The key observations were as follows. (1) The rate was strongly dependent on the hydrogen partial pressure. At low conversion levels, the rate was solely dependent on the hydrogen pressure. (2) The rate increased with the partial pressure of CO. (3) The partial pressure of CO2 did not have a direct effect on the rate. [It did affect the partial pressures of the other components via the water gas shift (WGS) reaction.] (4) The rate was markedly depressed by increasing the partial pressure of water. (5) The level of gaseous hydrocarbons had no apparent effect on the rate of the Fischer–Tropsch reaction. Based on these observations, a satisfactory rate equation for iron catalysts is Eq. (2), as it is in keeping with all the experimental findings. Note that this in effect ‘‘re-invented’’ the equation proposed by Anderson about 20 years prior to the Sasol investigation. When applying this equation in multi-step calculations and at each step also taking into account the simultaneously occurring WGS reaction, the conversion profile along the length of the reactors can be calculated. It was found to match the actually measured profiles in both the LTFT fixed-bed and the HTFT circulating fluidized-bed commercial reactors [21].

2979

The negative effect of water on the FT rate is common to all four equations above. This effect is probably due to the fact that iron is sensitive to oxidation by water vapor, and so the higher the H2 O partial pressure the higher is the occupancy of surface iron sites by O atoms/ions, which lowers the amount of active surface sites available for the FT reaction. A theoretical derivation of Eq. (2) is given elsewhere [6, 70] (in Ref. [6], see Chapter 7). Equation (2) predicts that if the total pressure is increased and the gas feed rate is increased by the same factor, i.e. the residence time in the catalyst bed remains the same, then the degree of conversion remains unchanged. This matches the experimental pilot-plant findings. This means that the production rate is increased in proportion to the increase in pressure. Based on this prediction, new pilot plants were constructed at Sasol and tests up to 6.0 and 7.5 MPa for the fixed-bed LTFT and fluidized-bed HTFT operations, respectively, were carried out. These tests confirmed the kinetic predictions. A 4.5-MPa fixed-bed commercial reactor was subsequently built, and it performed as per prediction. (The older commercial reactors operate at 2.7 MPa.) It has been estimated that HTFT fluidized-bed units can be viable up to 4.0 MPa [75]. The use of even higher pressures is constrained by economic considerations and heatexchange capacity requirements. Cobalt-Based Catalysts Currently, the Shell commercial plant in Malaysia is the only one using a cobalt catalyst. Both Sasol and Exxon have operated demonstration slurry bed FT units with cobalt catalysts. However, it appears that no kinetic information for these operations has been published. Various equations, based on laboratory experiments, have, however, been generated [6, 75–79]. A few examples are given below. 13.15.6.2

r = aPH2 2 /PCO [80]

(5)

r = aPCO PH2 /(1 + 1.82PCO )2 [76]

(6)

r = aPCO PH0.52 /(1 + bPCO + cPH0.52 )2 [81]

(7)

r = aPCO PH2 2 /(1 + bPCO PH2 2 ) [27]

(8)

0.2 r = a(PH0.52 /PCO )/(1 + 0.93PH2 O /PH2 ) [82]

(9)

r = aPH2 PCO /(PCO + 0.27PH2 O )

(10)

where r is the FT reaction rate, and a, b and c are the relevant constants. Despite the fact that Co is more resistant to oxidation than Fe, it has been reported that under FT conditions Co crystals smaller than 5 µm [62] are oxidized by water References see page 2992

13.15 The Fischer–Tropsch (FT) Synthesis Processes

vapor (see also Section 13.15.5.4). It is conceivable that the smaller the Co crystals, the lower is the PH2 O /PH2 ratio at which they will oxidize and so become inactive. If this process is reversible then it seems feasible that the effect of the water vapor can be incorporated into the kinetic equation. Thus Eqs. (9) and (10) could well be applicable at high conversions, i.e. under commercial conditions. Comparison of the Activity of Cobalt and Iron Catalysts Comparing the equations in the foregoing sections it can be seen that, contrary to the iron case, the majority of equations for cobalt do not contain a water partial pressure term. Since cobalt is much more resistant to oxidation than iron, it can be presumed that under FT conditions the occupancy of the cobalt surface by oxygen atoms/ions will be much lower than in the case of iron. This could account for the absence of a PH2 O term in some of the equations for cobalt. This can give cobalt a large activity advantage. To evaluate the situation the conversion profiles for three cases were calculated using Eq. (2) for iron and Eqs. (6) and (9) for cobalt (see Chapter 6 in Ref. [6]). In order that the Co and Fe catalysts start off on the same footing, the value of the constants a in these equations were chosen so that the 4% conversion levels were attained at the same position near the entrance of the catalyst bed. (Note that at low conversions the partial pressure of water is low and hence its influence on the kinetics is minimal.) This was taken to mean that the ‘‘initial’’ activities were the same. If Co is three times more active per surface site than Fe [69], the above means that there are effectively three times more active sites available per unit volume of catalyst bed in the case of Fe. For all the cases illustrated in Fig. 8, the chosen conditions were: feedgas PH2 /PCO ratio = 2/1, P = 3.0 MPa, the same feedgas flow, once-through basis and the same LTFT temperature. For the iron catalysts, allowance was made for the WGS reaction, but not for cobalt as it has a very low WGS activity under LTFT conditions. Included in the figure is a case for an Fe catalyst with an initial activity five times higher than that of the other Fe case (Fex5). As can be seen, even though Fe and Co have the same initial activity, the Co catalyst increasingly outperforms the Fe catalyst as the syngas moves through the reactor. For example, at a bed ‘‘length’’ of 30 units the conversion attained is over 80% for the Co catalyst using Eq. (6) and over 60% using Eq. (9) and just under 40% for the Fe catalyst using Eq. (2). If the Fe catalyst is made five times more active, then it does outperform the Co catalyst [Eq. (6)] up to about the 60% conversion level, after which the Co catalyst again is superior. Should Eq. (9) for Co apply, however, the five times more active Fe catalyst outperforms the Co catalyst. There is hence a big incentive 13.15.6.3

Conversion / %

2980

100 90 80 70 60 50 40 30 20 10 0

Co Fex5 Co/H2O

Fe

0

5

10

15

20

25

30

35

Bed length Comparison of the calculated conversion profiles of Co and Fe catalysts. The bed length is in arbitrary units. Plot ‘‘Fe’’ is for Eq. (2), plot ‘‘Co’’ is for Eq. (6) and plot ‘‘Co/H2 O’’ is for Eq. (9). For an explanation of Fex5, see text.

Fig. 8

to improve significantly the activity of the lower cost Fe catalysts. However, if the number of active sites for Co catalysts can also be increased, then the Co catalysts will remain superior. It is of interest to point out that when using Eq. (10) the reaction profile obtained is virtually identical with that predicted by Eq. (9). A further general point is that when applying the various kinetic equations for both Fe and Co catalysts, it is found that a change in the partial pressure of hydrogen has a greater effect on the reaction rate than a similar change in the partial pressure of carbon monoxide. This possibly explains why so many ‘‘apparently different’’ kinetic equations have been published. When using the kinetic Eq. (2) for Fe and Eq. (10) for Co (the two equations have the same structure), the response to changes in total operating pressure is the same. For instance, if the pressure and the syngas flow are doubled (thus the residence times in the catalyst beds are the same) the conversion stays the same but the production rate doubles. [In the case of iron catalysts this calculated prediction matches the observed pilot plant scale results (see Table 6).] For the cobalt Eqs. (6) and (9), however, while the production rate also increases with pressure, the actual conversions decrease [6]. Hence for the purpose of designing a commercial plant it is essential to know which kinetic equation to use. It would in any event be essential to determine on demonstration size FT reactors how the process reacted to changes in pressures. Although cobalt catalysts clearly have a large activity advantage, this does not mean that high conversions cannot be achieved with Fe catalysts. Using the commercial Fe LTFT catalyst, it has been shown that over 90% conversion can be attained. However, this requires using two reactors in series with water knockout between the two reactors and also each reactor

13.15.7 FT Product Selectivities

operating in the recycle mode (which involves feeding some tail gas, after water knock-out, together with fresh syngas feed to the reactor). This mode of operating lowers the average water vapor pressure in the reactors and thus increases the reaction rate. This scheme, however, increases construction and operating costs. For the fluidized HTFT process, conversions over 90% are obtained in single-stage reactors operating with recycle. It should be noted that Fig. 8 depicts the performance of fresh catalysts. With time on-stream both Co and Fe catalysts deactivate. If the FT reactors are slurry bed units then in the case of iron catalysts, because of their relatively low cost, high conversion levels can be maintained by regular on-line addition of fresh catalyst. The superior performance of cobalt catalysts nevertheless does mean that high conversions can be achieved in a single-stage reactor without the need to recycle tail gas. Hence fewer or smaller reactors would be required. In the case of cobalt catalysts, high conversion levels do result in high PH2 O /PH2 ratios in the reactor. This accelerates the rate of deactivation (see Section 13.15.5.4). To minimize these effects, a two-stage operation with water knock-out between the stages can be applied. Alternatively, a larger reactor with recycle of Tab. 7

a portion of the tailgas (after water knock-out) could be applied. The number or cost of the reactors will then be increased, but the overall costs should still be lower than those required when using iron-based catalysts. 13.15.7

FT Product Selectivities Primary and Secondary Reactions The FT product spectra are determined by mechanistic, catalytic and kinetic factors and are very different from what would be expected. Thermodynamics predicts that the main product should be methane and the amount of alkenes should be insignificant, whereas in practice the methane content is low and the alkenes content is high [6, 21] (in Ref. [6], see Chapter 3). Table 7 shows typical observed ratios of several products compared with the calculated ratios, if the reactions went to equilibrium at the actual partial pressures of the relevant gases in the reactor. It can be seen from Table 7, class (A), that the alkenes should theoretically be hydrogenated to alkanes, but the 13.15.7.1

References see page 2992

Equilibrium and actual ratios of product concentrations

Reaction class

Reactions

Temperature/K

Ratios Ratio

A

B C D

C2 H6 C2 H4 + H2  C3 H6 + H2

 C3 H8 C5 H10 + H2

 C5 H12 C10 H20 + H2

 C10 H22 C20 H40 + H2

 C20 H42 C2 H4 + H2 O

 C2 H5 OH C2 H5 OH + H2

 C2 H6 + H2 O C2 H5 OH + H2 O

 CH3 COOH + 2H2

E

C2 H5 OH

 CH3 CHO + H2

F G

3-Me-1-C4 H8 1-C5 H10  2CH3 COOH

 CH3 COCH3 + CO2 + H2 O CH3 COCH3 + H2

 2-C3 H7 OH

H

I a At

2981

C7 H14

 C7 H8 + 3H2

the relevant partial pressures at the reactor exit. catalysts in all cases.

b Iron-based

C2 H4 /C2 H6 C3 H6 /C3 H8 C5 H10 /C5 H12 C10 H20 /C10 H22 C20 H40 /C20 H42 C2 H5 OH/C2 H4

Expected at equilibriuma

Typically observedb

5.9 × 10−7 1.7 × 10−5 1.3 × 10−5 1.1 × 10−5 1.1 × 10−5 0.0027 0.024 3.6 × 10−9

2 8 7 6 3 0.0045 1.2 0.6

600 600 600 600 600 600 510 600

C2 H5 OH/C2 H6

610

CH3 COOH/C2 H5 OH

0.14

0.14

510 600

CH3 CHO/C2 H5 OH

0.0009 0.22

0.26 0.21

510 600 600

3-Me-1-C4 H8 /1-C5 H10 C3 H6 O/(C2 H4 O2 )2

610

2-C3 H7 OH/C3 H6 O

510 600

C7 H8 /C7 H14

0.0028 1.4 5.6 × 104

0.27 0.17 1.9 × 102

0.2

0.2

7 2000

0.1 0.15

2982

13.15 The Fischer–Tropsch (FT) Synthesis Processes

alkenes are in fact dominant. This indicates that alkenes must be primary products and, furthermore, that their subsequent hydrogenation must be slow. Even with cobalt, which is much more active for hydrogenation than iron, the alkene contents of the various cuts are significant (see Tables 8 and 9). Thermodynamically the hydrogenation of ethene should be more complete than that of propene, while for higher carbon numbers the ratios alkenes/alkanes should decrease. It can be seen that the observed ratios, although much higher, do follow the predicted trends. For the data in case (A), the H2 pressure was 0.6 MPa. At 510 K and a hydrogen pressure of 4.5 MPa the alkene to alkane ratios were lower but the trends were the same. If ethene is deliberately added to the syngas fed to an iron catalyst at 600 K, about 50% is hydrogenated to ethane [21]. However, if the liquid oil produced is recycled to the reactor then there is hardly any change in the alkene or aromatic content. Hence it is concluded that the higher the molecular mass of alkenes added to the syngas, the less likely they are to be hydrogenated in secondary reactions. The situation, however, for the growing primary FT species on the catalyst surface is different. The higher the carbon number, the longer is the time they have spent linked to the catalyst and hence the greater the likelihood that they are hydrogenated. For reaction class (B) in Table 7, the observed ratio of ethanol to ethene is higher than predicted. This means that alcohols are not formed by the hydration of alkenes in secondary reactions. The reverse reaction is, of course, possible. The data for reaction (C) show that the ratio of ethanol to ethane is very much higher than expected. From cases (B) and (C) it therefore appears that the alcohols must be primary products. For reactions (D) and (E), it can be seen that ethanol, acetic acid and acetaldehyde are in equilibrium at 600 K and so these reactions do take place. (This is confirmed by the observation that when ethanol or acetaldehyde or acetic acid, is added to the syngas, additional amounts of the other two compounds are always produced [21].) At 510 K, however, the ratios observed are much higher than expected. This implies that acetic acid is not formed by subsequent oxidation of ethanol by water, nor is acetaldehyde formed by dehydrogenation of ethanol. Hence it appears that aldehydes and acids too may be primary products. The data for reaction (F) show that the ratio of branched to linear pentenes is much lower than expected. This indicates that the FT mechanism favors the formation of linear products and also that the secondary isomerizations are slow. The branched FT hydrocarbons are predominantly methyl branches. With iron catalysts at 340 ◦ C, the degree of branching increases as the carbon number of the product increases [83]. At 220 ◦ C, however,

the degree of branching decreases with increasing chain length. Ketones and iso-alcohols only appear to be produced in significant amounts at higher temperatures. This could be because they are secondary products. From Table 7, class (G), the ketonization of acetic acid is feasible. This is supported by the fact that, as the temperature is increased from 310 to 380 ◦ C (see Table 10), the acid selectivity decreases and the ketone selectivity increases up to about 360 ◦ C, after which it decreases. The latter decrease could be due to the hydrogenation of ketones to isoalcohols. When at 340 ◦ C, acetic acid is deliberately added to the syngas, additional acetone is produced [21]. On feeding additional ethanol, there was also an increase in acetone production [21]. As reaction (D) is at equilibrium at 610 K, adding ethanol to the feed should increase the production of acetic acid, and ketonization of the latter could account for the increase in acetone. The FT ketones are predominantly methyl ketones. Since the dominant acid formed is acetic acid, it follows that, if ketonization of the mixture of acids takes place, then the methyl ketone would emerge as the dominant isomer for each carbon number ketone. For reaction (H), it can be seen that at 610 K the hydrogenation of acetone to 2-propanol has proceeded to equilibrium, but at 510 K the 2-propanol/acetone ratio is lower than expected. This suggests that the hydrogenation of acetone is slow at the lower temperature and so this may indicate that iso-alcohols are secondary products. Below about 240 ◦ C no aromatics are produced, but, as can be seen in Table 10, at higher temperatures the amount of aromatics in the gasoline progressively increases. It is possible that alkenes and aromatics are formed from common surface precursors. The thermodynamically unfavored monoalkylbenzenes predominate, and this fits the concept that aromatics are formed by linkage between the first and sixth carbon atoms of the carbon chain followed by dehydrogenation [6]. There are also significant amounts of unsaturated ring compounds present, which fits the proposed route to aromatics. Of the aromatics only about 1% is benzene. From Table 7, reaction (I), it can be seen that the expected ratio of toluene to heptene is very much greater than that obtained. If aromatization was a secondary reaction then it could be expected that recycling of the oil would increase the aromatic content but, as already mentioned, this did not happen. It is therefore possible that the formation of ring compounds and aromatics are primary reactions, which occur very slowly at FT temperatures. (Table 10 shows that at higher temperatures more aromatics are formed.) At the H2 pressures existing in FT reactors, hydrocracking of hydrocarbons is thermodynamically favorable but in practice it hardly occurs, if at all [6].

13.15.7 FT Product Selectivities

Product Distributions The selectivity, on a carbon atom basis, of methane can vary from 1 to 100%, whereas that of long-chain waxes can vary from 0 to over 70% [6, 21]. The spread of products is determined by the operating conditions (see Section 13.15.7.3), but there is nevertheless always a clear interrelationship between the individual products, as illustrated in Figs. 9 and 10 [6, 21] (in Ref. [6], see Chapter 3). The explanation is that FT involves a stepwise chain growth procedure of ‘‘monomers’’ (see Section 13.15.9). If it is assumed that the probability of chain growth (α) is independent of chain length, then the entire product spectrum can be calculated. These calculations show, as observed in practice, that all the intermediate carbon number cuts go through maxima as the probability of chain growth increases [6, 21]. There is a reasonably good agreement between the calculated and the experimental distributions. A clear exception is the C2 hydrocarbons: the calculations predict a maximum selectivity of about 30% whereas in practice it never exceeds 20%. This misfit is discussed in Section 13.15.9. A mathematical equation for the stepwise chain growth concept was developed by Herrington [84], Friedel and Anderson [85] and Flory [86] and is known as the ASF equation: 13.15.7.2

log(Wn /n) = n log α + constant 40 35

C5 -200 °C

Selecticity/ %

30 25 C2

20 15

C3

10 200–320 °C

5 0

0

10

20

30

40

50

60

70

CH4 selectivity / % Relation between the selectivities of the hydrocarbon cuts for the HTFT process. The selectivities are on a C atom% basis.

40 35 30

Selectivity/%

Overall, it can be concluded that once the primary hydrocarbons have desorbed they are not readily readsorbed, because they compete poorly with more strongly adsorbing species such as CO and H2 O [6]. The only significant exceptions are ethene and the lower molecular mass oxygenates.

2983

C5 - 200 °C

25 370 – 500 °C

20 15 10

200 – 320 °C

5 0

0

10

20

30

40

50

60

Hard wax selectivity/% Fig. 10 Relation between the selectivities of the hydrocarbon cuts for the LTFT process. The selectivities are on a mass% basis.

where Wn is the mass fraction of the species with carbon number n. From the slope of the plot of log(Wn /n) against n, the value of α is obtained. It is usually found that over the carbon number range from 3 to about 12 the plots are linear and so over this range α is constant. (Note that usually the C2 and C1 compounds do not fall on the line.) When considerable amounts of waxes are produced, it is found that in the vicinity of carbon number 12 the ASF plot shifts to yield another ‘‘straight’’ line which translates to a higher chain growth probability (α2 ). This ‘‘double alpha’’ effect is observed for both iron and cobalt catalysts. For the Fe commercial LTFT process α2 is typically 0.95 at the start of a run. The proposals as to why α changes to a higher value are varied. A review concluded that multiplicity of the chain growth probability was the main cause of the shift in the alpha values [87]. Based on the assumed longer residence times of the products in the wax-filled pores, alkene re-incorporation has previously been offered as the explanation of the ‘‘double alpha’’ phenomenon. Different reasons have been suggested to explain the longer residence times, namely slower diffusion rates, higher solubility in the liquid phase and stronger adsorption of the longer chain products [62, 79, 88–92]. Typical industrial product distributions and analyses are shown in Tables 8 and 9. As can be seen, LTFT produces large quantities of waxes but no aromatics. Co, being more hydrogenating, produces less alkenes and oxygenates than Fe. With the same LTFT Fe catalyst it is found that under the same operating conditions the smaller particles used in slurry bed reactors result in higher alkene/alkane ratios than do the larger particles used in fixed-bed reactors [6]. This suggests that as a result

Fig. 9

References see page 2992

2984

13.15 The Fischer–Tropsch (FT) Synthesis Processes

of the longer residence times inside the larger particles, some of the primary alkenes are hydrogenated to alkanes. The HTFT produces some aromatics, methyl-branched hydrocarbons and large amounts of alkenes, the majority of which are linear 1-alkenes. Of the alcohols, aldehydes and acids, the most abundant are the C2 compounds. A relatively large amount of methanol is produced in the LTFT process but not in the HTFT process. HTFT makes more acids and consequently more ketones and iso-alcohols (see Section 13.15.7.1). Control of Selectivity The main factors which alter the product selectivities are the temperature, the catalyst metal type, the promoters and the partial pressures of the gases inside the reactor. 13.15.7.3

13.15.7.3.1 Temperature Whatever the catalyst, the methane selectivity increases with increasing temperature. As desorption is endothermic, one would expect the probability of chain termination to increase with increasing temperature, thus resulting in a shift to lighter products. At higher temperatures the rates of all reactions increase and so the situation should shift towards thermodynamic predictions, namely the formation of the more stable products such as methane, branched alkanes

Tab. 9

Commercial FT hydrocarbon isomers

Cuta

Catalyst

C5 –C12 cut % Alkanes % Alkenes % Aromatics % Oxygenates C13 –C18 cut % Alkanes % Alkenes % Aromatics % Oxygenates C24 –C35 cut % Alkenes C4 cut % 1-Butene % Me-1-propene C6 cut % 1-Hexene % Me-1-pentenes C10 cut % 1-Decene % Me-1-nonenes

Cobalt 220 ◦ C

Iron 235 ◦ Cb

Iron 340 ◦ Cc

60 39 0 1

29 64 0 7

13 70 5 12

95 5 0 Low

44 50 0 6

15 60 15 10

Low

10

–

– –

– –

74 8

– –

– –

58 24

– –

– –

38 20

a Percentages

are mass%. bed operation (LTFT). c Fluidized bed operation (HTFT). b Slurry

Tab. 8

Commercial FT product spectra (pressure 2 MPa)

Selectivities (C atom basis)

CH4 C2 H4 C2 H6 C3 H6 C3 H8 C4 H8 C4 H10 C5 −C6 C7 −160 ◦ Ca 160−350 ◦ Ca >350 ◦ Ca Total water-soluble oxygenatesb ASF α-valuec a The

Catalyst Reactor type Temperature/ ◦ C Cobalt Slurry 220

Iron Slurry 235

4 0.05 1 2 1 2 1 8 11 22 47 1

3 0.5 1 2.5 0.5 3 1 7 9 17.5 51 4

0.92

0.95

Iron Fluidized 340 8 4 3 11 2 9 1 16 20 16 5 5 0.7

temperatures indicate the boiling points of the respective product fractions at 1 bar. b Alcohols, aldehydes, ketones and acids dissolved in the water phase. c ASF: Anderson–Schulz–Flory probability of chain growth.

and aromatics (see Table 7). Table 10 illustrates some of the effects of temperature on the HTFT process. As Co is more active than Fe, the methane selectivity increases more rapidly with increasing temperature. It was found in previous studies with Co that, when operating at 165 ◦ C, large amounts of oxygenates were formed (the oil fraction contained about 40% alcohols) whereas at 200 ◦ C only 1% was present [5, 93]. At about 170 ◦ C Ru produces high molecular mass waxes [54, 55] whereas at 400 ◦ C only methane is produced. 13.15.7.3.2

Catalyst Metals and Promoters

A Cobalt Catalysts Co, because of its high price, has to be effectively supported in order to obtain a high surface area per unit mass of Co. Unlike Fe, the selectivity of Co is less sensitive to the chemical nature of promoters or supports. The early German FT studies with Co/kieselguhr showed that, at atmospheric pressure, alkali promotion increased the wax selectivity. However, at higher pressures the effect of alkali was minimal [27]. Sasol investigated the effect of adding, individually, small amounts of the oxides of Cr, Mg, K and Th to alumina-supported Co and concluded that none of the additives had any beneficial

13.15.7 FT Product Selectivities

Tab. 10

Influence of temperature on an Fe catalyst in HTFT operation [21] Selectivity/% (C atom basis)

Temperature/ ◦ C CH4

310 330 350 360 370 380 a The

2985

10 14 17 20 23 28

Gasoline cut analysis

Water-soluble chemicals Alcohols

Ketones

Acids

2.3 2.3 1.6 1.1 0.8 0.5

0.4 0.8 1.1 1.3 1.2 0.8

0.3 0.4 0.2 0.2 0.1 0.1

Br numbera

% Aromatics

109 94 92 93 88 85

4 8 10 13 18 26

C3 H6 /C3 H8 ratio

6 10 9 8 6 4

bromine number is an indicator of the amount of alkenes present.

effects [48]. Iglesia studied the performance of Co supported on Al2 O3 , on SiO2 and on TiO2 . On the basis of unit metal surface area, all three catalysts had the same activity [62]. Hence the oxide supports had no chemical promoting effect. Exxon, using TiO2 or SiO2 as supports, found that the addition of low levels of ruthenium increased the selectivity of higher molecular mass hydrocarbons [46]. Goodwin and coworkers reported that Ru promotion of Co/alumina had no effect on selectivity [47]. It was reported that promoting Co/alumina or Co/silica with Pt did not affect the selectivity [94]. The selectivity of Co/silica was found to be enhanced by lanthanum oxide [95] and also by zirconium oxide [96]. B Iron Catalysts Unlike for Co, alkali promotion is vital for both the activity and selectivity of Fe catalysts [21]. Because the promoters K and Na have such a dominant influence, it is debatable whether other compounds or supports have, in their own rights, any marked influence. They can, however, have a marked influence on the effectiveness of the alkali promoter. Thus if there is any silica or alumina present, the alkali can react with these to form alkali metal silicates or aluminates, which will be less basic than the ‘‘free’’ alkali and hence the ‘‘basicity’’ of the iron surface would be lowered. Even if the supports do not interact chemically with the alkali, the alkali will not only be distributed on the iron surface but also on the surface of the support and thus the basicity of the iron surface will be lowered. Overall, this means that the amount of alkali that has to be added needs to be adjusted to take into account the chemical nature and the surface areas of the other components present. The Group I alkalis as promoters have been investigated [21]. As expected from their relative basicities, the wax selectivities increased in the order Li, Na, K and Rb. (The Li-promoted catalyst also had a much lower activity than the other cases.)

Table 11 shows the strong influence of K promotion on the activity and selectivity of the precipitated iron catalyst in the LTFT process [21]. It has been reported previously that, as the alkali level is increased, the activity increases up to a certain point and then decreases at higher alkali levels [27]. In Table 11, this peak in activity was only observed in the unsupported series. In the supported series, the alkali levels were apparently all beyond the peak level. As the wax selectivity increases, their average chain length increases, and this probably results in lower diffusion rates inside the wax-filled pores, which results in lower overall FT reaction rates. As for precipitated catalysts, the dominant promoters for HTFT catalysts are the strong Group I alkalis. Higher surface areas (due to structural promotion) coupled with a fixed amount of alkali lead to a lower surface basicity. If the promoters or other compounds present, chemically interact with the alkali, then the basicity of the alkali will be lowered. Increasing the silica/alkali ratio (at a fixed alkali amount) from 0.8 to 4.2 resulted in an increase in methane selectivity from 8 to 27% [6, 21]. Table 12 shows that, as the K level is increased, the selectivity shifts to longer chain products, as reflected by a decreasing methane selectivity and also an increase in the amount of alkenes and oxygenated products. The effect of sodium, which is a weaker alkali, is much less pronounced [6, 21]. The way in which the K promoter is added to the catalyst used in a fluidized bed appears to have little effect. When potassium silicate powder was added on-line to an alkali-free catalyst, the FT performance was similar to that of the catalyst prepared by adding the alkali during the fusion step [21]. This result is not surprising since the fused catalyst was in any event a heterogeneous mixture of alkali metal silicate and magnetite (see Section 13.15.5.1.2A). It appears that under FT conditions the alkali migrates over the surface from particle to particle. It nevertheless seems doubtful if References see page 2992

2986

13.15 The Fischer–Tropsch (FT) Synthesis Processes Influence of alkali content on synthesis performance of precipitated iron catalysts [21]

Tab. 11

Catalyst type

Unsupported Fe2 O3

Silica supported Fe2 O3

K2 O levela

Hard wax selectivityb

0 1.0 1.6 2.0 3.0 12 16 21 24 32

5 34 41 53 63 18 20 30 38 44

Water analysis/wt.% Alcohols

2.5 3.0 3.4 2.0 2.0

Activitya

Acids 500 ◦ C) selectivity as a function of the H2 /CO ratio of the feed gas (LTFT operation).

Fig. 11

Table 15 shows a set of results in which the various partial pressures were varied over wide ranges. Again it can be seen that there is no correlation between the selectivity and the H2 /CO ratio, but as illustrated in Fig. 12, there is a reasonable correlation with the partial /(PCO + 0.7PCO2 + 0.6PH2 O ). pressure factor PH0.25 2

mass products [27, 99, 100]. This is similar to Fe LTFT catalysts. However, contrary to Fe LTFT, but similarly to Fe HTFT catalysts, increasing the pressure over Co increased the probability of chain growth [101, 102]. In a more recent study with Co/alumina, it was found that the wax selectivity increased from virtually zero at 0.2 MPa to about 60% at 4.0 MPa [58]. The effect of pressure on Ru is similar [54]. The dominant factor controlling the wax selectivity appeared to be the CO partial pressure [103]. Combining the effects of PH2 /PCO ratio and of total pressure, the selectivity of Co catalysts could be proportional to a factor such as PHx 2 /PCO where x < 1. On feeding only CO2 and H2 over Co, methane was the main product, whereas under the same conditions large amounts of liquids were produced when CO and H2 were fed [104, 105]. Note that when operating with iron at 330 ◦ C a feed gas consisting of only CO2 and H2 produces a product spectrum similar to that of normal syngas (see Table 15). The reason is that Fe is an active WGS catalyst, and so CO2 is converted to CO, hence normal FT takes place. Alcohol Production The amount of alcohols produced is normally low under commercial FT conditions. The production of alcohols 13.15.7.4

B Cobalt Catalysts When the PH2 /PCO ratio was lowered. The spectrum shifted to higher molecular

13.15.8 Versatility of the FT Process and Product Quality Improvement

2989

FT reaction at 335 ◦ C [6]. For all runs the total feed consisted of fresh feed plus recycled tail gas after water and oil knock-out. The listed partial pressures are those after WGS and before FT synthesis

Tab. 15

CH4 selectivitya

PH2 /MPa

PCO /MPa

PCO2 /MPa

PH2 O /MPa

0.25 PH /(PCO + 0.7PCO2 + 2 0.6PH2 O )/MPa−0.75

PH2 /PCO

22.1 20.8 18.3 17.7 15.0b 14.5 13.8 13.7 13.5 13.4 11.4 10.9 10.8 10.5 7.1b 6.8 6.5b 5.5

1.155 0.49 0.427 1.159 1.322 0.931 0.853 0.964 1.345 1.175 1.148 1.766 1.129 1.422 0.914 2.46 0.766 3.073

0.145 0.098 0.102 0.164 0.104 0.184 0.195 0.19 0.112 0.296 0.064 0.366 0.241 0.358 0.171 0.526 0.183 0.67

0.006 0.049 0.085 0.023 0.207 0.051 0.15 0.056 0.234 0.01 0.277 0.117 0.162 0.282 0.672 0.379 0.903 0.455

0.002 0.008 0.015 0.006 0.094 0.01 0.024 0.01 0.1 0.001 0.159 0.021 0.025 0.035 0.134 0.061 0.124 0.071

6.89 6.10 4.74 5.65 3.51 4.35 3.06 4.21 3.21 3.43 2.93 2.50 2.79 1.89 1.35 1.51 1.05 1.28

8.0 5.0 4.2 7.1 12.7 5.1 4.4 5.1 12.0 4.0 17.9 4.8 4.7 4.0 5.3 4.7 4.2 4.6

aC

atom%. feed consisted only of CO2 and H2 .

Methane selectivity / carbon atom%

b Fresh

that, as the amount of alcohols decreases, the amount of alkenes increases, which suggests that some of the alkenes could be formed via dehydration of the alcohols (see Section 13.15.7.1).

25 20 15

13.15.8

Versatility of the FT Process and Product Quality Improvement

10 5

Gasoline The straight-run HTFT gasoline consists of predominantly linear hydrocarbons, and the aromatics content is only about 5%, therefore it has a low octane rating. To increase the octane number, it is hydrotreated to convert alkenes and oxygenates to alkanes. The C5 –C6 cut can then be isomerized to branched alkanes, and the C7 –C10 cut can be Pt-reformed to a mixture of branched alkanes and aromatics. To increase the gasoline yield, the C3 and C4 alkenes (about 20% of the total FT product) can be oligomerized over a non-selective acid catalyst to produce highly branched alkenes. The octane number of the final gasoline pool can be boosted further by adding ethanol or ethers such as TAME (tertiary amyl methyl ether) or DIPE (di-isopropyl ether). All of these additives are either direct FT products or can be produced from the FT products. 13.15.8.1

0 0

2

PH2

0.25

4

(PCO + 0.7PCO2 + 0.6PH2O

6

8

)−1/ MPa−0.75

Fig. 12 The selectivity of methane as a function of the partial pressure factor (HTFT operation). The unit of the partial pressure factor is MPa−0.075 .

can, however, be increased by operating at high pressures and space velocities but at relatively low temperatures (see Chapter 3 in Ref. [6]). For both Fe and Co the alcohol content of the oil fraction can vary between 30 and 65%. The alcohols are primary products and may be formed by CO insertion (see Section 13.15.9). The high pressure and low temperature probably ensure that there is a large amount of undissociated CO on the catalyst surface, which should enhance CO insertion. There are indications

References see page 2992

2990

13.15 The Fischer–Tropsch (FT) Synthesis Processes

Diesel Fuel The very factors which count against FT gasoline quality, namely linearity and low aromatic content, are very positive factors in favor of high cetane number diesel fuel. For maximum production of diesel fuel, the LTFT slurry bed operating in the high wax selectivity mode is recommended. The straight-run FT diesel makes up about 20% of the total FT product and has a cetane number of about 75. (At present the required cetane number of diesel fuels varies from 40 to 50, depending on the location.) Mild hydrocracking of the wax makes the largest contribution to the final diesel fuel pool. Hydrocracking with standard bifunctional catalysts was investigated at Sasol during the 1970s [106, 107]. The wax yielded about 80% diesel, 15% naphtha and 5% C1 –C4 gases. The above product spread appears to be the result of random β-scission of the wax chains. There is a big incentive to improve the selectivity of the wax hydrocracking in order to increase the diesel yield. Some chain branching occurs during hydrocracking but the final diesel pool nevertheless has a cetane number of about 70. Due to the mild FT and wax hydrocracking conditions used, the aromatic content of the diesel fuel is very low. The advantage of producing such a highquality diesel is either that it can be used in areas where there are strict constraints regarding automobile exhaust gases or it can be used as a blending stock to upgrade lower quality diesel. The naphtha produced in wax hydrocracking consists only of alkanes and so is an excellent feedstock for the production of ethylene by steam cracking. The yield is much higher than obtained from crude oil naphtha. 13.15.8.2

Chemicals The HTFT process with Fe catalysts produces large amounts of linear 1-alkenes. As the purified 1-alkenes are sold as chemicals they fetch much higher prices than when sold as fuels. The C2 –C8 alkenes are used in the production of a variety of polymers. The longer chain 1-alkenes can be selectively converted to high-value primary alcohols by hydroformylation. The LTFT processes make predominantly longer chain linear alkanes. After catalytic hydrofining, the oils and various grades of waxes are sold at high prices. 13.15.8.3

that, based on the ‘‘correct’’ mechanism, a catalyst has been developed that will markedly improve the selectivity, which is a key factor in the economics of the FT process. This therefore remains a major challenge. Development of Various Models The original proposal was that CO carbided the metal surface, which was then hydrogenated to methylene groups, which polymerized to hydrocarbons [113, 114]. Because this did not explain the formation of alcohols, the ‘‘enol’’ mechanism was proposed [5, 27]. The adsorbed CO was hydrogenated to adsorbed HCOH monomers. Neighboring species link up with the elimination of water. A brief version of the scheme is depicted in Fig. 13. The ‘‘CO insertion’’ (or ‘‘carbonyl’’) model could also explain the formation of oxygenated products [115, 116]. A version thereof is depicted in Fig. 14. The ‘‘alkyl’’ or ‘‘methylene’’ mechanism is an improved version of the original carbide one and for some time now has remained the most commonly accepted scheme. Studies in which either alkyl chlorides [117] or diazomethane [118] were utilized supported the concept that CH2 monomers and not oxygenated ones were involved in chain growth. A version of this mechanism is shown in Fig. 15. As can be seen, one of the chain termination steps is the β-abstraction of hydrogen to yield alkenes. This concept is surprising in view of the fact that, under industrial conditions, where large amounts of alkenes are formed, thermodynamics highly favor hydrogenation over dehydrogenation. Note that for the ‘‘adsorbed 13.15.9.1

R C

OH + H C

OH → R C

↓ +H2 RCH2CH2OH ← H2 +

RCH2COH + H2 → RCH2CH + H2O

RCH2CH3 ← H2 + RCH2CH → RCH

Fig. 13

CH2

An ‘‘enol’’ FT reaction mechanism.

RCH2CH

13.15.9

Possible Surface Reactions and Mechanisms

For 80 years, the details of the FT mechanism have been a controversial matter. Several mechanisms and variations thereof have been proposed and reviewed [5, 6, 21, 27, 108–112] (in Ref. [6], see Chapters 3 and 8). There is no doubt that the mechanism is of intriguing scientific interest. However, there is no convincing evidence to date

OH + H2O

C

C

O

↑ +CO RCH + CO → RCH C

RCH

Fig. 14

O + 2H2 → RCH2CH + H2O

CH2 ← RCH2CH + H2 → RCH2CH3

A ‘‘CO insertion’’ FT reaction mechanism.

13.15.10 Environmental Aspects

CO → C + O

C + H2 → CH2 CH2 + H → CH3 CH2R

CH2R + CH2 → CH2 → RCH

CH2R CH2

CH2 + H

+ H → RCH2CH3 + OH → RCH2CH2OH

Fig. 15

An ‘‘alkyl’’ FT reaction mechanism.

alkene’’ [109] scheme or for alternative schemes [21, 110, 119, 120], dehydrogenation is not required for the formation of alkenes. Initial Reactions and Formation of Monomers As mentioned, it has been variously proposed that the monomers could be adsorbed CO, HCOH or CH2 . In all cases the initial step must be the adsorption of CO and it is conceivable that this is followed by the formation of the other two monomer species: 13.15.9.2

CO + H2 −−−→ HCOH HCOH + H2 −−−→ CH2 + H2 O Hence all three monomers could be present on the surface to varying degrees, depending on the catalyst type and on operating conditions. The formation of methanol could occur via the hydrogenation of HCOH. In the ‘‘alkyl’’ mechanism it is assumed that the adsorbed CO dissociates to C and O and the C is then hydrogenated to CH2 and the O to H2 O. When only CO is fed to iron catalysts at normal FT temperatures, carbides are initially formed and so CO decomposition certainly occurs and the O is removed as CO2 . At lower temperatures the process ceases once the Fe is completely carbided, but at higher temperatures the deposition of ‘‘free’’ C continues [21]. The question arises of why CO, which is a powerful reducing agent, can remove O from the Fe surface but not from the Co surface (over Co catalysts little or no CO2 is formed during FT synthesis). Is it possible that the majority of the CH2 monomers are formed via the HCOH species and not via CO decomposition?

carbon number distributions. Even if two or more ‘‘different’’ mechanisms operate at the same time, the overall distribution would probably still be as per ASF. The mix of alkenes, alkanes, oxygenated and other products could then depend on the relative contributions of the various mechanisms. In Chapter 3 in Ref. [6] the possibility of how the surface species, CO, HCOH, CH2 , H, OH and H2 O, could be involved in the FT reactions is discussed. The Misfits of C1 and C2 Compounds in the ASF Distributions As mentioned in Section 13.15.7.2, the C2 compounds and methane do not fall on the ASF plots. The C2 compounds lie below the line and C1 , particularly so for Co catalysts, lies above the line. The ASF deviations are discussed in Chapters 3 and 8 in Ref. [6]. An isolated CH2 monomer (see Fig. 16a) is likely to be rapidly hydrogenated, and this alone could explain the higher than expected methane yields. One of the reasons for the low levels of C2 compounds could be their re-incorporation into growing chains. Alternatively, if the C2 surface specie is taken as adsorbed ethylene, the addition of another CH2 could occur on either side of it, which would increase the probability of chain growth (see Fig. 16b). For C3 and higher adsorbed alkenes, however, a CH2 , for steric hindrance reasons, is more likely to attach to the ‘‘open’’ side of the adsorbed alkene (see Fig. 16c). 13.15.9.4

13.15.10

Environmental Aspects (see Chapter 5 in Ref. [6])

In the production of syngas, all sulfur and nitrogen compounds are removed upstream of the FT reactors. The FT fuels are therefore S- and N-free. Hence the exhaust gases from combustion engines are free of SO2 . The NOx levels are lower because the FT fuels themselves make no contribution. Benzene is carcinogenic but the FT gasoline contains less than 1%, which is much lower than in crude oil-derived gasoline. The use of diesel is being encouraged because the efficiency of a diesel engine is about 44% as against about H2 → CH2 → CH4 (a)

(b)

Possibility of Various Mechanisms Operating Simultaneously It seems unlikely that each type of FT metal has its own unique mechanism. It is feasible that different mechanisms take place on the catalyst surface at the same time. Irrespective of the metal type or of the process conditions, the products always have ASF (see Section 13.15.7.2)

2991

13.15.9.3

(c)

CH2 → CH2

CH2 ← CH2

CH2 → CH2

CH CH2R

Possible reasons for the deviations of methane and of C2 hydrocarbons from the ASF product distribution. Fig. 16

References see page 2992

2992

13.15 The Fischer–Tropsch (FT) Synthesis Processes

24% for a gasoline engine. The diesel fuel produced in the LTFT process is of excellent quality (see Section 13.15.8.2). The aromatics content is less than 1% as against currently about 30% in diesel fuels produced from crude oil. The exhaust emission levels of hydrocarbons, CO, NOx and particulate matter are 56, 33, 28 and 21%, respectively, lower than with crude oil-derived diesel fuels [121]. In addition, FT diesel is readily biodegradable. In a coal-based plant there are two aqueous streams and one gaseous stream that have to be treated. The ‘‘raw gas liquor’’ stream is from the coal gasification step. After recovery of the ammonia and phenols, the water still contains low levels of organics, and these are destroyed in aerobic biodigesters. The FT synthesis water, after extraction of the desired oxygenated compounds, contains some organic acids, and this stream is also treated in the bioworks. The product water from the digesters is recycled to the plant for use as cooling water. The raw synthesis gas from the coal gasifiers is scrubbed to remove the CO2 and H2 S. The H2 S-containing CO2 stream is treated in a Sulfreen or similar process where the H2 S is oxidized to elemental sulfur. In a methane-based plant the situation is simpler. The low levels of sulfur impurities are removed upstream of the CH4 reformers. The FT water stream, after extraction of the oxygenates, contains only some organic acids, which biodegrade readily and can be purified in either aerobic or anaerobic digesters to produce good-quality cooling water. An advantage of using CH4 for syngas production is that it is much more carbon efficient than when using coal and hence much less CO2 is produced. As CO2 is one of the greenhouse gases, there is currently much emotional and political pressure to minimize CO2 emissions, even though the scientific evidence regarding the extent of the contribution of CO2 to global warming is as yet unresolved [8]. References 1. P. Sabatier, J. B. Senderens, C. R. Acad. Sci. 1902, 134, 514. 2. Badische Anilin und Soda Fabrik, German Patent 293 787, 1913. 3. F. Fischer, H. Tropsch, Brennst. Chem. 1923, 4, 276. 4. H. H. Storch, R. B. Anderson, L. J. E. Hofer, C. O. Hawk, H. C. Anderson, N. Golumbic, Synthetic Liquid Fuels from Hydrogenation of Carbon Monoxide, Bureau of Mines Technical Paper 709, US Government Printing Office, Washington, DC, 1948. 5. H. H. Storch, N. Golumbic, R. B. Anderson, The Fischer– Tropsch and Related Syntheses, Wiley, New York, 1951, 610 pp. 6. A. P. Steynberg, M. E. Dry (Eds.), Fischer–Tropsch Technology. Studies in Surface Science and Catalysis, Vol. 152, Elsevier, Amsterdam, 2004, 689 pp. 7. A. C. Vosloo, Fuel Process. Technol. 2001, 71, 149. 8. M. E. Dry, in Encyclopedia of Catalysis, I. T. Horvath (Ed.), Vol. 3, Wiley, Hoboken, NJ, 2003, p. 347.

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111. V. Ponec, in Handbook of Heterogeneous Catalysis, G. Ertl, H. Kn¨ozinger, J. Weitkamp (Eds.), Vol. 4, VCH, Weinheim, 1997, p. 1876. 112. P. M. Maitlis, R. Quyoum, H. C. Long, M. L. Turner, Appl. Catal. A: General 1999, 186, 363. 113. F. Fischer, H. Tropsch, Brennst. Chem. 1926, 7, 97. 114. S. R. Craxford, E. K. Rideal. J. Chem. Soc. 1939, 1604. 115. A. Steenberg, I. Wender, in Proceedings of International Conference on Coordination Chemistry, Chemical Society, London, 1959, p. 53. 116. H. Pichler, H. Schulz, Chem. Ing. Tech. 1970, 42, 1162. 117. W. A. van Baarneveld, V. Ponec, J. Catal. 1984, 88, 382. 118. R. C. Brady, R. Pettit, J. Am. Chem. Soc. 1981, 103, 1287. 119. M. E. Dry, Catal. Today 1990, 6, 183. 120. M. E. Dry, in Applied Industrial Catalysis, B. E. Leach (Ed.), Vol. 2, Academic Press, New York, 1983, p. 167. 121. P. W. Schwaberg, I. S. Myburgh, J. J. Botha, P. N. Roets, L. P. Dancuart, Presented at the 11th World Clean Air Congress, Durban, South Africa, September 1998.

13.16

Gas-to-Liquids Tom J. Remans, Gregor Jenzer, and Arend Hoek∗

13.16.1

Introduction to GTL

Gas-to-liquids (GTL) is the integrated process of the manufacture of synthesis gas (syngas) from natural gas, the subsequent conversion of the syngas into hydrocarbons via the Fischer–Tropsch (FT) reaction, followed by a liquid processing unit. For coal-to-liquids (CTL) and biomass-to-liquids (BTL), coal and biomass, respectively, are used in the syngas production step. The liquid processing unit typically consists of hydrogenation and hydroconversion reactors to give the final products the right properties and to yield the highest value product slate. The history of GTL started in 1925 with a patent by Franz Fischer and Hans Tropsch on the production of paraffin hydrocarbons from carbon monoxide and hydrogen by catalytic paths [1], later published in a journal article [2]. They built on the CO hydrogenation work on metallic catalysts by Sabatier in 1902 [3] and a 1913 BASF patent on hydrocarbon and oxygenate synthesis from carbon monoxide and hydrogen on oxide catalysts at very high pressure and temperature [4]. Since 1925, interest in large-scale application of GTL and CTL has mainly been triggered by specific political and economic situations. During World War II, 20% of the German gasoline demand was produced by CTL [5]. The first German ∗

Corresponding author.

plant was started up in 1936. By the 1950s, all nine CTL plants with a total production capacity of ca. 700 000 t a−1 [16 000 bbl day−1 (bbl = barrels)] were stopped. In South Africa, 31% of the fuel demand in the 1990s was produced by CTL and 10% by GTL [6]. Most of the old South African plants have been replaced by larger ones, now having a total production capacity of ca. 7 300 000 t a−1 (165 000 bbl day−1 ) by Sasol and ca. 2 100 000 t a−1 (47 000 bbl day−1 ) by PetroSA. These processes are based on high-temperature Fischer–Tropsch synthesis in fluidbed reactors. In 1993, Shell’s GTL plant, based on the Shell middle distillate synthesis (SMDS) technology, startedup in Bintulu, Malaysia, as a first step towards proving GTL to be an economically viable alternative to liquefied natural gas (LNG) at remote gas fields. The Bintulu plant applies low-temperature Fischer–Tropsch synthesis and currently has a production capacity of ca. 580 000 t a−1 (14 700 bbl day−1 ). A short-lived large-scale application of low-temperature Fischer–Tropsch chemistry was the ca. 300 000 t a−1 (7000 bbl day−1 ) ebullated bed reactor operated from 1948 to 1953 at Brownsville, TX, USA by a consortium led by Hydrocarbon Research [7]. An overview of all CTL and GTL plants is given in Table 1, showing a total current GTL + CTL capacity of ca. 11 500 000 t a−1 (260 500 bbl day−1 ). In 2007, Oryx, a joint venture of Qatar Petroleum (51%) and Sasol (49%), started its ca. 1 400 000 t a−1 (34 000 bbl day−1 ) GTL plant at Ras Laffan, Qatar. Oryx plans for a ca. 2 700 000 t a−1 (68 000 bbl day−1 ) expansion of their existing Ras Laffan GTL plant by 2009. It uses Sasol’s slurry phase distillate Fischer–Tropsch technology, which will also be applied in the ca. 1 400 000 t a−1 (34 000 bbl day−1 ) GTL plant currently under construction at Escravos, Nigeria, and scheduled to start-up in 2008. At Ras Laffan, Qatar Petroleum has a ca. 5 600 000 t a−1 (1 40 000 bbl day−1 ) GTL plant under construction based on Shell’s GTL technology, scheduled to start up in 2009–11. Plans for a ca. 6 200 000 t a−1 (154 000 bbl day−1 ) GTL plant based on ExxonMobil’s GTL technology were shelved in 2007. In 2005, plans for additional world-scale GTL plants in Qatar by Oryx (ca. 5 700 000 t a−1 or 130 000 bbl day−1 ), Marathon Oil (maximum 5 300 000 t a−1 or 120 000 bbl day−1 ) and Conoco-Phillips (ca. 7 000 000 t a−1 or 160 000 bbl day−1 ) have been put on hold for a minimum of 3 years [8]. Many other GTL projects are being studied in gas-rich regions such as Australia, Russia, Bolivia, Venezuela, Iran, Nigeria and Egypt. Countries with large coal reserves are also looking into CTL for additional energy security. Several CTL projects are being studied in the USA, holding the world’s largest coal reserves, and in China, positioned third in the world coal reserves ranking.

13.16.2 Basic Steps in GTL

Tab. 1

2995

Historical overview of large-scale CTL and GTL applications

Company

Location

Capacitya /t a−1

Reaction and catalystc

FT reactor (No. of reactors)

Operating period

Coal-to-liquids plants: current and past operations > ca. 45 000 t a−1 (1000 bbl day−1 ) – Sasol Sasol

Germany Sasolburg, South Africa Secunda, South Africa

700 000b 110 000 4 600 000b

LTFT by Co LTFT by Fe HTFT by Fe

Sasol

Secunda, South Africa

7 300 000

HTFT by Fe

Sasol

Sasolburg, South Africa

110 000

LTFT by Fe

ARGE multitubular fixed-bed (9) ARGE multitubular fixed-bed (5) Synthol, fluidized bed (16)

1936–1950s 1955–present (Sasol-I) 1980 and 1982–late 1990s (Sasol-II and -III) Sasol advanced Synthol, ebullated 1st in 1989, last in bed (8) 1999–present (replaced Synthol) Sasol slurry phase distillate (1) 1993–present

Gas-to-liquids plants: current and past operations > ca. 45 000 t a−1 (1000 bbl day−1 ) Consortium Brownsville, TX, USA PetroSA Mossel Bay, South Africa Shell Bintulu, Malaysia Oryx Ras Laffan, Qatar

300 000b 2 100 000

HTFT by Fe HTFT by Fe

Ebullated bed Fluidized bed

580 000 1 400 000

LTFT by Co LTFT by Co

Multitubular fixed bed by Shell (4) 1993–present Sasol slurry phase distillate (2) 2007–present

1948–1953 1991–present

estimated via known bbl day−1 at liquid density of 700 g L−1 . taken out of operation. c LTFT = low-temperature Fischer–Tropsch; HTFT = High-temperature Fischer–Tropsch. a Capacity b Unit

Biomass-to-liquids (BTL) is being scaled up by Choren Industries in Freiberg, Germany, together with Shell. The first commercial BTL plant is planned to have a capacity of ca. 160 000–200 000 t a−1 (4000–5000 bbl day−1 ), with start-up being scheduled for 2010 earliest. Table 2 presents an overview of the main large-scale GTL, CTL and BTL plants and locations being considered fur future construction. By 2015–2020, a total of ca. 500 million t a−1 (12 million bbl day−1 ) of additional GTL + CTL capacity could come on line, but as not all studies will result in actual plants being built, a more conservative estimate is an expansion by ca. 30 million t a−1 (750 000 bbl day−1 ) for GTL + CTL by 2015, yielding an installed world capacity of just above 40 million t a−1 (1 million bbl day−1 ). The product slate of GTL, CTL and BTL processes is similar as each process passes through a syngas production step followed by Fischer–Tropsch synthesis (Fig. 1). Syngas originating from feeds with low H/C ratio, especially from coal but also from biomass, requires adjustment of the H2 /CO ratio by H2 addition prior to the Fischer–Tropsch synthesis by ex situ water-gas shift (WGS) or promotion of the WGS reaction during the Fischer–Tropsch synthesis. The main product from high-temperature Fischer– Tropsch (HTFT), typically performed at 300–350 ◦ C, is naphtha/gasoline. For low-temperature Fischer–Tropsch (LTFT), typically performed at 200–260 ◦ C, the main products are gasoil for automotive applications, base oils for

high-quality lubricants and food-grade waxes. Furthermore, many product fractions are ideal feedstocks for chemical processes. HTFT consumes a lower H2 /CO ratio than LTFT. This implies that less H2 /CO adaptation has to be made in an HTFT process, resulting in less CO2 production. The HTFT product slate therefore contains more carbon, which raises the CO2 emission upon use. The economic reality of today shows that also for CTL, LTFT is to be preferred, essentially for its product slate and quality. An extensive development effort is being made worldwide to deal with the CO2 emission (sequestration, mineralization and other routes). 13.16.2

Basic Steps in GTL

Three basic process steps are common for all GTL, CTL and BTL processes. First, the hydrocarbon source is converted into syngas; second, the syngas passes through the Fischer–Tropsch synthesis step producing hydrocarbons; and third, the product is hydroprocessed, mainly via hydrogenation and/or mild hydrocracking into the final products (Fig. 2). All commercial bulk syngas manufacturing processes require oxygen from an air separation unit (ASU). Prior to the syngas manufacturing step, contaminants in the natural gas, especially sulfur compounds, are removed. References see page 3008

2996

13.16 Gas-to-Liquids

GTL (gas-to-liquids) Gasification or steam methane reforming

Natural gas

BTL (biomass-to-liquids) Gasification

Biomass

Syngas

FT synthesis & work-up

Naphtha Gasoil

CTL (coal-to-liquids) Coal gasification process

Coal

Common marketing and government/OEM engagement CO2

Fig. 1

Overall GTL, BTL and CTL process flows.

Naphtha, LPG CO + x H2 Natural gas

Syngas manufacture

−CH2− FT synthesis

Kerosene Hydrogenation/ hydrocracking Gasoil Specialities

O2

Fig. 2

H2O

Basic steps in GTL.

After its manufacture, the fresh syngas contains nitrogen compounds formed in the process and also sulfur compounds that may have slipped through. These are removed via washing and dedicated guard beds prior to introduction of the syngas into the Fischer–Tropsch reactor. After LTFT synthesis, the waxy product can be hydrogenated, yielding various wax grades and detergent feedstocks, or it can be hydrocracked, yielding LPG, naphtha, kerosene, gasoil and base oils. After HTFT, the lighter, more cyclic hydrocarbon product can contain up to 26% aromatics in its gasoline fraction and, depending on the product fraction, is routed to distillate hydrogenation, hydrocracking, olefin oligomerization, naphtha reforming and isomerization units. The three main steps are discussed in detail below. 13.16.3

Synthesis Gas Production and Purification

In principle, a wide range of hydrocarbon feedstocks such as natural gas, oil, coal and biomass may be converted to the intermediate syngas [17–19]. Remote or stranded natural gas is the feedstock of choice for current large-scale GTL projects.

The natural gas is converted to syngas, the feedstock for the Fischer–Tropsch synthesis and the source of pure hydrogen for the hydroconversion. The Fischer–Tropsch process requires syngas of the right quality: • an overall H2 /CO usage ratio for LTFT synthesis of about 2, with the ratio actually required depending on the catalyst characteristics • absence of catalyst poisons (see Section 13.16.3.3) • preferably low content of inerts (methane, nitrogen, carbon dioxide). Natural gas is converted to syngas in three steps: natural gas purification, conversion to syngas and its purification. The process is very similar to that for methanol synthesis, as both processes use syngas of comparable quality. The best choice of technology or combination of technologies for a specific GTL plant depends on the integrated concept of the GTL complex (Table 3). Purification of Natural Gas The process gas has to be conditioned for both the syngas generation and the Fischer–Tropsch synthesis. The raw feedstock usually contains many impurities and valuable natural gas liquids (Table 4). Often purification to very 13.16.3.1

13.16.3 Synthesis Gas Production and Purification

Tab. 2

2997

Plans and studies for new large-scale CTL, GTL and BTL plants

Company

Capacitya /bbl day−1

Location

Reaction and catalystb

FT reactor

Status and start date

Slurry technology from Rentech Slurry technology from Rentech Sasol technology

Study ongoing; start ca. 2007 [9] Study ongoing [10]

Sasol technology

Study ongoing

Shell technology

Study ongoing

Studies and plans for coal-to-liquids plants CCPRI Rentech Sasol + Shenhua Sasol + Shenhua Shell + Shenhua Headwaters + Shenhua Group

Vandalia, IL, USAQ Wyoming, USA

40 000 + electricity 10 000

Yulin, Shaanxi, China Yinchuan, Ningxia, China Yinchuan, Ningxia China China

80 000

LTFT by Fe LTFT by Fe

Study ongoing

– 80 000 – 70 000 – 50 000

Study ongoing [11] –

–

68 0000

LTFT by Co

34 000

LTFT by Co

140 000

LTFT by Co

154 000

LTFT by Co

Slurry phase distillate by Sasol Slurry phase distillate by Sasol Multitubular fixed bed by Shell Slurry

MOU signed; start ca. 2009 [12] EPC contract awarded; start ca. 2009 [13] EPC contract awarded; start ca. 2009 HOA signed; cancelled in 2007 [14]

Ras Laffan, Qatar

130 000

LTFT by Co

Slurry phase distillate by Sasol

On hold until ≥ 2008 [15]

Ras Laffan, Qatar

160 000

Ras Laffan, Qatar

91 000–120 000

Studies and plans for gas-to-liquids plants Oryx EGTLc Qatar Petroleum ExxonMobil (technology provider) Sasol and others (technology provider) Conoco-Phillips (technology provider) Marathon Oil technology provider) To be decided Several

Ras Laffan, Qatar Escravos, Nigeria Ras Laffan, Qatar Ras Laffan, Qatar

On hold until ≥ 2008 [15] LTFT by Co

Fluidized bed by Syntroleum

Tinrhert, Nigeria 30 000–36 000 Several –

–

–

LTFT by Co

Multitubular fixed bed by Shell

On hold until ≥ 2008 [15] Invitation to bid sent out [16]; cancelled in 2007 Studies ongoing for >10 other locations

Studies and plans for biomass-to-liquids plants Choren Industries

Freiberg, Germany

4000–5000

ca. 2010–12

the figures in column 3 by 41.1 gives the capacities in t a−1 . = low-temperature Fischer–Tropsch, HTFT = high-temperature Fischer–Tropsch. c EGTL = Escravos GTL, 25% owned by Nigerian National Petroleum Corporation (NNPC), 75% by Chevron Nigeria Limited (CNL). a Multiplying b LTFT

low levels is necessary to ensure acceptable FT catalyst lifetime. Sulfur Removal The sour natural gas is converted to sweet feed gas in three steps. First, the bulk of the H2 S is removed by chemical absorption (MDEA, Sulfinol) down to a level of about 10 ppmv [18, 33]. Next, sulfur compounds (COS, thiols, organic sulfur) are hydrogenated to H2 S over a CoMo or NiMo catalyst at 350–400 ◦ C with

13.16.3.1.1

addition of hydrogen upstream of the catalyst bed. Alternatively, COS is hydrolyzed by steam to CO2 and H2 S. The remaining sulfur is separated by adsorption on ZnO down to a level of 10 ppbv as required by typical reforming catalysts. Addition of promoters such as alumina to the ZnO catalyst enhance the adsorption of traces of COS. Cu and similar metals provide additional hydrogenation activity [18]. References see page 3008

2998

13.16 Gas-to-Liquids

Tab. 3

Syngas generation technologies used in GTL plants

Plant

Operation

Nominal GTL capacitya /bbl day−1

1991–present 1993–present 1993–1994 2001 2003 2003 Not reported 2003 2004 2004–present 2007/2009 2009 2009–11 cancelled

30 000 14 700 250 20 400 300 250 70 1000 Not reported 34 000/68 000 34 000 140 000 154 000

PetroSA, Mossel Bay, South Africa [20] Shell, Bintulu, Malaysia [35] Synhytech, Pueblo, USA [21] ENI/IFP, Sannazzaro, Italy [22] Conoco-Phillips, Ponca City, USA [23] BP Amoco, Nikiski, USA [24] ExxonMobil, Baton Rouge, USA [25] Syntroleum, Tulsa, USA [26] Statoil, Mossel Bay, South Africa [27] Sasol, Sasolburg, South Africac [28] Oryx, Ras Laffan, Qatar [29] EGTL, Escravos, Nigeria [30] Qatar Petroleum, Ras Laffan, Qatar [31] ExxonMobil, Ras Laffan, Qatar [32]

Technologyb

SMR and ATR (Lurgi), HTFT POX (Shell) and SMR Not reported Not reported CPO process Compact SMR Fluidized-bed ATR Air-blown ATR Not reported ATR (Haldor Topsøe) ATR (Haldor Topsøe) ATR (Haldor Topsøe) POX (Shell) and SMR ATR

the figures in column 3 by 41.1 gives the capacities in t a−1 . = steam methane reforming; ATR = autothermal reforming; HTFT = high-temperature Fischer–Tropsch; CPO = catalytic partial oxidation; POX = partial oxidation. c Revamp, originally applied coal feedstock. a Multiplying b SMR

13.16.3.1.2 Hydrocarbon Removal Autothermal reforming (ATR) requires natural gas free of higher hydrocarbons and olefins, in order to operate efficiently. The presence of higher hydrocarbons results in thermal cracking in the preheater coil and olefins tend to coke on the reformer catalyst. Pre-reforming quantitatively converts hydrocarbons and olefins in the feedstock with steam into a mixture of methane, steam, carbon oxides and hydrogen. Typically, pre-reforming runs adiabatically at 350–550 ◦ C over Ni/MgO or Ni/MgAl2 O4 .

Trace Components Removal Organic chlorine is removed by activated alumina between the hydrogenation and the sulfur adsorption to prevent the formation of HCl in the hydrogenation, causing absorption on the ZnO and ultimately deposition of sublimed ZnCl2 on downstream catalyst or heat transfer surfaces.

13.16.3.1.3

Tab. 4

Mercury in the raw natural gas is adsorbed on metal sulfide supported on alumina. Present Syngas Generation Technologies Currently applied syngas generation processes are summarized in Table 3. The two main ones applied are partial oxidation (POX) and ATR, although they are typically used in combination with other technologies [34, 35]. Typical syngas compositions from these reactions are given in Table 5. The main reactions are partial oxidation, steam methane reforming and shift, depending on the reaction conditions. Common secondary reactions include cracking, the Boudouard reaction, CO reduction and formation of ammonia and hydrocyanic acid. The syngas composition of these processes can, within limits, be manipulated by altering various process conditions: addition and removal of CO2 , H2 and steam, preheat and exit temperature and pressure. 13.16.3.2

Components removed from raw natural gas

Compound in raw natural gas

Typical removal options

H2 S COS, thiols, organic sulfur Monoethylene glycol (MEG) Natural gas liquids (NGL) Mercury Organic chlorides, HCl Higher hydrocarbons, olefins Moisture

Adsorption Hydrogenation, sulfur removal Condensation Cold fractionation Adsorption Hydrogenation and adsorption Pre-reforming Adsorption

Typical dry syngas composition (mol%) for main technologies [17]

Tab. 5

Component

POX

ATR

H2 CO CO2 CH4 N2

62.4 34.3 2.0 0.5 0.8

63.8 27.5 5.0 3.1 0.7

13.16.3 Synthesis Gas Production and Purification

13.16.3.2.1 Non-catalytic Partial Oxidation (POX) In partial oxidation, the oxidant and the hydrocarbon are mixed in the reactor, where they are allowed to react at temperatures of 1300–1400 ◦ C without any catalyst [36, 37]. Typically, entrained-flow refractory-lined reactors with top-mounted burners are used. Quenching of the raw syngas yields high-pressure steam. In contrast to other processes, partial oxidation is highly tolerant of sulfur. 13.16.3.2.2 Autothermal Reforming (ATR) Autothermal reforming refers to adiabatic oxidative reforming, where the heat for the reforming reactions is supplied by internal combustion with oxygen. The reaction is initiated homogenously and completed by heterogeneous catalysis [38]. In recent development, the required steam/carbon ratio was reduced from 1.5–2 to about 0.6, thanks to the development of new burners. Under these conditions, syngas with the ideal ratio of 2 may be produced and less oxygen is required. The ATR runs on a nickel-based reforming catalyst. The catalyst is exposed to high operating temperatures, resulting in sintering and low intrinsic catalyst activity. Alumina and magnesium alumina spinel are typical supports due to their high thermal stability and strength [17].

Developments Incremental improvements are continuously made on the above processes, particularly on catalysts and burners. New technologies aim at higher H2 + CO yield (at the correct ratio) and at better economics. Design challenges are mechanical design, material selection and scale-up. Alternative syngas generation technologies are in development: 13.16.3.2.3

• Catalytic partial oxidation (CPO) [39]: The production of syngas with solid catalysts is referred to as CPO. The oxidant and the hydrocarbon feedstock are premixed in a mixer before the feed enters the catalytic bed. Very short residence times are applied. Noble metal catalysts, in forms of pellets, monoliths and foams, are used. Concerns regarding pre-ignition limit the inlet feed temperatures of the hydrocarbon feedstock and oxidant and result in relatively high oxygen consumption. • Steam methane reforming (SMR) [40]: Conventional reforming on Ni- and alkali metal-promoted catalysts offers poor economy of scale. Present development focuses on compact reformers [41] and heat exchange reformers (e.g. gas-heated reforming), in which a portion of the heat of reaction is provided by heat recovery from the reformed gas, rather than burning the fuel.

2999

• Air-blown ATR eliminates the need for an oxygen plant [26]. In air-blown synthesis gas production, savings in air separation are traded for more inert gas in the downstream Fischer–Tropsch synthesis. • Oxygen membrane reforming (OMR) [42]: Oxygen production by means of membranes couples air separation and partial oxidation in one unit, thereby eliminating the need for a conventional cryogenic oxygen separation. The principle of these devices is based on the use of non-porous ceramic ion transport membranes.

Purification of Syngas The raw syngas contains impurities, some of which form in the syngas generation process. The syngas is polished to meet the requirements of the Fischer–Tropsch synthesis (Table 6). Scrubbing of the raw syngas removes NH3 , HCN, formic acid and soot particles. Small quantities of soot (ca. 50 ppm) form in partial oxidation of natural gas, although their formation is not favored thermodynamically. Alternatively, the soot may be removed by dry soot removal or with adsorbents. In partial oxidation, the syngas is advantageously desulfurized downstream of the syngas generation process, albeit a higher volume of gas to be treated: 13.16.3.3

• quantitative conversion of organic sulfur species in the natural gas to H2 S (and traces of COS) in partial oxidation, eliminating the hydrogenation step • inhibition of metal dusting with syngas at elevated temperatures • inhibition of spontaneous methanation of syngas at temperatures above 400 ◦ C. The technologies used for sulfur removal are the same as for natural gas. To limit the loss of efficiency in cooling syngas to the temperature required by current acid gas removal systems, warm gas cleanup technologies have been developed for temperatures above 250 ◦ C. Typical syngas specifications for the Fischer–Tropsch synthesis [43]

Tab. 6

H2 S + COS + CS2 NH3 + HCN HCN + HBr + HF Alkaline metals Solids (soot, dust, ash) Tars, benzene, toluene, xylenes Phenols and similar

References see page 3008

< 1 ppmv < 1 ppmv < 10 ppbv < 10 ppbv Essentially nil Below dew point < 1 ppmv

3000

13.16 Gas-to-Liquids

Optionally, CO2 is removed and sequestered. Typical absorbents and processes are MDEA, Flexsorb, Selexol and Rectisol [44].

Entrained fluidized-bed and fixed fluidized-bed reactors can only be applied at high temperatures as otherwise the products would condense too much. At high temperatures, only Fe-based FT catalysts can be applied. LTFT can be operated with both Fe and Co (see below). Entrained fluidized-bed and fixed fluidized-bed reactors have the same advantages and disadvantages as ebullated bed and slurry bubble column reactors for FT. An overview of the current commercially applied reactor designs for FT synthesis is presented in Table 7. A wide variety of alternative reactor designs have been and are being evaluated for Fischer–Tropsch synthesis. Structured reactors have recently attracted attention thanks to the capacity for high heat removal for most of these structures. Examples of structured reactors include ceramic monoliths [45], metallic monoliths or ‘‘micro-channel reactors’’ [46] and ceramic and metallic foams. Foams will display higher heat conductivity than monoliths thanks to mass transfer in the radial direction. For all structured reactors, good contact with the cooling wall under reaction conditions is essential to ensure efficient heat removal. For metallic monoliths and foams, much attention has to be paid to adhesion of the active phase to the support under reaction conditions. Other novel reactor designs focus on supercritical FT [47] and the use of membranes [48, 49]. Supercritical FT can solve the diffusion limitation in FT but adds the complications of using a solvent and operating at very low conversion-per-pass to sustain the supercritical state. Membrane reactors for Fischer–Tropsch synthesis can in theory generate the highest selectivity to preferred products but still suffer from a low productivity on a reactor volume basis.

13.16.4

Fischer–Tropsch Reaction Reactor Types The Fischer–Tropsch (FT) reaction (see also Chapter 13.15) is highly exothermic with an enthalpy change of about 165–180 kJ mol−1 CO converted, depending on the precise product composition. As a result, all FT reactors are designed to maximize heat removal. Insufficient heat removal would otherwise result in increased catalyst deactivation, decreased selectivity towards preferred products and at worst may induce a run-away. Furthermore, reaction products, especially H2 O, can deactivate the FT catalyst, calling for a limited conversion-per-pass, which is observed in all designs. Finally, diffusion limitation reduces the rate of reaction and affects the product distribution such that many designs aim to minimize this effect. The four typical reactor designs for FT synthesis are fluidized and ebullated bed reactors for HTFT and multitubular fixed-bed and slurry bubble column reactors for LTFT. The first multitubular fixed-bed reactor was developed by Arbeitsgemeinschaft Ruhrchemie (ARGE) and contained 2050 tubes of 12 m length and 5 cm diameter. Current state-of-the-art multitubular fixed-bed reactors as applied in the Shell GTL process contain more than 10 000 tubes per reactor. Intrinsic advantages of multitubular fixed-bed reactors over other commercial reactor designs are a higher C5+ selectivity and a lower CO2 selectivity, making the process more carbon efficient. Operation in slurry bubble column or ebullated bed mode avoids most of the heat removal issues. Because smaller catalyst particles are used, the FT reaction is also less diffusion limited versus standard multitubular fixed-bed operation. 13.16.4.1

Tab. 7

Catalysts The initial catalyst applied by Fischer and Tropsch was based on Co but yielded only small amounts of liquid product [1]. After additional catalyst R&D by Fischer and Koch, a good liquid product yield was achieved at 180 ◦ C applying 100 Co/18 ThO2 /100 kieselguhr as catalyst [50]. 13.16.4.2

Overview of commercial reactor types for Fischer–Tropsch synthesis

Process type

Catalyst

Technology

Slurry bubble column

LTFT LTFT LTFT

Co and Fe Co Co

Fixed fluidized bed Entrained fluidized bed

HTFT HTFT

Fe Fe

ARGE Shell GTL Sasol’s slurry phase distillate, ExxonMobil’s AGC-21 Sasol’s advanced Synthol Sasol’s Synthol (also used by PetroSAa )

Multitubular fixed-bed

a Company

Fischer–Tropsch reaction

formerly named Mossgas.

13.16.4 Fischer–Tropsch Reaction

In the early years, both Co- and Fe-based catalysts were studied for FT catalysis, but up to the mid-1970s, the commercial focus was on CTL with Fe. For the ARGE reactors, the initial catalyst formulation was 100 Co/5 ThO2 /8 MgO/200 kieselguhr during World War II in Germany at low operating pressure, later changed to 100 Fe/5 Cu/4.2 K/25 SiO2 , as applied by Sasol at higher operating pressure. Also Sasol’s ebullated bed units and Sasol’s and PetroSA’s fluidized-bed units apply Febased catalysts. From the mid-1970s onwards, there was renewed commercial interest in GTL and concomitantly in the use of Co as active metal. This culminated in commercial Co-based FT catalysts applied in Shell’s fixedbed reactor and Sasol’s slurry bubble column reactor. Over the years, much research has been done to understand the FT mechanism and the effect of carrier type, additives and contaminants on the performance of Co-, Fe- and also Ru-based FT catalysts. In this chapter, these will be discussed in more detail. Active Metals Many metals have been found active for Fischer–Tropsch synthesis. Activity-wise, they can be ranked in the following order for low-temperature FT synthesis: Pt, Pd, Ir, Zr (low) < Os (medium) < Fe, Ni (high) < Co (very high) < Ru (highest). Although ruthenium is most active for FT synthesis and also yields good selectivities [51], it is of lower commercial interest given its scarcity and high price [52, 53]. Still, Cosmo Oil has a ca. 280 t a−1 (7 bbl day−1 ) slurry phase pilot plant running on an Ru-based FT catalyst, operated at 30 bar (1 bar = 105 Pa) and 270 ◦ C [54]. Nickel has high activity but is characterized by low selectivity, typically producing high amounts of methane [7]. It also displays low stability due to readily formed Ni carbides [55, 56]. Furthermore, the highly toxic and volatile Ni(CO)4 can be formed in the presence of CO under some reaction conditions [5]. Cobalt is about three times more active than Fe per active site in LTFT application [57]. In addition to its higher FT activity, Co is superior to Fe in LTFT thanks to its (1) higher selectivity for longer chain hydrocarbons, (2) lower coke formation, (3) lower olefin selectivity, (4) lower oxygenate selectivity, (5) lower water gas shift activity and (6) higher resistance to oxidation by water [58–60]. These advantages yield a more efficient process, better product quality, a longer catalyst life and easier product processing. In contrast to cobalt, iron also displays good activity and selectivity in HTFT synthesis. As a result, Fe is the only active metal currently in commercial use for this application. Under HTFT conditions, Co displays low selectivity, generating much methane [7]. 13.16.4.2.1

3001

The higher WGS activity of Fe is said to be crucial for CTL applications as the initial H2 /CO ratio is lower relative to GTL. The WGS reaction of Fe can increase the H2 concentration in situ up to the required level. Still, one can always decouple H2 production from FT by using a separate hydrogen manufacturing unit (HMU) or steam methane reformer (SMR) and then apply the more selective Co as active metal. An additional advantage of a separate shift step is the generation of a more concentrated CO2 stream, which is more favorable for CO2 management. 13.16.4.2.2 Carriers Many carrier types have been evaluated to support the active phase in FT catalysts. The most common ones are silica, alumina and titania, but much work has also been done on zirconia, zinc oxides, various carbon structures and numerous mixtures of the above. The impact of a carrier on the performance of an FT catalyst can typically be divided into four categories:

• strength of interaction with active phase • propensity for overgrowth [strong metal–support interaction (SMSI); see Chapter 3.2.5.2] and crystallization with active metal • hydrothermal stability • diffusion effects. The strength of the interaction between carrier and active phase needs to be balanced, as it has to be sufficient to generate good dispersion but not too high in order to avoid decreased reducibility of highly dispersed crystallites. The latter negative effect can be counteracted by addition of noble metals to facilitate the reduction (see Section 13.16.4.2.3). For Co, the strength of the interaction with the carrier can be ranked in the order silica < zirconia < alumina < titania. For a given carrier type, the strength of the interaction can still be varied by the applied calcination temperature [61] or by changing the ingredients in the catalyst preparation [62] or the active phase precursor [60]. Although there are huge differences in the interaction between the carrier and the active phase, the carrier type does not affect the turnover number (TON) of the active phase. Only when the crystallite size of Co is 370 ◦ C, is recycled back to the hydrocracker. Alternatively, vacuum distillation can be applied to the >370 ◦ C fraction, yielding a waxy raffinate fraction with an atmospheric boiling point range of 370–540 ◦ C, and a fraction >540 ◦ C which is recycled back to the hydrocracker. The boiling point ranges given are for indicative purposes only. The waxy raffinate can be converted into base oils by removing the wax by either solvent dewaxing or catalytic dewaxing. When required, catalytic dewaxing can also be applied to improve the cold-flow properties of the gasoil. Generally, the operator is interested in maximizing the yield of heavier products and of gasoil and/or waxy raffinate in particular. The yield can be optimized by: 1. the choice of the process conditions such as pressure and space velocities of liquid feedstock and the hydrogen-containing gas phase 2. the choice of the catalyst or catalysts 3. the application of a recycle of unconverted Fischer–Tropsch paraffins. It is possible, in principle, to arrive at full conversion by once-through operation of the Fischer–Tropsch product. This generally results in overconversion or, in other words, in a very low yield of the desired products with the higher boiling point (waxy raffinate, gasoil) and high yields

of low boiling point products (LPG, naphtha). Applying a recycle of unconverted material solves this problem and leads to higher yields of the desired products. In this case the relevant parameter is the so-called conversion per pass of >370 or >540 ◦ C, i.e. the conversion taken over reactor inlet and outlet. On the other hand product specifications will demand a certain level of isomerization and, since isomerization increases with conversion per pass, the conversion per pas cannot be below a certain level. Very low conversions per pass are not attractive from a capital investment point of view either since this results in large recycle streams leading to large reactors, pumps, compressors, separator vessels, etc. In practice, the conversion per pass is in the range 30–80%. Other process configurations are possible: • pre-hydrotreating of the Fischer–Tropsch product in a separate reactor • using separate hydrotreating and hydrocracking catalyst beds in one reactor • production of normal paraffinic fractions by hydrogenation only, followed by distillation, so that solvents, detergent feedstocks and refined waxes can be produced • incorporating a catalytic dewaxing function in the hydrocracking step • have a certain boiling point fraction bypass the hydrotreating and/or hydrocracking step in order to have oxygenates in the final product. In the following sections, we will assume a single hydrocracking step for the full-range liquid product of the Fischer–Tropsch section. By the sheer nature of the Fischer–Tropsch product, hydrocracking in GTL is completely different from refinery hydrocracking. Of course, olefins and oxygenates are present, but sulfur- and nitrogen-containing molecules and (poly)aromatics are completely absent. The consequence is that we do not have to revert to high hydrogen pressure to protect the catalyst from poisoning or from coke laydown. This leads to a number of options and observations that are unique to GTL hydrocracking: 1. A relatively low hydrogen pressure, 973 K) of the OCM reaction. The importance of the nature of oxygen species is indirectly supported by the results of the OCM reaction, where different oxidizing agents were used [55–61]. Figure 2 clearly demonstrates an increase in C2 selectivity when O2 is replaced by N2 O. However, contradictory results were reported in [55]. For Na/CaO and Li/CaO catalytic materials with low ( 0.1 kPa Oxygen vacancies at p < 0.1 kPa active oxygen species at p > 0.1 kPa Impurity defects (transition metals)

S (C2) / %

80

60

40 Na(0.001 at.%)/CaO Na(1.2 at.%)/CaO Na(6.4 at.%)/CaO 20 0.00

0.05

0.10

0.15

0.20

Θo / Θo2 Selectivity of C2 products over Na2 O/CaO catalysts versus steady-state O / O2 ratio simulated at oxygen partial pressures without the presence of methane. Reproduced from Ref. [75].

Fig. 3

conductivity of solid materials. The electrical properties play an essential role in the OCM reaction [86]. Solid materials of n-type conductivity are usually non-selective catalysts. It appears that solids containing both p-type and oxygen-ion conductivity are desirable catalysts [74, 76, 80]. It must be stressed that the bandgap should be in the range 5–6 eV. The necessity for p- and ionic-type conductivities can be understood from the following discussion. As mentioned in Section 13.17.2.3, differently charged adsorbed oxygen species are considered to be responsible for methane activation. To form charged oxygen species, electrons from the solid should be transferred to oxygen. It is

well known that the concentration of electrons in the conduction band is a function of the bandgap: the lower the bandgap, the higher is the concentration of electrons. Therefore, if the bandgap is too small, high concentrations of active oxygen species are expected, which do not favor selective methane oxidation. If the bandgap is too high, reactive oxygen species can hardly be formed and no catalytic activity is observed. It may be expected that an optimized concentration of active oxygen species is required for selective catalyst performance. This may be one reason why all selective OCM catalysts possess a bandgap of 5–6 eV. The oxygen ionic conductivity promotes both dissociation of adsorbed molecular oxygen species, which favor the consecutive oxidation of C2 hydrocarbons, and fast exchange between surface atomic oxygen species and bulk anion vacancies. The latter process determines surface coverage by active oxygen species. Oxygen-ion and p-type conductivities are usually increased upon doping of solids with other compounds, if the oxidation state of the cation is lower than that in the host matrix. Catalytic materials of composition Na–W–Mn/SiO2 have been intensively characterized in order to establish structure–selectivity relationships [82, 88–90]. According to Palermo et al. [82], the crystalline structure of the support plays a very important role in designing a selective catalytic material. Amorphous silica led to a catalyst with poor selectivity, whereas α-cristobalite (crystalline SiO2 ) was a very favorable supporting material. The authors pointed out that Na plays a dual role: (i) crystallization of amorphous silica to the crystalline form and (ii) stabilization and dispersion of surface WOx species. WO4 was mentioned as a possible candidate. The importance of WO4 References see page 3021

3016

13.17 Oxidative Coupling of Methane

species was later highlighted by Green and coworkers [88–90], who catalytically tested and characterized a series of Na–W–Mn/SiO2 catalytic materials by means of XPS, Raman spectroscopy and XRD analysis. DFT calculations by Chen et al. [91] predicted that the tetrahedral [WO4 ] site with a single bridge oxygen is the most probable active center responsible for methane activation. Anshits and coworkers [56, 58, 59, 62, 92] demonstrated that the defect structure of OCM catalysts is strongly changed under reaction conditions. They used the following methodology. A sample was treated in a reactive feed (O2 , N2 O, He, CH4 –O2 or CH4 –N2 O) at 1023 K followed by rapid quenching of the pretreated sample from 1023 to 77 K. The quenched samples were γ irradiated. The irradiation did not create new defects but made them visible by EPR spectroscopy. These structural changes influence the catalytic performance. For Li/CaO and Na/CaO with low (50 nm), mesopores (2–50 nm) and micropores (400 ◦ C leads to rupture of bridges linking units within the macromolecular coal structure [74]. Aliphatic ether linkages are cleaved first and then methylene bridges of, in particular, diarylethane structures. Bond rupture forms free radicals, which must be scavenged by addition of a hydrogen atom. Solventmediated hydrogenolysis recently has been recognized as a second important mechanism for bond cleavage (see Section 13.18.3.3). Fragmentation to units of lower molecular mass then enhances the solubilization. Radicals which are not capped by hydrogen fast enough can lead to the formation of stable high-molecular mass products, producing ultimately char and coke. Thus, liquefaction critically depends on whether the supply of hydrogen atoms is able to follow the formation of free radicals. Sources of hydrogen atoms are the recycling solvent, the coal itself and molecular hydrogen [9]. However, uncatalyzed thermal reactions have been found to convert coal generally only to the stage of heavy fuels, whereas catalysts are necessary to remove heteroatoms efficiently and to lower the molecular size to the extent required for the production of gasoline and other fuels [74]. Coal Properties Generally, the behavior of coal in liquefaction depends on rank, petrographic composition and inorganic constituents. The macerals of the vitrinite group, accounting for >80% of the organic material of most coals, and of the exinite (liptinite) group are considered to be much more reactive than inertinite macerals. However, coal rank is the most crucial parameter for liquefaction. Figure 2 13.18.3.1

References see page 3034

3028

13.18 Catalysis in Direct Coal Liquefaction

Yields / % on daf coal

60

Distillate crude oil

40

20

Organic residue 25

20

15

10

5

0 Oxygen content / % daf

0 60

65

70

75

80

85

90

Carbon content / % daf

Dependence of liquefaction yields on coal rank (content of carbon and oxygen). Data obtained from the process development unit [150 kg day−1 , 30 MPa H2 , 1.2% Fe2 O3 catalyst (red mud), 465–470 ◦ C] of Deutsche Montan Technologie, Essen, Germany.
, •: Carbon content; , ◦: oxygen content; daf: dry and ash-free. Modified from Ref. [75].

Fig. 2

shows yields of distillate oil and vacuum bottom depending on coal rank, parameterized by the content of carbon and oxygen of the coals [75]. The yield of distillable liquids reaches a maximum for mid-rank coals between 80 and 85% carbon (high-volatile A bituminous coals in ASTM classification [14, 15]). Between 87 and 89% carbon (medium-volatile bituminous coals) the oil yield decreases sharply and high-rank coals of >90% carbon (low-volatile bituminous coals, anthracite) are almost inert. Low-rank coals, lignites, brown coals and sub-bituminous coals, in spite of high conversions, give lower oil yields due to the high content of oxygen which leads to formation of water and carbon oxides. Similar rank dependences have been reported by others [4–9, 76]. Generally, the low-rank coals behave differently in liquefaction than mid-rank coals and require different reaction conditions [9]. Primary Liquefaction Products Depending on the reaction conditions, coal liquefaction leads to a distribution of products which includes gases, distillable liquids, non-distillable liquids and solubilized coal fragments. The non-distillable products are usually characterized by their solubility in various solvents, using solvent fractionation methods which originally were developed in the petroleum industry. Oils are classified as materials soluble in pentane or hexane, asphaltenes as materials soluble in benzene or toluene but insoluble in pentane and preasphaltenes as materials soluble in pyridine but insoluble in benzene. The distinction between these fractions is based solely on solubility and is not directly related to chemical structures. Preasphaltenes have a similar content of heteroatoms to the parent coal and are believed to consist of relatively large organic fragments formed in the initial liquefaction stage by thermal cleavage of a few most reactive bonds. In comparison, asphaltenes are smaller and less polar fragments, but their chemical constitution 13.18.3.2

depends on the parent coal and the processing conditions. Long hydrotreatment increases aromaticity and reduces the amount of polar functional groups and molecular mass [9, 69, 77, 78]. On the basis of analytical data, average molecular structures have been proposed for preasphaltenes and asphaltenes, in order to visualize and to model chemical transformations in the liquefaction process [19, 69, 74]. However, just like coal structure models, they must be used with care, since solubility fractions are very complex mixtures of compounds. Solvent The solvent used in direct coal liquefaction takes a number of important functions. Its physical role is to provide the medium for coal transport, heat transfer and dispersion of liquefaction products [9]. Chemically it serves as an important source of transferable hydrogen (hydrogen donation) and as a transport vehicle for hydrogen within coal (hydrogen shuttling) [79–81]; under liquefaction conditions (i.e. 400 ◦ C), it even can take an active role in the cleavage of strong Caryl −Calkyl bonds (solvent-mediated hydrogenolysis, SMH) [82–85]. The solvent components responsible for these chemical effects are polycyclic aromatic hydrocarbons (PAHs) and their partially hydrogenated counterparts such as tetralin, 9,10-dihydroanthracene, 9,10-dihydrophenanthrene and 4,5-dihydropyrene [4,5-(H,H)Py]. The latter, as an example of a potent hydrogen donor solvent [86], is capable of supplying reactive hydroaromatic hydrogen for capping two radicals with hydrogen atoms [Scheme 1, Eq. (1)]. The pyrene formed is readily hydrogenated to 4,5-(H,H)Py during coal liquefaction and generally hydrogen donors can be generated in situ from PAHs [9, 69, 81]. Via repeated hydrogenation/dehydrogenation cycles, PAHs act as a ‘‘shuttle’’ between radicals and molecular hydrogen, which itself competes in the capping of the pyrolytically formed radicals only at higher pressures [87, 88]. 13.18.3.3

13.18.3 Fundamental Transformations in Direct Coal Liquefaction

H2

R

H2

R H

H H

R

3029

R H

(1) 4-(H)Py

4,5-(H,H)Py

Py

H H

[H ]

R

R H

(2) Py 1-(H)Py CH2

Py

1-(H)Py

Coal

Py H CH2 Coal

H +

CH2

Coal (3)

Scheme 1 Mechanisms of (1) hydrogen donation, (2) hydrogen shuttling and (3) solvent-mediated hydrogenolysis (SMH) of strong linkages in coal. Py = pyrene.

On the other hand, PAHs are acceptors for H-atoms forming resonance-stabilized cyclohexadienyl-like radicals [89]. Pyrene, for instance, can scavenge free H-atoms to form 1-H-pyrenyl radicals, 1-(H)Py• , which then act as hydrogen donor for capping radicals [Scheme 1, Eq. (2)]. However, hydrogen atoms [H• ] in Scheme 1 can also be delivered from hydroaromatic compounds such as 4,5-dihydropyrene by reversed radical disproportionation (RRD) [79, 83, 90]: • • −−

4,5-(H,H)Py + Py −− − −4-(H)Py + 1-(H)Py

(4)

However, just like hydroaromatic solvent molecules, hydroaromatic structures of coal can also be the source of hydrogen atoms [H• ] transferred by RRD to pyrene. In this case the hydrogen, which is shuttled by 1-(H)Py• and used in capping reactions of radical sites in the pyrolytically bursting coal structure, originates from the coal itself. Finally, in addition to thermal C−C bond homolysis and subsequent radical capping, a second mechanism via solvent-mediated hydrogenolysis (SMH) has been proposed on the basis of studies on model compounds [83–85, 90]. In this case the solvent radical such as 1-(H)Py• transfers an H-atom to the ipso-carbon atom of an aromatic ring connected to an aliphatic link [Scheme 1, Eq. (3)]. The resulting cyclohexadienyl radical is susceptible to cleavage of the Caryl −Calkyl bond, forming the unsubstituted aromatic ring and the radical of the detached link. Various PAHs, including phenanthrene and pyrene, have been known for a long time as disintegrating solvents to solubilize certain high-volatile bituminous coals at 350–400 ◦ C [9, 91–93]. Nowadays this is explainable by

the mechanisms given in Scheme 1 for bond cleavage reactions and hydrogen shuttling within coal from hydroaromatic structures to radical sites via PAH-derived radicals such as 1-(H)Py• . However, hydrogen donation, hydrogen shuttling and SMH require that solvent molecules must be able to diffuse into coal particles. As already discussed (Section 13.18.2.3), coal imbibes solvents and this phenomenon is strongly connected with swelling. Aromatic solvents cause considerably higher swelling effects of coal than aliphatic solvents, and this indicates that a sufficiently high mobility of aromatic and hydroaromatic solvent molecules within the macromolecular coal structure should be given. However, this is not possible for typical solid catalyst particles such as iron oxides and sulfides, no matter how finely dispersed they are. It has been proposed that hydrogen atoms [H• ] in Scheme 1 could be delivered from a catalyst surface loaded with hydrogen atoms, i.e. single H-atom transfer from Cat-H to pyrene could also lead to the formation of 1-(H)Py• radicals [83]. This proposal concurs with a recent report that the hydrogenolytic cleavage of low molecular mass model compounds with an ‘‘FeS’’ catalyst, which in coal liquefaction is commonly used for the first stage of the hydrotreatment in the liquid phase, is consistent with a reversible hydrogen atom transfer pathway from the catalyst [94]. On the other hand, in the case of coal liquefaction the hydrogenation of aromatic to hydroaromatic solvent molecules on the catalyst followed by RRD [Eq. (4)] would lead to the same solvent radicals as the single H-atom transfer from the catalyst to a PAH. In References see page 3034

3030

13.18 Catalysis in Direct Coal Liquefaction

by hydrogen addition, rearrangements or radical addition. The hydrogen transfer leads to stabilized soluble products, but competes with the last two reactions which can initiate retrograde processes forming high-molecular mass compounds [74, 103]. Only after coal fragments have been solubilized can the catalysts develop activity to promote directly the further breakdown of the coal structure to liquids. However, there still may exist severe constraints for large molecules to access the catalyst surface or pores. The activity of catalysts used in the coal dissolution stage is related to the catalyst dispersion and to the form and mode of catalyst application, which have an effect on the level of intimacy for the contact between coal and catalyst. Primary coal fragments and preasphaltenes still possess a similar content of heteroatoms to the parent coal, what makes them under unfavorable conditions particularly sensitive to retrograde reactions leading ultimately to the formation of char and coke (Fig. 3) [69, 97]. The conversion pathways of bituminous and subbituminous coals have been found to be very different. High intermediate yields of preasphaltenes and asphaltenes are produced from bituminous coals before oils begin to be generated in significant amounts. By contrast, the conversion of sub-bituminous coals is more direct, i.e. preasphaltenes, asphaltenes and oil are generated simultaneously. The use of a catalyst did not change the pathways, but solely increased the reaction rates [104]. Differences of the coal structure appear to be the reason for the distinct behavior in liquefaction [9].

either case, solvent molecules must be considered to play a very important role in transferring the H-atom activity generated at a catalyst surface to sites within the organic coal matrix where it is needed for hydrogenolytic cleavage of the macromolecular structure [83]. Generally coal-derived solvents are highly complex mixtures of compounds, some of which, for example heterocyclic nitrogen and phenolic oxygen compounds, can influence liquefaction beneficially or detrimentally [9]. Due to the severe reaction conditions at 450–500 ◦ C, solvent components also are exposed to side reactions such as isomerization, cracking, polymerization and incorporation in coal liquefaction products [9, 79, 89, 92]. Hydrogen exchange and transfer between coal, solvent and molecular hydrogen under liquefaction conditions has been studied using deuterium and tritium tracers [10, 95]. 13.18.4

Catalysis in Direct Coal Liquefaction

Efficient coal liquefaction processes require the use of catalysts to accelerate the various reactions involved in hydrogenation, cracking, hydrocracking (hydrogenolysis) and heteroatom removal. Several recent reviews of catalyst developments are available [9, 10, 96–102]. Catalysts are used either as solid metal compounds dispersed in the coal–solvent slurry or as supported metals and metal compounds in ebullated- or fixed-bed reactors. Generally, the former catalysts are used to promote coal solubilization, whereas the latter are mainly applied to upgrade the coal-derived crude oil. At the onset of the liquefaction process, where insoluble three-dimensional macromolecules mainly constitute the solid coal particles, typical solid catalysts can have only a very limited function (Fig. 3). As the coal particles disintegrate through thermal cracking and SMH, the fragments begin to disperse in the solvent and radical sites are stabilized

Dissolution Catalysts in the First Liquefaction Stage Generally, the oxides and sulfides of Fe, Co, Ni, Mo and W and halides of Zn, Sn and Pb have been found to be catalytically active for hydrocracking (hydrogenolysis) and hydrogenation reactions of coal and coal products. Iron compounds are only moderately active and have 13.18.4.1

Coke

Coal

Retrograde reactions

Thermal or solvent-mediated bond cleavage

[Coal fragments]

H-Transfer

Oil Fig. 3

H-Transfer

Solid catalyst

H-Transfer

Functions of solid catalysts in coal liquefaction. Modified from Ref. [9].

Preasphaltenes

H-Transfer

Asphaltenes

13.18.4 Catalysis in Direct Coal Liquefaction

to be used with higher concentrations, i.e. factors of 100–200 higher as compared with molybdenum. The costs make the highly active metals prohibitive for use as disposable (once-through) catalysts for the first stage of coal liquefaction in the liquid stage, and the recycling would require labor- and cost-intensive separation of the highly dispersed catalyst from mineral matter and coal residue. For this reason, only Mo, Sn and Fe dissolution catalysts have been used in commercial-scale developments and by far the highest preference has been given to Fe catalysts [3–9, 71–73, 97]. Bergius originally added to the coal feed 5% of red mud, the iron oxide waste from aluminum production from bauxite, as an absorbent for hydrogen sulfide evolved during coal liquefaction. Only later were iron oxides and sulfur recognized actually to serve as a catalyst and promoting agent, respectively. In 1924, Pier at BASF introduced oxides and sulfides of Mo, W, Fe, Co and Ni as sulfur-resistant hydrogenation catalysts. In the first commercial-scale plant that started operation in 1927 at Leuna, Germany, MoO3 together with sulfuric acid in low concentration were used in the liquefaction of brown coal and the hydrogenation of brown coal tar. A supported Mo catalyst formed by impregnating brown coal coke with aqueous ammonium molybdate was found to be effective at Mo concentrations below 1 wt.%. By impregnating the brown coal itself with ammonium molybdate and sulfuric acid, the Mo concentration could be reduced to 0.05 wt.%. Later the molybdenum was replaced by less active but much cheaper iron, which was added as Fe2 O3 in the form of 5–9 wt.% red mud or iron ore. In the case of low-sulfur brown coal, it was found necessary to add sulfur (1.2 wt.%). In the first two commercialscale liquefaction plants for bituminous coal, which began operation in 1935 at Billingham, England, and in 1936 at Scholven, Germany, 0.06 wt.% tin dichloride was used as the catalyst. In Germany, during the War, the tin was partly replaced by lead. In other German plants bituminous coal was liquefied with an iron catalyst (1.2 wt.% FeSO4 · 7H2 O, 2 wt.% red mud, 0.3 wt.% Na2 S), for which the iron sulfate was impregnated on the coal [3–7, 97]. Since 1950, mainly in the USA, new research on the catalysis of coal liquefaction was started [99], which at the end of the 1970s was intensified, especially in the USA [97, 98] and Japan [100–102]. Several developments were aimed at increasing the activity of iron-based dissolution catalysts as cheap, environmentally benign and disposable materials. As can be recognized from the abundant literature [9, 100, 101], three general strategies are pursued: (i) enhancement of the catalyst dispersion by preparing ultrafine powders of nanocrystalline iron oxides and sulfides with crystallite diameters 1400 ◦ C) to produce mostly CO and H2 . The venerable moving-bed Lurgi gasifiers [6, 7] do operate at lower temperatures and, if a low-rank coal is used, as at Great Plains, the well-dispersed inherent mineral matter (initially in the form of alkaline earth and alkali metal cations exchanged on the coal’s carboxyl groups) may have a substantial catalytic effect on both reaction kinetics [8] and product selectivity (e.g. towards CH4 ) [9, 10]. Clearly, a much greater benefit would be derived from a process based on the deliberate use of an active, selective and recoverable gasification catalyst. A large number of authoritative reviews on the various aspects of catalysis of coal and carbon gasification have been published, especially during the 1980s [1, 5, 11–25]. The Exxon Catalytic Coal Gasification process [7, 26] is a good point of departure both for its practical aspects – it was (and still appears to be) the closest to commercial development [27] – and as a basis for a more fundamental discussion. The (simplified) reactions of primary interest are the following: C + H2 O −−−→ CO + H2 ◦


H973 = +136 kJ mol−1 CO

(1)

2C + 2H2 O −−−→ CH4 + CO2 ◦


H973 = +11.7 kJ mol−1 CH4 References see page 3044

(2)

3038

13.19 Catalysis in Coal and Carbon Gasification

Coal (biomass, petroleum coke, waste) Gasification system

O2

CO + H2

Gas cleanup

H2 Fuel cell

CO, H2, CO2, H2S, NH3, CH4,…

Solid products

Gas turbine Exhaust

Fuels/ chemicals Electricity

Electricity Water

Steam Heat exchanger

CO2 removal Steam turbine Fig. 1

H 2O to stack Electricity

A typical coal gasification process scheme.

molten-salt processes have gone out of vogue. For example, the gasification process associated with M. W. Kellogg [1, 7, 30] is currently advertised as a ‘‘circulatingbed reactor concept that uses finely pulverized coal and limestone’’ [31]. Much uncertainty (and scope for improvements!) thus remains regarding the ‘‘winning’’ scheme for a commercial catalytic coal gasification facility.

Reaction (2) is the more attractive because of the higher heating value of its product (ca. 37 000 kJ m−3 for substitute natural gas versus ca. 11 000 kJ m−3 for syngas) and its much lower endothermicity. It represents a combination of reaction (1) and the exothermic water gas shift (CO + H2 O → CO2 + H2 ) and methanation (CO + 3H2 → H2 O + CH4 ) reactions (see Chapter 13.12). Figure 2 shows a simplified process flow sheet [4, 7, 26]. A more recent process development, at least at a conceptual level, is due, for example, to CeraMem (www.ceramem.com), which has attempted to develop a ‘‘novel nano-structured catalyst for steam gasification of carbonaceous feedstocks’’ [28]; see also Ref. [29]. Other

13.19.3

Catalysts

It is beneficial not only to analyze the catalytic behavior of carbonaceous substrates that are similar to coals and coal chars – such as graphite, activated carbons, petroleum

Coal Coal preparation

Recovered catalyst CO + H2 recycle

Steam Gasification Ash + ‘spent’ catalyst

Gas cleanup

Gas separation

SNG

Sulfur recovery

Sulfur

CO2 recovery?

(To stack)

Catalyst recovery Ash Fig. 2

The Exxon catalytic coal gasification process for the production of substitute natural gas (SNG).

13.19.3 Catalysts

cokes and carbon blacks – but also to appreciate the improvements in our understanding of related processes, such as regeneration of coked catalysts (see Chapter 7.1), the production of higher-value-added activated carbons for air and water purification, soot destruction in diesel engines, NOx emissions reduction from both mobile and stationary sources, oxidation protection of composite carbon materials used in aerospace and tribological applications, fabrication of more efficient electrocatalysts in fuel cells and better controlled growth of carbon nanotubes. Even an appreciation of the virtues of carbonsupported catalysts and the challenges in their preparation and use [32] is rewarding. The summary presented here should be viewed from such a perspective. An inescapable conclusion from even a cursory analysis of the published reviews is that a large portion of the periodic system is of interest for the catalysis of coal and carbon gasification [17, 19], as illustrated in Fig. 3. Alkali, alkaline earth and transition metals have been by far the most investigated, although the use of noble metals [11], lanthanides [33, 34] and even some actinides [34, 35] has also been explored. From a practical point of view, it is desirable to know which catalysts are most effective. However, their complexity typically does not allow a meaningful comparison, even among studies carried out under similar experimental conditions. Therefore, the results summarized in the reviews mentioned in Section 13.19.2 should not be viewed as more than illustrations of trends and of often dominant phenomena and concepts. The key issues of interest are, of course, catalytic activity, turnover frequency and selectivity, but these are difficult

to generalize because their relative values depend on the details of catalyst preparation and the nature of the carbon substrate. Thus, in some studies the order of catalytic activity of the alkali metal salts may be regular, e.g. Li < Na < K < Rb < Cs, whereas in others it may be completely different. For example, Harker [36] related the former trend, observed for a cellulose-derived char gasified in O2 , to the ionization potentials of the metals; similarly, the activity order Cs > Rb > K > Na in CO2 gasification has been interpreted in terms of a ‘‘dispersion effect’’ and the ability of the respective carbonates ‘‘to increase the steady-state concentration of oxygen at the carbon surface by increasing the total number of active sites’’ [37]. In contrast, Li has been reported [11] to be the most active among the carbonate catalysts of graphite gasification with CO2 , presumably because of its lowest melting point among alkali metal carbonates or because of the ‘‘instability of Li0 with respect to its oxide’’ [38]. In contrast to gasification reactions using steam and CO2 , which were the object of intense study through the 1980s, the catalyzed gasification of carbons with NO has been studied fairly thoroughly over the last 10–15 years. Even a pilot-scale study has been carried out [39], driven by the need to resolve the ammonia slip problem in selective catalytic reduction (see Chapter 11.3) and based on the demonstrated beneficial effect of O2 in the NOx -laden flue gas [40] and on the presence of abundant inherent catalysts in low-rank coal chars [41]. Figure 4 illustrates the typical results: high NO conversion is achieved at low temperatures by the highest-pyrolysis-temperature References see page 3044

H

Li

He

Be

Catalytic gasification with O2 Catalytic gasification with H2O or CO2

Na Mg

K

3039

Ca Sc Ti

Rb Sr

Y

V

B

C

N

O

F

Ne

Al

Si

P

S

Cl Ar

Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te

Cs Ba La Hf Ta W Re Os

Ir

I

Xe

Pt Au Hg Tl Pb Bi Po At Rn

Illustration of the (almost) ubiquitous catalytic effect of inorganic constituents on carbon gasification reactions. The bar height represents approximate catalyst effectiveness.

Fig. 3

3040

13.19 Catalysis in Coal and Carbon Gasification

100

% NOx reduced

1422-5 1422-7 1422-9 50

0

−50

0

100

200

300

400

500

T / °C

Temperature-programmed reaction profiles for NO reduction with coal chars prepared by lignite pyrolysis at 500 (1422-5), 700 (1422-7) and 900 ◦ C (1422-9).

Fig. 4

The most critical issue (as the Exxon CCG experience has demonstrated) – and one that is likely to be responsible for the often reported contradictory findings – is the nature of the interaction between the catalyst (or the catalyst precursor) and its reactive support. Indeed, this is a sui generis problem for both scientists and engineers: how does one preserve the active chemical state of the catalyst, the desirable high catalyst surface area, and also catalyst/substrate physical contact, while the ‘‘support’’ is being gasified away? Many attempts have been made to clarify the chemistry and physics of catalyst–support interaction, especially using carbons containing minimal inorganic impurities in order to avoid the complicating effects of catalyst interaction with mineral matter (whose discussion is beyond the scope of this chapter). Figure 6 illustrates the importance of the chemical state of the catalyst and the

70

0.9

60

0.8 0.7

50

Reactivity / h−1

CaO surface area / m2 g−1

char because of either the largest number of its reduced catalytic sites or its largest catalyst/carbon interfacial area; similarly, high conversion is also achieved at the higher temperatures, now by the lowest-pyrolysis-temperature char because of the highest concentration of its (re)active sites. The intermediate-temperature ‘‘window’’ of low catalytic activity is a consequence of the accumulation of oxygen on the catalyst surface, because of its inefficient transfer (‘‘spillover’’) to the carbon surface and this retards the subsequent desorption of CO2 and/or CO [41, 42] (see also Section 13.19.4). Meaningful fundamental comparisons of catalytic effectiveness are complicated by the fact that the catalyst surface area (or dispersion) is often not reported; for coal chars it is also often difficult to measure [8]. A notable early exception is illustrated in Fig. 5, using CO2 chemisorption to titrate alkaline earth metal oxide surface sites [43].

40 30 20

0.5 0.4 0.3 0.2

10

0.1 0

0 0 (a)

0.6

2

4

6

Calcium content (wt.%)

8

0

10 (b)

5

10

15

20

25

30

CaO surface area / m2 g−1

Illustration of measurement of catalyst dispersion (a) and its correlation with reactivity (b) in CaO-catalyzed gasification of carbon with CO2 at 1023 K (adapted from Ref. [43]).

Fig. 5

13.19.4 Mechanisms

Inactive phase (e.g., K2O, FeO)

3041

Active phase (e.g., KxOy, Fe)

Inactive phase H2O

Active phase

H2O Catalyst particle Carbon (support) surface Catalyst particle Carbon surface

Catalytic effect in carbon gasification as a function of catalyst’s chemical state, wetting behavior and interfacial contact area. In the bottom part, wetting propensity increases and the interfacial contact angle decreases from left to right.

Fig. 6

influence that the support exerts on it. Carbon is, of course, known to be a very good reducing agent [44] and the carbothermal (or carbothermic) reduction of a catalyst, or of its precursor, is an important chemical factor. Another practical example is the intriguing difference between the behavior of potassium and calcium gasification catalysts [45]: on a mass-of-carbon and mass-of-catalyst basis, potassium exhibited a higher initial activity than calcium in both O2 and steam, because of either higher initial dispersion or higher turnover frequency [43]; as gasification in O2 proceeded at 550–565 K, neither of the two catalysts was able to maintain contact with the reacting support and the reaction rate decreased, but at 900 K in steam there was no deactivation during potassiumcatalyzed gasification, presumably because of catalyst redispersion by virtue of the formation and regeneration of surface phenoxide (C−O−K) species [17]. The topography of catalyzed carbon gasification is also of interest [19, 46]. Some catalysts function by basal-plane penetration, thus forming pits on the support surface in the direction perpendicular to the basal planes (or graphene layer). Others form channels in the direction of one or several graphene layers [18] and yet others act by enhancing the recession of the edges of graphene layers. The effectiveness of the latter two modes of catalysis is related to the ability of the catalyst to wet the surface of the carbonaceous support: edge recession is exhibited by highly wetting (spreading) catalysts, while channeling is observed for larger catalyst–support contact angles (see Fig. 6). In this context, rationalization of the behavior of the practically attractive ‘‘composite’’ catalysts [47, 48] poses special challenges. Thus, for example, the reduced melting

points of eutectic phases are thought to be responsible for the higher catalytic activity of binary and ternary alkali metal halides, carbonates and sulfates [49], because they ‘‘facilitate contact between the catalyst and the carbonaceous substrate’’. Similarly, a synergistic effect has been illustrated for Ni/K catalysts in gasification of graphite with steam, wet H2 and wet O2 [50]. On the other hand, a complicating effect for binary catalysts has been illustrated in a study of Cu and Ni on a graphite substrate [51]: when the steam : H2 ratio was 100, the inactive NiO preferentially segregated on the surface. 13.19.4

Mechanisms

Significant progress toward commercialization of catalytic coal gasification will depend on further improvements in the control of catalyst dispersion, catalyst–support interaction and ultimately catalytic activity. These in turn are facilitated by progress in our understanding of the reaction mechanisms. Reflecting the broader scope of this topic adopted here, in addition to reactions (1) and (2), the following are also of interest: C + CO2 −−−→ 2CO ◦


H973 = +171 kJ mol−1 CO2

(3)

C + O2 −−−→ CO2 ◦


H973 = −395 kJ mol−1 References see page 3044

(4)

3042

13.19 Catalysis in Coal and Carbon Gasification

2C + 2NO −−−→ N2 + 2CO ◦
H973

−1

= −202 kJ mol

NO

according to the following mechanism: (5)

Of course, these global reactions are again simplified and it must be kept in mind that gasification with O2 almost always produces CO together with CO2 , especially at higher temperatures, whereas gasification with NO often produces CO2 together with CO, especially at lower temperatures. (The reaction heats are given at 973 K, a temperature that is thought to satisfy the conflicting demands of being sufficiently below the isokinetic point – see discussion below and Fig. 9 – and high enough to ensure a reasonable reaction rate. This is indeed the design temperature for the Exxon CCG process [26].) In both the presence and absence of a catalyst, reactions (1)–(5) have in common the oxygentransfer step and indeed the ‘‘mechanism of gasification for all oxidizing gases appears to be the same’’ [1], as illustrated below. Conspicuously absent from this list is the potentially desirable hydrogasification reaction (C + 2H2 → CH4 ). It will not be discussed here because its rate is orders of magnitude lower than that of the oxygen-transfer reactions, e.g. 105 times less than the uncatalyzed steam gasification reaction [52]; and, of course, in our current ‘‘hydrogen economy’’ mind-set, H2 is often considered to be a more valuable fuel than CH4 . Catalyzed steam gasification of carbon was among the first reactions to be studied in an attempt to distinguish between the so-called ‘‘electron transfer’’ and ‘‘oxygen transfer’’ mechanisms [24]. At a more fundamental level, even oxygen transfer can be interpreted as electron transfer when there is a distinguishable oxidation–reduction (redox) process occurring at the catalyst–support interface (see below). Indeed, the argument presented by Long and Sykes [53] against this ‘‘intermediate compound theory’’ and in favor of ‘‘an electronic interaction with the carbon’’ – in the sense that the ‘‘added [catalytic] substances interact with the [reactive] gases [such as CO2 and H2 O] to a much greater extent than do the natural impurities though both are efficient catalysts’’ – in retrospect seems unconvincing, although it is still often cited in the relevant literature. For all oxygen-transfer reactions – the case of steam gasification is used here as the most straightforward and practically most important example – the redox behavior of a gasification catalyst can be summarized as follows: Cat + H2 O −−−→ Cat−O + H2 Cat−O + Cf −−−→ Cat + C−O C−O −−−→ CO(+Cf )

(6) (7) (8)

Thus, for example, when an alkali metal (M) carbonate is used, the redox process is thought to proceed [22]

M2 CO3 + 2C −−−→ 2M + 3CO

(9)

2M + 2H2 O −−−→ 2MOH + H2

(10)

2MOH + CO −−−→ M2 CO3 + H2

(11)

The effect of a transition metal catalyst has been interpreted in two ways. The first is the analogous oxygentransfer redox cycle [24], for example [54] Fe + H2 O −−−→ FeO + H2

(12)

FeO + Cf −−−→ Fe + C(O)

(13)

C(O) −−−→ CO(+Cf )

(14)

and the second is the carbon dissolution mechanism [55–57]. The broader interest in the latter, summarized in Fig. 7, lies in the fact that vapor-grown carbon fibers and, more recently, multi-wall carbon nanotubes are thought to be produced according to the reverse process [58]. A critical step in this reaction sequence is ‘spillover’ (see Chapter 5.3.2), i.e. migration of surface oxygen from the catalyst to the gasifying support, as represented by reaction (7). Although there is broad consensus regarding its ability to account for many observed phenomena in catalytic carbon gasification, its kinetic quantification remains a challenge. A much more uncertain mechanistic issue is which one of the presumably elementary steps outlined above, if any, is the rate-determining step. Indeed, the concept of the rate-determining step (RDS) when applied to catalytic gas–solid reactions (as opposed to heterogeneously catalyzed gas–gas reactions) is arguably in need of some theoretical clarification because, even in the absence of catalyst deactivation, the number of catalytically active sites and the catalyst–‘‘support’’ interfacial contact are both temperature- and conversion-dependent parameters. Figure 8 illustrates the associated complexities [42]: chemisorption of NO is facile on Kx Oy sites (see Fig. 6), as reflected in the large increase in catalyst-bound surface oxygen at low temperatures; oxygen spillover from CaO to carbon is also facile, as reflected in the large increase in carbon-bound surface oxygen at low temperatures. Thus, the RDS is expected to shift from adsorption (G1) to spillover (G2) or vice versa, as temperature increases. Alas, for reasons discussed below, all this is difficult to prove by analyzing the changes in a reaction’s activation energy, as routinely done in catalytic studies. A key fundamental question that has direct impact on the practical issue of catalyst effectiveness is whether catalytic action is a consequence of a lower activation energy (implying a change in the RDS and a higher turnover rate) or a higher pre-exponential (or frequency) factor in the

13.19.4 Mechanisms

3043

Carbon gasification Carbon deposition H2O

2H2

CO + H2

C CH4 (1) Hydrocarbon dissociation on catalyst particle surface (2) Diffusion of carbon through catalyst particle (3) Precipitation and growth of carbon (nanofibers)

C C∗

(3) Carbon removal from catalyst particle surface (2) Diffusion of carbon through catalyst particle (1) Dissolution of carbon in catalyst particle

Dissolution–diffusion–precipitation mechanism of gasification or formation of carbon in the presence of transition metal catalysts (e.g. Fe, Co, Ni).

Fig. 7

Rate = k1 G1 = k2 G2 = k3 G3 [O]g

[O]g

G1 G1

[O]cat

G1

G1

Concentration

[O]cat

(a)

G2

G2

G2

G2

[O]c

Temperature

(b)

[O]c

Temperature

Temperature dependence of the rate-determining step (RDS) in the catalytic sequence: (a) potassium-catalyzed carbon gasification with NO and (b) calcium-catalyzed gasification with NO. The RDS is the one that has the largest concentration gradient (G).

Fig. 8

Arrhenius equation, implying an increase in the number of (re)active sites. Evidence for both can be found in the literature. For example, an increase in the number of active sites has been advocated by Radovic and coworkers [8, 59] and Kapteijn et al. [60]. Whether the true activation energy is also affected depends on several factors. Thus, for example, it is well documented that the activation energy is higher for oxidation of a more ordered (more ‘‘graphitic’’) carbon than for oxidation of a less ordered carbon. Also, the desorption-controlled reaction typically has a higher activation energy than the adsorption-controlled reaction; and as the reaction temperature decreases or reactant pressure increases there is often a transition from adsorption control to desorption control. Another complicating issue in this regard is the compensation effect [11, 61], a well-known phenomenon in

heterogeneous catalysis [62], although not necessarily a well-understood one. Even when catalyzed gasification has a lower activation energy than the concurrent uncatalyzed gasification (see Fig. 9), at temperatures close to the isokinetic point (Tiso ), and when the fraction of the reactive carbon surface (s) is small (a common occurrence), the observed activation energy corresponds to that of the uncatalyzed reaction; furthermore, under the most relevant reaction conditions (T < Tiso ) the observed activation energy decreases as the catalyst surface concentration increases. Therefore, mechanistic inferences based predominantly on observed activation energies must be viewed with caution. References see page 3044

3044

13.19 Catalysis in Coal and Carbon Gasification

13.19.5 s = 0.50 s = 0.10 s = 0.01 s = 0.0001

Reaction rate / arbitrary units

0.01

0.001

0.0001

k = s Acat exp − 0.00001

0.0007

Ecat RT

+ (1 – s) Auncat exp −

0.0008

Euncat RT

0.0009

0.001

0.0011

Reciprocal temperature / K−1

Compensation effect in catalytic carbon gasification: isokinetic temperature = 1273 K; Ecat = 100 kJ mol−1 ; Euncat = 200 kJ mol−1 .

Fig. 9

A final brief comment is in order about the inhibition of gasification reactions. This is sometimes referred to as ‘‘negative catalysis’’ and the practically important boron–carbon system [63] offers a nice example of why such a term is not as inappropriate as one might think. Only a few elements in the periodic table are known to be ‘‘intrinsic’’ inhibitors [64]; among them, boron and phosphorus dopants have been investigated the most. When boron is substituted for carbon in the sp2 -hybridized graphene lattice, it has two potentially conflicting electronic effects, apart from the reduced total electron density in the graphene layer [65]: (a) decreased contribution of delocalized π electrons to the electron density of the remaining carbon atoms; and (b) σ electron localization on carbon atoms in the lattice due to the lower electronegativity of boron. The dominant effect depends on the extent of boron doping and reaction conditions. The often observed inhibition has been interpreted as a consequence of suppressed chemisorption of the electrophilic reactant, due to the lower electron density at the graphene edges. However, the above-mentioned redistribution of π electrons can lead to a weakening of the C−C bonds in the graphene layer [24, 53], and this is consistent with a catalytic effect on the desorption of products, which has also been observed experimentally. Therefore, at least in this case, catalysis and inhibition cannot be viewed as mutually exclusive concepts and this is yet another example of a compensation effect [65]: concomitant with the decrease in the number of (re)active sites there is an increase in the gasification turnover frequency.

Conclusion

A broader view of catalyzed coal gasification has allowed us to highlight the progress and the challenges in this field of energy- and materials-related catalysis. Initial progress has been driven largely by the need to commercialize a process for producing methane from coal. While this remains an important goal, especially in coal-rich countries, the more recent trends include environmental applications (e.g. NOx removal) as well as the preparation of novel materials (e.g. specialty adsorbents and oxidation-resistant carbon composites). In the design and optimization stages of all these processes, the challenges for scientists and engineers are typically even greater than in the case of more conventional applications of heterogeneous catalysis. On the one hand, this means that much remains to be learned about the control of activity, selectivity and regenerability of gasification catalysts. On the other, as this review has strived to demonstrate, significant progress has been made in applying and developing some of the basic concepts in heterogeneous catalysis – e.g. catalyst dispersion, catalyst–support interfacial contact, spillover, rate-determining step, compensation effect, redox mechanism – to advance the understanding of how alkali, alkaline earth and transition metals, and even boron, catalyze (or inhibit) the gasification of coals and carbons with oxidizing gases. References 1. H. Heinemann, in Handbook of Heterogeneous Catalysis, G. Ertl, H. Kn¨ozinger, J. Weitkamp (Eds.), Wiley-VCH, Weinheim, 1997, p. 2074. 2. J. J. Morgan, in Chemistry of Coal Utilization, Vol. II, H. H. Lowry (Ed.), Wiley, New York, 1945, p. 1673. 3. M. A. Elliott (Ed.), Chemistry of Coal Utilization, Second Supplementary Volume, Wiley, New York, 1981, 2374 pp. 4. J. A. Moulijn, M. Makkee, A. van Diepen, Chemical Process Technology, Wiley, Chichester, 2001, 453 pp. 5. R. L Hirsch, J. E. Gallagher Jr., R. R. Lessard, R. D. Wesselhoft, Science 1982, 215, 121. 6. H. Perry, Sci. Am. 1974, 230(3), 19. 7. R. F. Probstein, R. E. Hicks, Synthetic Fuels, McGraw-Hill, New York, 1982, 490 pp. 8. L. R. Radovic, P. L. Walker Jr., R. G. Jenkins, J. Catal. 1983, 82, 382. 9. R. Meijer, R. van Doorn, F. Kapteijn, J. A. Moulijn, J. Catal. 1992, 134, 525. 10. C. A. Mims, J. J. Krajewski, J. Catal. 1986, 102, 140. 11. D. W. McKee, in Chemistry and Physics of Carbon, P. L. Walker Jr, P. A. Thrower (Eds.), Marcel Dekker, New York, 1981, p. 1. 12. H. J¨untgen, Fuel 1983, 62, 234. 13. J. R. Pullen, Catalytic Coal Gasification, International Energy Agency for Coal Research, London, UK, 1984. 14. B. J. Wood, K. M. Sancier, Catal. Rev. Sci. Eng. 1984, 26, 233.

13.20.1 Fuel Processors 15. J. A. Moulijn, F. Kapteijn, in Carbon and Coal Gasification, J. L. Figueiredo, J. A. Moulijn (Eds.), Martinus Nijhoff, Dordrecht, 1986, p. 181. 16. K. J. H¨uttinger, Erd¨ol Kohle Erdgas Petrochem. 1986, 39, 261. 17. C. A. Mims, in Fundamental Issues in Control of Carbon Gasification Reactivity, J. Lahaye, P. Ehrburger (Eds.), Kluwer, Dordrecht, 1991, p. 383. 18. R. T. Yang, in Chemistry and Physics of Carbon, Vol. 19, P. A. Thrower (Ed.), Marcel Dekker, New York, 1984, p. 163. 19. R. T. K. Baker, in Coal and Carbon Gasification, J. L. Figueiredo, J. A. Moulijn (Eds.), Martinus Nijhoff, Dordrecht, 1986, p. 231. 20. A. Tomita, Catal. Surv. Jpn. 2001, 5, 17. 21. J. R. Katzer, in Chemistry and Chemical Engineering of Catalytic Processes, R. Prins, G. C. A. Schuit (Eds.), Sijthoff and Noordhoff, 1980, p. 563. 22. D. W. McKee, C. L. Spiro, P. G. Kosky, E. J. Lamby, ChemTech 1983, 13, 624. 23. J. L. Johnson, Kinetics of Coal Gasification: a Compilation of Research, Wiley, New York, 1979, 234 pp. 24. P. L. Walker Jr., M. Shelef, R. A. Anderson, in Chemistry and Physics of Carbon, Vol. 4, P. L. Walker Jr. (Ed.), Marcel Dekker, New York, 1968, p. 287. 25. T. Baker, Chem. Ind. (London) 1982, 18, 698. 26. N. C. Nahas, Fuel 1983, 62, 239. 27. J. Haggin, Chem. Eng. News 1983, 61(43), 26. 28. www.science.doe.gov/sbir/awards− abstracts/sbirsttr/cycle19/ phase2/100.htm; accessed 15 December 2005. 29. A. M. Leas, US Patent 5 641,327 1997; A. M. Leas, US Patent 5 776 212, 1998; A. M. Leas, US Patent 5 855 631, 1999. 30. N. Berkowitz, An Introduction to Coal Technology, Academic Press, New York, 1979, 345 pp. 31. www.netl.doe.gov/coal/gasification/description/gasifiers. html; accessed 19 December 2005. 32. L. R. Radovic, F. Rodr´ıguez-Reinoso, in Chemistry and Physics of Carbon, Vol. 25, P. A. Thrower (Ed.), Marcel Dekker, New York, 1997, p. 243. 33. T. Suzuki, S. Nakajima, Y. Watanabe, Energy Fuels 1988, 2, 848. 34. S. Sampath, N. K. Kulkarni, M. S. Subramanian, N. C. Jayadevan, Carbon, 1988, 26, 129. 35. D. W. McKee, J. Catal. 1986, 97, 264. 36. H. Harker, in Proceedings of the 4th Conference on Carbon, American Carbon Committee, Pergamon Press, Oxford, 1959, p. 125. 37. F. Kapteijn, J. A. Moulijn, in Carbon and Coal Gasification, J. L. Figueiredo, J. A. Moulijn (Eds.), Martinus Nijhoff, Dordrecht, 1986, p. 291. 38. C. L. Spiro, D. W. McKee, P. G. Kosky, E. J. Lamby, Fuel 1983, 62, 180. 39. H. Gupta, S. A. Benson, L. S. Fan, J. D. Laumb, E. S. Olson, C. R. Crocker, R. K. Sharma, R. Z. Knutson, A. S. M. Rokanuzzaman, J. E. Tibbets, Ind. Eng. Chem. Res. 2004, 43, 5820. 40. H. Yamashita, H. Yamada, T. Kyotani, L. R. Radovic, A. Tomita, Energy Fuels 1993, 7, 85. 41. M. J. Ill´an-G´omez, C. Salinas-Mart´ınez de Lecea, A. LinaresSolano, L. R. Radovic, Energy Fuels 1998, 12, 1256. 42. M. J. Ill´an-G´omez, A. Linares-Solano, L. R. Radovic, C. Salinas-Mart´ınez de Lecea, Energy Fuels 1996, 10, 158. 43. A. Linares-Solano, C. Salinas-Mart´ınez de Lecea, D. CazorlaAmor´os, J. Joly, in Fundamental Issues in Control of Carbon

44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. 61. 62. 63. 64. 65.

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Gasification Reactivity, J. Lahaye, P. Ehrburger (Eds.), Kluwer, Dordrecht, 1991, p. 409. H. J. Grabke, Mater. Corros. 1998, 49, 303. L. R. Radovic, P. L. Walker Jr, R. G. Jenkins, Fuel 1984, 63, 1028. H. Marsh, K. Kuo, H. Marsh (Ed.), in Introduction to Carbon Science, Butterworths, London, 1989, Ch. 4. T. Haga, K. Nogi, M. Amaya, Y. Nishiyama, Appl. Catal. 1991, 67, 189. A. C. Sheth, C. Sastry, Y. D. Yeboah, Y. Xu, P. Agarwal, J. Air Waste Manage. Assoc. 2003, 53, 451. D. W. McKee, C. L. Spiro, P. G. Kosky, E. J. Lamby, Fuel 1985, 64, 805. J. Carrazza, J. J. Chludzinski, H. Heinemann, G. A. Somorjai, R. T. K. Baker, J. Catal. 1988, 110, 74. R. T. K. Baker, N. S. Dudash, C. R. F. Lund, J. J. Chludzinski Jr, Fuel 1985, 64, 1151. P. L. Walker Jr., F. Rusinko Jr., L. G. Austin, in Advances in Catalysis, Vol. 11, D. D. Eley, P. W. Selwood, P. B. Weisz (Eds.), Academic Press, New York, 1959, p. 133. F. J. Long, K. W. Sykes, Proc. R. Soc London, Ser. A 1952, 215, 100. K. J. H¨uttinger, J. Adler, G. Hermann, in Carbon and Coal Gasification, J. L. Figueiredo, J. A. Moulijn (Eds.), Martinus Nijhoff, Dordrecht, 1986, p. 213. L. S. Lobo, D. L. Trimm, J. Catal. 1973, 29, 15. J. L. Figueiredo, D. L. Trimm, J. Catal. 1975, 40, 154. J. L. Figueiredo, Mater. Corros. 1998, 49, 373. L. P. Bir´o, C. A. Bernardo, G. G. Tibbetts, P. Lambin (Eds.), Carbon Filaments and Nanotubes: Common Origins, Differing Applications, Kluwer, Dordrecht, 2001, 366 pp. L. R. Radovic, P. L. Walker Jr, R. G. Jenkins, Fuel 1983, 62, 209. F. Kapteijn, O. Peer, J. A. Moulijn, Fuel 1986, 65, 1371. F. S. Feates, P. S. Harris, B. G. Reuben, J. Chem. Soc., Faraday Trans. 1 1974, 70, 2011. A. K. Galwey, Adv. Catal. 1977, 26, 247. L. R. Radovic, in Encyclopedia of Materials: Science and Technology, K. H. J. Buschow (Ed.), Pergamon Press Elsevier, Oxford, 2001, p. 975. D. W. McKee, in Chemistry and Physics of Carbon, Vol. 23, P. A. Thrower (Ed.), Marcel Dekker, New York, 1991, p. 173. L. R. Radovic, M. Karra, K. Skokova, P. A. Thrower, Carbon 1998, 36, 1841.

13.20

Fuel Cell Related Catalysis 13.20.1

Fuel Processors Ralf Peters∗

Introduction Fuel cell technology is of special interest from the present perspective due to its higher efficiency in energy 13.20.1.1

References see page 3076 ∗ Corresponding author.

3046

13.20 Fuel Cell Related Catalysis

conversion and less harmful emissions, especially if hydrogen is used as energy carrier. However, in the medium term, pure hydrogen generated from renewable or sustainable energy sources will not be extensively available, which is a requirement if fuel cells are to penetrate the market. Additionally, the present hydrogen storage systems for automotive applications are not able to match the operating range of conventional technologies based on internal combustion engines and fossil fuels. In order to open the market for fuel cell systems, a significant cost reduction must be achieved. This is only possible with continuous component development including innovative steps. A stepwise market entry via niche markets is a viable way to increase production volumes, thus achieving significant cost reductions and paving the road to mass markets and production. In order to introduce fuel cells systems successfully, a maximum number of possible applications must be addressed. These applications are based on different energy carriers. Only hydrogen, methane and methanol can be converted to electricity by the electrochemical processes of the fuel cell. Other fuels must be converted to hydrogen-rich fuel gases by chemical processes, so-called fuel processing. Therefore, all energy carriers currently used with an existing infrastructure have to be taken into account. Fuel processor development started in the mid-1980s with methanol reformer systems for fuel cell drive systems. It must be stated that the combination of fuel processor technology and low-temperature fuel cells working at temperatures between 333 and 353 K is a rather complex issue. The automotive industry has thus shifted the research focus towards hydrogen as energy carrier. One reason was the lack of infrastructure for methanol and the complexity of the systems developed. In the meantime, automotive applications concentrate on gasoline and diesel. The first activities targeted drive systems, whereas later research and development activities were focused on the so-called ‘‘auxiliary power units’’ (APUs). The past decade has seen rather fast developments in technology for natural gas reforming centered on stationary systems mainly for domestic heating systems. Unfortunately, many of these activities have been stopped for reasons of cost and reliability. New trends are the development of APU systems for aircraft driven by the aviation industry [1, 2], the use of small methanol reformer systems for portable electronic devices [3, 4] and the development of high-temperature PEFCs with an operating temperature between 433 and 473 K [5]. This chapter will explain the major aspects of fuel processor development from the viewpoint of heterogeneous catalysis, taking into account several engineering tools for constructing fuel processor units for different applications based on a variety of fuels. With

respect to simple fuels such as methane, methanol and ethanol, detailed investigations have been made in the past to understand the processes on catalytic surfaces. Shifting to fossil-based fuels, the reacting species consist of thousands of different molecules with an unlimited number of chemical reactions. Here, the principles of reforming must be analyzed to overcome degradation effects, carbon deposition and so on. Finally, technical data must be available to construct components in the range from 100 We up to 400 kWe fuel cell power. Energy Carriers for Fuel Cells In order to analyze the functionality of fuel processors, several thermodynamic properties must be determined. First, it must be known whether the fuel consists of a single chemical component or whether it is a mixture of hundreds of different hydrocarbon species. Before the chemical reaction can start the reacting species must be heated or evaporated. The chemistry of the reaction can be affected to a certain extent by the composition at the reactor entrance, the conditions for the reaction and the catalyst used. Some reactions in fuel processing are controlled by the kinetics of the reaction, others by the chemical equilibrium. The next two sections will discuss the most important conditions for different fuels and various reforming processes. The different fuels which have been discussed for fuel processing in the past 10 years are alcohols such as methanol, ethanol and glycerin, ethers such as dimethyl ether, fossil fuel products and renewable fuels. Fossil fuels include natural gas as a mixture of short-chain alkanes starting with methane, ethane, propane, etc., which are also the light products of oil refining. Fuels such as liquefied petroleum gas, gasoline, kerosene, diesel and heating oil consist of long-chain paraffins, olefins, naphthenes and mono-, di- and tri-aromatics. An important issue with regard to catalyst poisoning is the type of sulfur species in different fuels. Figure 1 shows typical sulfur compounds in fossil fuels [6]. Some of them are typically found in diesel or gasoline according to their boiling temperatures [7]. In kerosene, benzothiophenes are found with various bonded methyl groups [8]. Furthermore, various fuels contain additives in order to assure special properties for specific applications. The specific properties for different fuels are guaranteed by the data given in regulations, i.e. DIN EN 228 and DIN EN 590 for gasoline and diesel, respectively. Given are the ranges of boiling temperature and density, the flame point, the contents of water, ash, oxygen, etc., and the cetane number for diesel fuel and the octane number for gasoline. Considering the use of these fuels in fuel processors, other properties are important. These are the residual amount that does not evaporate and the contents 13.20.1.2

13.20.1 Fuel Processors

Fig. 1

RSH Thiol

RSSR′ Disulphide

RSR′ Sulphide

Thiophene

1-Benzothiophene (BT)

Thianthrene

Dibenzothiophene (DBT)

4-Methyldibenzothiophene (4-MDBT)

4,6-Dimethyldibenzothiophene (4,6-DMDBT)

3047

Sulfur-containing hydrocarbons in fossil fuels [6].

of aromatics, i.e. mono-, di-, tri- and polyaromatics, olefins and naphthenes. Figure 2 shows typical boiling curves for different fuels, i.e. gasoline, kerosene, diesel and heating oil. Fuels such as gasoline and naphtha evaporate immediately at ambient pressure and at temperatures higher than 473 K. Middle distillates such as diesel and heating oil evaporate only at higher temperatures and leave a residue of up to 5% even at 673 K. It is evident for diesel and heating oil that during the evaporation process a residue remains, which can lead to carbon deposition. Methanol–water mixtures evaporate only in a narrow temperature range due to the similar thermodynamic properties of methanol and water. If the boiling temperature of a rather complex mixture of hydrocarbons is to be calculated, this implies the description of the real gas state by an equation of

state, for example Redlich–Kwong–Soave [15], and the usage of models for the liquid-phase behavior. Deviations from the ideal liquid-state behavior of mixtures can be calculated by models for the Gibbs excess enthalpy GE . Unfortunately, these models involve high complexity for multicomponent mixtures. Detailed information is given by Smith et al. [16]. Moreover, the calculation methods for the Gibbs energy allow the determination of chemical equilibria data. Most often an ideal mixture of ideal gases is assumed. With the aid of an equation of state [15], one can check this assumption under the conditions corresponding to a typical inlet state for reforming, with temperatures between 300 and 673 K at elevated pressures. Long-chain paraffins, olefins and diaromatics References see page 3076

750 Kerosene Diesel Heating oil Gasoline MeOH-water, 1:1.5

700

Boiling temperature / K

650 600 550 500 450 400 350 300 250 0

10

20

30

40

50

60

70

80

90

100

Evaporated fuel / %

Boiling curve for several fossil fuels and for a methanol–water mixture at 0.1 MPa showing the evaporated part of the original fuel as a function of boiling temperature. Data from Refs. [11–14].

Fig. 2

3048

13.20 Fuel Cell Related Catalysis

clearly show non-ideal behavior. Also, polar components, such as steam and methanol, show slight deviations from the ideal gas state. Under the operating temperatures of the reforming processes, in the range of 973–1273 K and at moderate pressures, the ideal gas model can be used. Keeping the limited resources of fossil fuels in mind, renewable energy carriers will attain increasing importance. Today, renewable energies contribute about 5.8% to total primary energy use in the European Community [17]. In the sector of transportation fuels, the share of biofuel is about 1.4% and therefore still far less than the 2% target which was envisaged for 2005 by the Biofuels Directive of the European Commission [18]. In Germany, the main source of biodiesel (RME, rape oil methyl ester) is the rape plant. In France, biodiesel is produced from sunflower and rape seed. The origin of the plant oil are three organic acids bonded to the triglycerin molecule. In the case of rape oil, about 90% of these branches are oleic acid. An esterification process starts with the addition of methanol, resulting in the formation of glycerin and RME. The reaction must be catalyzed by sodium or potassium hydroxide. Bioethanol as fuel is produced from sugar cane or sugar beet by a biological fermentation process via yeast: −−

C6 H12 O6  −− − − 2C2 H5 OH + 2CO2

(1)

A much better process route for fuel production from biomass is a combination of gasification and Fischer–Tropsch synthesis. This method is attractive because other energy carriers such as coal and natural gas can be converted to a synthetic gas to feed the gasto-liquid process (GTL). If biomass is used as feedstock, the method is frequently described as biomass-to-liquid (BTL). Major plants for GTL are located in South Africa, Qatar and Malaysia, with a capacity of about 107 t a−1 with an increasing trend. The properties of diesel fuels produced by Fischer–Tropsch synthesis result in excellent fuel specifications with a cetane number between 73.5 and 75, a density of about 775–785 kg m−3 , a flame point between 340 and 383 K and an aromatic content between 0 and 2.7%. The sulfur content is lower than 10 ppm. Fischer–Tropsch diesel is an ideal fuel for fuel processors. 13.20.1.3

Thermodynamic Calculations

Reforming Processes for Hydrogen Production One possible way of producing hydrogen from carbonaceous substances, such as hydrocarbons or alcohols, using a so-called fuel processor, is steam reforming. Fuel is converted to a hydrogen-rich gas according to the following 13.20.1.3.1

reaction: −

Cn Hm Ol + (n − l)H2 O −  −− − − nCO + (m/2 + n − l)H2

(
HR > 0)

(2)

For alkanes, the indices n, m and l amount to m = 2n + 2, l = 0, resulting in −−

Cn H2n+2 + nH2 O  −− − − nCO + (2n + 1)H2

(
HR > 0)

(3)

Except for methane and methanol, the steam reforming reaction can be regarded as irreversible at the temperatures commonly used for reforming, that is, higher than 773 K. In addition, two side reactions can be found, i.e. the water gas shift reaction (WGS) and methane synthesis, which is the opposite of methane steam reforming: −−

CO + H2 O  −− − − CO2 + H2

(
HR < 0)

(4)

−−

CO + 3H2  −− − − CH4 + H2 O

(
HR < 0)

(5)

According to the principle of Le Chatelier, more CO and H2 and less CH4 are formed with rising temperature. An increase in the ratio of water vapor content to the carbon content also causes a decrease in carbon monoxide content according to Eq. (4), a decrease in methane content and an increase in carbon monoxide content according to Eq. (5). Equation (2) leads to a decomposition reaction if methanol is used as fuel (n = l = 1; m = 4): −−

CH3 OH  −− − − CO + 2H2

(
HR > 0)

(6)

Further, steam reforming of methanol can be described by the formation of carbon dioxide: −−

CH3 OH + H2 O  −− − − CO2 + 3H2

(
HR > 0) (7)

As a side reaction, the reverse WGS reaction occurs: −

CO2 + H2 −  −− − − CO + H2 O

(
HR < 0)

(8)

Methanation according to Eq. (5) can be suppressed by the catalyst in methanol reforming. The reforming reaction is endothermic in the equilibrium state when low-CH4 gas mixtures are produced, i.e. the reactor must be heated. The required heat for steam reforming can be provided by the total oxidation of hydrocarbons:   m m l −−

Cn H m O l + n + − O2  −− − − nCO2 + H2 O 4 2 2 (9) Apart from the supply of heat for endothermic steam reforming, the total oxidation must also ensure the complete chemical conversion of residual gases of the

13.20.1 Fuel Processors

fuel cell system. In order to guarantee the proper electrochemical function of each cell and to overcome strong mass transfer limitations, a certain hydrogen partial pressure must still be present at the outlet of each anode gas channel. Hydrogen utilization in the fuel cell is not complete – hydrogen utilization rates between 75 and 92.5% are normal. Therefore, the anode exhaust gas contains hydrogen, but also carbon monoxide and possibly methane or other organic substances. In order to be able to comply with the required emission standards, catalysts are used in the burner. The composition of the gas flowing into the burner is dependent on the design and operating conditions of the fuel cell system. For the operation of a methanol steam reformer, for example, residual gases and fresh methanol must be burned free of emissions, i.e. NOx , CO and unconverted hydrocarbons. It should be noted that there is no limit on CH4 emissions according to present regulations. The amount of added methanol depends on the energy balance between steam reformer and the catalytic burner [19]. As an alternative to the steam reforming reaction, it is also possible to produce a hydrogen-containing gas mixture via the partial oxidation of hydrocarbons: Cn Hm Ol +

(n − l) −−

O2  −− − − nCO + mH2 2

(10)

via an exothermic reaction. In the partial oxidation of gasoline with air, 17% of the lower heating value is converted for the enthalpy increase of the feedstock, which corresponds to a temperature increase of the gas from 298 to approximately 1143 K [20]. The reaction can be performed as a homogeneous gas-phase reaction at temperatures higher than 1273 K and with the aid of a catalyst as a heterogeneous catalyzed reaction at temperatures below 1273 K. Another possibility for producing hydrogen is the pyrolysis reaction, which is described for LPG, i.e. propane and butane. This method has not been fully established yet, apart from some research projects [22]. C3 H8 −−−→ 3CO + 4H2

(11)

C4 H10 −−−→ 4CO + 5H2

(12)

Steam reforming according to Eq. (2) can also be performed at lower temperatures between 823 and 973 K and high pressures above 22 MPa using supercritical steam with the so-called supercritical steam reforming process without catalysts [23, 24]. It is advantageous that using this process even fuels with high levels of sulfur can be fed to the reformer. Considering fuel cell specifications and system design, a subsequent desulfurization unit to remove the H2 S must be installed behind the reformer. Unfortunately, the apparatus requires a high mass and a rather high reaction volume at high pressures due to the

3049

homogeneous reaction and therefore leads to low power densities. Other technologies are focused on combined processes. So-called autothermal reforming constitutes a compromise between heated steam reforming and partial oxidation [21]. The autothermal reforming process comprises the combined reaction of hydrocarbons with oxygen according to Eq. (10) and steam according to Eqs. (2), (4) and (5). In order to provide the required heat for steam reforming and to reduce the effort of heat transport, part of the fuel in the feed gas can be converted directly by total oxidation [Eq. (9)]. This method of fuel processing is referred to as ‘‘hybrid steam reforming’’ [25]. The reactor does not need a catalyst, which means that every possible fuel can be used. Analogous to supercritical steam reforming, H2 S removal is required and power densities are, in principle, lower than those for autothermal reforming. In another technology, heat is delivered to the reforming process by a plasma-supported flame [26–28]. This method can be combined with a catalytic reaction zone. Plasma reforming suffers from the high electricity demand of the plasma. The so-called ‘‘Glid-Arc’’ process tries to transfer only the electrons in a plasma state, which lowers the electricity demand to an acceptable level [27]. Thermodynamic Equilibrium Calculations Thermodynamic equilibrium calculations are often used to determine the preliminary reaction conditions and to evaluate the product gas quality. Figure 3 shows the temperature dependence of H2 concentration for different fuels and reforming processes. In addition, the sum of H2 and CO is given because a downstream water gas shift reactor could produce extra hydrogen by Eq. (4). As one can see, fuels such as methane, isooctane and higher alkanes demand reforming temperatures between 873 and 1073 K to reach an appreciable conversion indicated by higher H2 concentration under comparable conditions. Steam reforming results in higher H2 concentrations with 63% compared to 37.5% at 1073 K for the autothermal process. Higher alkane chains lead to lower H2 concentrations due to the worse C/H ratio of the original fuel. Lower reforming temperatures are possible for methanol. Assuming that carbon formation can be suppressed by a catalyst, hydrogen production occurs at temperatures below 473 K. In practice, commercial Cu/ZnO/Al2 O3 catalysts are operated between 523 and 573 ◦ K. Figure 4 shows the temperature dependence of hydrogen product gas concentration for methanol reforming under various conditions. If carbon formation 13.20.1.3.2

References see page 3076

3050

13.20 Fuel Cell Related Catalysis

100

80

90

70

80

Concentration / %

Concentration / %

60 70 60 50 40 30

50 40 30 20

20

10

10

0 250

0 250

500

750 1000 Temperature / K

1250

1500

Temperature dependence of H2 and (H2 and CO) concentration for different reforming processes of hydrocarbons at chemical equilibrium at 0.2 MPa. Conditions steam reforming (SR): H2 O/C = 2. Conditions autothermal reforming (ATR), H2 O/C = 2; O2 /C = 0.45. (white squares) SR CH4 (H2 and CO); (dark gray squares) SR CH4 ; (white diamonds) all ATR: CH4 (H2 and CO); (pale gray squares) ATR CH4 ; (black squares) C8 H18 ; (pale gray diamonds) C14 H30 ; (dark gray diamonds) diesel (H2 and CO); (black diamonds) diesel. Fig. 3

is taken into account, higher temperatures are required. The decrease in hydrogen concentration at higher temperatures is caused by the shift reaction. Therefore, the sum of H2 and CO is nearly constant at temperatures higher than the lowest temperature resulting in complete fuel conversion. 13.20.1.3.3 Carbon Formation and Deposition The most important side reaction which must be avoided in all these technologies is the formation of carbon, often wrongly described as soot formation. Carbon deposition is a very complex mechanism, which still has not been completely understood or described in physical terms [29, 30]. Carbon deposits in a reformer lead to catalyst deactivation and must therefore be avoided. Polycyclic aromatic hydrocarbons (PAHs) seem to be the cause of carbon formation. The first step is nucleation, the formation of particle structures from PAHs with molar mass greater than 500 g mol−1 by their accumulation or by the formation of new aromatic rings with the addition of five carbon atoms. The carbon particles reach a size of ca. 50 nm by surface growth and coagulation. This process is described in Ref. [31]. The most important precursor for the formation of aromatic rings is acetylene. Figure 5 shows the schematic reaction steps in the formation of the first aromatic compounds such as benzene, naphthalene

500

750

1000

1250

1500

Temperature / K

Temperature dependence of H2 and (H2 and CO) concentration for different reforming processes of methanol at chemical equilibrium at 2 MPa. All SR: (black squares) H2 O/C = 1.5 (H2 and CO); (gray squares) H2 O/C = 1.5; (white squares) H2 O/C = 1.5 with carbon formation; (black diamonds) H2 O/C = 2; (gray diamonds) H2 O/C = 2 with carbon formation; (black circles) ATR, H2 O/C = 2, O2 /C = 0.45 (H2 and CO); (white circles) ATR, H2 O/C = 2, O2 /C = 0.45. Fig. 4

and phenanthrene, which are decisive for the reaction pathway, the formation of higher PAHs and the growth of carbon particles. Figure 6 shows that carbon formation in the steam reforming of diesel is especially intensive at low temperatures (ca. 573–873 K). The air/fuel ratio and the amount of steam are also important. High air ratios are therefore preferable. Since the fuel should only be partially oxidized in the reformer and not completely be burned, the air ratios λ are always less than 1 (λ < 1). Palm et al. [35] used an air ratio between 0.22 and 0.30, which corresponds to an O2 /C ratio between 0.34 and 0.47 according to Eq. (13) for a hydrocarbon species Cn Hm :   nO2 n (13) λ= nC n + m 4 The parameters in the mixing chamber of the reformer are much more favorable for carbon deposition than in diesel engines and combustors. Steam, probably adsorbed in the form of hydroxyl groups (OH), shows a positive impact on the prevention of carbon formation. In order to avoid carbon formation and deposition, the reactors for diesel processing are operated at high steam/carbon ratios. In addition to the process of carbon formation described above, which occurs in the reforming reactions of higher

13.20.1 Fuel Processors

3051

Reaction time

d < 50 nm Coagulation Surface growth and coagulation Nucleation Particle zone

d < 50 nm

H2 CO2

CO

O2

H2 O

Molecular zone

Mixture of fuel and oxidant

Schematic reaction mechanism of carbon formation in homogeneous mixtures and premixed flames. From Ref. [31], p. 287.

30

70

Concentrations / %

60

25

CH4

50

CO

H2 40

CO2

H2O 30

20 Diesel Heated steam reforming p = 0.1 MPa H2O:C = 15

15 10

20

Concentrations / %

Fig. 5

soot 5

10 0 600

0 800

1000

1200

Temperature / K

Fig. 6

Temperature dependence of the gas composition in heated steam reforming of diesel.

hydrocarbons, other processes of carbon formation and deposition have also been identified, especially in the area of fuel injection and evaporation. At the Los Alamos National Laboratory (LANL), diesel reforming experiments were carried out, particularly with regard to carbon deposition. According to Borup et al. [36], diesel shows a strong tendency to pyrolyze. The nozzle temperature is an important parameter here. This will be discussed in the next section. There are different mechanisms of carbon formation, all of which affect the carbon morphology. The most common structures are:

• carbon whisker • encapsulated carbon • pyrolytic carbon. The characteristics of the individual carbon forms are described in Ref. [38]. In contrast to other carbon structures, carbon whisker did not cause catalyst deactivation. The following side reactions of the reforming process are referred to as the main sources of carbon formation: −−− −− → 2CO ← − CO2 + C References see page 3076

(Boudouard reaction) (14)

3052

13.20 Fuel Cell Related Catalysis

−− −− → CH4 − ← − 2H2 + C

(methane decomposition) (15)

−−− −− → CO + H2 ← − H2 O + C

(CO hydrogenation) (16)

Cn Hm −−−→ ‘‘carbon deposit’’ + m/2H2

(17)

Reactions (14)–(16) are reversible and Eq. (17) is irreversible for n > 1. For n = 1 reaction (17) is identical with Eq. (15). Steam is mostly considered to be a means of preventing carbon formation. The addition of steam directly reduces carbon production in reaction (16). Steam is not present in reactions (14) and (15) and only affects carbon formation indirectly here. The variation of educt and product concentrations caused by reaction (16) has an impact on the reaction pathway of Eqs. (14) and (15), because all three reactions take place simultaneously. Theoretically, carbon deposition can be expected below a certain H2 O/C ratio. It is very important to determine the minimum amount of steam needed to avoid carbon formation. Based on the thermodynamic equilibrium, this amount can be calculated using the principle of minimization of the Gibbs free enthalpy. More details are given in Ref. [38]. For a specific gas composition, the potential for carbon formation corresponds to the free enthalpy −
Ge according to the following equation: −
Ge = RT ln(Kp /QR,e )

(18)

where Kp is the equilibrium constant under standard conditions (0.1 MPa, 298 K) and QR,e the reaction quotient of reactions (14) and (15), respectively, after equilibrium of the steam reforming and partial oxidation reactions; see Ref. [38]. Carbon formation is a complex network of heterogeneous reactions. Solid carbon ideally has an activity ac = 1 in equilibrium and therefore does not appear in the equation for the determination of the equilibrium constant K:  ν yi i (19) K= i

where yi is the molar fraction and νi the stoichiometric coefficient of the ith reaction component. This means that the carbon concentration does not have a direct effect on the reaction equilibrium. At a defined gas composition, there is a temperature TB below which there is a thermodynamic potential (
Ge < 0) for the exothermic Boudouard reaction (14) and a temperature TM above which carbon formation takes place by endothermic decomposition: see the example for methane in Eq. (15) and in Ref. [38]. Thermogravimetric studies allow the determination of a critical educt composition, above which carbon formation does not appear. These considerations are restricted to methane.

Even when the thermodynamic conditions favor carbon formation, the reaction kinetics can be so slow that carbon formation does not take place. A hydrocarbon can decompose instead of reacting with steam, even if the thermodynamic equilibrium does not favor carbon formation. For example, this occurs on the surface of a catalyst pellet which shows an affinity for carbon formation (
G < 0). A hydrocarbon radical as an intermediate product of reforming reactions can react on the catalyst surface to give either product gas or carbon whisker. Carbon nucleation appears when the activity (concentration) of the carbon dissolved in the catalyst exceeds the saturation activity. Real gases have a potential for carbon formation unless higher hydrocarbons are present. Consequently, carbon formation depends on the kinetics and on local reaction equilibrium [38]. Therefore, determination of carbon formation based on reaction kinetics is more relevant than that based on equilibrium. More than 95% of carbon is formed in the second process step, surface growth. The greatest challenge in modeling and calculation of this process is that surface growth is not a gas-phase reaction but a heterogeneous process. Mass growth typically rises asymptotically and approaches a limit value [31]. The variations in fuel processor design result from the number of different fuels, the choice of fuel cell technology, the kind of application and the selected reforming process. There is a strong dependence between fuels and reforming processes leading to promising fuel processing routes. Therefore, the reforming processes must be analyzed on the basis of the special properties required for the targeted application. It might be that the special features of a given technology mean that it cannot be used in certain cases, e.g. the necessity of a steam supply would not be appropriate in automotive applications [32–34]. Generally, partial oxidation leads to compact reformer units and a lower efficiency compared with steam reforming. Steam reforming is characterized by high system efficiency and mostly by rather complex systems and lower power densities. This results in a larger reactor volume, more thermal mass and a poorer dynamic performance. The next sections will be focused on the most successful combinations of reforming processes, fuels and fuel cells. Steam reforming will be discussed for methanol for a polymer electrolyte fuel cell (PEFC) and as a solid oxide fuel cell (SOFC) internal reforming process for methane. Autothermal reforming is a preferred option using diesel for PEFCs and high-temperature PEFCs. In addition, partial oxidation (POX) is a possible fuel processing step for the SOFC using fuels such as liquid petroleum gas (LPG), gasoline, kerosene or diesel.

13.20.1 Fuel Processors

Fuel Pre-Processing Fuel pre-processing units must perform two main tasks: 13.20.1.4

• catalyst protection against harmful substances in the fuel • evaporation of liquid educts. 13.20.1.4.1 Removal of Harmful Species from Methanol and Water Different catalyst materials imply a variety of species which must be separated from the fuel. The most important catalyst poison is sulfur. Non-precious materials such as Ni, Co, Fe and Cu are destroyed at sulfur concentrations higher than several tens of ppb. Precious metal-based catalysts can tolerate sulfur concentrations in the range 1–10 ppm. If copper-containing catalysts are used, as is especially the case in methanol steam reforming, it should be noted that the feedstock may only contain extremely small quantities of chlorine, sulfur and its components; see Section 13.20.1.4. A typical chlorine concentration in ultra-clean water purified by reverse osmosis membranes amounts to 10 ppb Cl− and for methanol about 100 ppb Cl− . Ion exchanger cartridges can reduce the chlorine load to one-third. Therefore, ion exchangers should be used in the liquid phase in combination with filters to protect the catalyst in the gas phase. Unfortunately, they are not optimized for methanol–water mixtures. The aim is to lower the ion concentration to values corresponding to a conductivity of 1200 h) 2

1

0 0

20

40

60

80

100

Methanol conversion / %

CO selectivity as a function of methanol conversion. Water:methanol molar ratio 1.5:1; mixture density 0.905 kg L−1 (298 K); temperature of heating fluid 553 K; operating pressure 0.38 MPa; methanol feed flow 3.8 mol h−1 each tube [100].

Fig. 12

A typical mixture exhibits low ion concentrations, i.e. −1 Si and 80 µg L−1 Cl− , 30 µg L−1 SO2− 4 , 50 µg L −1 + 108 µg L Na , with a conductivity of 0.35 µS cm−1 . In addition, it should be considered that during a catalyst operation time of 1000 h and an educt flow-rate of 250 mL h−1 , the educt volume in contact with the catalyst amounts to 250 L. The chlorine, for example, contained in this volume weighs 0.02 g, leading to a concentration of 280 ppm relative to a catalyst mass of 71 g. The traces of chlorine determined by means of chemical analysis amount to 570 ppm Cl in the inlet section, 137 ppm in the middle and 99 ppm at the end of the catalyst bed. Sulfur is located more or less exclusively in the inlet section at 1120 ppm compared with 100 ppm from the − SO2− 4 balance. Poisoning of the catalyst with Cl leads to accelerated sintering of the active Cu species on the catalyst surface. Ultimately, an ion-exchange cartridge must be used due to the Cl sensitivity of the catalyst in order to prevent sintering. If this is done, stable operation without degradation effects is possible for 300 h of operation checked by experiments. One might ask whether information on the active sites of the catalyst can be obtained by technically oriented measurements. The start-up behavior is of special interest for the STR and the CMR test runs [99]. Figure 13 shows the CO concentration during start-up depending on the amount of hydrogen purged before start-up. The start-up behavior can be separated into three different phases. In the first phase (0–10 min), just hydrogen from the purging process leaves the reactor until the fresh methanol–water mixture passes into the evaporating zone and evaporates immediately. In the second phase (10–23 min after start-up), the catalytic reaction starts and hydrogen is formed in the reforming

reaction. This is indicated by the simultaneous formation of carbon oxides, i.e. CO2 and CO. The concentration of carbon monoxide in the product gas is of special interest especially for the gas treatment unit. The shape and width of the CO peak depend on the amount of hydrogen previously purged. At high purge rates, a high CO concentration of about 30% occurs for several minutes and then the concentration suddenly drops to 10%. The steady-state equilibrium of 1% CO is then reached almost via a step change between 14 and 23 min. The H2 content remains almost constant at 67% during this phase. It is striking that the sum of carbon dioxide and carbon monoxide concentrations in the start-up phase is almost constant. Only the ratio of CO to CO2 changes dramatically. Such processes are not observed during load changes. The highest amount of CO formed during the start-up of the reforming process is 8 LN CO after purging 71 g of the catalyst with 12 000 LN H2 at a constant operation temperature of 553 K. The best option for overcoming the high CO level is to use much less hydrogen for purging. Dispensing with a hydrogen purge before starting up the methanol reformer can reduce the CO peak. Unfortunately, condensed water is usually harmful to the catalyst so that purging is recommended by the catalyst manufacturers. Long-term experiments, however, have shown that the hydrogen purge has only a minor influence on the degradation rate. The simulation of breakdown of the heating system, in which a liquid water–methanol mixture is in close contact with the catalyst, did not reveal any immediate decline of the catalyst activity [65]. Catalyst degradation is mainly determined by tiny amounts of chlorine. On-going hydrogen purging during start-up does not influence CO production. In such a case, the hydrogen concentration in

13.20.1 Fuel Processors

3061

30 25 CO concentration (dry) / %

12 000 NI 6800 NI 2000 NI No purge

Time period for reaching steady-state concentration

20 15 10 5 0 0

5

10

15 Time / min

20

25

30

CO peak during start-up for the STR depending on the amount of H2 purged. Catalyst loading 50% (71 g); TSteam = 553 K; p = 0.38 MPa; methanol feed flow 3.8 mol h−1 at water:methanol molar ratio 1.5:1 [103].

Fig. 13

the reformer off-gas merely changes more slowly during start-up. Furthermore, neither the addition of steam nor a change from hydrogen to nitrogen purging affects the occurrence of CO peaks. Considering all the results described in this section, it can be stated that a purge time of 1 h with an amount of 560 LN kgcat −1 removes the wet synthesis gas from the catalyst. In this case, the CO peak is damped during start-up. A comparison between different catalysts shows that a high catalyst activity results in larger CO peaks. D¨usterwald [65] tried to localize the zone where the carbon monoxide was produced. The CO peak seems to be built up at the end of the catalyst bed, where during operation the main reaction zone of the shift reactor can be localized by experiments [65] and reactor models. There are different explanations for this behavior in the literature. It seems that carbon dioxide is reduced to carbon monoxide with the formation of oxides of some metal components on the catalyst surface. A change between reducing and oxidizing atmospheres varies the structure of catalyst particles, as discussed by Oveson et al. [87]. Details of catalyst development and the layout and operation of reformers on a laboratory scale are described in the literature. For obvious reasons, detailed preparation data for commercial catalysts are not discussed in the open literature or in patent applications. Analogously, industrial research and the development and manufacturing of reformers will not be disclosed to the scientific community. Example of Methanol Steam Reformer Development In this section, the performance of a methanol steam reformer in the 25-kWe class will be discussed based

13.20.1.5.7

on papers by Emonts and coworkers [61, 104–107]. This example clearly reveals the general problems connected with steam reforming. Figure 14 shows the so-called compact methanol reformer. In order to illustrate the successful scale-up from single tube reformers to a compact reformer design, data from both experimental set-ups were compared [61, 105]. The deviation between STR and CMR results can be explained by the effect of chemical equilibrium for different pressures, i.e. 0.38 and 2.1 MPa, respectively. Before the liquid methanol–water mixture can be fed into the reformer, the CMR has to be heated to its operation temperature, that is, 533–553 K, for the steam circuit and the reformer section (see Fig. 15). Therefore, methanol is fed into the catalytic burner to provide the heat required. For the catalyst, 3 kg of Cu−ZnO−Al2 O3 has to be heated from ambient temperature to 533 K. Assuming a heat capacity of 0.8 kJ (kg K)−1 , it can be calculated that an energy difference of 564 kJ leads to a warm-up time of about 47 s at a maximum thermal power of 12 kW from the catalytic burner. Unfortunately, the mass of the steel casing with a heat capacity of 0.477 kJ (kg K)−1 also has to be heated up. For the worst case, 130 kg of the CMR must be heated to 533 K. The liquid methanol flow controller can begin operation between 36 and 39 min after the catalytic burner, at an inlet reformer temperature of 523 K. The temperatures of the catalytic burner and the exhaust gases reach their equilibrium values of 706 and 456 K, respectively, after 1 h. The heat provided by the burner amounts to about 6.5 kW. As can be seen, the initial assumptions about start-up are consistent with the experimental data. The disadvantage of the steam References see page 3076

3062

13.20 Fuel Cell Related Catalysis

Compact methanol reformer (CMR) for 61 kWth (H2 , LHV) corresponding to 25 kWel PEFC system and one of six catalytic burner units of 2.7 kWth installed in the bottom of the catalytic burner. Technical data: pressure, 2.1 MPa; temperature, 533–573 K; water:methanol molar ratio, 1.3–1.5; 3 kg catalyst from HTAS; CMR weight, 130 kg; volume, 142 L.

Fig. 14

900

8000 Volume flow, product gas/LN h−1

Temperature / K

Catalytic burner outlet temperature / K

6000

700 Q Steam temperature / K

600

4000

Reformer inlet temperature / K

Mass flow / gh−1; volume flow / LN h−1

800

Q

500

Mass flow, methanol input / g h−1

2000

400

300

0 0

Fig. 15

10

20

30 Time / min

40

50

60

Starting up a compact methanol reformer for 61 kWth (H2 , LHV).

circuit is a further retardation of 12 min due to the high thermal mass which is demanded by the high-pressure design. An improvement should be possible by reducing the thermal mass. The start-up strategy is discussed in more detail in Ref. [61]. In order to find preset values as an input for activating the mass flow controller in the drive test rig, the profile of the reformer reactant flow was determined by dynamic

simulation according to the New European Drive Cycle (NEDC). Relative to the methanol supply, it can be seen that the time lag between preset value feed and adjustment of the desired mass flow can range between 2 and 10 s depending on load. This is a consequence of the time constants of mass flow controllers for respective load steps, so that in a real drive system the flow controller must be replaced by a faster proportioning unit. A more

13.20.1 Fuel Processors

extensive analysis of the heat exchanger system was discussed by Tschauder et al. [107]. Finally, the time constants for a steam reforming system must be improved by an apparatus with a lower thermal mass. Nevertheless, such a system can be operated in a hybrid mode as an alternative option. It seems to be better to select the autothermal reforming process to achieve high dynamic operation and wide partial load ranges. Autothermal Reforming of Gasoline and Diesel By analyzing the state of the art in the reforming of hydrocarbons, it can be seen that huge progress has been made in the last decade. The application of hydrocarbon reforming has been modified due to a decision made by the US Department of Energy (DOE) in autumn 2004 [110]. Research and development activities in the USA have shown that the targets defined for gasoline reforming in fuel cell propulsion systems cannot be reached. Therefore, these activities have been shifted to electricity production with auxiliary power units for cars, trucks, ships and aircraft. Activities in the USA were focused on multi-fuel operation with fuels ranging from gasoline to diesel, kerosene (as JP-8 quality) and navy marine diesel. In Europe and Japan, developments were targeted more or less on special fuels. The results have been presented mostly at conferences and in proceedings reporting on progress in component development and installation. Catalytic aspects are published to only a minor extent. Special emphasis should be given to catalyst degradation, complete fuel conversion, power density and dynamic performance. In general, power density and dynamic aspects such as start-up and response time for load changes are coupled by the thermal mass. The power density of catalysts is very high, ranging from 20 kWe L−1 [111] for regular gasoline, i.e. 60 kWth L−1 , to 120 kWth L−1 for isooctane [112]. The target values are given by system aspects. Therefore, they will be discussed in later sections. Unfortunately, the collection of data for all items is far from complete for all reactor systems developed so far. A detailed benchmark cannot be presented. It is only possible to find relevant problems which must be solved in the future. It should be noted that the level of aromatic compounds in gasoline amounts to 40% and that of olefins 15% compared with 25 and 1%, respectively, for diesel. Kelly and Meyer [113] exclude gasoline steam reforming as an option due to the large amounts of aromatics and olefins. Borup et al. [114] showed by XPS investigations that both oxidative and steam reforming lead to carbon deposition on the catalyst surface, which increases with the complexity of the hydrocarbon bondings, i.e. 13.20.1.6

3063

paraffins < olefins < naphthenes < aromatics. Docter et al. [111] published experimental results for several brands of gasoline and to a certain extent for diesel and naphtha. Gasoline can be reformed up to a residual amount of non-methane hydrocarbons of 0.01% consisting mainly of C2 hydrocarbons. Benzene and toluene are analyzed at the detection limit of 2 ppm. For diesel they emphasized a preheating temperature of 823 K for the evaporation process. Mauzey et al. [115] reported on 268 start-up/shut-down cycles with their autothermal gasoline reformer for 175 h. The efficiency was 75–80% for gasoline and 73% for diesel, which is somewhat lower than the results of Docter et al. [111] of 80–85%. Berlowitz et al. [116] reported on dynamic studies leading to response times of 8 s for flows and 30 s for the final gas composition. Only tiny amounts of methane and propene could be observed during load changes for isooctane as model fuel. Kopasz et al. [10] discussed the reactivity of several species, i.e. isooctane, n-octane, methylcyclohexane, trimethylbenzene and toluene, at 873, 973 and 1073 K. Generally, reforming of isooctane is easiest, followed by n-octane > methylcyclohexane > toluene > trimethylbenzene. At low temperatures, the aromatic bonding cannot be cracked. Furthermore, the effects of sulfur and catalyst degradation were studied using isooctane as fuel. After 1000 h of operation with 26 start-up and shut-down cycles a degradation of 6% in flow and of 10% in hydrogen yield was observed. The BET surface area decreased from 13.5 to 3.6 m2 g−1 . Carbon deposition was assumed due to an increase from 0.1 to 2% C. An improved catalyst was tested with a model gasoline consisting of 74% isooctane, 20% dimethylbenzene, 5% methylcyclohexane and 1% pentene. It must be noted that the level of aromatics and olefins does not correspond to that of typical gasoline brands described in Ref. [111]. The number of start-up/shut-down cycles corresponds to five times per week. The hydrogen production rate drops by about 5% in 1000 h and the BET surface area decreases from 16.5 to 2.6 m2 g−1 . The addition of sulfur as benzothiophene leads to somewhat strange results. Without the addition of sulfur the hydrogen concentration decreased from 39 to 35% after 1000 h whereas the hydrogen concentration increased from 35 to 37.5% when sulfur was added. Finally, it should be stated that complete gasoline reforming was possible, but strong degradation effects were still observed. Work on propulsion systems based on gasoline has been terminated by the DOE decision [110]. As described above in Section 13.20.1.3, the most important issue in diesel reforming is the homogeneous mixing of diesel, steam and air. Carbon deposition can be avoided in combination with a uniform flow field References see page 3076

3064

13.20 Fuel Cell Related Catalysis

FZJ - IWV 2005

Fig. 16

EHT - 2.00 kV

Detector = SE2

WD = 7 mm

2 µm

SEM image of deposition in diesel fuel line next to the nozzle.

and a controlled temperature profile [45]. Attention must be given to the deposition of ashes in small channels. Figure 16 shows crystallite depositions from K2 SO4 as an SEM image in the inlet device of a reformer, and in Fig. 17 the corresponding EDX plot is given. Moreover, a further aspect is the reactivity of aromatic compounds in diesel. In general, hydrocarbon conversion increases with increasing oxygen to carbon ratio. At moderate gas hourly space velocities (GHSV) between 25 000 and 30 000 h−1 equilibrium concentrations can be reached, i.e. 9–12% CO, 0.2–0.5% CH4 , etc., depending on the thermodynamic conditions. Nearly complete conversion can be reached at O2 /C ratios between 0.4 and 0.5 and H2 O/C ratios of 1.7–2. The influence of fuel additives has been discussed by Konrad et al. [118] for diesel, i.e. corrosion inhibitor, antioxidants and defoamer, and by Lenz and Aicher [119] for kerosene, e.g. antioxidants, metal deactivator, static dissipater, corrosion inhibitor, lubricity improver and anti-icing additive. Most critical are additives based on silicon oil for defoaming and sulfur compounds to increase lubricity. Most developers observed degradation effects during their experiments [118–125] with decreasing hydrogen concentrations for partial oxidation, steam reforming and autothermal reforming. Typically, for an autothermal process carbon monoxide concentration increases whereas hydrogen concentration starts at 36–37% and decreases to 31–33% after several hundred hours. Reformer efficiency decreases from 80 to 70%. Lenz [120] reported on 300 h of degradation tests using low-sulfur kerosene at a high GHSV of 50 000 h−1 . Typically during their

experiments, several by-products of reforming increased, i.e. 0.1–0.35% methane, C6 ). Perna et al. [125] reported conversion values of 99.9% at the beginning and 90% after 1000 h. The most successful experiments were presented by Porˇs et al. [45, 126]. They achieved a conversion of 99.997% after 500 h and 99.7% after 1000 h using an advanced reactor with a commercial low-sulfur diesel containing 16% aromatics. To determine such high conversion levels the residual species must be analyzed in detail. After cooling of the product gas, the gas phase did not contain any non-methane hydrocarbons. The condensate had to be analyzed and finally Porˇs [126] found 0.5–2 µgorganic carbon L−1 , corresponding to a conversion better than 99.999% (Table 1). In order to avoid degradation effects in the subsequent components, a conversion of 99.99% should be aimed for. Residual carbon in the product gas of the reformer, bonded in hydrocarbons, not in CO, CO2 and CH4 , should be lower than 100 ppm C. In Japan, kerosene steam reformers for domestic heating systems were developed by Maeda et al. [127] and Saito et al. [128]. They claimed 100% conversion for 10 000 and 30 000 h, respectively. They both used an Ru-based reforming catalyst. Maeda et al. [127] promoted the catalyst with CeO2 to avoid carbon deposition by decomposition of H2 O into adsorbed OH and H groups. Unfortunately, attention was not given to residual substances and there is no detailed analysis. After the discussion of the technically oriented results, one might ask how heterogeneous catalysis can help to improve reactor performance. Figure 18 shows the

13.20.1 Fuel Processors

K C Ca Ca

3065

Pt S

O Cr Cr Fe

K

Na Zn

Cr

Zn Zn Mg Pt Fe

Pt

S

Ca K

Cl Cl

1 2 3 Full scale 446 cts cursor: 0.090 (9 cts) Fig. 17

Ca 4

Cr 5

Cr 6

Fe Pt Zn

Fe 7

8 KeV

EDX measurement of deposition in diesel fuel line next to the nozzle showing K and S. XRD indicates S–O bonding as K2 SO4 or

KSO3 . Diesel conversion of an autothermal reformer at 0.15 MPa: O2 /C = 0.47; H2 O/C = 1.9; GHSV ≈ 30 000 h−1 and residual concentrations

Tab. 1

Operating time/h

TOC(l)/µg L−1

Organic phase/L

C3 H7 (g)/ppmv

C4 H8 (g)/ppmv

C6 H6 (g)/ppmv

TOC (g)/ppmv

nC,out /mol h−1

nC,in /mol h−1

X/%

100 250 500 600 750 800 900 1000

2 3 17 33 62 72 120 130

– – – – – – – –

0 0 0 0 14 23 43 55

0 0 0 0 4 9 19 25

0 0 0 0 0 5 9 13

0 0 0 0 116 250 520 680

0.0003 0.0006 0.0030 0.0062 0.058 0.12 0.21 0.27

7.03 7.26 6.84 7.19 6.95 7.23 6.57 6.83

99.9997 99.9994 99.997 99.993 99.9 99.9 99.8 99.7

degradation results of Porˇs [126]. There is an overlapping effect of decreasing water gas shift activity and increasing by-products in the product gas. This can probably be explained in terms of the catalytic effects. Qi et al. [129] reported on the autothermal reforming of n-octane on Ru-based catalysts, i.e. Ru/K2 O–CeO2 /γ Al2 O3 . During their experiments they chose mild operating conditions, i.e. O2 /C = 0.3–0.45, H2 O/C = 1.2–.3, 723–1123 K and a GHSV of 1000 h−1 for n-octane, which corresponds to a GHSV for all inlet gases in the range 24 000–40 000 h−1 . They observed a continuously increasing CO/CO2 ratio in the product gas from 0.7 to 1.5 and explained this result by an increase in reverse water gas shift reaction. After 800 h of operation they observed a sudden decrease in octane conversion with a final conversion of 95% after 1000 h. They found a decrease

in BET surface area from 69.5 to 54.5 m2 g m2 g−1 and an increase in average pore size from 16.0 to 18.3 nm. Furthermore, a loss of Ru on the Al2 O3 support from 0.27 to 0.18 wt.% was found with XRF measurements accompanied by a loss of CeO2 and K2 O, i.e. from 2.94 and 2.63 wt.% to 0.85 and 2.02 wt.%, respectively. Qi et al. [129] interpreted the Ru loss by the strong etching and corrosive effect of K2 O on the Al2 O3 support. Additionally, carbon deposition occurred. Krumpelt et al. [130] described the development of new catalysts based on SOFC technology, where a transition metal is supported on an oxide-conducting substrate, such as ceria, zirconia or lanthanum gallate, which had been doped with a small amount of a non-reducible References see page 3076

3066

13.20 Fuel Cell Related Catalysis

40

H2 + CO (wet)

35 H2 (dry) Degradation in water-gas shift activitiy

Concentration / %

30 25 20 15 10 5 0 0

100

200

300

400 500 600 700 Time on stream / h

800

900 1000 1100

Long-term measurement of hydrogen production from diesel by an autothermal reforming process with educt mixtures with respect to O2 /C = 0.47; H2 O/C = 1.9; GHSV = 18 000–30 000 h−1 .

Fig. 18

element such as gadolinium, samarium or zirconium. The experimental conditions were O2 /C = 0.46, H2 O/C = 1.14, 773–1073 K and a GHSV of 3000 h−1 . They tested the transition metals Ru, Pd, Fe, Cu, Pt, Ni, Co and Ag, but only Ru resulted in a complete conversion for the whole temperature range. Pd required temperatures higher than 923 K for complete conversion. The highest selectivity H2 :H2,max was reached with Pd at 1073 K followed by Ru at 923 K. With increasing GHSV, the hydrogen product yield decreased drastically. The addition of up to 1300 ppm (mole) sulfur in the form of benzothiophene did not influence Pt as much. The Pt-supported Gd-doped Ce catalyst was also able to reform commercial gasoline. Cheekatamarla and Lane [131] reported on efficient bimetallic catalysts for hydrogen generation from diesel fuel. Their experimental results revealed that the impregnation of Ni or Pd, in addition to Pt on Al2 O3 and on CeO2 , increases activity and sulfur resistance. Surface analysis showed that the enhanced stability could be linked to strong metal–metal and metal–support interactions in the catalyst. This improved performance of the bimetallic catalyst is related to structural and electronic effects. Finally, the degradation effects are not fully understood. The complexity of diesel fuel leads to overlapping of a huge number of effects. Diesel and kerosene reformers have been designed on the basis of experimental experience and with the aid of CFD calculations. Kinetic considerations, however, give only an indication of proper reformer size. Pacheco et al. [132] proposed a kinetic model for isooctane reforming based on a Langmuir–Hinshelwood–Hougen–Watson formulation.

13.20.1.7 Internal Steam Reforming of Methane for SOFC Systems SOFCs operate at temperatures between 973 and 1273 K using hydrogen-containing gas mixtures as fuel and oxygen in the air as oxidant. At the Forschungszentrum J¨ulich (FZJ), an anode-supported planar substrate concept has been developed [133] consisting of thin electrolyte films (10 µm) deposited on thick substrates (1.5–2 mm). The thin electrolyte allows operating temperatures between 973 and 1073 K and thus the use of cheaper materials for the stack and the peripheral components compared with conventional systems with temperatures between 1223 and 1273 K. The most interesting fuel for SOFC systems is natural gas consisting mainly of methane, i.e. 80–95% CH4 . The H2 /CO-rich gas used in the SOFC can be produced by heated steam reforming (HSR) [38, 134, 135] or by partial oxidation (POX) [136–138]. The POX process has advantages of fast start-up, quick load changes and simplicity of the reformer. System efficiency and the H2 content in the reformate are drawbacks in comparison with steam reforming. Therefore, it is useful to apply the POX process in small systems such as portables and the steam reforming process in stationary systems where high efficiencies are required. The reforming process has to be integrated into the energy balance as part of the SOFC system. It is meaningful to perform the reforming process in two steps in different components, i.e. in a pre-reformer and within the anode chamber of the SOFC stack (internal reforming) [135]. The endothermic steam reforming of methane in the anode chamber is usefully applied to provide additional cooling of the cell and to reduce the expense of a pre-reformer. Because of higher

13.20.1 Fuel Processors

hydrocarbons in natural gas, i.e. 5.5% (vol.) C2 H6 , 1.3% C3 H8 , 0.5% C4 H10 , 0.1% C5 H12 , etc. (remainder: 88.3% methane, 3.0% nitrogen, 1.3% carbon dioxide) and several problems related to complete internal reforming, a small pre-reformer for higher hydrocarbons will have to be provided in future SOFC plant concept developments. More detailed information about the pre-reforming of natural gas is given by Peters et al. [139]. In order to optimize the process engineering design of the system, various process configurations were proposed by Riensche et al. [140]. This also includes recycling of anode exhaust gas to the pre-reformer to provide the steam required for reforming through the anode exhaust gas and thus making the steam generator redundant. The measurement results for a 10-kWe pre-reformer (SOFC) with simulated anode loop shows a considerable influence of CO2 on catalyst activity [139]. As a result, omitting the vaporizer leads to a significantly larger pre-reformer due to reduced catalyst activity with this anode loop. It was found that, in particular, carbon dioxide leads to greatly reduced activity of the anode cermet [141]. Higher partial pressures of hydrogen show a slightly higher methane conversion compared with measurements without the addition of hydrogen. Carbon monoxide displays neutral behavior in the experiments. Moreover, possible changes of the anode cermet during the reaction were examined. Two side reactions can occur: the formation of carbonaceous deposits on the catalyst surface due to methane decomposition: CH4 −−−→ C + 2H2

(34)

and the discharge of the catalytically active Ni2+ species via NiOH: −

Ni + 2H2 O −  −− − − Ni(OH)2 + H2

(35)

Figure 19 shows the effect of extremely high loads and flow velocities on the surface structure of the anode cermet. Figure 19a shows an SEM image of an unreduced

3067

specimen (NiO) for comparison. The anode cermet is subjected to a defined reduction procedure in the builtin state, adapted to that of the SOFC stack. After a continuous 4-week test, the specimens are removed again and examined by X-ray spectroscopy. Additional spectral analysis reveals that the catalytically active nickel species is depleted in the cermet at the gas inlet (Fig. 19b), whereas the nickel accumulates again as fairly large clusters in the rear region of the anode cermet (Fig. 19c). This migration effect of nickel can be suppressed by an equilibrium shift towards nickel by adding hydrogen. The effect is further slowed by the reduction of flow velocities in the channel. Generally, in a solid oxide fuel cell (SOFC) stack, the temperature distribution resulting from the current density distribution, the gas flow distribution and the heat transfer has to be homogeneous both within the cell plane and perpendicular to the cell plane. Large temperature gradients in either direction can cause damage to one or more of the components or interfaces due to thermal stresses. Experiments on internal methane steam reforming have shown that the strong activity of nickel results in a marked temperature reduction directly in the reaction zone. Since the associated temperature distribution cannot be tolerated in the SOFC stack, the reforming activity must be reduced by suitable measures. Several measures have been examined by experiments to optimize the catalyst activity and the flatness of the resulting temperature profile [142]. The SOFC stack performance concerning temperatures, current densities, reaction zones and mechanical stresses on a local scale must be calculated by suitable models [143]. Kinetic models for internal steam reforming of methane over nickel/zirconia have been evaluated by different groups [144–148]. Hecht et al. [148] described methane steam reforming with 42 reaction steps including reverse reactions. Their model incorporates channel flow, convective and diffusive porous media transport and elementary heterogeneous reactions. References see page 3076

Nickel deposition

10 µm

(a)

10 µm

(b)

10 µm

(c)

Fig. 19 SEM images of an anode cermet (30 × 10 × 1.5 mm). Flow velocity in the channel (30 × 1.5 × 1.5 mm) over the cermet, 17 m s−1 ; operating temperature, 973 K; load, 330 N m3 CH4 (m2 h)−1 (14 723 mol CH4 (m2 h)−1 ). (a) Unreduced cermet; (b) nickel depletion in entrance region; (c) cluster-type Ni deposits in the rear region of the cermet.

3068

13.20 Fuel Cell Related Catalysis

Several aspects of methane steam reforming over new materials such as Ni on CGO have been discussed [149–152].

• reduction of the concentration of carbon monoxide from 10–12% to 10–50 ppm • reduction of the concentration of sulfur in the form of H2 S from 1 ppm to a few ppb • emission reduction by catalytic combustion of residual fuel cell gas, i.e. H2 , CH4 and CO. In addition to CO and H2 S, other impurities or byproducts may also be present in the fuel gas which could degrade the performance of the PEFC. A rough specification of the fuel gas and a compilation of the influences of further possible impurities in the fuel gas on the performance of the PEFC were discussed by Amphlett et al. for methanol steam reforming with respect to formic acid (HCOOH) and formaldehyde (HCHO) [153]. In the case of hydrocarbon reforming, it can be assumed that the concentration of unconverted hydrocarbons might not exceed several tens of ppm. That leads to a conversion of better than 99.99%. 13.20.1.8.1 Water-Gas Shift Reaction The advantage of the water gas shift reaction, called CO conversion in the following, in comparison with all the other gas purification processes, is the fact that not only carbon monoxide is removed but also additional hydrogen is formed. Therefore, when more and more CO can be converted according to the conversion reaction, it will become more favorable for the overall balance of the fuel gas production system. The reaction equation for CO conversion is given in Eq. (4). The weakly exothermic reaction is aimed at decreasing the CO content. Thus lower temperatures promote CO conversion according to the chemical equilibrium (Fig. 20). It is therefore decisive to find catalysts that exhibit sufficiently high activity at low temperatures. Suitable catalysts are to be found within the Group Ib metals and also in Group VIII metal oxides and metal sulfides. The demand for sufficient stability under process conditions restricts the choice to metallic copper (Cu), iron oxide (Fe3 O4 ) and iron sulfide (FeS) [154, 155]. Fe3 O4 is a typical high-temperature catalyst for temperatures above 623 K. In the low temperature range, metallic copper is still catalytically active below 473 K, but the finely dispersed copper particles are sensitive to thermal sintering at higher temperatures. Grenoble et al. [156] investigated all Group VIII transition metals and Cu on different supports (Al2 O3 , SiO2

90 80 Conversion Xco / %

Gas Treatment for PEFC Systems The gas treatment units in fuel processor systems must mainly fulfill three different functions: 13.20.1.8

100

70 60 50

Equil. M = 1.3

40

Equil. LTS

30

HTS: GHSV = 25 000 h

20

LTS: GHSV = 10 000 h

−1

−1 HTS & LTS : GHSV = 7100 h

10 0 350

−1

450

550

650

750

850

950

Temperature / K Fig. 20 Performance chart of water gas shift reaction. HTS, high-temperature shift; LTS, low-temperature shift; HTS<S: XCO, NTS, in = 3.5%, X-axis shows TNTS , GHSV values in L h−1 [163].

and C). They found an optimal heat of adsorption of CO on Cu (about 80 kJ mol−1 ), leading to an increased activity of Cu by a factor of 50 in comparison with the elements in Group VIII. Detailed information about industrial water gas shift catalysts is given in Chapter 13.12. The challenges of automotive applications are fundamentally different from those of industrial units [157]. In order to develop WGS catalysts for fuel cells for mobile applications, several companies have developed new catalysts [158–161]. Wieland et al. [112, 158] reported that their catalyst does not require any in situ activation, it is non-pyrophoric and it is tolerant to temperature peaks. Furthermore, it can be coated on different substrates such as foams, monoliths and metallic surfaces and therefore fulfills the basic requirements of a WGS catalyst for mobile applications. The same group also reported that this catalyst does not undergo any methanation [162]. With respect to mobile applications, they pointed out that with their catalyst rapid start-up can be achieved by injecting air into the reformate stream. H2 and CO can be catalytically converted without any loss in activity for the WGS reaction. Balakos and Wagner [159] pointed to the improved tolerance of their precious metal catalyst to condensation, poisons and oxidation, which might occur during the numerous start-ups and shut-downs in mobile applications. Ruettinger et al. [160] developed a base metal non-pyrophoric particulate catalyst with a very promising catalytic behavior with respect to the requirements of fuel cell applications. This catalyst shows a very slight temperature increase of only 40 K without any deactivation if accidentally exposed to air. Lost activity due to liquid water exposure can be regenerated in situ or ex situ. They also coated this material on monolithic substrates which

13.20.1 Fuel Processors

can be operated at space velocities of 30 000 h−1 at temperatures between 533 and 573 K, resulting in only 1% CO in the product gas. Swartz et al. [161] investigated a Pt/CeO2 catalyst and discovered that it is non-pyrophoric and more active than commercial Cu-based catalysts at temperatures above 523 K. However, it showed significant deactivation rates caused by carbonaceous deposits. A deactivated Pt/CeO2 catalyst can be regenerated by annealing in air. A kinetic study of the WGS reaction on a Pt/MgO catalyst was performed by Wolf et al. [164]. They developed a very complex reaction scheme with five individual steps. CO2 and H2 adsorb on the MgO support whereas CO is bound to the Pt particles. These adsorption processes on both surfaces lead to the formation of several oxidized and reduced surface complexes which determine the reaction characteristics. Koryabkina et al. [165] determined the kinetic parameters of different Cu-containing catalysts. They found that Cu is the active site for catalysis since the addition of CeO2 and ZnO did not increase the rate per unit Cu surface area. The results of their kinetic measurements indicate strong inhibition of the forward reaction rate by the partial pressures of H2 and CO2 . They presented a redox mechanism with the reduction of surface oxygen by adsorbed CO as the rate-determining step. Baumann et al. [162] performed kinetic measurements which showed positive reaction orders for CO and H2 O but a slight inhibition by the partial pressures of CO2 and H2 . Furthermore, Choi and Stenger [166] tested several kinetic models. Finally they found good accord with experimental results for a commercial Cu−ZnO/Al2 O3 catalyst for adsorptive and redox mechanisms. The CO content can be reduced to 0.5–1% by the water gas shift reaction. Considering the size of such a reactor unit, an outlet concentration of 1% is meaningful. With respect to the influence of chemical equilibrium [see Eq. (4) and Fig. 20], it is obvious that the WGS reaction must be applied for autothermal reforming of hydrocarbons such as gasoline and diesel and that the WGS is not useful for methanol steam reforming. At 553 K and 0.38 MPa, the equilibrium concentration of CO amounts to 1.8% for methanol steam reforming, corresponding to an equilibrium constant of 55. At 473 K the equilibrium constant for the WGS reaction is 235, leading to a CO concentration of 0.5%. For autothermal reforming of tetradecane (C14 H30 ), a value for the WGS equilibrium constant of 0.8–1.6 is obtained, depending on the educt composition. Due to the much higher driving force expressed by the different equilibrium constants and with the aid of added steam, the WGS reaction can contribute to the major part of CO reduction in the product gas from a diesel reformer [167, 168].

3069

13.20.1.8.2 CO Fine Cleaning The possible processes of CO fine cleaning, i.e. for CO contents of less than 1%, can be basically divided into chemical and physical processes. The aim of the chemical processes is to lower the CO content in the fuel gas to a value for the PEMFC of below 10 ppm (Pt anode catalyst) and 100 ppm (Pt−Ru anode catalyst) with the aid of the catalytically supported reactions. The chemical processes include selective CO methanation and selective CO oxidation.

A H2 Separation Membranes Among the physical processes, in particular adsorption processes and membrane processes are of importance. Adsorption processes are used, among other things, for gas purification in stationary units for hydrogen production. Analogously to adsorption, the aim of the membrane processes is the separation of all impurities, especially of CO, from the product gas in order to obtain the purest hydrogen possible. The inlet mass flow is called feed, the mass flow passing through the membrane is the permeate and the residual mass flow is the retentate. The advantage of the membrane processes in comparison with the chemical conversion processes is that not only carbon monoxide but also all other components remain in the retentate. The reverse shift reaction towards CO is suppressed due to the lack of CO2 . The process step of a membrane process provides the favorable possibility of compressing liquid energy carriers and water in the liquid state using a fuel pump and of operating the reformer without appreciable compression losses at elevated pressure. This advantage is turned into a disadvantage in the case of gaseous energy carriers such as natural gas or air supply in POX reforming due to compression energy losses. Hydrogen of purity of ca. 99.9995% (vol.) can be produced with non-porous metal membranes, especially Pd/Ag membranes. It is therefore possible to obtain hydrogen of the required purity for operating a PEM fuel cell in a single gas purification step. The Pd/Ag membrane works in the temperature range 533–623 K. The attainable hydrogen yield ranges between 85 and 95%. However, the Pd/Ag membranes commercially available today have the disadvantage that their permeation rate is very low. Starting again from the standard fuel cell with a fuel cell power of 50 kWel , a membrane area of 7.3 m2 is needed. An important approach to cost reduction is the membrane thickness δ. By reducing the membrane thickness by half, material expenditure is halved with the same membrane area, whereas permeation doubles. For the separation of the same amount of hydrogen, the material price is thus reduced to one-quarter. Recent developments move towards coating porous ceramic membranes [104, 169], References see page 3076

3070

13.20 Fuel Cell Related Catalysis

metallic membranes [171] or metallic intermediate substrates with separating layers [170] of Pd or Pd/Ag alloys. Several factors currently prevent the wide application of ceramic-supported Pd/Ag membranes. During the activation process, cracks are often formed, leading to undesired leakage flows of feed gas. Moreover, carbon monoxide and methanol can reduce the permeation rate through the membrane. The solution behavior of hydrogen in the metal lattice again requires flushing with inert gas at lower temperatures. An excellent overview of the variety and the progress of manufacturing technologies was given by Uemiya [172]. B Preferential CO Oxidation Preferential oxidation is the process most often used for CO fine cleaning. Strongly exothermic, preferential oxidation into CO2 proceeds on condition that pure oxygen is supplied as the oxidant according to the following stoichiometric relation:

CO + 1/2O2 −−−→ CO2
HR(T = 473 K) = −283.6 kJ mol−1

(36)

If air is used as the oxidant and ideally assuming that this air is composed of 79% N2 and 21% O2 , N2 as an inert gas must be taken into account in fluid dynamics and in the thermal balance (in stoichiometry): CO + 1/2O2 + 79/42N2 −−−→ CO2 + 79/42N2 ,
HR(T = 473 K) = −283.6 kJ mol−1

(37)

The essential side reaction consists of the oxidation of hydrogen – as the main constituent in the fuel gas – to water: H2 + 1/2O2 −−−→ H2 O
HR(T = 473 K) = −243.5 kJ mol−1

(38)

In this case, too, a catalyst of high activity and selectivity for CO oxidation is required to suppress the undesirable H2 oxidation. This is made difficult by the very high H2 content in the fuel gas, amounting to about 67% (vol.), but it is facilitated by the fact that CO preferentially adsorbs on the catalyst surface. The oxidant supply requires, on the one hand, additional process and regulation expenditure, but permits, on the other, control of the reaction by restricting the O2 educt. Candidate catalysts are above all noble metals such as Pt, Pd, Rh or Ru, but Cu/Co, Cu/Cr and Ni/Co/Fe mixtures and also Ag, Cr, Fe and Mn are additionally being studied as catalyst materials [173]. Depending on the catalyst, not only H2 but also the CO2 contained in the fuel gas can have an influence on the reaction [174, 175]. Considering the product gas composition of a methanol

steam reformer and the product gas composition of a water gas shift reactor behind an autothermal reformer for hydrocarbons, a CO content of 1–2.5% must be preferentially oxidized. For PEFC applications, a final concentration in the range 10–100 ppm CO must be achieved. The preferential oxidation will be performed in isothermal or adiabatic reactors with intercooling devices in a single or more stage design [177, 178]. The catalyst can be divided into precious and non-precious metals. As precious metals Pt [179–196], Ru [197–201], Pt−Ru [202, 203] and Au [204–209] play a major role. The metal loadings are known from patents and papers in the open literature, and vary from 0.5 to 10%. The GHSVs vary between 5000 h−1 [197] and 500 000 h−1 [188]. Considering Eqs. (37) and (38), the selectivity for CO oxidation should be optimized. The selectivity for CO oxidation varies between 30 and 60% and during some single experiments values of 99% could be achieved [186, 214]. The optimum operating temperature amounts to 423–473 K for Pt, Ru and 323–353 K for Au catalysts. Promotion by Ce allows a higher activity [194]. The presence of steam increases conversion and selectivity [179]. Han et al. [200] reported on the disadvantage of Ru catalysts, i.e. their methanation activity. In addition, a low selectivity of 33% leads to a high hydrogen conversion of 17.5% and hot spots [201]. Typical supports for precious metals are Al2 O3 [179, 192] and zeolites [186, 188]. Liu et al. stated that TiO2 is a better support [73]. Some special compositions were developed by a few groups. A potassium-promoted Rh/USY-zeolite catalyst was investigated by Tanaka et al. [211]. Zhang ¨ et al. [212] reported on a Pt/Au catalyst on ZnO. Ozdemir et al. [187] applied Pt-SnO2 /Al2 O3 catalysts. The investigations of Choi and Stenger [189] give an explanation of the apparently inconsistent effects of steam during preferential CO oxidation. The addition of steam had a positive effect at temperatures below 493 K. At higher temperatures this effect remains for concentrations of 15 mol% H2 O with respect to conversion and selectivity. Low concentrations of 7 mol% H2 O, however, result in a poor performance. Choi and Stenger introduced the reverse water gas shift reaction in their models and found a good correlation with their observations. Liu et al. [192] reported on a Pt/γ -Al2 O3 catalyst which was promoted by FeOx . The addition of FeOx creates a non-competitive dual site adsorption enhancing the CO activity for the promoted catalyst. Moreover, active oxygen can be provided by FeOx . FeOx is located on or immediately adjacent to the Pt particles. Related to the CO adsorption on Pt, the CO adsorption on FeOx is weak. FeOx reduces the adsorption surface for CO and shows a strong interaction with the metallic Pt particles, leading to a stretching effect of the CO bonding during CO adsorption.

13.20.1 Fuel Processors

Kahlich et al. [191] preferred a Pt catalyst for the first stage and an Au catalyst for the second stage of their PROX reactor design studies. The advantage of the Au catalyst is the low activity for the reverse water gas shift reaction, the drawback the high necessary Au loading of 3.15 wt.% [180]. The optimal particle size for the Au catalyst is 5–10 nm [205, 207]. Also MnOx as promoter could improve the activity relating to a regular Au/Al2 O3 catalyst [208]. In addition to precious metal catalysts, different metal oxides have been examined with regard to their catalytic activity for preferential CO oxidation [213–220]. Teng et al. [213] investigated NiO, MnO, Mn3 O4 , Mn2 O3 , MnO2 , Cr2 O3 , Fe2 O3 , CuFe2 O4 , NiFe2 O4 , BaFe12 O19 and ZnFe2 O4 . The best results were achieved with CoO, MnOOH and Co3 O4 . The most commonly used nonprecious catalyst is a CuO−CeO2 mixed oxide. The main challenges of a PROX system are the development of an active and selective catalyst, effective process control and a good reactor design. Different concepts have been published [58, 178, 221–224]. Lee et al. [178] show that the CO concentration can be lowered from 1.8% down to 20 ppm during start-up within 3 min in the 10-kWe class. During load changes a slight increase of the CO outlet concentration could be observed. Recupero et al. [222] observed strong CO peaks during load changes which should be smoothed by a subsequent adsorption process [223]. Ahluwalia et al. [224] proposed a two-stage adiabatic reactor design with cooling stages. The reaction starts at 373 K and the GHSV amounts to about 30 000 h−1 . They emphasized a third stage at CO inlet concentrations higher than 1.75%. Finally, the PROX process is rather complex and should be avoided by an improved CO tolerance of new fuel cell types. Besides, it shows interesting aspects for heterogeneous catalysis. Kinetic models for the PROX reaction were published by, for example, Kahlich et al. [191] and Choi and Stenger [189]. C Selective CO Methanation Like CO shift conversion, CO methanation can be performed with the aid of a special catalyst without supplying additional components because the required CO and H2 educts are already contained in the fuel gas. In contrast to CO conversion by the water gas shift reaction, however, selective CO methanation does not proceed by forming hydrogen, but by consuming hydrogen. CO hydrogenation into CH4 is described by the following stoichiometric relation:

−

CO + 3H2 −  −− − − CH4 + H2 O
HR(T = 473 K) = −214.1 kJ mol−1

(39)

3071

The reaction is highly exothermic so that temperature control is of particular significance. A major undesirable side reaction is the slightly less exothermic CO2 hydrogenation into methane according to the stoichiometric relation −

CO2 + 4H2 −  −− − − CH4 + 2H2 O
HR(T = 473 K) = −174.7 kJ mol−1

(40)

Considering thermodynamic equilibrium, CO methanation is favored in comparison with CO2 methanation. Mills and Steffgen [225] and Allen and Yen [226] reported that CO2 methanation could not be observed in the presence of CO. Cambell et al. [155] limit this behavior to CO contents of 200–300 ppm in the syngas. In addition, the well-known shift reaction [Eq. (4)] has an important role in this reaction system since due to insufficient temperature control the shift equilibrium is shifted towards CO. The following equation is a combination of Eqs. (4) and (39): −−

2CO + 2H2  −− − − CH4 + CO2
HR(T = 473 K) = −253.5 kJ mol−1

(41)

In contrast to the use of methanation in industrial process technology [155, 225–229], e.g. in ammonia synthesis [229], the special challenge of gas purification by methanation in fuel gas production for PEFCs is to hydrogenate CO exclusively into CH4 without consuming additional hydrogen due to CO2 hydrogenation. The CO2 content of the product gas from methanol steam reforming, however, amounts to about 20% (vol.) and is thus one order of magnitude higher than the CO content, so that a correspondingly CO-selective catalyst is required. The lower performance of the fuel cell induced by the hydrogen consumption of methanation, however, can be neglected in comparison with the performance loss due to CO in the fuel gas, if sufficient selectivity for CO methanation is provided. Suitable catalysts can be found, in particular, among the Group VIII transition metals and the noble metals. According to Pichler [230], Ru, Ni, Co, Fe and Mo can be specified as significant in order of decreasing activity. In industrial applications, nickel is used rather than ruthenium on account of the price situation and its availability. Nickel is comparatively cheap, sufficiently active for large surfaces and the most selective metal with respect to methane formation among all those specified. According to Vannice [231], Ni/Al2 O3 catalysts are about 85% less active, but about 35% more selective than Ru/Al2 O3 catalysts. Nickel is rapidly poisoned by sulfurcontaining components [232]. Moreover, apart from the formation of solid carbon and nickel carbide, there is the References see page 3076

3072

13.20 Fuel Cell Related Catalysis

danger of nickel tetracarbonyl formation, which, however, can be suppressed by suitable process control and is neglected in the selective CO methanation described here. Depending on their application, commercial nickel catalysts consist of 25–77% (mass) nickel, finely dispersed on substrate materials such as aluminum oxide or kieselguhr [225]. Further information can be found in Refs. [233–235]. Wagner and Takeda [236] reported a new catalyst especially for fuel cell applications. Catalytic Combustion of Fuel Cell Tail Gases In fuel cell systems, in general, different residual gas streams occur which still contain burnable components. These are in most cases catalytically converted into water and carbon dioxide largely free of emissions. The heat that is generated in these combustion reactions can be used for pre-warming educts, for evaporation and for the provision of reaction heat via the process heat source. Considering fuel cell systems, two different principles of combustion are preferred, i.e. catalytic combustion for low-temperature fuel cell systems and flame burners for high-temperature fuel cell systems. This difference is a result of the high entrance temperature of the afterburner behind an SOFC leading to homogeneous gas-phase reactions. At the high combustion temperatures involved, intermediate products with catalytic properties arise during conversion. These gaseous radicals accelerate conversion during the combustion reactions in subsequent chain reactions or chain branching. During combustion processes with air, NO (nitrogen oxide) and NO2 (nitrogen dioxide) are formed from the nitrogen of the air at high temperatures. Pronounced NOx formation only occurs at ca. 1300–1800 K and doubles with every temperature increase of 40 K. There are also developments to achieve low NOx emissions using flame burners. However, this is achieved at the expense of relatively poor flame stability and elevated carbon monoxide emissions. In the case of flame burners, therefore, a minimum of NO and NOx emissions must be balanced. Catalytic combustion provides important advantages over flame combustion with regard to low-emission conversion and high efficiencies. These are: practically no formation of NOx due to low reaction temperatures, low CO and HC emissions, high efficiency at high fuel conversion of over 99.9%, good stability of conversion with varying gas supply, large variation range of the air ratio and low operating temperatures. In the last three decades, the development of catalytic combustion concepts has been pursued with particular intensity. However, combustion in fuel cell systems must be simultaneously capable of both heat supply by complete conversion of energy carriers such as methanol, hydrogen or propane and also 13.20.1.8.3

Fig. 21 Performance chart of total methane oxidation in microstructure. Data taken from Pasel et al. [257].

low-emission disposal of residual gases containing much lower concentrations of burnable components. Pasel et al. [257] reported on the catalytic combustion of methane in microstructures coated with commercial catalysts. Figure 21 shows the methane conversion as a function of temperature. The requirement of low emissions leads to high conversions exceeding 99%. Therefore, temperatures should be higher than 723 K for Pd and 773 K for Pt catalysts. It is well known from the literature [258–268] that Pd catalysts are much more applicable for the catalytic combustion of methane than Pt-based catalysts. Finally, methane conversion requires a longer contact time with the catalyst compared with hydrogen or carbon monoxide. Therefore, the size and the operating conditions of a catalytic burner in a PEFC fuel cell system must be determined regarding the slowest combustion process, i.e. methane conversion. A distinction can be made between four different concepts for catalytic combustion: the self-limiting catalytic burner [237], the diffusion burner [238], the catalytic hightemperature burner [97] and the premixed catalytic radiant burner [240, 241]. Oxidation reactions are catalyzed most strongly by metals of the platinum group, which are more or less expensive. In the search for cheaper catalysts, Hermann [242] investigated mixtures of noble metals with metal oxides ([Sr0.4 La0.6 ][Co0.9 Pt0.1 ]O3 on LaAlO3 , La2 O3 /Cr2 O3 /Pt), zeolite with noble metal, perovskite (La0.6 Sr0.4 MnO3 ), binary mixtures of metal oxides (Cr2 O3 /CO3 O4 ) and pure metal oxides (Cr2 O3 ). Finally, these catalysts cannot reach the activities of noble metals. In order to achieve larger surface areas with smaller material quantities, the catalyst

13.20.1 Fuel Processors

Steam Diesel Air

Reformer

Steam Tail gas

H2S Trap

Air Gas cleaning HTS

LTS

PROX

Air HeatBurner Exch.

PEFC

Air

Water

Process diagram of a fuel cell system including H2 production by autothermal reforming of diesel fuel. Fig. 22

materials are finely dispersed on substrates. The substrates consist of activated carbon, kieselguhr, aluminum and aluminum oxide, some of them with large inner surfaces and high porosities. Component and System Design In order to design a fuel processor for a fuel cell, the different components must be combined in an efficient system. The targets of such a system depend on the application and are set by end users, origin equipment manufacturers (OEMs), industrial companies and government organizations. For example, automotive propulsion systems demand a power density of about 650 We L−1 and 650 We kg−1 defined as a target value in 2010 of the DOE based on hydrogen excluding the hydrogen storage system; see status 2003 and 2006 [269, 270]. Considering gasoline with 30 ppm sulfur as fuel, the targeted power density of a 50-kWe system is somewhat lower at 325 We L−1 . The fuel processor for such a system must gain a power density of 800 We L−1 . As mentioned before, investigations on propulsion systems with gasoline as fuel were stopped in 2004 [110]. Auxiliary power units demand lower power densities of about 150 We L−1 [110, 269], leading to a fuel processor system with a power density of 500 We L−1 [111]. The DOE revised the targets in 2006 for APUs in the power range 3–30 kWe , resulting in a power density of 100 We L−1 by 2010 [270]. From these values, it is possible to calculate a fraction concerning the size of reactors, i.e. reformer, desulfurization unit, WGS reactor, PROX reactor and catalytic burner, heat exchanger, evaporators, compressor, pumps and finally electronics and insulation [244]. To achieve a value of 700 We L−1 (DOE target 2005), a power density of about 9 kWe L−1 can be postulated for the reformer. By integration measures, all the required heat exchangers, i.e. water vaporization, steam superheating, air and fuel preheating, can be taken into account and lead to a value of about 3.6 kWe L−1 [243]. Present data show a power density of 1.8 kWe L−1 for a 5-kWe autothermal diesel reformer with integrated heat exchangers [163]. 13.20.1.9

3073

How could such a power unit be further improved? The GHSV amounts to 30 000 h−1 , leading to a specific hydrogen production of 18 kWe L−1 [243], which is comparable to gasoline at 20 kWe L−1 [111]. By recalculating the demanded volumes it can be found that the catalyst needs only 10–20% of the total reactor volume. Part of the total volume is required by the integrated heat exchanger (10–20%), by the mixing zone (5–15%) and finally by the steel of the pressure vessel and the inner insulation (50–75%). Under the restrictions of the technical rules for apparatus design and construction, there is still potential for improvements. The set of targets for each specified component depends strongly on the system chosen. A SOFC system combined with a POX process for gasoline is extremely reduced with respect to components. The POX process requires only vaporized gasoline and preheated air. The afterburner behind the electrochemical cell uses the residual cathode air. An advantage of the SOFC is the possibility of converting CO electrochemically in the cell. In contrast to the SOFC system, the PEFC system demands a lot of components at different temperature levels. The ATR process needs three different flows including the separate evaporation of gasoline and water. The flows behind the PEFC are still separated to guarantee proper functioning of the catalytic burner. The number of heat exchangers for the PEFC system is much higher in comparison with the SOFC system. Although all the components of the SOFC system work at a uniformly high temperature level with a reduced effort for heat transfer between them, the high operating temperature itself is a disadvantage concerning start-up time and dynamic performance. 13.20.1.9.1 Heat Exchanger Networks To elucidate the size of the heat exchanger network, the pinch-point methodology of Linnhoff [246] can be applied. Figure 23 shows a typical pinch-point diagram for a PEFC process based on autothermal reforming of diesel. In a pinchpoint diagram, the enthalpy flow of the complete system is sketched as function of temperature. The heat sources and the heat sinks are combined in two separate graphs without recognizing the real apparatus and system design but leading to one giant heat exchanger. Without analyzing each heat exchanger, a rough number for the required area of the heat exchanger net work can be estimated using 
H2 ˙ dQ AWT = (42) kdϑ
H1

A higher temperature gradient leads to smaller heat exchangers. How can the temperature gradient be References see page 3076

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13.20 Fuel Cell Related Catalysis

PEFC: Q °BZ + Q °cond.

1400

Gas treatment Educt pre-heating

1000

ATR process

1200

∆Tln,min = 35 K

T /K

800 600

100 ∆Tln / K

1000

Water- and heat management

10

400 200 1

0 0

2

4

6

8

10 ∆H / kW

12

14

16

18

20

Pinch-point diagram showing the heat exchanger network for APU fuel cell systems based on autothermal reforming of diesel for a 5-kWe PEFC.

Fig. 23

increased? Fuel utilization in the fuel cell plays a crucial role for the heat management of the system. Mass transport limitation in the fuel cell does not allow complete utilization of hydrogen. The residual fuel is converted in the catalytic burner delivering heat to the system. In the case of methanol steam reforming, a defined utilization must be fulfilled to close the heat balance of the system. According to the second law of thermodynamics, a somewhat lower fuel utilization is demanded. Disregard of the second law of thermodynamics can be observed in the pinch-point diagram if the sum of all hot enthalpy flows indicates lower temperatures than the so-called cold curve of the heat sinks. The pinch-point is given if the hot curve hits the cold curve in on defined state, i.e. the pinch-point. During the optimization process, one must decide on the targeted performance data, i.e. optimized efficiency vs. lowest heat exchanger size or a suitable mixture. Applying this method to ATR processes, the heat input is given by the addition of air to the reformer and the fuel utilization. Most parameters are fixed by catalytic constraints, e.g. the O2 /C ratio by complete fuel conversion in the reformer and the operation temperature of the shift stages by the kinetic data for the catalyst and the chemical equilibrium. Apart from the theoretical approach, the engineering aspect plays a major role in the design and construction of each apparatus. Furthermore, experimental results on a pilot scale provide support for the component design. One must consider that the weakest component in the system determines the power density of the whole system. Innovative Reactor Concepts The apparatus design determines the size, mass and internal

13.20.1.9.2

construction elements of chemical reactors and heat exchangers. The power density can be improved by integration measures combining different functions in one apparatus, e.g. microreactors [251–253] and membrane reactors [271, 272]. Unfortunately, various authors [247–254] have reported that the projected advantages of this technology have not yet been fully achieved. Mass transfer limitations of chemical reactions can be significantly reduced at channel dimensions of less than 100 µm. Narrow channels lead to high pressure drops, which are not acceptable from the system perspective. Additionally, the required mass of catalyst must be coated on the small passages, which is a rather complex issue and sometimes the catalyst mass is too high. Other concepts use turbulence inserts to improve catalyst efficiency [255]. Microstructures are an excellent tool for evaporation processes [251, 256]. Finally, it must be remembered that the small channels can be blocked by dust particles and products from side reactions of the fuel, such as pyrolysis of higher hydrocarbons. 13.20.1.9.3 Dynamic System Modeling Beside the size of components, the dynamic performance of the fuel cell systems is of special interest. During a dynamic simulation of the complete system, further improvements concerning the start-up strategy can be verified. Dynamic simulations of fuel cells must include the performance of the complete system with all required reactors and heat exchangers. Such models allow the development of improved start-up strategies [273–275]. Sommer et al. [276] analyzed the response behavior of a fuel processor system consisting of ATR and WGS on a load change from 10 to 90%. The dynamic performance

13.20.1 Fuel Processors

of the PROX reactor was not considered. Because of the lack of water in one heat exchanger during the first 20 s, the reformate cannot be cooled as much as planned in the design stage. Consequently, a temperature peak of about 50 K in the high-temperature shift stage occurs, causing a CO peak of about 20%. The low-temperature shift stage is not able to decrease this concentration significantly. Such a high CO concentration cannot be treated in the PROX reactor. Finally, the calculated response time of the fuel processor system amounts to 1 s. Porˇs [126] calculated a mean residence time between 70 and 130 ms for an ATR including evaporation and mixing zones. Taken these data and other technical information into account, one can calculate by a Matlab/Simulink model a response time of 1 s to achieve 95% of the proposed load for the line-up of ATR, WGS and PROX sections. These data are in good accord with our own calculations using the technical data given by Severin et al. [277] for all inner apparatus volumes of a PEFC system, resulting in 1.5 s. The calculated response times show that even for a highly dynamic reformer the system will be too slow to work as a propulsion system without any energy storage. In order to estimate possible start-up times and to develop start-up strategies, the thermal mass of all apparatus must be defined. In this context, the design of the heat exchanger network plays an important role. Montel [273] proposed a start-up strategy leading to start-up times of about 6 min for the complete PEFC system based on diesel fuel as energy carrier under the condition that the ATR is preheated. The CO concentration could be lowered to an acceptable level for the fuel cell after 100 s. During the first few minutes the shift reactor did not have the full performance due to low temperatures. At the beginning, a catalytic burner produces a hot tail gas that heated the ATR to 673 K. Experiments with electrical devices for preheating the catalyst [44, 277] offer a heating time between 10 and 15 min with the potential to reduce this time to 5 min. Springmann et al. [274] postulated preheating by an electrically heated catalyst to 673 K within 60 s. Reforming should be started in POX operation mode for 60 s. The reformer start-up could be performed in 15 s while the water gas shift reactor required 6 min to reach the final state. Springmann et al. [274] discussed several measures to shorten startup times to 3 min while the shift section ignited after 60 s. Possible measures are the oxidation of reformate in the high-temperature shift reactor section, electrical preheating of the shift reactor, variation of air ratios and the availability of steam. Outlook Some system developments have been very successful, such as the PEFC stack development for 13.20.1.10

3075

DaimlerChrysler’s NeCar III using methanol by reforming and also the gasoline fuel processor development by General Motors [278]. At present, the automotive industry is concentrating on hydrogen for propulsion tasks in the 75-kWe class and gasoline or diesel for APUs in the 5-kWe class. Further improvements must be made to achieve market entry. Hybridization is a possibility for qualifying fuel cells and fuel processor systems in mobile applications, leading to several options for electric and mechanical drive systems. New applications are being investigated by industry such as APUs for trucks, ships and aircraft or domestic heating systems. How can heterogeneous catalysis help to achieve these ambitious goals? Reliability is an important issue for fuel processors. From the chemical point of view, catalyst degradation must be fully understood. Important analysis methods are long-term test under stationary conditions and in dynamic operation mode and a subsequent investigation of the catalyst surface with XPS and EDX. A correlation between catalytic activity and the XPS measurements for fresh and aged catalysts should be drawn. In order to analyze the surface phenomena, in situ Raman spectroscopy is a promising method. Unfortunately, fewer investigations are published on autothermal reforming for hydrocarbons compared with methanol steam reforming. The loss of active sites results from different overlapping effects, which include poisoning by different substances such as sulfur and chlorine, ash and carbon deposition, sintering effects and the direct loss of active catalyst material. Furthermore, catalyst activity must be improved at lower temperatures, especially for the water gas shift reaction. The most difficult task is the preferential CO oxidation catalyst. Most catalysts have a narrow operating window, low activity, high sensitivity to unconverted fuel and unstable operation at load changes. Today, catalyst coatings for microchannels, membranes and other construction elements are limited to a defined structural size. To overcome mass and heat transfer limitations, further developments must be made. Acknowledgments

The author thanks the fuel processing group at the Forschungszentrum J¨ulich, Institute of Energy Research – Fuel Cells (IEF-3), for their excellent cooperation and the important results obtained during the last 10 years of fuel cell research, and R. C. Samsun, Z. Porˇs and J. Pasel for valuable discussions and technical assistance in preparing the manuscript. References see page 3076

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13.20.2 Fuel Cells

13.20.2

Fuel Cells .. Hubert A. Gasteiger and Jurgen Garche∗

3081

and cathodic reduction (i.e. O2 ) reactions occurring in PEMFCs and DMFCs; where appropriate, findings will be compared with the phosphoric acid fuel cell literature. Working Principles of Fuel Cells Fuel cells are electrochemical power sources (ECPS), similar to primary and secondary batteries (also referred to as accumulators). Their working principle is shown schematically in Fig. 2. These electrochemical cells consist of two electrodes separated by an ionically conducting electrolyte. Conversion of chemical energy into electrical energy occurs at the interface between the electronically conducting electrodes and the ionically conducting electrolyte. In principal, the electrodes can operate in two modes, either transforming electrical into chemical energy (electrolysis cell mode): 13.20.2.2

Introduction Fuel cells directly convert chemical energy into electrical energy, without the intermediate formation of heat and mechanical energy as in conventional energy conversion devices based on the Carnot cycle. The direct electrochemical energy conversions of H2 at the anode and O2 at the cathode occur spatially separated on two electrodes connected by electrolyte and, in principle, allow for higher energy conversion efficiencies as shown in Fig. 1. The largest efficiency differences between Carnotbased and fuel cell-based energy conversion are found in the low power range of CO2 + H2O + 2e−

Fig. 5

O2,CO2 1/2 O2 + CO2 + 2e− = > CO32−

Schematic view and working principles of a MCFC.

Lix Ni1−x O (0.022 < x < 0.04) in the electrolyte. A current problem is NiO corrosion, leading to the formation of metallic nickel via reduction of Ni2+ by H2 ; ultimately this may lead to electronic short circuiting via Ni dendrites growing into the electrolyte: NiO + CO2 −−−→ Ni2+ + CO2− 3 H2 + Ni2+ + CO2− 3 −−−→ Ni + CO2 + H2 O

(8) (9)

The MCFC is used in stationary applications as a combined heat-and-power station for electric power ranges of >200 kW. The high temperature exhaust heat (550–600 ◦ C) can be utilized for process heat, absorptions chiller, drying and sterilization and can even drive a steam turbine for electricity generation. MCFCs are developed/produced by CFC Solutions (formerly MTU) (Germany), FC Energy (USA), Ansaldo (Italy) and several Japanese companies (Mitsubishi, Fuji, IHI, Hitachi). An overview of MCFCs is given in Refs. [20, 21]. 13.20.2.3.5

Solid Oxide Fuel Cell (SOFC)

Electrolyte: Electrodes:

Operating temperatures: Reactants: Cell power density: Cell electrical efficiency:

solid-state ZrO2 −Y2 O3 membrane; conducting ion, O2− anode, Ni cermet; cathode, doped perovskites (LaMnO3 , LaSrMnO3 , LaCoO3 ) 800–1000 ◦ C CH4 , CO, H2 ; oxygen, air ∼0.2–0.3 W cm−2 60–65% (with respect to lower heating value)

The commonly used oxide ion conducting electrolyte is ZrO2 with about 10% Y2 O3 to form oxygen vacancies, referred to as yttrium-stabilized zirconia (YSZ). YSZ has a conductivity of about 0.1 S cm−1 at 1000 ◦ C [22]. Conventionally, the YSZ thickness is ∼100–200 µm, but in order to extend the operating temperatures to below

1000 ◦ C, design concepts with 10–20 µm thin YSZ are being developed. Alternative electrolytes with better lowtemperature conductivity are doped CeO2 (Ce0.9 Gd0.1 O2 ) and doped LaGaO3 [(La1−x Srx )(Ga1−y Mgy )O3 ], but their stability and reliability are lower. Their application is envisaged especially for ‘‘low-temperature SOFCs’’ at ∼600–800 ◦ C [23]. As anode, an Ni−ZrO2 cermet (35% N in a porous YSZ matrix) is used. The cathode generally consists of doped lanthanum manganite (La1−x Srx MnO3 , 0.1 < x < 0.15). While about half of the overall SOFC voltage losses are due to limited electrolyte conductivity, the remaining losses mainly occur at the cathode. Therefore, large efforts are being made to develop alternative cathode materials, as Ce- and Ca-doped LaMnO3 , solid solutions of LaMnO2 with LaCoO3 (LaCo1−x Mnx O3 , x ≈ 0.2) and high-temperature superconductors such as La2 (BaSr)x CuO4 and YBa2 Cu3 O7−x . Normally, one sub-cell element (anode, cathode or electrolyte) is chosen as the substrate to provide mechanical stability and to serve as carrier of the subsequently deposited remaining sub-cell components. Two main types of cell design exist: tubular and planar. In the tubular concept, the cathode is formed by extrusion as tube (up to 1.5 m length) with one opening. Then the electrolyte and the cell interconnects are deposited. The anode is finally formed on the electrolyte layer by slurry deposition. One advantage of the tubular type is that it eliminates the need for reactant manifolding at the high operating temperature. The disadvantage is a relatively long current pass around the tube circumferences to the interconnects, significantly limiting cell power densities. The planar or bipolar concept used in PEMFCs and PAFCs would permit a simple serial electrical connection with low ohmic resistance between the cells, but gas leaks and mechanical problems with large electrode/electrolyte areas are impediments to using this approach for SOFCs. Natural gas is used as fuel for SOFCs, being reformed either internally or externally. The high SOFC operating temperature of 800–1000 ◦ C leads to material stability issues, particularly with regard to the thermally induced stresses during startup and cooling. Similar to PAFCs, the SOFC lifetime strongly depends on the number of startup and cooling cycles. Seals are the most critical components and currently high-cost materials are used to minimize these problems. In recent years, the development of lowtemperature SOFCs operating at ∼600–800 ◦ C has been investigated to overcome lifetime and cost problems. The research focus toward developing low-temperature SOFCs is electrolytes and cathode materials. SOFCs are primarily used for combined heat-andpower (CHP) stations at >100 kW. The high-temperature exhaust heat on the order to 800 ◦ C can be utilized for highquality process heat or electricity generation using a steam

13.20.2 Fuel Cells

turbine. First demonstrations have shown, that a total system electrical efficiency of 60–65% could be reached. A further envisaged application are low-power (∼5 kW) auxiliary power units (APUs) for electricity generation on-board vehicles, but thermal cycling durability remains a critical issue. SOFCs for CHP are being developed by, for example, Siemens–Westinghouse (Germany/USA), Ceramic Fuel Cells (Australia), Ztek (USA), Mitsubishi (Japan) and Global Thermal Electric (Canada). SOFC APU developers are, for example, Delphi (USA), Webasto (Germany) and BMW (Germany). An overview on SOFCs is given in Ref. [24]. Fuels for Fuel Cells From the electrochemical point of view, hydrogen is the ideal fuel. In addition to hydrogen, methanol also is used directly as fuel, but the achievable power densities and cell efficiencies are much lower due to poor electrode kinetics. Other organic fuels such as ethanol, glycol and ethers (dimethyl ether, dimethoxymethane) have been tested, but their reactivity with currently known anode catalysts is too low for practical applications and C−C bonds are generally not broken, thereby reducing the available energy content of the fuel (a possible exception is discussed in Section 13.20.2.7.2). Although hydrazine was used with reasonable success in the 1960s in direct-fuel fuel cells, its development was stopped due to its carcinogenicity. Daihatsu Motor Co. however has announced in 2007 that it has developed a hydrazine driven fuel cell. Currently, hydrogen is stored either compressed at pressures up to 70 MPa or liquefied. Alternative H2 storage materials, for example metallic, covalent and ionic hydride 13.20.2.4

3087

materials, zeolites, carbon nanotubes and metal–organic frameworks, are under development [25–29]. Figure 6 shows the state-of-the-art values of different storage technologies in comparison with US Department of Energy transportation targets [30]. Since reaching these targets is challenging, hydrocarbon reformers are being developed in parallel for transportation applications, but with lower R&D intensity. For stationary applications, pure hydrogen is not available in decentralized applications, necessitating H2 production from hydrocarbons via reforming (e.g. from natural gas, propane/butane, gasoline, diesel, methanol and ethanol). For the conversion of these chemical hydrogen carriers to an H2 -rich reformate, different processes are used: steam reforming, autothermal reforming and partial oxidation. The basic equations for the most commonly used processes of steam reforming and autothermal reforming are shown below. Steam reforming process:

− CO + 3H2 CH4 + H2 O ← −−− −− →


R H(25 ◦ C) = +206 kJ mol−1

(10)

← − CO2 + H2 CO + H2 O − −− −− →


R H(25 ◦ C) = −41 kJ mol−1

(11)

← − CO2 + 4H2 CH4 + 2H2 O − −− −− →


R H(25 ◦ C) = +165 kJ mol−1

(12)

References see page 3117

100 Current cost estimates (based on 500 000 units)

Volumetric capacity / g L−1

80

10 000 psi 5 000 psi Liquid H2 Complex hydride Chemical hydride

60

2015 targets

$0 $5 2015 targets 2010 targets

$10 $ / kWh

$15

$20

2010 targets 40

Chemical hydride

Liquid hydrogen 10 000 psi

20

5 000 psi Complex hydride

0 0

2

4

6

8

10

Gravimetric capacity / wt.%

DOE targets and practical hydrogen storage values of different storage technologies and materials [30]. Complex hydride, thermal H2 generation; chemical hydride, H2 generation by reaction with water; costs exclude regeneration/processing.

Fig. 6

3088

13.20 Fuel Cell Related Catalysis

Tab. 1

Advantages and drawbacks of different reforming processes

Process Steam reforming

Partial oxidation

Autothermal reforming

Advantages

Disadvantages

High H2 content in reformate, H2 anode exhaust is used for heating up the steam reformer Fast start, highly dynamic operation Fast start, dynamic operation

Autothermal reforming process: In the autothermal reforming process, the endothermic steam reforming reaction [Eq. (13)] is combined with the exothermic partial oxidation reaction of methane [Eq. (14)] and the water gas shift reaction [Eq. (15)] in order to achieve overall thermal equilibrium in the reactor. Endothermic reaction:

← − CO + 3H2 CH4 + H2 O − −− −− →
R H(25 ◦ C) = +206 kJ mol−1

(13)

Exothermic reactions: ← − CO + 2H2 + (2N2 ) CH4 + 12 O2 + (2N2 ) − −− −− →
R H(25 ◦ C) = −247 kJ mol−1

(14)

← − CO2 + H2 CO + H2 O − −− −− →
R H(25 ◦ C) = −41 kJ mol−1

Low H2 content in reformate, highly exothermal process, H2 anode exhaust could not used H2 anode exhaust cannot be further utilized in the process

still too high for PEMFCs, which require an additional selective CO oxidation (referred to as SELOX or PROX) reactor to reduce the CO content to 0 V vs. RHE and hydrogen evolution at secondary > tertiary. Basic catalysts cover a wide range that includes basic oxides, alkali-metal loaded zeolites, organometallic compounds such as alkylsodium and metallic sodium or potassium. Synthesis of Styrene from Toluene with Methanol Toluene undergoes side-chain alkylation with methanol in the presence of Cs-exchanged zeolite X to produce ethylbenzene and styrene [58]. The reaction involves first a base-catalyzed dehydrogenation of methanol to formaldehyde, which undergoes an aldol-type condensation with toluene to form styrene. Ethylbenzene is formed by hydrogenation of styrene. The incorporation of excess alkali metal ions produces alkali metal oxide/hydroxide in the zeolite pores and increases the rate [59]. The selectivity of the reaction is further increased by promotion with borate [60]. 14.3.2.2.2

14.3.2.2.3 Isobutylbenzene This compound cannot be produced by acid-catalyzed alkylation. Reaction of benzene with isobutyl chloride invariably yields tertbutylbenzene. By contrast, the reaction of toluene with propylene over base catalysts yields isobutylbenzene [56, 61], with minor amounts of n-butylbenzene. The reaction

occurs via an anion chain reaction (Fig. 4). The addition step (a), which produces a more stable primary carbanion, is much faster than the competing step (b), which results in a less stable secondary carbanion. Isobutylbenzene is an intermediate in the synthesis of the analgesic ibuprofen. It is commercially produced via base catalysis by Albemarle and by Phillips in the USA and by Synthetic Chemicals, a Shell subsidiary, in the UK. 14.3.2.2.4 tert-Amylbenzene The base-catalyzed reaction of isopropylbenzene with ethylene produces tertamylbenzene (α-ethylcumene). The reaction occurs with 99% conversion and with 99% selectivity. The catalyst is a solid superbase, KOH/K/γ -alumina [62]. The reaction occurs at room temperature and is commercially operated by Sumitomo Chemical.

Butenylation of Alkylaromatics With base catalysts, alkylbenzenes react readily with butadiene [63] and other dienes. The reaction of a xylene with butadiene has been developed by Amoco Chemical into an industrial process for the production of an intermediate for the synthesis of 2,6-dimethylnaphthalene (2,6-DMN). o-Xylene is reacted with butadiene in the presence of sodium on potassium carbonate to yield l-methyl2-n-pentenylbenzene (Fig. 5), which is then further converted by phosphoric acid-catalyzed ring closure, dehydrogenation and isomerization to 2,6-DMN.

14.3.2.2.5

Initiation CH3

CH2− Na+

R−Na+

+

RH

CH3

Addition (1a) CH2−

+

CH 2 C CH2−

H2C=CHCH3

Addition (1b)

−

CH2−

+

CH2 C CHCH3

H2C=CHCH3

Proton Transfer CH3

CH3 CH2

Fig. 4

C CH2−

+

CH3

CH2 C CH3

–

+

CH 2

Mechanisms of base-catalyzed side-chain alkylation of toluene with propylene via a carbanion chain reaction.

14.3.2 Alkylation of Benzenes, Phenols and Anilines

+

Base catalyst

Ring closure Dehydrogenation 2,6-dimethylnaphthalene (DMN)

Methyl group isomerization

Base-catalyzed butenylation of o-xylene and subsequent reactions for the synthesis of 2,6-dimethylnaphthalene (2,6-DMN).

Fig. 5

Alkylphenols Alkylated phenols are important intermediates in the chemical industry for a great diversity of products, including pharmaceuticals, dyes and antioxidants. The major markets are non-ionic detergents, phenolic resins, polymer additives and agrochemicals [64]. The variety of alkylated phenols and synthesis methods are covered by several excellent reviews [64, 65]. Therefore, only a few general and novel aspects of this field will be discussed. 14.3.2.3

14.3.2.3.1 Methylphenols Cresol (methylphenol) and xylenol (dimethylphenol) are produced by the reaction of phenol with methanol using acid and basic catalysts. The initial product is probably anisole (methyl phenyl ether), which can rearrange intramolecularly to form o-cresol, which can undergo a second ortho alkylation to 2,6-xylenol. In the absence of acid catalysts, these ortho products are obtained in high purity. In the commercial process, phenol reacts with methanol in a fixed-bed flow reactor at 475–600 ◦ C over MgO [66] as catalyst. With an excess of methanol, 2,4,6-trimethylphenol is obtained. An improved phenol methylation process has been introduced by Asahi Chemical Industry [67]. It uses a fluidized-bed reactor with a silica-supported promoted iron–vanadium catalyst. Its high activity allows lower temperature operation (300–350 ◦ C). o-Cresol and 2,6xylenol are obtained in high purity of 99.9%. In the presence of acid catalysts, such as silica–alumina, MFI and USY, o- and p-cresol are formed under kinetically controlled conditions. The xylenols contain over 80% 2,4and 2,5-dimethylphenol. Since the thermodynamically favored meta isomer of cresol is essentially absent, these products are formed by electrophilic substitution in the ortho and para positions, rather than by subsequent isomerization of an initial ortho-alkylated product. At higher temperature, isomerization to m-cresol does occur. 14.3.2.3.2 Higher Alkylphenols Alkylphenols containing 3–12 carbon alkyl groups are obtained by alkylating

3163

phenol with the corresponding alkenes using acid catalysts. Conventional acid catalysts such as sulfuric acid, phosphoric acid, activated montmorillonite clay, boron trifluoride or aluminum chloride are used. In view of the greater nucleophilicity of phenol relative to benzene, reaction conditions are milder, and weak acid catalysts can be used. For example, sulfonated polystyrene ion-exchange resins, which are stable only up to 140 ◦ C, have found increased use as solid catalysts. The alkyl group enters the nucleus ortho and para to the OH group of phenol, according to the rules of electrophilic substitution and subject to minor steric effects in the ortho position. Under mild conditions, etherification of the OH group can occur. Aluminum trisphenoxide, generated in situ by reaction of phenol and aluminum, is also a frequently used catalyst. In contrast to the Brønsted acid catalysts, it causes ortho alkylation selectively. A possible mechanism to explain this special ortho effect has been proposed [68]. Shape-selective zeolites have been proposed to obtain alkylated phenols with altered product distribution. For example, the reaction of phenol with ethylene at 400 ◦ C with MFI catalyst produced ethylphenols with a para:meta:ortho ratio of 17 : 52 : 31. MFI, modified by treatment with tetramethoxysilane and subsequently calcined, gave an ethylphenol distribution of 94 : 5 : 1, i.e. predominantly para [69]. Another example is the synthesis of the difficult to obtain 3-methylphenol (m-cresol) [70]. The reaction proposed starts with the alkylation of toluene with methanol over zeolite MTW. Under optimized conditions, it produces an isopropyltoluene mixture containing 64% of the meta and 5% of the ortho isomer. This mixture is then shape-selectively cracked over a modified MFI, which cracks all the para and some of the ortho isomer, leaving unaffected essentially all of the m-cymene. Conventional oxidation produces m-cresol in high isomeric purity. Alkylanilines The alkylation of aniline by acid catalysts is difficult, due to the high basicity of aniline. For example, attempts to catalyze the reaction of aniline with ethylene with AlCl3 led mostly to resinous materials. However, the reaction can be carried out satisfactorily by using as catalyst aluminum anilide, Al(C6 H5 NH)3 , which is formed in situ by heating of aniline with aluminum powder, with the evolution of hydrogen. At 300–340 ◦ C, 20 MPa and a short reaction time, the ortho substitution product, 2-ethylaniline, is obtained; longer reaction times yield 2,6-diethylaniline [65a]. This ortho selectivity is analogous to the phenol system discussed above. However, the reactivity of alkenes decreases in the order 14.3.2.4

References see page 3166

3164

14.3 Alkylation of Aromatics

ethylene > propylene > isobutene, the reverse of what is observed in the phenol alkylation with Al(OC6 H5 )3 as catalyst. Zeolites as Brønsted acids are effective catalysts for aniline alkylation with methanol [71]. At 275 ◦ C, 100% N -alkylation occurs, yielding N -methylaniline and N ,N dimethylaniline. At 450 ◦ C, C-alkylated toluidine is produced, with only 0.3% N -methylaniline. The toluidines contain mostly the para isomer at short contact times, but reach thermodynamic equilibrium at longer times, with the meta isomer as the major component. The isomerization of the three toluidines has been studied separately with ZSM-5 at 350–500 ◦ C [72]. Alkylation of aniline by alkenes catalyzed by zeolites has also been carried out. For example, with dealuminated MOR or Y-zeolite catalysts, propylene will alkylate aniline at 250 ◦ C to yield 84% o-isopropylaniline, 4% p-alkylate and 15% N -alkylate [73]. Large-pore zeolites, amorphous silica–alumina and alumina have also been studied by Burgoyne and coworkers in the alkylation of aniline by propene and isobutene [74]. Again, the product distribution is a strong function of the reaction temperature. 14.3.3

Alkylation of Binuclear Aromatics

The alkylation of bi- and polynuclear aromatics has not attracted broad interest in the past. However, this field has recently attracted attention, as discussed below. Alkylation of Naphthalene Alkylnaphthalenes are important monomers for the production of advanced aromatic polymer materials and are used in fine chemical synthesis as chemical intermediates. Song and coworkers [75, 76] have reviewed this area. For example, commercially produced vitamin K is prepared via 2-methylnaphthalene as intermediate, and 2,6-dialkylnaphthalene is an important precursor for the production of polyethylene naphthalate (PEN), polybutylene naphthalate (PBN) and liquid crystal polymers [77, 78]. The use of solid catalysts to synthesize alkylnaphthalenes has attracted attention [76, 79–85]. The great difficulty in preparing alkylnaphthalenes with high selectivity is caused by the great number of isomer positions available. Compared with the dialkylbenzenes, for which three isomers (ortho, para and meta) are possible, there are 10 possible isomers for a dialkylnaphthalene system. Therefore, a catalyst chosen for specific reactions must have special properties in terms of activity and selectivity. Zeolites, with their uniform pore sizes, have been found to provide such catalysts. 14.3.3.1

As reported by Cusumano [81], alkylation of naphthalene with propene over zeolite catalysts yields predominantly 2,6-diisopropylnaphthalene, an important precursor to specialty polymers. This pathway is illustrated in Fig. 6. It has long been recognized that electrophilic substitution in the 1-position of naphthalene is kinetically preferred and that the 2-isomer is the more stable. There is strong steric hindrance for bulkier molecules at the 1-position, as a result of antibonding interaction with the peri-hydrogen in the 8-position. This may affect the rate of alkylation and will affect the thermodynamic stability of the compound. For example, it has not been possible to introduce a tert-butyl group at that position. The reported equilibrium data for various alkyl substituents are given in Table 6. While 1-methylnaphthalene is only slightly less stable than the 2-methyl isomer, even the ethyl group experiences a greater steric repulsion by the adjacent peri-hydrogen than from an ortho-methyl in benzene. In the alkylation of naphthalene with conventional Friedel–Crafts catalysts, it is very difficult to obtain good yields of dialkylation products, since further alkylation leading to tri- and tetrasubstituted naphthalene occurs very rapidly. Due to their sterically constrained pore system, zeolite catalysts have been found that limit such polyalkylation, while at the same time showing high selectivity for one or two of the possible 10 isomers, including the most desirable 2,6-disubstituted alkylnaphthalenes. Data for alkylating naphthalene with methanol and 2-propanol are given in Table 7. Methylation occurs selectively at the 1-position when the larger pore MTW (12-ring) is the catalyst (kinetic control). With dealuminated MOR and the medium-pore MFI (10ring), high selectivity to 2-methylnaphthalene can be obtained. The selectivity exceeds that expected from thermodynamic control (Table 6) and must be a result of diffusion control. Dimethylnaphthalene containing 55% of the most desired 2,6-isomer has also been obtained, Equilibrium composition of alkylnaphthalenes. Experimental data with AlCl3 catalyst [77] and equilibrium isomer distributions calculated from thermodynamics [78]

Tab. 6

Alkyl group Methyla Methyla Methyl Ethyl Isopropyl tert-Butyl a Adapted

Temperature/ ◦ C

1-Alkyl/%

2-Alkyl/%

27 227 46 46 0–25 0

35.3 38.2 24.5 9.5 1.5 0

64.7 61.8 75.5 90.5 98.5 100

from Song and Kirby [76].

14.3.3 Alkylation of Binuclear Aromatics

3165

2,6 isomer AlCl3 + 2,7 isomer + tri- and tetrasubstituted products H-Mordenite

CO2H

Specialty polyesters

HO2C CO2H HO OH

Liquid crystal polymers

Specialty polyesters

HO

Fig. 6

Alkylation of naphthalene with propylene: selective production of 2,6-diisopropylbenzene and its oxidation products.

Tab. 7

Isomer distribution in the alkylation of naphthalene with methanol and 2-propanol

Reaction

Methylation

Isopropylation

Catalyst

MTW MOR MFI MFI HY MOR MTW

Temperature/ ◦ C

280 400 400 300 250 250 280

although in low yield [86]. Similar 2,6-isomer selectivities were found in the methylation of 2-methylnaphthalene [82]. Alkylation with 2-propanol and the large-pore HY catalyst produces monoalkylate with a strong preference for the 1-position and a low selectivity for 2,6-dialkylate. Dealuminated mordenite and ZSM-12 alkylate first with a preference for the 2-position and produce the desired 2,6-diisopropylnaphthalene with a high selectivity of up to 82% [76] (Fig. 6). Using propylene as feed leads to similar selectivity [87]. These results are clearly due to the subtleties of shapeselective reactions made possible by zeolite catalysis.

Monoalkylate/% 1-

2-

77 20 13 – 69 23 23

23 80 87 – 31 77 77

2,6-Dialkylate/%

Ref.

– – – 55 30 82 82

[84] [83] [84] [86] [76] [76] [84]

Although 2,6- and 2,7-diispropylnaphthalene appear to have the same critical molecular diameter, computerassisted catalyst screening showed that the 2,6-isomer has a somewhat more linear structure than the 2,7-isomer and therefore can diffuse faster in the zeolite channels [88]. Alkylation of Biphenyl The interest in novel, high-performance polymers has driven the search for new linear bifunctional monomers. One area of research has focused on biphenyl derivatives with substituents at both the 4- and 4 -positions. Using 14.3.3.2

References see page 3166

3166

14.3 Alkylation of Aromatics

Zeolite

+

Catalyst 4,4′-diisopropylbiphenyl HO2C

O2

CO2H

[1,1′-biphenyl]-4,4′-dicarboxylic acid Fig. 7

Shape-selective isopropylation of biphenyl.

Tab. 8

Effect of SiO2 :Al2 O3 molar ratio of mordenite on product distribution in the propylation of biphenyl [89]

SiO2 :Al2 O3

14 38 230 a MIPD,

Conversion/%

49 71 23

Product distributiona /mol%

MIPD isomer distribution/%

DIPD isomer distribution/%

MIPD

DIPD

TIPD

2-

3-

4-

3,3 -

3,4 -

4,4 -

74 54 67

25 42 32

0.5 3.0 0.5

9.2 10.2 2.0

24 28 18

66 62 80

3.9 2.7 0.9

17 13 10

66 72 87

13.8 12.2 2.5

DIPD, TIPD = mono-, di-, triisopropyldiphenyl.

shape-selective zeolites as acid catalysts, it was possible to develop a synthesis for 4,4 -diisopropyldiphenyl (DIPD) by reacting biphenyl with propylene. The alkylation product can be oxidized to (1,1 -diphenyl)-4,4 -dicarboxylic acid, which is a monomer for the production of heat-resistant polymers (Fig. 7). A large number of patents by different companies and some scientific papers reflect the competitive nature of this development [89–97]. In a typical example, biphenyl (50 parts) and propylene are heated with dealuminated mordenite catalyst (1 part) in an autoclave for 20 h at 525 K. 4,4 -DIPD is obtained with 73% selectivity [89]. Data obtained by Schmitz and Song [98] show the effect of different silica:alumina ratios of the mordenite catalyst on the ‘‘product distribution’’ (Table 8). With the highsilica catalyst, the yield of recyclable mono- plus dialkylate exceeds 99%. The isomeric purity of 4,4 -DIPD is 87%. 14.3.4

Conclusions

Alkylation of aromatics is a very rich field, with a wide variety of aromatics (e.g. substituted benzenes, naphthalenes, biphenyl), alkylating agents (e.g. alkenes, alcohols, chlorides) and catalysts (e.g., acidic, basic and radical, in selective and amorphous forms) participating. Not surprisingly, this diversity has led to a wealth of literature ranging from detailed mechanistic studies to the application of specialized catalysts for carrying out the transformation of choice. In this last respect,

the interplay of mechanistic studies and advances in molecular sieve catalysts, in terms of both classical shape selectivity and catalyst designs which energetically favor the desired reaction, has had a remarkable impact on the field. As a result, process yields and product purities well above 99.5% and 99.98%, respectively, are realized commercially for commodity petrochemicals production, demonstrating the movement of inorganic catalysts into the range normally associated with enzymatic selectivity. The future of aromatic alkylation also holds exceptional promise, driven by the strong demand growth for basic petrochemicals and the economic requirements of feedstock utilization and energy efficiency. Advances in polymer science will also continue to create the need for increasingly complex, tailored, high-purity aromatics, serving as monomer precursors for specialty polymers used in the materials and electronics industries. The application of advances in material science to prepare ever more versatile and selective designer catalysts will continue to push the pace of innovation in this field to meet cost-effectively the demand for these increasingly specialized products and processes. References 1. G. A. Olah (Ed.), Friedel–Crafts and Related Reactions, Vol. II, Interscience, New York, 1963, 656 pp. 2. G. A. Olah (Ed.), Friedel–Crafts Chemistry, Wiley, New York, 1973, 581 pp. 3. H. Pines, The Chemistry of Catalytic Hydrocarbon Conversions, Academic Press, New York, 1981, 305 pp.

References 4. W. Keim, M. R¨oper, in Ullmann’s Encyclopedia of Industrial Chemistry, 5th Ed., Vol. A1, VCH, Weinheim, 1985, p. 185. 5. B. V. Vora, J. A. Kocal, P. T. Barger, R. J. Schmidt, J. A. Johnson, Kirk-Othmer Encyclopedia of Chemical Technology, 4th Ed., Wiley, New York, 2001, pp. 2: 169–203. 6. P. B. Venuto, Microporous Mater. 1994, 2, 297. 7. G. A. Olah (Ed.), Friedel–Crafts and Related Reactions, Vol. I, Interscience, New York, 1963, p. 205. 8. N. Y. Chen, W. E. Garwood, R. H. Heck, Ind. Eng. Chem. Process Des. Dev. 1987, 26, 706. 9. R. H. Allen, L. D. Yates, J. Am. Chem. Soc. 1961, 83, 2799. 10. J. E. Szulejko, T. B. McMahon, J. Am. Chem. Soc. 1993, 115, 7839. 11. A. Corma, V. Martinez-Soria, E. Schnoeveld, J. Catal. 2000, 192, 163. 12. L. M Stock, H. C. Brown, J. Am. Chem. Soc. 1959, 81, 3323. 13. C. W. McCarthy, Y. Okamoto, H. C. Brown, J. Am. Chem. Soc. 1955, 77, 3037. 14. P. B. Venuto, P. S. Landis, in Advances in Catalysis and Related Subjects, D. D. Eley, H. Pines, P. B. Weisz (Eds.), Vol. 18, Academic Press, New York, 1968, p. 259. 15. S. M. Csicsery, J. Catal. 1970, 19, 394. 16. P. B. Weisz, Chem. Technol. 1973, 3, 498. 17. W. O. Haag, N. Y. Chen, in Catalyst Design, L. L. Hegedus (Ed.), Wiley, New York, 1987, p. 163. 18. W. O. Haag, in Zeolites and Related Microporous Materials: State of the Art 1994, J. Weitkamp, H. K. Karge, H. Pfeifer, W. H¨olderich (Eds.), Studies in Surface Science and Catalysis, Vol. 84, Part B, Elsevier, Amsterdam, 1994, p. 1375 19. W. O. Haag, R. M. Lago, P. G. Rodewald, J. Mol. Catal. 1982, 17, 161; C. D. Chang, W. H. Lang, R. L Smith, J. Catal. 1979, 56, 169. 20. V. N. Romannikov, K. G. Ione, J. Catal. 1994, 146, 211. 21. J. C. Cheng, T. F. Degnan, J. S. Beck, Y. Y. Huang, M. Kalyanaraman, J. A. Kowalski, C. A. Loehr, D. N. Mazzone, in Science and Technology in Catalysis 1998, H. Hattori, K. Otsuka (Eds.), Studies in Surface Science and Catalysis, Vol. 121, Elsevier, Amsterdam, 1999, p. 53. 22. H. Du, D. H. Olson, J. Phys. Chem. B 2002, 106, 395. 23. G. Girotti, O. Cappellazzo, E. Bencini, G. Pazzuconi, C. Perego, US Patent 6 034 291, assigned to Enichem, 2000. 24. B. Maerz, C. Morris Smith, in Handbook of Petrochemicals Production Processes, R. A. Meyers (Ed.), McGraw-Hill, New York, 2005, Chapter 5.3, p. 5.23. 25. F. G. Dwyer, P. J. Lewis, in Encyclopedia of Chemical Processing and Design, J. J. McKetta, W. A. Cunningham (Eds.), Vol. 20, Marcel Dekker, New York, 1989, p. 77. 26. R. R. Coig, V. A. Welch, S. Ram, J. Singh, in Ullmann’s Encyclopedia of Industrial Chemistry, 5th Ed., Vol. A10, VCH, Weinheim, 1987, p. 35. 27. H. U. Hammershaimb, T. Imai, G. J. Thompson, B. V. Vora, in Kirk-Othmer Encyclopedia of Chemical Technology, 4th Ed., Vol. 2, Wiley, New York, 1992, p. 94. 28. B. Maerz, C. Morris Smith, in Handbook of Petrochemicals Production Processes, R. A. Meyers, (Ed.), McGraw-Hill, New York, 2005, Chapter 5.3, pp. 5.23–5.38. 29. H. W. Gote, Oil Gas J. 1958, 56(13), 73. 30. F. G. Dwyer, P. 1. Lewis, Chem. Eng. 1976, 83(1), 90; P. J. Lewis, F. G. Dwyer, Oil Gas J. 1977, 75(40), 55. 31. M. C. Clark, R. J. Cimini, C. M. Smith, B. Maerz, US Patent 6 995 295, 2006. 32. M. C. Clark and B. Maerz, EBMax SM Process Design, Research and Development, Badger L.L.C. Technology Conference, TC2005, Vancouver, BC, 2005.

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14.4

Isomerization and Transalkylation of Alkylaromatics David L. Stern∗ , Stephen H. Brown, and Jeffrey S. Beck

14.4.1

Introduction

The commodity aromatics industry can be said to have started with the development of synthetic fibers early in the 20th century [1] using feedstocks from coaltar distilleries that reach back to the beginnings of organic chemistry. Simple aromatic molecules are easily functionalized, yielding aromatic diacids (terephthalic acid, isophthalic acid), phthalic anhydride, styrene, phenol and adipic acid/caprolactam. These monomers are prepared, respectively, from the aromatic feedstocks p-, m- and o-xylene, ethylbenzene, cumene and benzene; hence the importance of developing economic processes for their production. Polymerization or copolymerization of these functionalized monomers yields synthetic fibers with properties that are attractive for a variety of consumer products. The isomerization and transalkylation of alkylaromatics and the synthesis of alkylaromatics via alkylation evolved as an efficient means of producing purified base aromatics as the demand for synthetic fibers grew. BTX feedstocks (benzene, toluene, xylenes) became readily available from refineries following World War II [2] with the widespread ∗

Corresponding author.

14.4.2 Catalysts

introduction of reforming to produce high-octane motor gasoline via naphtha aromatization. Significant progress was made in the catalysis and separations processes needed to recover purified BTX isomers from reformate in the mid-20th century, concurrent with improvements in the conversion of purified aromatics to monomers. Advances in polymer science and spinning techniques later in the 20th century allowed the production of tailored fabrics, plastics and films from these functionalized monomers, fueling a steadily rising demand for aromatics that continues today. Aromatics alkylation processes synthesize heavier aromatics such as ethylbenzene and cumene via the addition of olefins to light aromatics; this related subject is reviewed in Chapter 14.3. Transalkylation processes redistribute alkyl side-chains on aromatics to more desirable products, while disproportionation is a special case of transalkylation that transfers an alkyl group between identical molecules. Isomerization of alkylaromatics increases the yield of a desired aromatic isomer by the redistribution of side-chains without a change in molecular weight. Isomerization and transalkylation processes are acid catalyzed and are typically conducted in fixed-bed reactors with solid catalysts. This chapter provides an up-to-date review of the catalysts, processes and mechanisms of these alkylaromatic transformations. 14.4.2

Catalysts

Most often the catalysts are microporous, crystalline aluminosilicate zeolites or silicoaluminophosphates (SAPOs). The chemistries of commercial aromatics isomerization and transalkylation processes are similar enough to make some solid acid catalysts effective candidates for more than one of the transformations. These catalysts share a number of basic features. They contain 10- and/or 12-ring channel systems, may be onedimensional or have intersecting channel systems (i.e. three-dimensional) and have Brønsted acid sites. Threedimensional molecular sieves have intersections larger than each individual pore, and can accommodate larger bimolecular transition states, and may contain pockets to accommodate larger molecules and reactions. The zeolites are silica rich (Si:Al > 10), with acid strength tailored to carry out the desired reaction and minimize undesirable side-reactions. Catalyst Acid Strength Zeolites have both strong and weak Brønsted acid sites and Lewis acid sites. The hydroxyl group adjacent to aluminum in the silicoaluminate matrix of a zeolite, which 14.4.2.1

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satisfies the valence requirement in Si−O(H)−Al−O−Si bonding, is a strong Brønsted acid site, while the surface hydroxyl groups from surface silica in zeolites are weak Brønsted acid sites. The exposed cations, such as the extra-framework Al3+ on an alumina binder, are weak Lewis acid sites. The strong Brønsted sites are most important in catalytic reactions to produce and transform alkylaromatics [3], since most of these reactions require at least moderate acidity; in transformations requiring high selectivity, attention must be paid to all sources of acidity which can lead to less desirable products. The overall acid strength of a hydrogen-form aluminosilicate zeolite depends on the aluminum distribution. Acidity associated with aluminum tetrahedra is stronger with a smaller number of aluminum atoms as next-nearest neighbors [4–6]. Loewenstein’s rule forbids the formation of Al−O−Al bonds in zeolite structures [7], therefore the potential number of acid sites equals the number of aluminum atoms in any reference unit of a zeolite crystal. Careful studies with H-ZSM-5 demonstrate that acid sites with 0 and 1 next-nearest neighbor aluminums are very close in acid site strength [8–9]. Most of the acid sites in zeolites with Si:Al ratios >10 have only a small number of their sites with more than one aluminum nextnearest neighbors, leading to uniform acid site strength. The strength of this site has been well characterized by NMR and IR probes of simple sorbates, allowing the conclusion to be reached that the acid site strength is similar to that of 70% sulfuric acid [10–12]. Model compound reaction studies uncomplicated by mass transfer limitations provide further support for uniform acid site strength [13–16]. The relative rate of conversion of nhexane at 538 ◦ C versus an amorphous silica–alumina cracking catalyst [17], designated the alpha-test, is a useful probe for characterizing the acidity of a zeolite catalyst. The rate of n-hexane cracking (or α value) correlates well with the aluminum content (and thus the SiO2 /Al2 O3 ratio) in H-ZSM-5, with samples ranging over three orders of magnitude in aluminum content [18]. Similar correlations have been found for a number of acid-catalyzed reactions over H-ZSM-5, including toluene disproportionation, hexene cracking, methanol conversion and ethylbenzene dealkylation [19]. Silicoaluminophosphates [20] are also used in petrochemical processing. These crystalline sieves of aluminum phosphate [21], which are isoelectronic with the silica in zeolites, derive their acidity by isomorphous substitution of silicon into an aluminophosphate framework via substitution of silicon for aluminum [substitution mechanism 1 (SM1)], substitution of silicon for phosphorus (SM2) or substitution of two silicons for an Al−P pair (SM3). The SM2 and SM3 pathways leads References see page 3192

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to the creation of Brønsted acid sites, via the hydroxyl group in −O−Si−O(H)−Al. SAPOs are active acid catalysts for reactions, as illustrated by their ability to transalkylate methylaromatics [22], but find most use in acid-catalyzed conversions requiring mild acidity, such as methanol conversion to olefins [23] and the hydroisomerization of n-octane [24]. As can be expected, the activity of SAPOs increases with the silicon content, as has been found for the hydroisomerization of ndecane over SAPO-41 [25]. The nature of these sites has been examined in some detail via NMR, IR and amine sorption [26]. Catalyst Structure Up-to-date information about molecular sieve structures is available from the International Zeolite Association (IZA), with the structure atlas in both written and online form [27, 28]. At the time of writing, the on-line atlas contained 165 known structure types, of which fewer than 20 meet both the acid site density and pore dimension criteria necessary to qualify as potentially useful transalkylation or isomerization catalysts. These catalysts are listed in Table 1, which includes the IZA structure code, some common names for the sieves and the dimensionality of the pores. FAU, BEA, MWW, MTW, MOR, MFI and AEL are all believed either to have been or are being used commercially for alkylaromatics transformations. Although a straightforward exercise, no document is known that compares well-characterized samples of all 14.4.2.2

Zeolite structures with Si:Al between 5 and 50 and with 10- and/or 12-ring channel systems and some aluminophosphate molecular sieves

Tab. 1

IZA structure code

EUO (ZSM-50) MTT (ZSM-23) TON (ZSM-22) AEL (SAPO-11) AFO (SAPO-41) NES FER (ZSM-35) MFS (ZSM-57) MWW (MCM-22) MEL (ZSM-11) MFI (ZSM-5) MTW (ZSM-12) LTL (zeolite L) MOR (mordenite) BEA (beta) EMT FAU (faujasite)

Ring size

10 10 10 10 10 10 × 10 10 × 8 10 × 8 10 × 10, plus 12-ring 3D pocket 10 × 10 × 10 10 × 10 × 10 12 12 12 × 8 12 × 12 × 12 12 × 12 × 12 12 × 12 × 12

Dimensionality of pore system 1D 1D 1D 1D 1D 2D 2D 2D 2D (10 × 10), plus 3D pocket 3D 3D 1D 1D 2D 3D 3D 3D

candidate catalysts in Table 1 for xylene isomerization or transalkylation of methyl or ethyl/propyl aromatics. The IZA Catalysis Commission has established a standard test to determine the rate of ethylbenzene transalkylation over the zeolite LaNaY [29] (see Chaper 2.6.2). Results from testing at five independent laboratories are in reasonable agreement. Presumably, a standardized test may be widely used in the future to provide more rigorous comparisons of activity amongst different zeolite and SAPO catalysts. FAU, EMT, FER, LTL and MOR are most easily produced with high aluminum content (Si:Al < 5) and were first synthesized at Union Carbide. Union Carbide also pioneered the discovery and initial development of aluminophosphate and silicoaluminophosphate molecular sieves, such as AEL and AFO. ExxonMobil pioneered the synthesis of high-silica zeolites (Si:Al ratios >10), including the zeolites MEL, MFI, MFS, TON, MTT, MTW, BEA and MWW. Subsequent to ExxonMobil’s commercialization of MFI for xylene isomerization in the 1970s, extensive academic work appeared. This 1980s literature is dominated by work using MFI (ZSM-5). Like MFI, MEL and NES are also multi-dimensional 10-ring zeolites. The available literature using these catalysts for xylene isomerization and toluene disproportionation indicates similar selectivities. Lack of more intensive work suggests that catalysts cannot be made from these materials with significantly improved performance over MFI. The conversion of polyethyl- and polyisopropylbenzenes does not readily occur over the 10-ring only zeolite structures, since they are too small and so only the 12-ring structures and MWW are candidates for these reactions. In the 1990s ExxonMobil commercialized a revolutionary process for cumene production using a structurally unique zeolite (MWW) [30]. Subsequent to ExxonMobil’s success, extensive academic literature emerged reporting on the uses and characteristics of MWW [31]. MWW is a unique zeolite structure [32], which can be described as a two-dimensional zeolite, with an intersecting twodimensional 10-ring channel system in the plane and 12-ring ‘‘pockets’’ covering the surface (Fig. 1). For transethylation and transisopropylation reactions, only the 12-ring pockets are active.

12 ring surface pockets

25.2 Å

Fig. 1

The unique structure of MWW (MCM-22).

14.4.2 Catalysts

EUO, MTT, TON, FER and MFS generally have low acid activity. Each of these structures contains a onedimensional 10-ring pore system. Both m- and o-xylene are too large to diffuse readily. These zeolites are also not used for toluene disproportionation because they do not have internal spaces with dimensions large enough to accommodate the aromatic dimer intermediates believed to be required. MTW, MOR, EMT, MWW and FAU all have sufficient activity and large enough pore systems to be considered for xylene isomerization and aromatics transalkylation applications. All methylbenzenes up to hexamethylbenzene can readily diffuse in and out of these 12-ring materials. For this reason, they are effective catalysts to bring the methylbenzenes to their calculated equilibrium distribution. As a result, significant quantities of C9+ polymethylbenzenes are produced independent of the composition of a polymethylbenzene feedstock as equilibrium is approached. Poor stability is known to be a problem for MOR and FAU [33, 34]. Hydrothermally synthesized MOR is typically an aluminum-rich, one-dimensional structure. Post-processing is required to yield a lower aluminum content, high-activity zeolite catalyst; lower aluminum content is associated with reduced coke selectivity. This is an important property for the onedimensional 12-ring channels of MOR because a single coke species at a pore mouth will deactivate multiple acid sites. FAU and EMT have large internal supercages which promote aromatic oligomerization and conjunct polymerization reactions. The high coke selectivity of FAU and EMT has prevented them from becoming widely used transmethylation catalysts. Shape Selectivity and Diffusional Restriction Assuming acidity is of the proper level, the pore size and dimensionality of the molecular sieve are typically chosen so as to govern the selectivity of the reaction, termed shape selectivity. Three main types of shape selectivity (see Fig. 2) have been utilized to explain the regioselectivity relevant to alkylaromatics transformations: 14.4.2.3

A. Transition-state selectivity or spatioselectivity, which refers to a limitation in the formation of the transition state, due to the size and dimensionality of the sieve pores and intersections. In the example shown, xylene disproportionation is limited due to the difficulty of forming the sterically bulky, bimolecular intermediate, whereas unimolecular xylene isomerization is facile. B. Reactant shape selectivity occurs when the reactants diffuse at significantly different rates through the sieve pores. In the example shown, n-hexane diffuses faster through the ZSM-5 pores and thus cracks at a fast rate relative to the bulkier 3-methylpentane, which has much lower rates of cracking.

2

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H+

(a)

(b)

(c)

Some examples of shape selectivity. (a) transition-state selectivity or spatioselectivity; (b) reactant shape selectivity; (c) product shape selectivity.

Fig. 2

C. Product shape selectivity occurs when the products diffuse at significantly different rates through the sieve pores. In the example shown, isomerization of trimethylbenzenes in the large intersections of ZSM-5 yields 1,2,4-, 1,2,3- and 1,3,5-trimethylbenzene (TMB), but since 1,2,4-trimethylbenzene is the sterically least demanding of the TMB isomers, it more readily escapes the 10-ring pores and hence has higher than equilibrium concentrations outside the zeolite pores. Diffusion control refers to the relative rate of diffusion of one molecule over another in a particular, constrained environment. Experimental measurement of diffusivities is commonly carried out by gravimetric sorption using probe molecules. Using this technique, the diffusivities of p-, m- and o-xylene in ZSM-5 [35] and of ethylated and propylated aromatics in MCM-22 [36] can be measured, providing insight into sieve pore size and shape and gaining predictive power on reaction selectivities. A particularly useful test for understanding the pore systems in molecular sieves is a reactive probe called the constraint index (CI) [37]. In this test, the rates of n-hexane and 3-methylpentane cracking (kH and kMP , respectively) are measured at conditions where References see page 3192

3172

14.4 Isomerization and Transalkylation of Alkylaromatics

bimolecular hydride transfer is typically the rate-limiting step [38]. The observed ratio kH /kMP , or CI, correlates well with the cross-sectional area for a variety of small- and medium-pore zeolites. The measured CI is independent of the crystal size and acidity, at least for ZSM-5.

3. transalkylation of C9+ alkylaromatics with toluene and/or benzene 4. selective toluene disproportionation to p-xylene and benzene 5. alkylation of toluene with methanol to make either equilibrium xylenes or to selectively make p-xylene.

14.4.3

Xylenes

Aromatic Feedstocks and Reformate Separations Aromatic streams are produced in the refinery via reforming of straight-run and catalytic naphtha for the production of aromatic-rich, high-octane gasoline. Aromatics, primarily benzene, toluene and ethylbenzene, are also produced as a by-product from ethylene manufacture via pyrolysis, called pyrolysis gasoline (pygas). The volumes of C8 aromatics recovered from pygas are much smaller than from reforming (less than 3%). Reformate contains both aromatics and non-aromatics, the composition of which varies considerably since it is a function of naphtha feedstock composition and reformer conditions. Although not prescriptive, the ratio of aromatics to non-aromatics in reformate is usually higher than 3:1. The carbon number distribution of the reformate, as expected, is a product of the molecular weight distribution of the feed: with higher C6 −C7 content, more benzene and toluene are formed, while with higher C8 −C10 content, more xylenes and higher aromatics are formed [42]. The non-aromatic fraction 14.4.3.1

Purified p- and o-xylene isomers are used on a large scale as intermediates for many derivatives, some of which include major petrochemicals. Demand for purified mxylene is believed to be between 100 and 1000 kt a−1 , the volume level separating commodity and specialty chemical products. The 2005 industrial demand for p- and o-xylene is ca. 32 000 kt a−1 , split roughly 7:1 [39]. p-Xylene demand is so large that industry has developed processes to produce it from other aromatic molecules. Routes for separating and converting the common feedstocks for the production of p-xylene are shown in Fig. 3. They include equilibrium processes (1, 2 and 3) and para-selective routes (4 and 5): 1. isomerization of xylenes from a mixture of C8 aromatics 2. toluene disproportionation to make a mixture of equilibrium xylenes and benzene

(1)

(2)

(3)

+

+

Para-selective processes Equilibrium processes (4)

(5)

Fig. 3

+

+ MeOH

Equilibrium and para-selective processes for the production of p-xylene.

14.4.3 Xylenes

of reformate decreases with increasing carbon number, since the reformer can more easily dehydrocyclize heavier paraffins and olefins than lighter ones, so that C6 and C7 paraffins have the highest concentration, followed by much smaller amounts of C8 non-aromatics, and C9+ non-aromatic components are typically of insignificant concentration. The boiling points of aromatics and non-aromatics overlap sufficiently so that distillation does not separate them – for example, benzene and cyclohexane have boiling points within 1 ◦ C (boiling points of 80.1 and 80.7 ◦ C, respectively). The non-aromatics can either be converted in aromatics processes via cracking to light gases or they can be removed via extraction or extractive distillation processes using a polar solvent, such as sulfolane, ethylene glycol, dimethyl sulfoxide, dimethylformamide or pyrrolidones [40]. Typically, the benzene and toluene streams are extracted, whereas the C8+ aromatics are not. Reformate is distilled to produce benzene, toluene and C8+ reformate. Additional distillation towers and extraction or extractive distillation are employed to produce chemical-grade benzene, which typically requires >99.9% purity. Toluene can sometimes be used directly in further aromatics processing and is referred to as non-extracted toluene. The C8+ reformate is distilled to produce mixed xylenes and C9+ reformate. Both extracted and non-extracted toluene and xylenes are suitable feedstocks for toluene disproportionation and xylene isomerization processes, respectively. The purpose of commercial methylbenzene isomerization and transmethylation processes is to use the fractionated and/or extracted toluene, xylenes and C9+ reformate streams as feedstocks to produce more benzene, p-xylene and o-xylene. Xylene Separations The four C8 aromatics, ethylbenzene and p-, m- and oxylenes, are fairly close in boiling point (Table 2). Because of their close boiling points, all four C8 aromatic isomers are not fractionated to yield the high purity required commercially (typically >99%). 14.4.3.2

Melting and boiling points of C8 aromatics (at atmospheric pressure)

Tab. 2

C8 aromatic Ethylbenzene m-Xylene o-Xylene p-Xylene

Melting point/ ◦ C −95.0 −47.9 −25.2 13.3

Boiling point/ ◦ C 136.2 139.1 144.4 138.3

3173

o-Xylene, however, is separated from the C8 heart cut by distillation. This separation requires a well-trayed tower and high reflux (known as the ortho tower) due to the small difference in boiling point and hence is capital and energy intensive. In complexes employing an ortho tower, purified C8 aromatics depleted in o-xylene are recovered as overhead, whereas o-xylene and C9+ are recovered as bottoms; a second column then recovers o-xylene in high purity. p-Xylene purification is typically accomplished by a multistage, countercurrent sorptive process employing sieves as sorbent beds, typically cation-exchanged zeolite X, coupled with desorption using p-diethylbenzene [41] or toluene (termed heavy and light desorbents, respectively). Distillation is used to separate the desorbent and p-xylene, with resulting p-xylene purity exceeding 99.8% and with yields of >97%. The two most widely utilized sorptive processes were developed by UOP (the Parex process [42]) and IFP (the Eluxyl process [43]) and introduced in 1971 and 2000, respectively; the Parex process is more widely used commercially than the Eluxyl process. These two processes simulate a countercurrent, moving bed sorption process in multiple fixed beds and differ primarily in their configuration. The Parex process employs either single or dual vessels with multiple beds and a highly engineered rotary valve to control flows and switch between sorption and desorption, whereas the Eluxyl process employs multiple separate beds with associating valving. Because it has a higher freezing point than the other C8 aromatic isomers, p-xylene may also be removed via crystallization. A eutectic composition limit restricts p-xylene recovery from an equilibrium xylene mixture to ca. 65% per pass and at 90%). m-Xylene may be removed by two processes: a sorptive process introduced by UOP (MX Sorbex) or in a process introduced by Mitsubishi Gas Chemical Company (MGCC) via complexation with HF−BF3 . UOP’s MX Sorbex process [44] is believed to be an extension of their countercurrent sorptive Parex process in which the characteristics of the sorptive sieve are modified. This process, which was introduced in 1995, has been reported to achieve 95% per pass m-xylene recovery at 99.5% purity. The MGCC process is based on the significantly higher basicity of m-xylene than either p- or o-xylene (20-fold higher basicity for m- than for p- or o-xylene). In this process [47], mixed xylenes are treated with strongly acidic HF−BF3 . Two layers are References see page 3192

3174

14.4 Isomerization and Transalkylation of Alkylaromatics

formed – a 1:1 molecular complex, m-xylene–HBF4 layer, and an organic layer containing the remaining xylenes; m-xylene is then recovered in 97–99% purity by thermal decomposition of the m-xylene–HBF4 complex. This process has been commercially practiced by MGCC and its licensees, however, the difficulties in dealing with the highly corrosive HF−BF3 and the need for high-purity mxylene make this process intrinsically less attractive than sorptive recovery. Ethylbenzene can be isolated from the C8 heart cut by superfractionation in a column with several hundred distillation plates. This was commercially practiced over 40 years ago, but because the purity of ethylbenzene required is lower than that afforded by alkylation of ethylene with benzene, and the expense of operating superfractionation equipment is high, isolation of ethylbenzene from mixed C8 streams by superfractionation is believed to no longer be used commercially. Toluene Conversion to Benzene Toluene can be hydrocracked to benzene and methane in hydrodealkylation (HDA) processing. HDA, which currently accounts for approximately 15% of petrochemicalgrade benzene production, can be carried out either catalytically or thermally. The catalytic process utilizes acidic, supported Group VIII metal or metal oxide supported on alumina [46, 47] and operates at temperatures below 500 ◦ C in order to minimize metal sintering. Thermal processing employs high temperatures (typically >650 ◦ C) and high concentrations of hydrogen (H2 :toluene is typically 2–6). Thermal HDA, such as the UOP THDA process [48], is currently the dominant route for new process units using HDA, since it produces high yields of benzene and avoids the use of a catalyst which needs to be replaced periodically. HDA processes have been tuned with plantwide control design [49] so that a nearly quantitative molar yield of benzene can be recovered (99%), but the formation of methane is a net aromatic weight yield loss 14.4.3.3

and hence the ultimate net weight yield is limited to ca. 84%. The relative pricing of benzene and toluene varies over the course of petrochemical cycles and thus dictates economics but, as a rule of thumb, HDA is favorable when the spot pricing of benzene is greater than 1.25 times that of toluene. HDA processing is also used to convert heavy reformate and FCC cycle oils to higher value naphthalene. 14.4.4

Xylene Isomerization

The small volumes of p-xylene required in the early days of this industry were readily obtained by crystallization from distilled mixed xylenes. The C8 aromatic by-products were typically used as gasoline blendstock. By the 1960s, the market for p-xylene grew beyond the available supply obtained by single pass crystallization. Producing incremental p-xylene from the aromatic by-products of crystallization was an attractive option. Ethylbenzene is either catalytically isomerized to incremental mixed xylenes or hydrodealkylated to benzene and ethane and m- and o-xylene are converted to additional p-xylene by xylene isomerization. Fundamentals of Xylene Isomerization The product distribution of xylene isomerization processes is governed by thermodynamic equilibrium. The equilibrium concentrations of the three xylene isomers as calculated from thermochemical data are shown in Table 3. m-Xylene is the most abundant of the xylenes and of the C8 s. The ethylbenzene fraction of the C8 s increases with temperature – it is lowest at typical xylene isomerization and reforming conditions (12–15% at 425 to 540 ◦ C), and is higher under steam cracking conditions (>35% at 850 ◦ C). The experimentally determined equilibrium concentration [50] of p-xylene (of the xylenes) is higher than the calculated equilibrium value at typical xylene isomerization conditions. This is likely due to inaccuracies in the thermodynamic data. 14.4.4.1

Equilibrium concentrations of the three xylene isomers, as calculated from thermochemical data and experimentally determined xylene concentrations at equilibrium

Tab. 3

Compound

p-Xylene m-Xylene o-Xylene

Equilibrium concentration/mol% From thermodynamic data

From experimental data

623 K

623 K

23.7 53.0 23.3

673 K 23.5 52.4 24.1

24.4 54.0 21.6

673 K 24.0 53.2 23.8

14.4.4 Xylene Isomerization

Ethylbenzene Conversion in Xylene Isomerization Processes Three types of xylene isomerization processes that differ in the way ethylbenzene is converted have been commercialized: [Eq. (1)] de-ethylation of ethylbenzene to benzene and ethylene and saturation of the ethylene formed, [Eq. (2)] ethylbenzene hydroisomerization to produce additional xylene or [Eq. (3)] transalkylation of the ethyl group, followed by distillation to remove the heavies formed: 14.4.4.2

ethylbenzene + H2 −−−→ benzene + ethylene + H2 −−−→ benzene + ethane

(1)

ethylbenzene + H2 −−−→ ethylcyclohexane −−−→ dimethylcyclohexanes −−−→ xylenes + H2

(2)

ethylbenzene + xylenes −−−→ benzene + diethylbenzenes + ethylxylenes

(3)

The first route [Eq. (1)] does not have equilibrium limitations for all practical purposes due to the exhaustive saturation of ethylene formed from dealkylation. It can therefore operate at much higher conversions, typically 60–85% per pass, which allows for reductions in the size of the xylene isomerization loop equipment. In the second route [Eq. (2)], ethylbenzene conversion is limited

to 30% by the equilibrium between ethylbenzene and xylenes. The third route [Eq. (3)] operates at moderate ethylbenzene conversion (30–40%) since transalkylation of ethylbenzene with xylenes results in a net xylene yield loss. Ethylbenzene Hydrocracking Processes in Xylene Isomerization Ethylbenzene hydrocracking [Eq. (1) above] is the most common xylene isomerization process and co-produces p-xylene and benzene. The benzene yield obviously depends on the concentration of ethylbenzene in the feedstock. This process utilizes a bifunctional catalyst: an acidic zeolite to crack the ethylbenzene into ethylene plus benzene and a metal function to exhaustively saturate the olefin thus formed. Typically, a medium-pore (10-ring) multi-dimensional zeolite, such as ZSM-5, is the acid catalyst. The choice of metal function is critical and must effectively saturate the olefin without hydrogenating the aromatics present in the feed and must also hydrogenate dimeric coke precursors so as to limit catalyst aging in service. Figure 4 shows a typical process flow diagram. The C8 and C9+ feedstock is first fractionated to produce the C8 heart cut by removing C9+ aromatics in the 14.4.4.3

References see page 3192

Para-Xylene product

Para-Xylene recovery

Stabilizer H2

Xylene isomerization reactor Lights (C5−)

C8+ aromatics (from light reformate splitter)

Ortho-Xylene product

Separator drum Benzene and toluene

Xylene splitter

Ortho tower

Fig. 4

3175

C9+ aromatics

Process flow diagram for a typical xylene isomerization process.

Light aromatics splitter

3176

14.4 Isomerization and Transalkylation of Alkylaromatics

C9 splitter; optionally an o-xylene distillation tower is employed if o-xylene is desired (see Section 14.4.3.2). p-Xylene is recovered from the C8 heart cut in a sorption process; since the sorption process recovers >95% of the p-xylene in the feed, the raffinate has a low p-xylene content of ca. 1%. The p-xylene-depleted raffinate from the sorption is then treated in the xylene isomerization reactor, in which ethylbenzene is hydrodealkylated, the xylenes are isomerized to equilibrium and co-boiling paraffins are hydrocracked to light gas. The isomerate is then distilled in successive towers: the stabilizer removes light gases, whereas benzene and toluene are removed as overhead in the splitter and are subsequently distilled to produce benzene and toluene. The C8 isomerate is then combined with fresh feed and recycled back to the C9+ splitter for further processing. Since the equilibrium concentration of p-xylene is ca. 24%, the C8 stream undergoes approximately three to four recycles if only p-xylene is recovered. Because of the high recycle, avoiding undesirable reactions in the xylene isomerization reactor is critical. Undesirable reactions include methyl [Eq. (4)] and ethyl [Eq. (5)] aromatic transalkylation to yield C9+ aromatics, saturation of aromatics including both benzene and xylenes to yield cyclohexane and dimethylcyclohexanes [Eq. (6)] and hydrocracking of aromatics to produce light gas [Eq. (7)]. All four reactions reduce the ultimate yield of p-xylene and are referred to as xylene loss. Reactions (6) and (7) reduce the useful product yield, since they result in net aromatic yield loss. Aromatic saturation [Eq. (6)] must be avoided, since the distilled benzene is typically sold as chemical grade, requiring high purity. 2 xylenes −−−→ toluene + trimethylbenzenes

(4)

xylenes + ethylbenzene −−−→ ethylxylenes + benzene

(5)

benzene (or xylene) + H2 −−−→ cyclohexane (or dimethylcyclohexanes)

(6)

xylenes + H2 →→→ light gas (C2 −C6 paraffins) (7) ExxonMobil’s vapor phase isomerization process (MVPI), introduced in 1975, was the first-generation technology. The next-generation technology was ExxonMobil’s high-temperature isomerization process (MHTI), which was introduced in 1981. This process was developed for unextracted feeds containing C8+ aliphatics and operates at higher temperatures (ca. 425 ◦ C). The ExxonMobil high-activity isomerization process (MHAI), advanced MHAI process (AMHAI) and XyMax process are more recently commercialized technologies for xylene isomerization, introduced in 1990, 1999 and 2000, respectively. In all three processes, a unique

Top bed Ethylbenzene and non-aromatics conversion

H2

+

Bottom bed Xylene isomerization

Dual-bed catalyst system used in ExxonMobil’s MHAI, AMHAI and XyMax xylene isomerization processes.

Fig. 5

dual-bed catalyst system is used (Fig. 5) that takes advantage of the low p-xylene content (ca. 1%) of the feed to the isomerization reactor. In these processes, the top-bed catalyst has high activity for ethylbenzene hydrodealkylation and olefin saturation, thus reducing xylene losses through transalkylation, such as in reaction (5) above. The ZSM-5 catalyst also limits xylene disproportionation, another route for xylene loss [reaction (4) above] and is active for non-aromatics conversion via hydrocracking to lighter products, which are removed as light gas. The lower-bed catalyst has high xylene isomerization activity to ensure that a very high p-xylene approach to equilibrium is obtained – up to 104% versus the calculated equilibrium values. The result of this stacked-bed catalyst system is the ability to achieve very high ethylbenzene conversion (60–85%), a very high p-xylene approach to equilibrium (up to 104% versus the calculated equilibrium values) and high non-aromatics conversion (typically 30–40%). High ethylbenzene and non-aromatics conversion are important in order to prevent these molecules from building up in the recycle loop. Typical process conditions are 400–470 ◦ C, 14–21 bar, a 1–5 molar hydrogen:hydrocarbon ratio and a WHSV of 5–20 h−1 C8 aromatic feedstock [51, 52]. Ethylbenzene conversion and paraffin cracking are more demanding reactions than xylene isomerization and thus benefit from higher temperature operation. Even under severe conditions, per pass xylene losses are kept well below 1% using the XyMax catalyst and process. Minimizing xylene losses at high ethylbenzene conversion has substantial economic value at locations where mixed xylenes feedstock availability bottlenecks plant production. Finally, the benzene produced can

14.4.5 Methylaromatics Transalkylation

exceed the >99.85% purity specification needed for direct chemical sales. UOP also offers technologies for ethylbenzene dealkylation, denoted Isomar. Their latest offering (I-300) has modest xylene losses at somewhat lower ethylbenzene conversions and utilizes a single catalyst bed for both ethylbenzene conversion and xylene isomerization.

3177

one-dimensional SAPO such as SAPO-31 or SAPO-41 or a medium-pore zeolite with mild acid strength. The per site acid activity of SAPOs is typically lower than that of aluminosilicates, which may account for the higher selectivity for EB isomerization and reduced hydrocracking. Ethylbenzene Disproportionation Processes in Xylene Isomerization ExxonMobil’s vapor-phase isomerization process (MVPI) utilizes ethylbenzene transalkylation as the primary means of ethylbenzene removal. This process operates at relatively low temperature (about 315 ◦ C), uses hydrogen co-feed and converts approximately 30% of the ethylbenzene contained in the C8 stream via disproportionation. The process takes advantage of shape selectivity provided by the H-ZSM-5 catalyst, so the primary route for EB conversion is via disproportionation, hence transalkylation of ethylbenzene to xylenes is reduced. The 10-ring pores in H-ZSM-5 also suppress aging by limiting the formation of coke within the pores relative to larger pore zeolites. 14.4.4.5

Ethylbenzene Hydroisomerization Processes in Xylene Isomerization Ethylbenzene (EB) hydroisomerization catalysts are directed towards interconverting all the C8 aromatic isomers and producing p-xylene without benzene as a co-product [53]. This process requires the formation of cycloolefinic intermediates in order to interconvert the C8 aromatics, which are prone to acid-catalyzed hydrocracking to light gases. Ultimate aromatic yields based on feed aromatics are always less than 100% and typically a net loss of 2–10% of feed aromatics to light gas is observed. Fewer and/or weaker acid sites are thus required. Higher temperatures are also required to achieve sufficient EB conversion to xylenes. Typical process conditions are 400–470 ◦ C, 14–21 bar, a 1–10 molar hydrogen:hydrocarbon ratio and a WHSV of 1–2 h−1 C8 aromatic feedstock [54]. The catalysts for this process have high hydrogenation/dehydrogenation activity and low acidity, with either weaker or a small number of acid sites. The acid sites isomerize the intermediate naphthenes with minimal cracking. Naphthene isomerization is of significantly lower activation energy than isomerization of the aromatic ring. The earliest process for ethylbenzene isomerization, called the Octafining process, was developed and commercialized by Arco and Engelhard in 1960 [55] and used a dual-functional Pt on amorphous silica–alumina catalyst. In the 1970s, the acidic function of the aluminosilicate was replaced with mordenite in commercial catalysts [56]. Since then, several sieves have been shown to be effective for this service [57]. UOP introduced two new catalysts for this process, called I-9 and I-210 [57d]. I-210 may be platinum supported on SAPO-11, a medium pore-sized, one-dimensional SAPO. UOP reports that the latest version of this technology, I-400, achieves 91% ultimate p-xylene yield from fresh feed when employed in a modern petrochemical complex, a significant advance over the 84% ultimate p-xylene yield reported for their earlier I-9 technology [58]. Axens introduced the Oparis process for ethylbenzene isomerization in 2001, which is currently used in several commercial units. Axens [59] reports up to 93% ultimate p-xylene yield using the Oparis catalyst and an optimized reactor/monitoring system. Presumably, these catalysts are based on an improved version of SAPO-11, another medium-pore, 14.4.4.4

14.4.5

Methylaromatics Transalkylation

Toluene, benzene and the C9+ aromatics in reformate are low-cost feedstocks for incremental benzene and p-xylene production via processes which transalkylate methylaromatics to produce mixed xylenes and benzene as the major products [60]. The role of methylaromatics transalkylation processes in the modern petrochemical complex is increasing. This is driven by several factors: • insufficient xylenes yield and the increasing availability of C9+ aromatics from the reformer; • the realization that toluene and C9+ aromatics transalkylation can be a low-cost route for xylenes in both revamp and grass root designs; and • the need to reduce the T95 (temperature to boil 95%) of motor gasoline. Fundamental Aspects of Methylaromatics Transalkylation In all methylaromatics transalkylation reactions, conversion is limited by equilibrium. The product distribution is determined by the alkyl:ring molar ratio (e.g. the alkyl:ring molar ratio for toluene is 1). When every methylbenzene out to hexamethylbenzene is an allowed product [Eq. (8)] and the average molar methyl:aromatic ring ratio is 1, the toluene concentration at equilibrium is 14.4.5.1

References see page 3192

3178

14.4 Isomerization and Transalkylation of Alkylaromatics Tab. 4

Toluene disproportionation: product distribution at thermodynamic equilibrium

Temperature/K

600 800

Product/mol% Benzene

Toluene

Xylenes

Trimethylbenzenes

Tetramethylbenzenes

31.5 32.0

41.7 40.6

22.7 23.1

3.8 3.9

0.3 0.4

near 41 wt.%; Table 4 provides the calculated equilibrium methylaromatic product distributions. For toluene disproportionation, which allows xylene formation but not C9+ aromatics as products [Eq. (9)], the toluene concentration at equilibrium conversion is 43 wt.%, but if the xylenes are limited to p-xylene [Eq. (10)], the toluene concentration at equilibrium is 63 wt.%. toluene −−−→ benzene + xylenes + tri-/tetra-/penta-/ hexamethylbenzenes (8) toluene −−−→ benzene + mixed xylenes toluene −−−→ benzene + p-xylene

(9) (10)

Methylaromatics Transalkylation Catalysts and Process Chemistry Several reactions are desirable in methylaromatics transalkylation: transfer of methyl groups from C9+ aromatics to lighter aromatics [Eq. (11)], hydrocracking of ethylated and propylated aromatics to produce lighter aromatics and saturates [Eq. (12)] and hydrocracking of feed non-aromatics to light gas [Eq. (13)]. Undesirable reactions include the hydrocracking and demethylation of methylaromatics, aromatic hydrogenation and coking. 14.4.5.2

trimethylbenzenes + toluene −−−→ 2 xylenes

(11)

ethyltoluenes + H2 −−−→ toluene + ethane

(12)

non-aromatics+H2 →→→ light gas (C2 −C6 paraffins) (13) Zeolites with added metal function are used to catalyze this reaction. These zeolites typically have a 12-ring pore system to allow fast diffusion of sterically more demanding feed molecules, including trimethyl- and dimethylethylbenzenes, to the acid sites of the zeolite. Larger molecules, such as penta- and hexaalkylbenzenes, are not usually admitted. Since these larger aromatics easily form coke, this reactant shape selectivity reduces catalyst aging. The internal pore system must also be large enough to accommodate dimeric intermediates formed in methylaromatic transalkylation reactions, which requires an internal 12-ring pore system and/or the intersections of multi-dimensional pore systems. The transmethylation

of aromatics [Eq. (11)] requires strong acid sites. MWW, MTW, LTL, MOR, BEA, EMT and FAU are a nearly complete list of suitable candidates. The metal function in these catalysts is chosen so as to saturate the light olefins formed from dealkylation of ethyl and propylaromatics [Eq. (12)]. This increases the per pass C9+ conversion and allows for subsequent reaction of the lighter alkylaromatics, such as the toluene formed from methylethylbenzene hydrodealkylation, to form additional xylenes. Aromatic saturation and hydrocracking must be avoided. Hydrocracking of paraffins [Eq. (13)] is desirable to prevent their contamination of product and to avoid build-up of their concentration in the recycle loop. Methylaromatics Transalkylation Processes Commercial processes for methylaromatics processing [61] can be roughly divided into two types, depending on the feedstock employed: transalkylation of toluene with some benzene, with C9 /C10 aromatics and conversion of heavier feedstocks, such as C9 /C10 aromatics, without added toluene or benzene. Both process types employ the methylaromatic and aromatic de-ethylation/depropylation chemistry described above. These processes operate at 400 – 470 ◦ C, 14–28 bar, a 1–5 molar hydrogen:hydrocarbon ratio and a WHSV of 0.5–5 h−1 depending on the feedstock aromatic. Several vendors offer commercial transalkylation processes for license. ExxonMobil’s TDP-3 [62, 63] process utilizes a high-activity ZSM-5-based catalyst to transalkylate toluene to equilibrium and operates at 45–50% toluene conversion. An advantage of the H-ZSM-5 catalyst is the relatively high selectivity for xylene formation. The smaller 10-ring zeolite pores of ZSM-5 limit the amount of C9+ aromatics that may be coprocessed to about 25%. UOP’s Tatoray process [67], ExxonMobil’s TransPlus process [68] and UOP’s TAC-9 process all employ larger pore-sized zeolites for heavy aromatics transalkylation processing and differ in how much toluene and benzene are utilized in the fresh feed. All these processes operate with extinction recycle of aromatics, with product xylenes fed to a xylenes isomerization loop. Figure 6 shows how alkylaromatics transalkylation processes may be integrated into an aromatics processing complex employing transalkylation processing and xylene isomerization. 14.4.5.3

14.4.5 Methylaromatics Transalkylation

3179

Benzene, Toluene, and heavies towers Toluene product

Benzene product

Toluene

Non --aromatics product to Gasoline Pool

C9/C10

Aromatic extraction

H2

C10+ to fuels

C5−

Transalkylation reactor

Reformate

C8+

Separator & stabilizer H2

Reformate splitter

Para-Xylene product

C8+

Xylene Isom. reactor

Para-Xylene recovery

Ortho-Xylene product

Benzene / toluene

Light aromatics splitter

Xylene splitter

Ortho tower

H2 C5− H2

Separator & stabilizer

C9+

Fully integrated petrochemical complex for aromatics production, including aromatics extraction, distillation to separate C6 −C10+ aromatics, xylene isomerization and methylaromatics transalkylation for increased xylene production. Fig. 6

Selective Toluene Disproportionation (STDP) The transmethylation processes described thus far all produce equilibrium xylene mixtures. Selective disproportionation of toluene to chemical-grade benzene and high-purity p-xylene (>80% p-xylene selectivity) is a breakthrough process pioneered by ExxonMobil. The process can convert low-cost distilled toluene or extracted toluene directly into chemical-grade benzene and high-purity p-xylene. The ZSM-5 catalysts that exhibit high para-selectivities have been modified by chemical treatment of the ZSM-5 with organics such as coke or other surface modifications. The experimental observations and model to describe them have been well studied and described in detail elsewhere [35, 66–69] and are presented in brief here. 14.4.5.4

14.4.5.4.1 Shape Selectivity in STDP The selective disproportionation of toluene in ZSM-5 is depicted conceptually in Fig. 7. Toluene disproportionation (TDP) occurs on the acid sites in the pores of the ZSM-5 crystal. The primary products of TDP are probably p- and o-xylene, since these are the kinetic products. Experimentally, an equilibrium mixture of xylenes is observed in an unmodified ZSM-5 catalyst, because the rate of xylene isomerization (kISOM ) is much higher than that for toluene disproportionation (kDIS ) – in ZSM-5 the intrinsic kISOM /kDIS ratio can be 7000. Catalyst modification, with treatments such as coke or inorganic modifiers such as magnesium, phosphorus, silicon, boron and antimony, can block off a significant References see page 3192

3180

14.4 Isomerization and Transalkylation of Alkylaromatics

Relative D

Relative D ZSM-5 crystal

1

104 1

+

104

Model of selective toluene disproportionation in modified H-ZSM-5.

Fig. 7

fraction of the pores of the zeolite, thus creating a much more tortuous path for the xylenes produced within the ZSM-5 crystal. The resulting diffusivity from the pores of the crystal is significantly reduced for all three xylene isomers, and also the toluene feed. However, since p-xylene is less sterically demanding than either m- or o-xylene, the relative diffusivity of p-xylene to the other isomers is ca. 1000 as measured experimentally [35]. Therefore, p-xylene escapes from the ZSM-5 crystal 1000 times faster than m- and o-xylene. Isomerization within the crystal rapidly converts the remaining meta- and orthoisomers toward equilibrium, creating additional p-xylene. When the rate of isomerization in the crystal is much faster than the rate of m- and o-xylene diffusion out of the

pores, the result is a modified catalyst where high paraselectivity is observed or the apparent kISOM /kDIS 1. In addition to creating a diffusional barrier, these modifications also limit secondary isomerization of p-xylene once it has exited the crystal. Since the intrinsic kISOM /kDIS is high, achieving high para-selectivity is critically dependent on ensuring that the modified catalysts do not have appreciable activity for secondary xylene isomerization. This can occur on acid sites on the exterior of the modified zeolite crystal, the binder used to formulate the zeolite into a finished catalyst or on unmodified acid catalyst. The rate of sorption of o-xylene is an excellent probe for characterizing the diffusion barrier in unmodified and modified ZSM-5. Measuring the time required to adsorb 30% of the o-xylene adsorbed at infinite time, t0.3 , is a direct measure of the mass transfer property r 2 /D, where r is the radius of the zeolite crystal and D is the diffusivity of o-xylene, or t0.3 ∝

r2 D

(14)

A good correlation can be drawn between the measured t0.3 for o-xylene (measured at 120 ◦ C in a thermogravimetric analysis experiment) and the observed para-selectivity in toluene disproportionation (measured at 20% toluene conversion and 550 ◦ C). Figure 8 shows

100 S = Small crystal

90

L = Large crystal

Si–L Coked–L

Para-Xylene content (percent of Xylenes)

80 70

Coked–S Sb–S

60 B–S

50

Mg–S Si–L

H–L

40

Ca–S 30

Zn–S

Si–S H–S

U–S Mg–S

20 1

10

100

1000

10 000

Diffusion time, t0.3 / min

Relation of o-xylene diffusion time (t0.3 ) to the para-selectivity achieved in selective toluene disproportionation (STDP) for a number of modified ZSM-5 samples. Conditions: toluene conversion 20%, 550 ◦ C, 41 bar; t0.3 = time required for 30% uptake of o-xylene at infinite time.

Fig. 8

14.4.5 Methylaromatics Transalkylation

this relationship for samples of small- and large-crystal H-ZSM-5 and modifications of these crystals with a variety of surface modifications. The para-selectivity is governed not only by the relative diffusivity, D, of the three xylenes from the pores of the crystal, but also by the relative rate of xylene formation (from toluene disproportionation) and xylene isomerization (to produce additional p- from m- and o-xylenes). To think of this another way, reduced para-selectivity for a catalyst with the same diffusional restriction can result if the relative diffusivity for all three xylenes is held constant but the catalyst activity is lower. If xylene isomerization is slow, m- and o-xylene can diffuse out of the crystal before they are isomerized to p-xylene in the interior of the catalyst, and if TDP is slow, diffusion competes more favorably with the rate of xylene formation. The Thiele modulus, 2 , includes the intrinsic catalyst activity k and fundamentally describes the selectivity obtained in this diffusionally limited STDP reaction [Eq. (15)]. As described in Section 14.4.2.1, the alpha value α, or n-hexane cracking activity, is proportional to the intrinsic activity of the ZSM-5 catalyst for TDP and xylene isomerization. Rearrangement of Eqs. (14) and (15) results in Eq. (16), which shows that the critical parameter in determining para-selectivity, r 2 k/D, can be replaced by two easily measurable quantities, the alpha value α and the o-xylene diffusion time t0.3 :

r 2k  = or  = r D




2

k D

3181 (15)

kr 2 ∝ (αt0.3 ) D

(16)

Empirically, a plot of αt0.3 or the pseudo-Thiele modulus, shows a much better correlation with paraselectivity for catalysts with widely differing catalyst activity. This is illustrated for a series of magnesiummodified catalysts with 0–5% added MgO, which have widely differing activities (α values range over an order of magnitude). A graph of o-xylene diffusiuon time (t0.3 ) versus para-selectivity does not fall on the smooth curve, but when the pseudoThiele modulus (αt0.3 ) is plotted against para-selectivity, a smooth curve is obtained which includes catalysts differing in activity by an order of magnitude (Fig. 9). Selective Toluene Disproportionation Processes As described above, the diffusion barrier is critical to obtaining high p-xylene selectivity in STDP technologies. Addition of the barrier to reduce diffusion and secondary xylene isomerization is known as selectivation. The diffusion barrier may be added to the catalyst either while 14.4.5.4.2

References see page 3192

100 90 SiO2

Para-Xylene content (percent of xylenes)

80

Coke

70 60 50

SiO2 4% MgO

40 3% MgO 30 20 102

2% MgO

H

H

103

104

Equilibrium 105

106

107

at0.3

Relation between the pseudo-Thiele modulus (αt0.3 ) and the observed para-selectivity in STDP for diffusionally modified catalysts with varying activity. (Data from Ref. [68].)

Fig. 9

3182

14.4 Isomerization and Transalkylation of Alkylaromatics

in the reactor (in situ selectivation) or during catalyst manufacture (ex situ selectivation). In the first variant, the MSTDP process, the catalyst is selectivated by coking the catalyst on stream to provide the diffusion barrier discussed above. Following selectivation, the p-xylene selectivity obtained is over 80%. The MSTDP process was introduced commercially in 1990 at EniChem [70]. Commercial experience found that the catalyst was stable, easily air-regenerated, reselectivated and streamed for a nearly identical second cycle, with each cycle lasting about 500 days. The MSTDP technology has since been licensed and used in over nine additional plants. A very attractive feature of the process is that the benzene produced is obtained in over 99.98 wt.% purity and is thus of nitration grade. UOP offers a competing technology for license called the PX-Plus process [71]; however, it is not clear that this technology has been applied commercially. A variant of this technology may also have been developed and commercialized at a location in Asia [72]. In the second variant, the PxMax process, the catalyst is selectivated during the catalyst manufacturing process [73]. Because of the complex interplay between reaction and diffusion, catalyst performance responds to aluminum content, aluminum distribution, morphology and selectivating agent. The first application for PxMax was a retrofit at the ExxonMobil Chalmette petrochemicals plant in 1996, and the first grass-roots application of this technology was at ExxonMobil’s Beaumont refinery in 1997. The catalyst is still in its first cycle and has produced over 200 000 t of products per ton of catalyst over 9 years, demonstrating superior catalyst stability. PxMax is currently used in over seven plants. There are many operational advantages for the PxMax process. The catalyst does not require the complex coke selectivation procedures used for the MSTDP or PX-Plus processes, since the catalyst is selectivated prior to loading. Likewise, it does not require the high temperatures needed for coke selectivation and, hence, grass roots units can utilize reactors with less expensive metallurgy. The p-xylene selectivity and purity (p-xylene as a fraction of the C8 hydrocarbons) is well over 90%, which allows for the use of lower cost crystallizers for high p-xylene recovery at high recovered purity in grass roots units. For retrofits, improved utilization of the sorption units is possible with the high p-xylene purity. The catalyst has high activity and is stable at lower hydrogen to hydrocarbon ratios. This allows the PxMax process to enjoy a temperature advantage of 50 ◦ C over STDP, leading to long run lengths, and results in high p-xylene yield as a fraction of toluene converted. However, when the price of benzene is high, the PxMax process may be operated in the high-severity mode to yield additional benzene.

The STDP process is carried out at 400–470 ◦ C, 14–35 bar, 0.5–5 molar hydrogen to hydrocarbon ratio and a WHSV of 2–5 h−1 C8 aromatic feedstock. Feedstock toluene is extinction-recycled, and the products are benzene, C8 aromatics enriched in p-xylene and ethylbenzene, light paraffins and C9+ aromatics. Per-pass toluene conversion is limited to about 32% if high yields of p-xylene are desired, close to the 36% toluene conversion at equilibrium when p-xylene is made exclusively. At higher conversion, hydrocracking increases as evidenced by higher light gas yields. When paraffins that distill with toluene are present in the feedstock (non-extracted toluene), they crack over the acidic catalyst if they are less branched and hence their diffusion does not sufficiently limit entrance to the zeolite pores. Paraffins such as n-octane and 3-methylheptane diffuse readily into the catalyst and crack at a rate faster than toluene disproportionates, but highly branched paraffins and cycloparaffins do not crack, indicating that they are severely diffusion limited. Bulky, toluene coboiling paraffins, such as 3-ethylhexane, crack at slow rates and hence can build up in the recycle loop. Selective toluene disproportionation is the lowest cost method to produce benzene and p-xylene for those producers with large enough reformers to produce both products on a world scale. Light reformate can be distilled and extracted to produce benzene and some extracted toluene. The benefits of this strategy are that half of the benzene is produced without extraction, reformer xylenes are never distilled overhead, crystallizers can be used instead of sorptive separation and no by-product toluene is produced. The C8+ reformate remains available for p-xylene growth steps using xylene isomerization and C9+ processing. 14.4.6

Ethylbenzene and Cumene Transalkylation

As mentioned in Chapter 14.3, polyisopropylbenzenes and polyethylbenzenes are by-products of the alkylation of benzene with propylene and ethylene, respectively. These heavies represent 5–15% of the per-pass product from cumene and ethylbenzene synthesis. The goal of transalkylation is to convert these heavies, along with incremental benzene, to additional cumene and ethylbenzene, while minimizing the production of byproducts, so that overall yields of above 99% cumene and ethylbenzene are achieved from alkylation. Catalysts for Ethylbenzene and Cumene Transalkylation FAU, BEA, MWW and MOR are all believed to have been developed as transalkylation catalysts. Some comparisons 14.4.6.1

14.4.6 Ethylbenzene and Cumene Transalkylation

of zeolite selectivities are available [74, 75]. Lower cost catalysts are often used for benzene/diethylbenzene transalkylation because selectivity to trace by-products is not among the most important performance features. Selectivity, however, differentiates cumene transalkylation catalysts. The unique morphology and 12-ring pockets of ExxonMobil’s proprietary MWW (Fig. 1) can accelerate the conversion of larger o- and m-diisopropylbenzenes and triisopropylbenzenes since the transalkylation chemistry is believed to occur in the 12-ring surface pockets. ExxonMobil has developed a suite of catalysts for this application that take advantage of the unique performance characteristics of its proprietary zeolites to achieve higher conversion and selectivity in these sterically demanding transalkylations [76]. Other technology providers have developed commercial catalysts with optimized formulations of FAU, BEA or MOR. Catalyst cost is of minimal economic importance because the catalysts produce on the order of 20 000 kg of ethylbenzene or cumene per kilogram of catalyst each cycle and are regenerable. Older alkylation and transalkylation units utilized AlCl3 in ethylbenzene synthesis and SPA in cumene synthesis (solid phosphoric acid or phosphoric acid supported on kieselguhr), but will not be discussed here. Feeds and Products of Ethylbenzene and Cumene Transalkylation The feedstock to the reactor is chemical-grade benzene and polyalkylated aromatic. Treatment of fresh benzene feed with clay, zeolites or other adsorbents is recommended to remove trace components which can poison the zeolite catalyst. The polyethylbenzene and polyisopropylbenzene transalkylation processes occur at sufficiently low temperatures that negligible amounts of cracked (C3− ) products are produced. Losses in both processes are most dependent on process design capability to capture heavies for recycle to the transalkylator. Heavies comprise 0.99 is reached for an averaged hydrogenation rate of 5000 g COD g−1 Pd h−1 ! In 1,3-cyclooctadiene, the double bonds show a parallel orientation. Both are therefore able to interact easily with the Pd surface. Cyclododecadienes, however, show no parallel orientation of double bonds in the state of energy minimum. An approach to a more parallel orientation resulting in stronger adsorption seems possible only 14.10.1.2.4

14.10.1 Selective Hydrogenation of Hydrocarbons

100

100

90

80

A B C D

70 60

Selectivity / %

90

Yield CDE / %

3271

80 70

Pd 263 K Pt 383 K p C H = 6.7 kPa

60

2 2

50 50

40 0

10

40 40

50

60

70

80

90

Cyclododecene yield versus degree of triene + diene conversion [32]. (A) pH2 = 1.5 bar, and stepwise decreased hydrogen pressure; (B) 1.5 bar (X < 30%), 0.3 bar (X > 30%); (C) 1 bar (X < 30%), 0.5 bar (30% < X < 60%), 0.15 bar (X > 60%); (D) 0.5 bar (X < 25%), 0.25 bar (25% < X < 65%) 0.13 bar (X > 65%), pCO = 0.1 bar. Fig. 6

by distortion in the state of the adsorbed molecule as the energy of distortion is compensated by part of the energy of adsorption. Therefore, it is expected that the equilibration of the adsorption layer needs much more time than in the case of cyclooctadiene with parallel double bonds from the start. A decreasing hydrogenation rate favors the equilibration of the adsorption layer with high coverage of the cyclododecadienes [32]. Notwithstanding the very low hydrogenation rate, this selective hydrogenation can be easily carried out on an industrial scale and also in a continuous or batch procedure [32, 34]. Hydrogenation of Ethyne and Other Alkynes Small amounts of ethyne in ethene from naphtha crackers are removed by selective hydrogenation, down to less than 1 ppm. This is because ethyne deactivates the catalyst for ethene polymerization. Therefore, many kinetic studies of ethyne hydrogenation have been carried out for mixtures with high ethene : ethyne ratios. Palladium catalysts show the highest selectivities towards ethene. However, a small amount of ‘‘green oil’’ is formed in parallel with the conversion of ethyne, which is a complex mixture of highly unsaturated C4+ hydrocarbons with an even number of carbon atoms [1]. Moreover, benzene is formed in small amounts. The kinetics of ethyne hydrogenation on Pd catalysts are very similar to those of 1,3-butadiene. Reaction orders in hydrogen reported in the literature lie between 1.0 and 1.6 and ethyne reaction orders range from −0.5 to 0 [35–39]. 14.10.1.3

30

40

50

p H2 / kPa

100

Conversion (CDT + CDD) / %

20

Dependence of selectivity with respect to ethane on hydrogen pressure for ethyne hydrogenation over Pd and Pt (adapted from Ref. [12]).

Fig. 7

Very often, the combination of 1 (hydrogen) and 0 (ethyne) is found, but a combination of reaction orders greater than 1 (hydrogen) and smaller than 0 (ethyne) have also often been observed [39]. The interpretation of these reaction orders is the same as given for the hydrogenation of dienes in Section 14.10.1.2.2. Whereas in the hydrogenation of 1,3-butadiene or other alkadienes and of 1-butyne on supported Pd catalysts the selectivities with respect to alkenes are 100% provided that the diene exceeds a certain concentration, a small fraction of ethane is inevitably formed in the hydrogenation of ethyne. The ethane fraction increases with increasing hydrogen pressure and decreases with increasing temperature [12] (Fig. 7). Margitfalvi et al. [35] have shown by use of 14 C2 H2 and 14 C H that on Pd black ethane is exclusively formed from 2 4 ethyne provided that ethyne is present. This was shown also for the hydrogenation of ethyne on Pt by Arafa and Webb [40]. These results reveal that on Pd and Pt catalysts a strong thermodynamic factor prevents the readsorption of ethene and, hence, the consecutive hydrogenation of readsorbed ethene. The interpretation of the kinetics and of the properties of catalysts is made more difficult by carbonaceous deposits formed in the starting period of the hydrogenation, which cause a decrease in activity in the first hours on-stream [1, 36]. Also, the selectivity changes drastically. In the case of ethyne hydrogenation on Pd, the selectivity towards ethene usually increases during this period. Furthermore, in gas-phase hydrogenation, as usually applied in the hydrogenation of C2 H2 −C2 H4 mixtures, high-boiling fractions of green oil accumulate in the References see page 3282

3272

14.10 Hydrogenation Reactions

(C2H4)g

Green oil (C2H2)g

HC

CH

+H∗

+H∗ HC

∗

∗

H2C

CH2

∗

(C2H2)∗

CH2 ∗

+H∗ +H∗ (C2H)∗

∗ ∗ ∗

C

CH3

+H∗

∗ ∗

CH

CH3

+H∗

∗

CH2

CH3

+H∗

(C2H6)g Carbonaceous residues Scheme 4

pores of the catalyst pellet [37, 41, 42], so that periodical regeneration of the catalyst is necessary. Scheme 4 is based on considerations by Margitfalvi et al. [43]. However, as there is no experimental proof of most intermediates, such a reaction scheme should be treated with caution. The upper reaction sequence shows the stepwise hydrogenation of ethyne to adsorbed ethene, which finally desorbs. Chemisorbed ethyne or the semihydrogenated ethyne are expected as the starting compounds of dimerization and oligomerization to form the so-called green oil. The formation of carbonaceous residues may occur via intermediates of low hydrogen content which are formed by hydrogen abstraction [44]. Furthermore, the formation of ethane is discussed in terms of stepwise hydrogenation of ethylidyne (see Section 14.10.1.5) because the route via consecutive ethene hydrogenation has been excluded. In the selective hydrogenation of 1-butyne [20, 21] only 1-butene is formed. As expected, 2-butyne is mainly hydrogenated to cis-2-butene for the same reason as the cis preference in the hydrogenation of 1,2-butadiene. The hydrogenation of vinylacetylene over supported palladium proceeds consecutively through the formation of 1,3-butadiene and a mixture of 1-butene and cis- and trans-2-butene and then butane [45, 46]. An important feature of vinylacetylene hydrogenation is the formation of a vinylacetylene−palladium complex which is soluble in hydrocarbons. Therefore, at low temperatures palladium is eluted from the catalyst [1]. Numerous metal additives are applied to improve the catalysts for the hydrogenation of alkynic compounds. The Lindlar catalyst [47], a lead-promoted palladium on

calcium carbonate, is well known for the selective reduction of triple bonds to cis-configured double bonds. The partial selective catalytic hydrogenation of alkynes, as used for the preparation of sex pheromones over the Lindlar catalyst in the presence of quinoline, generally gives high yields of the corresponding cis-alkene (>98% cis-isomer) [48]. A detailed characterization of the Lindlar catalyst has been given by Schl¨ogl et al. [49]. Effect of Carbon Monoxide in Selective Hydrogenation The addition of small amounts of carbon monoxide to the mixture of reactants strongly reduces the consecutive hydrogenation of the alkenes produced. This has been clearly demonstrated by Weiss et al. [50] (Fig. 8) for the selective hydrogenation of a simulated tail-end mixture (0.3% C2 H2 , 0.4% H2 being balanced with C2 H4 ) on a 0.04 wt.% Pd/γ -Al2 O3 catalyst. The selectivities with respect to ethene and ethane are defined by the ratio of rates of formation of ethene and ethane, respectively, and the rate of ethyne consumption. At about 50 ppm CO, a marked increase in SC2 H4 is observed without a significant reduction in ethyne conversion. The authors interpreted this CO effect as the blocking of sites for ethene hydrogenation. However, Al-Ammar and Webb [51, 52] attributed this effect to a competition between CO and hydrogen, thus lowering the hydrogen concentration at the surface. This effect of CO is also applied in industrial processes for the selective hydrogenation of 1,3-butadiene [53, 54]. 14.10.1.4

14.10.1 Selective Hydrogenation of Hydrocarbons

C C2H2

100

S C2H6

Selectivity / %

S C2H4 80

80

60

60

40

40

20

Conversion / %

100

3273

20

0 0

500

1000

0 5000 1000015000

1500

pCO / ppm Dependence of ethyne conversion and selectivity towards ethane and ethane on carbon monoxide pressure in ethyne hydrogenation (adapted from Ref. [50]).

Fig. 8

1.0

0.8

k / k CO = 0

The influence of CO has been studied for the hydrogenation of 1,3-cyclooctadiene and also for the hydrogenation of cyclooctene in the absence of the diene [55]. The dependence of the relative rates (referenced to the experiment without CO addition) on the CO content in the feed shows that the rate of cyclooctene hydrogenation is more strongly affected than the rate of 1,3-cyclooctadiene hydrogenation (Fig. 9). This result suggests a competition between CO and hydrogen and also between CO and cyclooctene. The effect of CO on cyclooctadiene hydrogenation decreases as the temperature is raised, whereas CO still strongly affects the hydrogenation rate of cyclooctene.

0.6

0.4

k 1 COD k 2 COE

0.2

0.0 0.0

14.10.1.5

Catalysts for Selective Hydrogenation

14.10.1.5.1 General Aspects of Catalyst Selection Of the Group VIII noble metals, palladium shows by far the highest selectivity with respect to (cyclo-)alkenes in (cyclo-)alkadiene and alkyne hydrogenation. Because of the relatively high reaction rate and therefore the strong influence of mass transfer processes on the selectivity catalysts of the egg-shell type with a very thin active layer (20–100 μm) of palladium on alumina are generally used. Knitted metal fabrics as thin-layer catalysts with active layers of L < 0.01 μm have been introduced as improved alternatives [56]. Additives, for example Ag, are applied in order to suppress or reduce undesired side-reactions such as the green oil formation in the selective hydrogenation of acetylene [12]. A detailed study on the morphology of Pd deposits on alumina has been carried out by Freund and co-workers [57]. Figure 10 shows an STM image of Pd grown on Al2 O3 /NiAl (110) at 300 K. The particles adopt

COE COA

0.2

0.4

0.6

0.8

p CO / Pa Influence of carbon monoxide on the rate constants in COD and COE hydrogenation [55].

Fig. 9

the form of small truncated cubo-octahedron crystals limited mainly by planes with a {111} orientation (and a smaller fraction of {100} planes). 14.10.1.5.2 Comparison of Pd and Pt Catalysts 1,3Butadiene is hydrogenated on supported Pd catalysts with a selectivity of 100%, whereas on supported Pt the hydrogenation under comparable conditions leads to a marked fraction of butane [13, 58]. The same differences in selectivity have been observed for 1,3-butadiene hydrogenation on samples of Pd{111}, Pd{110}, Pt{111}, Pt{110} and Pt{100} [59–61]. Therefore, it is of interest References see page 3282

3274

14.10 Hydrogenation Reactions Dependence of adsorption energy (kJ mol−1 ) of di-σ -1-butene on coverage  (adapted from Ref. [62])

Tab. 1



Pd{111} Pt{111}

STM 500 × 500 nm

STM image of Pd grown on Al2 O3 /NiAl (110) at 300 K and the model of an idealized crystal (adapted from Ref. [57]).

Fig. 10

to compare these catalysts in order to find the factors which are responsible for the exceptionally high selectivity obtained when Pd catalysts are employed. The selectivity of alkadiene hydrogenation to alkenes is determined by the ratio of the rates of alkene desorption and the consecutive hydrogenation, provided that the surface is densely covered by diene or alkyne. We may assume that the differences in selectivity on Pd and on Pt depend on different rates of desorption. The hypothesis that the rate of alkene desorption determines the selectivity finds support by the strong increase in alkene selectivity in the hydrogenation of 1,3-butadiene on Pt as the temperature is raised, shown in Fig. 11 [12]. The same dependence was found for the hydrogenation of ethyne. A comparative density functional theory (DFT) study of the adsorption of 1,3-butadiene and butene isomers on the Pt{111} and Pd{111} surface by Valc´arcel et al. [62] has shown that the adsorption of butenes on the Pd{111} surface compared with Pt{111} shows a smaller 100

Selectivity / %

90 80 70 Pd p C4H6 = 6.7 kPa Pt p H2 = 13.3 kPa

60 50 40 300

350

400

450

500

550

600

Temperature / K Fig. 11 Temperature dependence of selectivity in 1,3-butadiene hydrogenation over palladium and platinum (adapted from Ref. [12]).

1/3

1/4

1/6

1/9

−25 −45

−57 −83

−70 −91

−87 −93

molecular distortion upon di-σ -adsorption, which means an increased tendency towards π-adsorption. Also for 1,3-butadiene on Pd{111} there is a marked tendency for di-π-adsorption with less molecular distortion. Furthermore, the calculations have shown that the difference in the strength of 1,3-butadiene and 1-butene adsorption is nearly equal for Pd and Pt so that the higher selectivity of 1,3-butadiene hydrogenation towards butene on Pd compared with Pt cannot be explained by this means. However, the results offer another interpretation if the rate differences of desorption and consecutive hydrogenation are regarded as crucial for the differences in selectivity. The DFT study gave adsorption energies of di-σ -butenes which are markedly lower for Pd{111} than for Pt{111}. This difference is even more pronounced if π-bonded 1-butene is considered. Furthermore, the adsorption energy of di-σ -1-butene shows for Pd a stronger decrease with increasing coverage , as demonstrated in Table 1. It should be noted that the crucial arrangement in the selective hydrogenation of 1,3-butadiene is an adsorbed butene molecule, just formed, which is densely surrounded by butadiene molecules. Considering the above-mentioned adsorption energies of the DFT study and in particular their dependence on the coverage, an enhanced rate of desorption of butenes on Pd compared with Pt can be expected. In this context, DFT studies of ethene adsorption on Pd {111} by Neurock et al. [63] are of interest. They showed that at higher coverage lateral repulsive interactions between adsorbates destabilize the di-σ -adsorption and result in a less strongly bound π-adsorption. The preference for π-coordination of the intermediate alkene on Pd may favor its desorption, while the stronger di-σ -adsorption on Pt increases the chance of subsequent hydrogenation. Numerous studies have led to the assumption that π-coordination is generally favored on palladium [64–70]. Recent studies by in situ surface-sensitive methods operating in a pressure range from UHV to ambient conditions have contributed to the knowledge of surface species in the hydrogenation of unsaturated hydrocarbons. In a review by Rupprechter [71], recent studies of surface/vibrational

3275

14.10.1 Selective Hydrogenation of Hydrocarbons

14.10.1.5.3 Particle Size Effect Selective hydrogenations are usually carried out by Pd dispersed over silica and alumina supports with an average particle size of about 2–5 nm. However, the same high selectivity is also obtained with Pd black or electrolytically deposited Pd crystallites with a particle size of about 25 nm [14]. Chou and Vannice [78] evaluated the adsorption enthalpies of H2 and CO on Pd for various supports and found these to be independent of the particle size down to 3 nm. Below this diameter, both adsorption enthalpies rise rapidly with decreasing size. Boitiaux et al. [79] observed a decrease in the turnover frequency (TOF) of the hydrogenation of 1,3-butadiene and of 1-butyne on supported palladium by a factor of 12 when the degree of metal dispersion was increased from 20 to 80%. The degree of dispersion is defined as the number of surface metal atoms divided by the total number of metal atoms in the particle. These studies were confirmed by Tardy et al. [80] for 1,3-butadiene hydrogenation on Pd/C. Infrared studies of

100 1,3-butadiene

80

trans -2-butene

y/%

60 40

cis -2-butene

1-butene 20

n -butane

0 0

400

800

1200

t / min

(a) 100 1,3-butadiene

80

y/%

spectroscopy and DFT calculations are discussed in detail. These studies concern the hydrogenation of ethene mainly on Pt{111} and Pd{111} single-crystal surfaces and on model catalysts of well-defined nanoparticles on Al2 O3 . The following species are considered: π-ethylene, di-σ ethylene and ethylidyne. IR/Vis sum frequency generation (SFG) indicated that on Pt{111} both ethylidyne and di-σ -bonded ethene were present during hydrogenation (10–20 mbar ethene, 40–120 mbar hydrogen, 300 K) [72]. However, these strongly bound species had to be considered as spectator molecules because the reaction rate could not correlated with the relative concentrations of these species. The same conclusion was drawn by Beebe and coworkers [73, 74]. Therefore, π-bonded ethylene and ethyl were suggested as active intermediates even though they could not be unambiguously identified. SFG studies of ethylene adsorption and hydrogenation on Pd{111} have revealed that compared with Pt{111}, ethylene shows a smaller tendency to produce ethylidyne on Pd{111} [75]. However, under reaction conditions on Pd{111} and on Pd nanoparticles on Al2 O3 , no adsorbed ethylene species could be identified [76, 77]. Unfortunately, π-bonded ethylene shows almost no signal due to its parallel orientation to the surface. Nevertheless, the authors again considered π-bonded ethylene to be the most likely active species. The conclusion drawn from this comparison is that there may be some starting points to find an explanation for the outstanding properties of palladium, for example a higher rate of alkene desorption compared with platinum. However, the development of a sound theory of selective hydrogenation still remains a difficult task.

trans -2-butene

60 40

n -butane

1-butene

cis -2-butene

20 0 0 (b)

400

800

1200

t / min

Product distribution y versus reaction time t for 1,3-butadiene hydrogenation at 373 K on Pd−Al2 O3 model catalysts with mean particle diameter (a) 2 and (b) 8 nm [25].

Fig. 12

CO adsorption showed a shift of the bands towards lower frequencies as the dispersion increased, thus indicating a strengthening of the metal−CO bond [81]. Silvestre-Albero et al. [25] reported atmospheric pressure kinetic studies of 1,3-butadiene hydrogenation on well-defined Pd/Al2 O3 model catalysts for a wide range of mean particle sizes from 2 to 8 nm. An exact characterization of the Pd nanoparticle morphology and surface structure was obtained by scanning tunneling microscopy (STM) and a number of surface-spectroscopic techniques under ultra-high vacuum. Furthermore, reaction rates on Pd{111} and Pd{110} single crystals were measured for comparison [82]. Figure 12 shows the product distribution versus the reaction time for the catalysts of mean particle diameter 2 and 8 nm, both samples showing about (1 ± 0.2) × 1015 surface Pd atoms. In both cases a 100% butene selectivity is obtained even at a very low concentration of butadiene. For particles of 8 nm, the rate of butadiene hydrogenation at a partial pressure of 5 mbar is about 10-fold higher than for 2 nm. For the latter, the reaction rate increases markedly with decreasing partial pressure of butadiene, indicating a negative reaction order of References see page 3282

3276

14.10 Hydrogenation Reactions

butadiene while the reaction order is zero for catalysts of 8 nm particle. The TOF increases linearly with increasing particle size when the total number of surface Pd atoms is considered. However, based on the kinetic measurements for a Pd{111} single crystal and based on the STM characterization of the Pd particles, the authors showed that the TOF is independent of the particle size when the TOF is scaled only to the number of Pd atoms in incomplete {111} terraces. With decreasing particle size, the fraction of {111} terraces decreases in favor of higher index sites (and edges), which obviously show a stronger adsorption of butadiene and consequently a lower hydrogenation rate due to a decreased activation of hydrogen. Furthermore, the diagrams show that the rate of isomerization and hydrogenation of butenes after complete consumption of butadiene is independent of the particle size. A stronger bonding of 1,3-butadiene with decreasing particle size (2.8–1.2 nm) was also suggested by Bertolini et al. [83]. Role of Mass Transfer in Selective Hydrogenation For the gas- and liquid-phase hydrogenation of 1,3cyclooctadiene, which can be regarded as representative for selective hydrogenations of dienes and alkynes, the resulting selectivity of the intermediate 1,3-cyclooctene is calculated as a result of the interdependence of reaction and mass transfer [84–86]. 14.10.1.6

14.10.1.6.1 Gas-Phase Hydrogenation Even if catalysts of the egg-shell type with a very thin active layer of palladium on alumina are used for selective hydrogenations, mass transfer in the pores of this layer strongly influences the selectivity because of the high reaction rate of all hydrogenations considered. The study concerns the consecutive hydrogenation of 1,3-cyclooctadiene: r1

1,3-cyclooctadiene −−−→ COD

cyclooctene COE

The intrinsic reaction rate r1 is approximately first order with respect to hydrogen and independent of the partial pressures of COD and COE: r1 =

(8)

The intrinsic reaction rate r2 is represented by a Langmuir−Hinshelwood rate expression: r2 =

k2 KCOE pCOE CH2 1 + KCOE pCOE + KCOD pCOD

d2 (CCOD /CCOD.s ) = ϕ2 d(x/L)2 with the Thiele modulus  k1 cH2 ϕ=L Deff cCOD.s

(9)

(10)

(11)

The boundary conditions are:   cCOD (i) =1 cCOD.s x/L=0

(12)

(ii) The gradient of cCOD at x = 0 is obtained from the steady-state condition that the flow of COD into the active layer is equal to the amount of converted COD:   dCCOD = x  k1 CH2 (13) Deff dx x=0 corresponding to   x d(cCOD /cCOD.s ) = − ϕ2 d(x/L) L x/L=0

(14)

where x  /L is the position in the active layer where cCOD = 0. Therefore, the active layer is covered by COD in the range 0 < x/L < x  /L. Integration leads to x  x 2 1 2  x 2 cCOD =1− ϕ + ϕ cCOD.s LL 2 L

r2

−−−→ cyclooctane COA

1 k1 CH C0 2 COD

These reaction rates are related to the volume of the palladium-containing layer. The ratio of the adsorption constants is KCOD /KCOE > 104 , so that in the presence of COD the reaction rate r2 is almost zero. The concentrations of the components in the pores of the active layer are calculated from the mass balance. Since the diffusion coefficient of hydrogen is about 10 times larger than that of the C8 components, a nearly constant concentration of hydrogen in the active layer is assumed. Thus, for the calculations the hydrogen concentration of the gas phase is substituted. With this simplification, the system is described only by the mass balance equation of COD and the corresponding boundary conditions where the coordinate along the active pores is x, with x = 0 at the outer rim and x = L at the inner rim:

and hence √ 2 x = L ϕ

(15)

(16)

In Fig. 13, calculated concentration profiles along the active pores are shown. The apparent reaction rates which are related to the mass of the catalyst are given by COD −−−→ COE: r1∗ =

x  VPd 1 pH2 k1 L VP ρP RT

(17)

14.10.1 Selective Hydrogenation of Hydrocarbons

4

r * / mol h−1 kgcat.−1

1.00

0.75

C / CS

3277

p H2 = 50 Pa

0.50

0.25

r1* r2* L = 7 × 10−5 m T = 356 K p COD = 150 Pa p COE = 1500 Pa

2

1

p H2 = 75 Pa

p H2 = 150 Pa COD hydrogenation

3

0

COE hydrogenation

0

0.00 0.0

0.2

0.4

0.6

0.8

100

1.0

200

300

400

500

600

p H2 / Pa

x/L Dimensionless concentration profiles across the active shell for different hydrogen pressures [14]. Fig. 13

  x  VPd 1 COE −−−→ COA: r2∗ = 1 − k2 L VP ρP KCOE pCOE pH2 RT + KCOE PCOE RT

(18)

Hydrogenation of COE is only possible by readsorption of COE on that fraction of the active layer which is not covered by COD, i.e. The fraction (1 − x  /L), thus COD and COE hydrogenation take place in different zones of the active layer. These are indicated in Fig. 13. Two cases can be distinguished: x cCOD ≥ 0 at = 1, cCOD.s L

ϕ2 ≤ 2

(19)

The palladium-containing layer is completely covered by COD (x  /L = 1). Therefore the apparent reaction rates are r1∗ =

VPd 1 pH2 k1 VP ρP RT

1 VPd 1 1 √ 2Deff k1 pCOD.s pH2 (22) L VP ρP RT   KCOE pCOE pH2 VPd 1 x r2∗ = 1− k2 (23) VP ρP L RT + KCOE PCOE RT √ with x  /L = 2/ϕ. The calculated apparent reaction rates are shown in Fig. 14 as a function of the hydrogen pressure. The onset point of consecutive COE hydrogenation corresponds to the condition ϕ 2 = 2 (cCOD = 0 at x = L) and divides the diagram into two parts. The left part is characterized by ϕ 2 ≤ 2 and 100% COE selectivity and the right part by ϕ 2 > 2 and reduced COE selectivity. This model was validated by experiments in which all parameters were varied, namely L, k1 (via Pd loading), Deff (via the dependence on the total pressure) and finally pCOD and pH2 [9, 83, 84]. r1∗ =

where VPd /VP is the volume of the Pd-containing layer related to the volume of the pellet and ρP is the density of the pellet.

(i)

Fig. 14 Dependence of the apparent reaction rates for COD and COE hydrogenation on partial pressure.

(20)

r2∗ = 0 corresponding to 100% selectivity with respect to COE. x cCOD (ii) = 0 at 0 < (21) < 1, ϕ 2 > 2 cCOD.s L The active layer is only covered by COD in the range 0 < x/L < x  /L. The apparent reaction rates are

The condition φ 2 = (1/L2 )(Deff /k1 )(pH2 /pCOD,s ) < 2 for 100% selectivity is the key for the successful design of an egg-shell catalyst. For complete COD conversion the ratio pH2 /pCOD,s must be chosen such that a minimal excess of hydrogen is guaranteed. All other parameters can be adjusted by appropriate preparation of the catalyst. For example, k1 is approximately proportional to the Pd content of the active layer. Even Deff can be adjusted to a certain extent by the total pressure and the pore diameter, so that mass transfer is less determined by Knudsen diffusion. The limiting case is very sensitive to the thickness of the active layer, L. Therefore, a highly selective catalyst must have a uniform thickness of the Pdcovered layer. If Pd is deposited in deeper regions of the References see page 3282

3278

14.10 Hydrogenation Reactions

pellet, for example in narrow fissures, then mass transfer limitation resulting in cycloalkene hydrogenation occurs for a small fraction of Pd at much lower pH2 /pCOD,s ratio than for the main part of the catalyst. It is a difficult task to avoid such uneven patches in the active layer. The development of catalysts providing a very high selectivity depends strongly on improved preparative methods for such ideal egg-shell catalysts [22, 23, 55, 113]. If the thickness of the active layer, L, is obtained by microscopy, the experiment r1∗ = f(pH2 ) at constant pCOD,s ≈ pCOD can serve as appropriate method for the determination of Deff on the basis of the condition for the onset point of consecutive COE hydrogenation. ϕCOD 2 = 2 −−−→ Deff =

∗ pH L2 2 k1 2 pCOD

(24)

∗ is the hydrogen pressure at the onset point where pH 2 and k1 is obtained from the plot r1∗ = f (pH2 ) in the range ϕ 2 < 2.

Liquid-Phase Hydrogenation This section follows a study by Wuchter [9] on the liquid-phase hydrogenation of COD. The liquid-phase hydrogenation differs from the gas-phase hydrogenation first of all by a lower effective diffusivity in the liquid filled pores of the catalyst. With regard to the hydrocarbon, the low diffusivity is usually compensated by a much higher concentration and consequently higher gradients of concentration. However, the Thiele modulus with regard to hydrogen is much higher in the case of liquidphase hydrogenation because of the low diffusivity in the liquid-filled pores. The drastic fall of the hydrogen concentration towards the pellet center in this case favors the selectivity towards the intermediate products. Usually the concentration of hydrogen in the liquid phase is of the same order of magnitude as in the case of gasphase hydrogenation. Such concentrations correspond to a hydrogen pressure of a few bar. For the intrinsic reaction rate r1 we can use the same expression as for gas-phase hydrogenation that is approximately first order with respect to hydrogen and independent of the partial pressures of cyclooctadiene and cyclooctene: 14.10.1.6.2

r1 = kCH2

(25)

In contrast to the gas-phase hydrogenation, the mass balance equation of hydrogen must also be taken into account, so that the liquid-phase hydrogenation is represented by the following set of mass balance equations and the corresponding boundary conditions: k1 d2 CH2 = C 0 C 1 ; mass balance H2 2 dx DH2 ,eff COD H2

(26)

d2 CCOD k1 = C 0 C 1 ; mass balance COD dx 2 DCOD,eff COD H2 (27) with the boundary conditions   dCH2 =0 CH2 (x = 0) = CH2 ,s dx x=x    dCCOD =0 CCOD (x = 0) = CCOD,s dx x=x 

(28) (29)

The boundary condition for hydrogen at x = x  is based on the simplification that the consumption of hydrogen by hydrogenation of cyclooctene in the back part of the active layer x  < x < L is negligible. Otherwise the evaluation of x  where CCOD (x  ) = 0 needs a rather complicated iteration. Since the rate of cyclooctene hydrogenation is lower than the rate of diene hydrogenation and reaction states of a relatively high rate of cyclooctene formation are not of interest, this simplification is justified particularly because it renders possible an algebraic solution. The solutions are as follows:

  k1 cosh (x  − x) DH2 ,eff

 (30) CH2 (x) = CH2 ,s k1  cosh x DH2 ,eff CCOD (x) = CH2 ,s

DH2 ,eff DCOD,eff 

· (x 

− x) DH2 ,eff

 k1 cosh · x DH2 ,eff

cosh ·

k1

+ CCOD,S − CH2 ,s

DH2,eff DCOD,eff

with CCOD (x  ) = 0 it follows that ⎛  ⎜ DH2 ,eff 1 x = arccosh ⎜ ⎝ CCOD,s DCOD,eff k1 1− CH2 ,s DH2 ,s

(31)

⎞ ⎟ ⎟ (32) ⎠

The calculation of the effective reaction rate, taking into account the influence of mass transfer from the liquid phase to the outer surface of the catalyst pellet, and finally the numerical integration of the material balance of a batch experiment:  CCOD  dCCOD =t (33) 0 rA,COD CCOD

14.10.1 Selective Hydrogenation of Hydrocarbons

Fraction C8 / %

100

3279

20

80

CCOD = 0.53 mol l−1 V R = 0.45 l m cat = 760 mg T = 353 K p H2 = 5 bar

60 40

15 10 5

20 0 0

1

2

3

4

5

6

2

Time / h

3

4

5

Time / h

Fraction of COD (), COE () and COA (◦) in the hydrogenation of 1,3-cyclooctadiene versus reaction time at constant hydrogen pressure (5 bar).

Fig. 15

give the concentration of COD and approximately also that of COE as a function of time. The parameter of the model k1 and DH2 ,eff are obtained by fitting the model to experimental data. Figure 15 shows the concentration curves of COD and COE for a batch experiment at constant hydrogen pressure. For this batch experiment and the gas-phase hydrogenation carried out for comparison, an egg-shell catalyst of 0.1% Pd on alumina with an active layer of L ≈ 20 μm and a pellet diameter of 2–3 mm were employed. The estimated reaction rate constant k1,li ≈ 90 s−1 is of the same order of magnitude as the rate constant k1,g ≈ 200 s−1 determined for the gas-phase hydrogenation. Taking into account that the COD hydrogenation is of slightly negative order (∼ −0.15) with respect to COD, an extrapolation to the saturation pressure of COD would give k1,g ≈ 130 s−1 . Therefore, the conclusion is drawn that in both cases, liquid- and gas-phase hydrogenation, the access of hydrogen through the dense adsorption layer to the catalyst surface and the subsequent dissociation can be regarded as the rate-determining step of the reaction. Partial Hydrogenation of Benzene to Cyclohexene Benzene is hydrogenated to cyclohexane on an industrial scale. In most processes, platinum catalysts are employed. Minor amounts of cyclohexene are obtained if the degree of benzene conversion is kept below 1%. In 1963, Hartog and Zwietering [87] investigated the liquid-phase hydrogenation of benzene and found that ruthenium is the only catalyst of Group VIII metals on which at least a small yield of cyclohexene of about 0.1% at 20% benzene conversion was obtained. The cyclohexene selectivity could be markedly increased when an alcohol, such as methanol or butanol, was added to the liquid benzene [88, 89]. However, the most important step for 14.10.1.7

improved selectivity was to carry out the hydrogenation in an agitated two-liquid-phase system. The two phases were benzene−cyclohexene and water, to which salts of transition metals were added [90]. Several companies have developed processes for the production of cyclohexene from benzene [91–94]. At present, the best result is a cyclohexene selectivity of 70% at a degree of benzene conversion of 85% [95]. A tremendous number of catalysts have been tested. It turned out that supported ruthenium catalysts are most suitable. These catalysts are usually prepared by impregnation of the carrier with an aqueous ruthenium chloride solution. Oxides of lanthanides such as La2 O3 were found to be particularly suitable [92, 96]. A special type of supported ruthenium catalyst is prepared by mixing a dissolved ruthenium chloride−alcohol complex and tetrahydroxysilane or aluminum tri-sec-butoxide (sol−gel process). Very small ruthenium crystallites with a diameter below 2 nm are obtained. The exceptional feature of this catalyst is that addition of salts to the aqueous phase is not necessary [97–99]. Most other supported catalysts are promoted by salts of transition metals such as Ni, Fe, Co and Zn, which are added directly by impregnation or coprecipitation or added to the aqueous phase [90, 92, 94, 100]. In order to reach reasonable selectivities to cyclohexene, the catalyst particles must be suspended in the aqueous phase [101, 102]. Hydrogen and benzene are dispersed as bubbles and droplets. The solubility of benzene in water is about eight times higher than that of cyclohexene. Therefore, it is assumed that water enhances the removal of cyclohexene from the catalyst surface. Furthermore, the access of hydrogen to the catalyst particles is moderated because of the low solubility of hydrogen in water. The diffusion of hydrogen through the adherent water References see page 3282

3280

14.10 Hydrogenation Reactions

layer of the particles is regarded as the rate-determining step [102]. Access of benzene and fast removal of cyclohexene from the catalyst surface are determined by the specific structure of the hydrogen bond network in the adherent layer. It is assumed that this structure depends on the cations which are adsorbed on the surface and, hence, make the ruthenium particles more hydrophilic [100, 102]. A particular effect was found by Nagahara and Konishi [103], who performed the hydrogenation in a system containing supported ruthenium particles suspended in an aqueous salt solution to which fine particles of oxides, such as oxides of Zr, Ti, Nb and Ta, were added. Cyclohexene yields of about 60% were achieved at benzene conversions of 80–90% [103, 104]. To date, kinetic studies are scarce because of the complexity of the multiphase reaction. In the work by Struijk et al. [102], a reaction order of unity with respect to hydrogen was determined. This result is in accord with the assumption that the diffusion of hydrogen through the adherent water layer of particles is the rate-determining step. The question as to whether benzene is converted following a stepwise Horiuti−Polanyi mechanism or whether there exists also a direct route from benzene to cyclohexane [101, 105] is controversially discussed. Recently, it was shown for the system Ru/La2 O3 suspended in aqueous ZnCl2 solution that the selectivity to cyclohexene approaches unity with decreasing degree of benzene conversion, so that a direct route from benzene to cyclohexane can be excluded [106]. The first and until now the only plant for partial hydrogenation of benzene to cyclohexene (60 000 t a−1 ) was constructed by Asahi Chemical Industry in Japan in 1990. In this process, benzene is converted into cyclohexene using a ruthenium−zinc catalyst. The aqueous phase contains salts and oxides as additives. The partial hydrogenation of benzene is carried out at 420–450 K and 5–7 MPa in an agitated reactor or in a series of agitated reactors [107]. The reactor contains a thoroughly dispersed organic−aqueous−catalyst mixture through which hydrogen is blown. In order to separate the multiphase reactor liquid into an organic phase and an aqueous phase, a stationary zone inside the reactor is provided or the separation occurs in a decanter. Some research work has been undertaken in order to substitute the four-phase procedure (gas, aqueous liquid, non-aqueous liquid, solid catalyst) by a simpler one. For example, D¨obert and Gaube [106] and Gescheidle [108] showed that the presence of a non aqueous phase is not crucial to achieve high selectivity. Gaseous benzene is blown through an aqueous suspension of the catalyst, e.g. ruthenium supported on La2 O3 . The gaseous product flow leaving the reactor contains cyclohexene, cyclohexane

and unconverted benzene. It is not necessary to separate the catalyst unless the deactivated catalyst has to be removed and replaced by fresh one. A series of tank reactors can be easily arranged in order to realize a high degree of benzene conversion and to approach the performance of a tubular reactor that is characterized by the highest selectivity with respect to the desired intermediate product cyclohexene. The simplest procedure would be the selective gas-phase hydrogenation of benzene. Unfortunately, selectivities obtained with this procedure are very low. On coated Ru with methanol as modifier, selectivities up to 45% at conversion degrees of 5%, i.e. a yield of 2.3%, and on an Ru/SiO2 sol−gel catalyst with water as modifier, selectivities up to 10% at a conversion degree of 30%, i.e. a yield of 3%, are observed [109]. Obviously an adsorption layer of the modifier is not sufficient to achieve a high selectivity of cyclohexene. Selective Hydrogenation Processes in the Downstream Treatment of Naphtha Cracking Several selective hydrogenation processes are constituent elements of the downstream treatment of naphtha cracking, as shown in Fig. 16. In the C2 fraction, ethyne is removed by gas-phase hydrogenation. Propyne and propadiene in the C3 fraction are preferably converted in a liquid-phase hydrogenation. If 1,3-butadiene is produced, the first step in C4 processing is the extraction of butadiene using polar solvents. In the course of butadiene separation and purification, vinylethyne, 1-butyne and 2-butyne are hydrogenated to butenes. The next step is the recovery of isobutene, for example by reaction with methanol forming MTBE, a valuable gasoline additive. The raffinate contains more than 70% butenes and still small amounts of butadiene which have to be removed by selective hydrogenation in the liquid phase. With the rapid increase in ethene and propene production, butadiene has become a surplus product. In order to adjust the excess butadiene supply to the demand, processes have been developed for the selective hydrogenation of the crude C4 cut to convert butadiene to butenes [110]. Many alternative processes have been developed, depending on the specific fractionation concepts of the reaction product of naphtha cracking. A comprehensive review on hydrotreatment was presented by Derrien [1]. Contributions published by groups at IFP [111], BASF [22], Degussa [23], CDTECH [112] and Leuna Werke [2] give excellent information about new processes and trends. In the following, descriptions of some selected processes are given. The most common method of C2 H2 removal is the so-called tail-end hydrogenation of the C2 fraction containing only C2 H4 , C2 H6 and 0.5–2 vol.% C2 H2 [22]. 14.10.1.8

14.10.1 Selective Hydrogenation of Hydrocarbons

H2, CH4, CO

C2

C3

Hydrogenation

H2, CH4

Selective hydrogenation

C2H4, C 2H6

Selective hydrogenation

C3H6, C 3H8

3281

Cracking Selective hydrogenation C4 Butenes Butadiene extraction Selective hydrogenation C5+

Selective hydrogenation Full hydrogenation

Fig. 16

MTBE extraction

Selective hydrogenation

1,3 –Butadiene

Aromatics gasoline

Scheme of downstream treatment of steam cracker effluents.

The selective hydrogenation for an approximate C2 H2 content of 1.5 vol.% is usually carried out in an adiabatic reactor with two separate catalyst beds and intercooling. The hydrogen is added stepwise. The reaction temperature is kept below 373 K and a H2 : C2 H2 molar ratio of 1.5 or lower is maintained to achieve a high selectivity towards ethene. To meet the requirements of a drastic reduction in the C2 H2 content ( catalyst structure (morphology, metal location) > reaction conditions. While catalyst activity can with some precaution be related to metal surface area and metal accessibility, selectivity is usually not correlated with easily measurable parameters [13]. For a specific hydrogenation reaction, most important for catalyst selectivity are the intrinsic selectivity properties of the metals. These selectivities, however, can be changed significantly by the use of modifiers (inhibitors, mediators, promoters). Examples of some remarkable effects are given in Section 14.10.2.5 (see also Chapter 14.16.1). Although there is no clear-cut strategy available for selecting a modifier for a specific selectivity problem, the in-process modification of metallic surfaces with organic modifiers uses concepts familiar to the chemist educated in organic chemistry. An interesting approach to transferring organometallic mechanistic principles to surface chemistry has been described by Augustine and O’Hagan [14]. Other concepts such as molecular recognition, proximity and haptophilicity are sometimes fruitful. The solvent has a very significant effect on both the activity and selectivity of catalytic hydrogenation [15, 16]. In some cases, the solvent effect can be interpreted as some sort of surface modification. The proper choice of the reaction conditions for a heterogeneous catalytic hydrogenation is more complicated than for normal organic reactions because at least three phases (gaseous, liquid, solid) are involved. Mass and

3287

heat transport phenomena not only influence the apparent activity, but can also alter significantly the selectivity and stability of the catalyst [13, 17, 18]. In some cases, the use of specially designed, highly efficient reactors strongly influences the overall performance of the process [19]. The pore and particle structure of the catalyst, mainly influenced by the morphology of the support, displays its effect on activity and selectivity of the reaction by restricting mass and heat transport at a microscopic scale [13, 18]. General Strategy and Experiment Planning As pointed out, selectivity in heterogeneous catalytic hydrogenation depends on many factors but the hierarchy is, in general, catalyst > reaction medium > reaction conditions. As many aspects are not fully understood or cannot easily be quantified, a purely rational approach in planning and optimizing a selective transformation is not feasible. Therefore, the methods used extend from theoretical considerations, literature and patent information, empirical correlations, experience and intuition to trial and error. Nevertheless, the following certain guidelines can be advantageous: 14.10.2.3.3

(i) Following the result of a literature and patent search for the desired reaction, exploratory experiments are first carried out. These experiments should be carried out in the kinetic regime, i.e. in the absence of external and, if possible, internal transport limitations. This can be achieved by the appropriate choice of reaction conditions and catalyst type. To find out what causes undesired side-reactions, the following questions have to be answered: first, is the side-reaction catalyzed by the hydrogenation catalyst or is it a homogeneous reaction in solution?; second, if the side-reaction is catalyzed by the catalyst, is it catalyzed by the metal itself or the support (acidic and basic centers)?; and finally, what is the type (decomposition, condensation, isomerization) and mechanism (single- or multiple-step reactions, parallel reaction or consecutive reaction) of both the hydrogenation and the actual side-reaction? (ii) In a second step, when the above questions have been answered, a specific strategy to improve the selectivity of the desired reaction can be developed. The strategy should be based on the general concept of speeding up the rate of the desired hydrogenation reaction versus the competing reaction, taking into account all steps of the process, i.e. kinetics and thermodynamics of adsorption/desorption, surface reaction of adsorbed species and mass transport. References see page 3307

3288

14.10 Hydrogenation Reactions

Catalyst Selection A Metal Selection Metals have inherent selectivity profiles (see below). Such a profile arises from the different adsorption/desorption enthalpies of the adsorbates on the various metals and their activation energies for the surface reactions. Both rate and selectivity depend on these surface processes. Although a rational approach based on the understanding of the various elemental steps could help, in practice empirical correlations and experience are used for the selection of the appropriate metal for a desired hydrogenation. Some of the characteristic properties of the most widely used catalytic metals for the hydrogenation of organic compounds can be summarized as follows [1–8]:

14.10.2.3.4

(i) Palladium is the metal of choice for hydrogenolyses (C−halogen, benzylic and allylic C−O and C−N and also N−N, O−O and N−O bonds) and for the saturation of C=C, C=O and C=N double bonds and C≡C and C≡N triple bonds conjugated to an aromatic ring. Palladium has also a strong tendency to catalyze isomerizations and double bond migrations. (ii) Platinum is very active for the hydrogenation of C=C, C=N and C=O double bonds. In contrast to palladium, isomerizations and double bond migrations almost never occur. C−halogen, allylic and benzylic C−O and C−N bonds are normally retained. (iii) Rhodium is used for the hydrogenation of aromatic rings under mild conditions and is very active for the hydrogenation of C=C and C≡C bonds. (iv) Ruthenium is applied for the hydrogenation of aromatic rings and of carbonyl and carboxyl functions, normally under higher pressure and elevated temperature. (v) Nickel, particularly Raney nickel, is especially suited for the hydrogenation of carbonyl functions (ketones, aldehydes), of nitriles to amines and of phenols to cyclohexanols. (vi) Copper catalysts are used for the hydrogenation of esters to the corresponding alcohols under rather harsh conditions. B Catalyst Support On the one hand, the support can interact with the substrate (polarity, lipo- or hydrophilicity, acidity/basicity), thereby influencing the course of a chemical reaction, or it can lead to side–reactions. On the other hand, the support interacts with the catalytic metal, thus changing its structural and electronic properties, which in turn has an effect on the hydrogenation process, i.e. adsorption/desorption (enthalpies and activation energy) and surface reaction (activation energy). In our experience, the catalyst activity decreases in the following order for the most useful and most prevailing supports: C (active carbon) > Al2 O3 > CaCO3 > BaSO4 .

C Catalyst Particle Size and Morphology For porous catalysts, the degree of internal mass transport limitation (and thus rate and selectivity) can depend on the size of the catalyst particle. The particle size also has an influence on the dispersion of the catalyst in solution, on the aggregation of gas bubbles in solution and on the separation of the catalyst from the reaction solution by filtration. The form of the catalyst particle has an influence on the physics of the process, particularly in fixed-bed applications (catalyst effectiveness). D Porosity of the Support and Location of the Metal The location of the metal on porous supports can have a strong impact on the mass and heat transport, thereby affecting not only the rate but also the selectivity of the reaction, which might be different from the selectivity in absence of diffusion control. The following aspects should be remembered when working with porous catalysts:

(i) If the metal is dispersed equally throughout the support, mass transport limitations are more likely to have an influence on the selectivity. This is particularly the case in consecutive hydrogenation reactions; an example is the partial hydrogenation of alkynes (Section 14.10.2.5.2). (ii) A porous catalyst with uniform metal distribution has advantages in terms of higher dispersion of the active metal and hence of activity. However, this is true only if the pores are large enough to allow easy access to the substrate molecules and when the rate of the reaction is not too high relative to the diffusion rate. (iii) In order to prevent hydrogen transport limitation in the pores, it can be essential to work at elevated pressure, particularly when catalysts with a uniformtype distribution of the metal are used. (iv) A porous support may have advantages in terms of catalyst lifetime, but only if deactivation preferentially occurs at the most accessible centers. E Metal Dispersion and Metal Crystallite Size Dispersion and crystallite size of the metal particles depend on the preparation and activation procedures used and on the metal loading of the catalyst support. The number of surface metal atoms and therefore the activity of the catalyst increases with increasing dispersion. However, ˚ a decrease in the specific for very small particles ( Br > Cl > F. f Rate of double bond hydrogenation: mono- > di- > tri- > tetrasubstituted. g Y = N, O. Selectivity profiles for the hydrogenation of aldehydes and ketones. R2 C=O + H2 → R2 CH−OH

Tab. 5

Group/metal

Function to be retained Ar−Halc

Pd Pt Ru Rhf Ni

± ±b +a +a +

[2, p. 210] [3, p. 307] [2, p. 210] [3, p. 307] [24]

+ +

[2, p. 210] [3, p. 307]

C≡Cf

C=Cd

C≡N

±a

[2, p. 224]

+ +a

[2, p. 218] [2, p. 218]

+a ±

[26] [2, p. 219]

±

±

[3, p. 305]

[3, p. 305]

ArNOf2

Y–benzyle ±

[3, p. 306]

+ +

[3, p. 306] [2, p. 213]

± −

[3, p. 306] [2, p. 213]

Key: +, selective; ±, partially selective; −, unselective. a Modified with second metal. b Non-metallic modifier. c Rate of dehalogenation: I > Br > Cl > F. d Rate of double bond hydrogenation: mono- > di- > tri- > tetrasubstituted. e Y = N, O. f No examples found.

which metal (listed vertically) has been shown to be suitable. In many cases, the inherent selectivity of a metal is not sufficient and modified catalysts have to be employed. To make the literature more easily accessible, Rylander [1], Houben-Weyl [2] or Freifelder [3] references are given where possible. The hydrogenation of functionalized nitroarenes, which is of importance for the intermediate manufacture for dyestuffs and agrochemicals and the semihydrogenation

of alkynes used especially for vitamins, are discussed in more detail in the following sections. 14.10.2.5 Selective Hydrogenation of Functionalized Aromatic Nitro Compounds 14.10.2.5.1 Introduction Substituted anilines are intermediates of importance in the manufacture of a large References see page 3307

3292

14.10 Hydrogenation Reactions

Selectivity profiles for the hydrogenation of C=C double bonds. R2 C=CR2 + H2 → R2 CH−CHR2 a

Tab. 6

Group/metal

Function to be retained Ar−Halb

Pd Pt

+e ± +

[18] [3, p. 160] [3, p. 159]

Ru Rh Ni

+

[3, p. 160]

C≡C

+

C=O

[29]

C≡N

+a + ±

[1, p. 40] [2, p. 161] [2, p. 161]

+ + ± +

[1, p. 41] [2, p. 168] [2, p. 164] [2, p. 161]

ArNO2 c

Y–benzyld

+ + + +

[2, p. 168] [3, p. 157] [2, p. 168] [3, p. 157]

+ + + +

[27] [28] [30] [31]

−

[2, p. 168]

+ +

[3, p. 158] [32]

Key: +, selective; ±, partially selective; −, unselective. a Rate of double bound hydrogenation: mono- > di- > tri- > tetrasubstituted. b Rate of dehalogenation: I > Br > Cl > F. c No Examples found. d Y = N, O. e Non-metallic modifier.

variety of fine and specialty chemicals. Agrochemicals, pharmaceuticals, dyestuffs and pigments equally depend on their supply. Some compounds are produced on a large scale. In recent years, older manufacturing processes, e.g. reduction with iron in aqueous acidic media (B´echamp reduction) or sulfides, have been replaced by catalytic hydrogenation, mainly for ecological reasons [33]. Catalytic hydrogenation is also a more economical alternative for the production of haloanilines because B´echamp reduction produces large quantities of iron oxides contaminated with toxic aromatics, the disposal of such wastes becoming increasingly difficult and expensive. The technology of selective catalytic hydrogenation has been under continuous development since the early 1940s and has been improved to a remarkable state of perfection, especially for chlorinated nitroarenes. Whereas the hydrogenation of simple nitroarenes poses few selectivity problems and is indeed carried out on very large scale, the situation is different if other reducible functional groups are present in the molecule. A summary of the state of the art has been presented in Table 4. Here we discuss the following topics in further detail: the reaction mechanism of nitroarene reduction, practical aspects such as the choice of catalyst, reaction medium and reaction conditions or questions of safety and equipment, the selective hydrogenation of halogenated nitroarenes, new modified catalyst systems and the problem of hydroxylamine accumulation.

the relatively complex set of reaction intermediates that can occur when reducing nitroarenes. More than 100 years ago, Haber [34] proposed the reaction network depicted in Scheme 1 to explain the results of the electrochemical reduction of nitroarenes. Since then, the suggested intermediates have all been verified and it has been generally accepted that catalytic hydrogenation reactions proceed via the same routes. However, it must be stressed that due to the very strong adsorption of several species, the mechanistic investigation and interpretation of heterogeneous catalytic reactions are much more complex. The reduction to anilines occurs in three individual steps with the overall formation of two molecules of water, either by the direct route via nitroso and hydroxylamine intermediates or via the condensation route, which is favored under basic conditions. Many of the intermediates can react further and byproduct formation resulting in poor product quality or, even worse, accumulation of metastable intermediates such as hydroxylamines and the possibility of runaway reactions (see below) are inherent problems due to the complexity of the reaction. When R is a function which can also be reduced or cleaved by hydrogen, these reactions can in principle occur on any of the various intermediates. However, due to the preferential adsorption of the species with a higher oxidation state, the unwanted side-reaction often occurs as a consecutive reaction after the substituted aniline has been formed. Practical Aspects A Catalysts The ‘‘classical’’ and most frequently used hydrogenation catalysts for the hydrogenation of nitro

14.10.2.5.3

Mechanism of Nitroarene Reductions In order to facilitate the discussions, we first present 14.10.2.5.2

14.10.2 Selective Hydrogenation of Functionalized Hydrocarbons

nitro

3293

NO2 R reduction

− H2O

nitroso NO R condensation reduction

− H2O

N R

+

O−

azoxy R

N − H2O reduction

hydroxylamine NHOH

azo

R

R

N reduction

R

− H2O

N reduction

aniline NH2

reduction

R

hydrazo

H N R

R N H

Scheme 1

groups are the noble metals Pt and Pd, supported on active carbon, Raney nickel and Ni supported on kieselguhr. Sometimes a second metal is added in order to modify the catalytic performance (see below). Since the active metal is present in the form of very small particles on a support or as a skeletal material, the metal specific surface area is generally very high. Many problems can be solved adequately using standard catalyst types, but more demanding processes often require tailored catalysts. Typical examples are a 1% Pt/C catalyst developed by Johnson Matthey [35] and a promoted 5% Ir/C catalyst of Degussa [36] for the selective hydrogenation of halonitroarenes and/or the suppression of by-product formation. The development of two modified catalysts with a much broader selectivity pattern and of a vanadium promoter to suppress hydroxylamine accumulation is described below in more detail. In recent years, several lines of research to produce catalysts with improved selectivities have been described. Colloids [37–39] were reported to give remarkably selective catalyst systems for the hydrogenation of chloronitroarenes. Reusable Pd complexes on different supports were described for the selective hydrogenation of nitro aromatics in the presence of C=O [40] and C−Cl functions [41]. In our view, despite good laboratory results, most of these catalysts are not (yet) ready for technical application. Progress has also been made in the use

of chemoselective transfer hydrogenation systems, most notably with iron hydroxide catalysts in combination with hydrazine hydrate as reducing agent [42], and first technical applications have been reported [43, 44]. B Reaction Medium and Conditions Catalytic hydrogenations of nitro groups are usually carried out in solution. The choice of the solvent not only affects the solubility of the reactants and products but can also very strongly influence the activity and selectivity of a catalyst. Solvents should not be hydrogenated under the particular reaction conditions. At the laboratory stage, only high-purity solvents should be used in order to minimize poisoning of the catalyst. Most often used are alcohols (MeOH, EtOH, iPrOH, BuOH), ethyl acetate, aromatic and aliphatic hydrocarbons, ethers such as tBuOMe, THF, dioxane (care has to be taken with Raney nickel at high temperatures), water, ketones and acetic acid. In special cases, amides such as dimethylformamide (DMF), dimethylacetamide or N -methylpyrrolidone and dichloromethane are also used. Especially for the production scale, it is important to optimize carefully all parameters of the catalytic system: catalyst, reaction medium and reaction conditions. The quality of the optimization will strongly affect the costs References see page 3307

3294

14.10 Hydrogenation Reactions

of the hydrogenation step! The following parameters can be influenced and will affect the process performance: Hydrogen pressure (which usually has an influence on the rate of reaction, and sometimes also on selectivity); temperature (because of the very high exothermicity of the reaction and in order to avoid accumulation of intermediates, it is of advantage to carry out the hydrogenation at temperatures >80 ◦ C); substrate concentration (which determines volume yield); catalyst/substrate ratio (which depends on the catalyst activity and determines reaction time and catalyst cost); agitation (which affects gas–liquid diffusion and is especially important because many nitroarene hydrogenations are very fast and, in addition, under hydrogen-starved conditions metal leaching can be very pronounced); and catalyst pre-treatment (e.g. pre-reduction is sometimes necessary in order to improve catalyst activity and to diminish corrosion of precious metals). In some cases, the continuous addition of unstable or dangerous substrate(s) should be considered. C Safety and Health Aspects The catalytic hydrogenation of aromatic nitro compounds is a potentially hazardous reaction. The safety of a specific operation depends both on the nature of the nitroarene and on the operating conditions. a Hazards Related to Hydrogen Hydrogen–air mixtures are explosive over a wide range of concentrations (4–75 vol.%) and have a very low ignition energy (0.02 mJ). b Hazards Related to the Catalyst Dry hydrogenation catalysts such as Raney nickel and palladium or platinum on charcoal are pyrophoric. One measure is the use of wet catalysts; this is usually not a problem since the hydrogenation reaction produces water anyway. c Hazards Due to the Hydrogenation Reaction The reduction of nitro-aromatics is a very exothermic reaction (560 kJ mol−1 ). In the case of loss of reaction control, decomposition of the aromatic nitro compound or of partially hydrogenated intermediates could easily be triggered. d Hazards Due to the High Decomposition Energy of Aromatic Nitro Compounds and some Intermediates Aromatic nitro compounds have high decomposition energies (about 2000 kJ mol−1 ). Even though the activation energy is generally also high, the decomposition often follows a self-accelerating mechanism. In solution, in presence of a metallic catalyst and especially under basic conditions, decomposition may start at much

lower temperatures. The catalytic hydrogenation of nitroarenes involves a series of intermediates, especially N -arylhydroxylamines, which can decompose exothermically and trigger the decomposition of the reaction mixture. This can be an acute safety problem under heat accumulation conditions such as cooling failures or when the circulation pump fails in loop reactors because these reactions do not consume hydrogen and therefore cannot be stopped by controlling hydrogen feed or stirring. Therefore, the accumulation of hydroxylamines must be kept at a minimum at any time of the reaction. More details and a discussion of the tools used for risk evaluation have been summarized by Blaser et al. [33]. e Toxicity Aromatic nitro compounds and anilines, in addition to their acute toxicity, are in some cases known carcinogens and must be handled accordingly. Toxicological studies suggest that arylhydroxylamines are metabolites produced by the organism involved in carcinogenesis [45, 46]. D Hydrogenation Equipment The catalytic hydrogenation of nitroarenes in the fine chemicals industry is usually carried out in the liquid phase and in the batch mode. For successful implementation, the following demands have to be met: very good dispersion of the hydrogen gas and the suspended solid catalyst in the reaction solution (efficient gas–liquid mixing and stirring), very effective heat removal (reaction control) and safe handling of the sometimes pyrophoric catalyst and the sometimes toxic starting materials and products. In practice, two reactor types have proven to be capable of meeting these requirements and also the needs for high reliability in operation and ease of control: the stirred autoclave and the loop reactor (see Table 7). The loop reactor provides very efficient hydrogen dispersion and the heat exchanger surface is almost unlimited and is especially useful when the space–time yield is very high (fast reaction, high substrate concentration) or when a low reaction temperature is required. The stirred autoclave is probably more versatile and better suited for multipurpose installations; it has an advantage when substrate slurries have to be used or when the starting material is added continuously. In addition, it is usually easier to clean and less space and lower investment costs are required. Selective Hydrogenation of Halogenated Nitroarenes A Introduction Halogenated anilines are the most important class of functionalized aromatic compounds and most early efforts to find selective hydrogenation catalysts have been concentrated here. The field has been 14.10.2.5.4

14.10.2 Selective Hydrogenation of Functionalized Hydrocarbons

Tab. 7

3295

Comparison of the loop reactor and the stirred autoclave Loop reactor

Gas dispersion Efficiency Heat removal Problem areas Recommended for

Stirred autoclave

Mixing nozzle (Venturi principle) High >1300 W m−2 K−1 , very high exchange area Circulating pump (viscous slurries), continuous feed addition High performance, dedicated plant

reviewed extensively by Kosak [47, 48] and by Baumeister et al. [49]. Most R&D effort has been directed towards the selective hydrogenation of monohaloanilines (F, Cl, Br), dichloronitrobenzenes, dinitrochlorobenzenes and aminochloro- and aminodichlorophenols. The reactivity of the halogen substituents towards hydrogenolysis decreases in the order I > Br > Cl > F [48]; the activity for the hydrogenolysis of haloaromatics for different metals is Pd Ni > Pt. The nature of the substituents has a strong effect on the dehalogenation rate but the selectivity pattern for dehalogenation seems to change with the metal and the acidity of the reaction medium. Since the main path to the loss of halogen is hydrogenolysis of the halogenated anilines [50], it appears that amino groups are activating substituents for the hydrogenolysis of the C−X bond. The activation decreases in the substituent order ortho- > para- > meta [48]. Substitution with more than one amino group further enhances the ease of displacement by hydrogenolysis. Protonation or acylation of the aniline weakens the activating character of the amino group and improves the selectivity in some cases. B Catalysts Platinum and nickel are the catalyst metals most widely used for the hydrogenation of halonitroarenes. Metals with favorable selectivity patterns, but for economic reasons of only limited and specialized use, are rhodium, ruthenium and iridium. Palladium, the catalytic metal of choice for the cleavage of carbon–halogen bonds in aromatics, can be used only for the hydrogenation of fluoronitroarenes. However, the classical catalysts show insufficient selectivity for most of the important transformations and various procedures

Mechanical agitator (hollow-shaft turbine) Medium to high ∼900 W m−2 K−1 , limited exchange area Heat exchange capacity Multi purpose plant

have been developed over the years, either for the preparation of modified catalysts or for the modification of the process conditions. a Modified Catalysts One of the first successful modified catalysts was developed by Greenfield and Dovell [51, 52], who used platinum sulfides on carbon for the selective hydrogenation of several chloronitrobenzene derivatives. The metal sulfides are highly selective but show acceptable activity only at rather high temperatures (140–150 ◦ C) and hydrogen pressures between 30 and 50 bar. Later, a variety of different in situ methods for the sulfidation of platinum and base metal catalysts were published. The effect of co-deposition of metallic modifiers on the selectivity of platinum catalysts was investigated both by catalyst manufacturers and by academic research groups. Special mention should be made of Pt catalysts modified with Pb, Bi, Sn, Ge, Zn, Al or Ag compounds [53, 54] or with copper [55] with high selectivity in the hydrogenation of various mono- and dichloronitrobenzenes. b In Situ Modification This is obviously the method of choice for the chemist working in development since the modification procedure can be adapted to the individual requirements of a specific reaction. Over the years, a number of effective modifiers for a variety of halonitroarenes have been developed which give not only a highly selective reaction but also acceptable catalytic activity. Pt/C catalysts modified with morpholine or phosphorous acid were developed by Kosak with References see page 3307

3296

14.10 Hydrogenation Reactions

selectivities of >99% for the hydrogenation of 3,4dichloronitrobenzene [47, 48]. Baumeister et al. [25] used Raney nickel or Pt/C modified with either dicyanodiamide or formamidine acetate for the selective hydrogenation of 1-chloro-2,4-dinitrobenzene, one of the most difficult selectivity problems in this field. New Catalyst Systems for the Hydrogenation of a Nitro Group in the Presence of Reducible Functional Groups Although the modified catalysts described above gave a satisfactory performance for selective chloronitroarene hydrogenations, functional groups involving C=C, C=O, C=N or C≡C bonds were either not tolerated or the catalytic activity was too low, leading to high concentrations of intermediates or side-products. Such selectivity problems could only be solved by using stoichiometric methods such as B´echamp or sulfide reduction, with the consequence of large amounts of waste. For these reasons, the Ciba-Geigy catalysis group developed two novel modified Pt catalyst systems with surprisingly good performance for a variety of selectivity problems [23]. 14.10.2.5.5

A Lead-Modified Platinum Catalysts: the ‘‘Lindlar The idea behind this approach was Approach’’ influenced by the successful Lindlar catalyst, namely to modify the selectivity of a catalyst by the addition of a second metal. The final catalyst was developed during the course of a process development for a herbicide intermediate carried out in collaboration with Degussa. In summary, the studies showed that: NH2

(i) CaCO3 is the best carrier material, other supports (e.g. charcoal or alumina) leading to catalysts with low selectivity; the lead content is very important, the optimal level being 1% w/w. (ii) To achieve reasonable catalytic activity, a reaction temperature ≥120 ◦ C is necessary. (iii) The catalysts can be prepared according to the Lindlar procedure [9, 10], but reproducibility can sometimes be a problem. (iv) The polarity of the solvent has a strong influence on both catalyst activity and yield. The best results are obtained with polar solvents. (v) The addition of small amounts of FeCl2 and tetramethylammonium chloride (TMAC) has a beneficial influence on the hydrogenation rate and to a lesser extent also on the yield. The preparative scope of the catalytic system is shown in Fig. 1. Except for the C≡C bond, all other functions are not reduced. The sometimes low yield is due to side-products and intermediates. Noteworthy is the high yield of a substituted 2-aminonitrile without formation of benzamide by-products due to oxygen transfer from the hydroxylamino to the cyano group. B H3 PO2 -Modified Pt Catalysts in the Presence of Vanadium Promoters For the second catalyst system, the basic idea was to change catalyst selectivity by the addition of organic or inorganic modifiers to the reaction solution. The starting point for the novel catalyst was a paper by

NH2

R R

N N

NH2

OH CN O A) 75% B) 95%

X

B) 95%

A) 82% I

NH2

NH2

R O

Ar-NO2

Cl

O

A) 65% B) 99% (prop/allyl 4:1) (allyl < 1%)

H2N

NH2

N (X) R

S A) 80%

N OH H2 N

CN

A) 51% B) >50% A) Pt-Pb/CaCO3

Fig. 1

R R A) > 90%

R A) 75%

B) Pt/C – H3PO2 – V

Scope of the new modified Pt catalysts (with few exceptions, the reaction conditions have not been optimized).

14.10.2 Selective Hydrogenation of Functionalized Hydrocarbons

NO2 SO2 N

5% Pd/C 2-propanol

3297

NH2 SO2 N

20 bar H2

400

time / min

40

ArNHOH / %

30 20

200

ArNHOH / %

Time / min

600

10

0.03% 0.02% V NH4VO3 (on Pd/C)

Fig. 2

0.02% CuCl

0.09% unmodified Tl(NO3)3 system

Effect of promoters on hydrogenation time and maximum hydroxylamine accumulation.

Kosak [48], who described Pt/C modified with H3 PO3 for the highly selective hydrogenation of iodonitrobenzenes, albeit with a strong accumulation of hydroxylamines. The challenge was to find a co-catalyst that would decrease the accumulation of hydroxylamines without interfering in a negative way with the catalytic hydrogenation. After extensive screening, a novel combination of modifier and promoter was found and tested for various chemoselective nitro reductions. In summary, the studies showed that: (i) The best modifier is H3 PO2 ; other additives such as H3 PO3 , (PhO)2 P(O)H and HPPh2 also showed an excellent selectivity. P(OPh)3 and especially PPh3 were less efficient. (ii) The H3 PO2 concentration has a strong effect on rate and selectivity; the critical level was at 2.5%, and optimum results were obtained with 5% relative to the Pt catalyst. (iii) The addition of V promoters is essential for low hydroxylamine accumulation. (iv) Toluene or toluene–water mixtures are favorable reaction media; in contrast to Pt−Pb−CaCO3 , polar and protic solvents were less suitable for Pt/C modified with H3 PO2 . Figure 1 shows selected results for various functionalized nitro compounds with this new catalyst system to give the corresponding anilines. No reduction of the second function was ever detected; even a C≡C bond remained completely unreduced! Compared with the lead-modified catalyst, the H3 PO2 /vanadium-modified Pt catalyst system is much more versatile and, especially important for technical applications, normal commercial Pt catalysts can be used. 14.10.2.5.6 Hydroxylamine Accumulation As pointed out above, hydroxylamines are problematic due their

potential for decomposition with a strong exotherm. In addition, they are also known to be strong carcinogens and therefore are hazardous in case of interrupted or incomplete hydrogenation [46]. Furthermore, hydroxylamine accumulation can lead to poor product quality because reaction with the nitroso compound gives colored condensation products. In many cases, the hydrogenation of aromatic nitro compounds is very fast at the beginning but becomes sluggish after 60–80% conversion. Sometimes it is even difficult to define the end-point of hydrogen uptake, generally an undesirable situation in production and unacceptable when a consecutive hydrogenation step reduces selectivity. The maximum concentration of hydroxylamines can vary and is notoriously difficult to predict; therefore, the product quality may differ from batch to batch. As a consequence, the suppression of hydroxylamine accumulation is highly desirable and a topic of industrial importance. Two recent publications describe the addition of small amounts of metal salts, especially vanadium compounds, to commercial Pt, Pd and Ni catalysts [56, 57], leading to a dramatic decrease in the maximum hydroxylamine accumulation often below the detection limit (for an example, see Fig. 2). The results can be summarized as follows: (i) The addition of promoters, especially vanadium salts, had a dramatic effect on the course of the hydrogenation of several nitroarenes. In some cases, the maximum level of hydroxylamine accumulation decreased from >40% to tertiary > secondary > primary N -benzyl group [2, 3]. If two or more benzyl groups are attached to a single nitrogen, then stepwise removal is often possible [97, 98]. This allows the synthesis of mixed secondary and tertiary amines by debenzylation/alkylation sequences [97]. Also here, differentiation between two benzyls attached to the same amide nitrogen is possible and N -benzylamines can selectively be cleaved in the presence of N -benzylamides [2]. E Selective Removal of N-Benzyl Groups in the Presence of other Reducible Functions As for O-benzyl groups, the removal of N -benzyl groups is possible in the presence of aromatic halides (especially Cl and F), C=O groups (aliphatic and aromatic) and nitriles [2, 3]. By adding acidic modifiers such as HCl or acetic acid, the selectivity for N -debenzylation can be improved in the presence of halogens. On the one hand, the reaction is speeded up by protonation of the nitrogen (‘‘quaternization’’), and on the other, the removal of the halogen X is slowed by the lack of an acceptor for H−X. Solvents can also play a major role, as mentioned before. Since N -debenzylation is usually more difficult than O-debenzylation, selective deprotection in the presence of C−C triple bonds or NO2 groups is even more difficult and no examples were found in an extensive literature search. C=C bonds are known to survive N -debenzylation conditions only if they are highly substituted and/or conjugated. Bornmann and Kuehne [99], for instance, described the selective deprotection of a molecule containing an α, β-unsaturated ester. F New Protecting Groups There is strong interest in protective groups which can be removed selectively and easily. Figure 7 depicts a series of recently published References see page 3307

HO

O OBn

HO

Si(C6F13CH2CH2)3

CF3 HO

BOB

CFTB

fluorous benzyl X

X HO OMe

Fig. 7

X = O, N

HO

MPM

N 1-NAP

3303

2-NAP

4-QUI

Structures and abbreviations of new protective groups that can be removed via selective hydrogenolysis.

3304

14.10 Hydrogenation Reactions

structures with interesting properties. 1-NAP and 4-QUI esters have been cleaved with a homogeneous Pd complex and a formate donor; benzyl esters, alkenes, Ar−Br and other functional groups are tolerated [100]. The highly selective removal of 1-NAP from N - and O-functions without affecting Bn and CFTB groups was reported for Pd/C−H2 [101]. MPM−OAr groups survived the deprotection of Bn–OAr and CbzNH with Pd/C modified with pyridine, but could easily be removed in the absence of the pyridine modifier [102]. Tagging with fluorous benzyl groups allowed a clever combination of protection and fluorous phase chemistry with easy subsequent removal of the auxiliary group [103]. BOB-protected hydroxy groups were deprotected via hydrogenolysis/lactonization compatible with a number of fatty acid esters [104].

on the catalyst surface whereas for aliphatic benzyl ethers the adsorption is blocked. (v) Very little is known about the solvent effects in debenzylation. In the case of halogen-containing solvents [88], small amounts of HCl formed by dehalogenation of the solvent could make the reactions faster and more selective. Other factors such as solvation/adsorption properties and hydrogen bond formation may also be important. (vi) Catalyst manufacturers claim that oxidic or unreduced egg-shell-type catalysts with a high Pd loading give the best selectivity for debenzylation (see above). So far, no explanation for this recommendation has been given. 14.10.2.8

Mechanism of Debenzylation and Mode of Action of Modifiers As mentioned above, there is a lack of fundamental studies on the hydrogenolysis reaction. A possible mechanism [105] involves the adsorption of X–CR R –Ph on the catalyst surface through the phenyl group and, if not too bulky, also through X. This species is thought to form a π-benzyl complex with the Pd after cleavage of the X−C bond. Reaction with adsorbed hydrogen finally forms the X−H and X−CR R −Ph species. Even though this mechanism seems to be generally accepted [84], it does not help in understanding how selectivity can be obtained. There are some hypotheses, however, which plausibly explain some of the observed effects and are worth mentioning, if only for practical reasons:

Stereoselective Reactions

14.10.2.7.5

(i) Removal of benzyl groups is accelerated by acids. In the case of benzylic alcohols, it is reasonable to assume that the reaction proceeds through the corresponding ester, when a carboxylic acid is added [3]. Kieboom and van Rantwijk [5] postulate that the debenzylation reaction occurs via a nucleophilic attack of a hydride on the C−X bond and explain the effect with the increased nucleophilicity of a protonated X group. (ii) For N -benzyl compounds, the acceleration by acids can be explained by protonation of the nitrogen, which eliminates the inhibition by amines. This is a non-catalytic mode of action and 1 equiv. of a strong acid is necessary for an optimal effect. (iii) For molecules containing aromatic halogens and basic amines, acids remove a possible acceptor for H−X. Therefore, the speed of dehalogenation decreases relative to the debenzylation. (iv) The effect of bases is not well understood. Deprotection of aliphatic [77, 81] but not aromatic [93] benzyl ethers is blocked by aliphatic nitrogen bases. One explanation is that the latter are still able to adsorb

14.10.2.8.1 Diastereoselective Hydrogenation This topic is discussed in more detail in Section 14.16.3, but since the asymmetric hydrogenation of (hetero) aromatic rings would be an attractive route to chiral (hetero) cyclic compounds but so far has not been feasible with homogeneous catalysts, we will briefly summarize some recent results. Several diastereoselective hydrogenations of carbocyclic or heterocyclic systems coupled to chiral auxiliaries such as proline or related compounds have been reported and diastereomeric excesses (des) up to 98% were achieved [106–110]. Selected examples are depicted in Fig. 8. It is worth noting that for optimal results, the structure of chiral auxiliary has to be adapted to the aromatic systems. In general, the chiral auxiliary has to be removed in an additional step, but recently Glorius et al. [109] have reported a procedure for the hydrogenation of 2oxazolidine-substituted pyridines where the auxiliary is cleaved during the hydrogenation. The observed enantiomeric excesses (ees) are very high and both substrate synthesis and auxiliary recycling are very efficient. The issue of cis selectivity in the hydrogenation of disubstituted heterocyclic and carbocyclic rings was addressed by several groups and classical catalysts such as Rh and Raney nickel were able to give satisfactory cis selectivities, but in some cases bimetallic systems proved to be superior [110, 111]. A remarkable example of the synergism of bimetallic catalysts is the hydrogenation of pyridine-2-carboxylic acid derivatives. Surprisingly, a 4.5% Pd–0.5% Rh/C catalyst is twice as active as a 5% Rh/C catalyst and, in addition, shows better cis selectivity [110]. 14.10.2.8.2 Enantioselective Hydrogenation A metal surface can be made chiral by the addition of suitable modifiers, leading to potential enantioselective catalysts (see Section 14.16.1). Here we briefly recapitulate the synthetic scope of this methodology (for a recent review, see

14.10.2 Selective Hydrogenation of Functionalized Hydrocarbons

Reaction

O

O R = Alkyl

O

EtOH, r.t. 50 bar

de up to 96% yield >90%

O

COOMe

O

N COOMe

O

N H

MeOH, 25 °C. 20 bar H2

R3

de 27%

R4 N

O

N R

O

Pd, Rh or Pt catalyst

R2

AcOH, 20−70 °C 100 bar H2

R1

[108]

O R4 +

N H

HN

O

ee up to 98% yield 60−94%

[109]

cis: trans 97:3 yield 95%

[110]

R PO3H2

PO3H2 4,5%Pd-0.5%Rh / C

Fig. 8

[107]

COOMe

O

R3

R2

N

de up to 95%

Rh / C

N

[106]

N * N

COOMe MeOH, 25 °C. 20 bar H2

O

N

R1

COOR

N

Rh / C

N N

Ref.

Rh / Al2O3 amine

COOR

N

3305

COOH

COOH

N H

Selected diastereoselective hydrogenations of aromatic rings.

Ref. [112]). The three most important asymmetric catalyst types are Raney nickel modified with tartaric acid, effective for β-functionalized ketones with ees up to 98.6%, platinum catalysts modified with cinchona alkaloids and related modifiers, successful for α-functionalized ketones with ees up to 98%, and palladium catalysts modified with cinchona alkaloids, which achieve ees up to 85% for selected activated C=C bonds. The most suitable substrates for the three catalytic systems are depicted in Fig. 9. The first technical application of a cinchona-modified Pt catalyst was reported in 1986 for the synthesis of an intermediate for the angiotensin-converting enzyme inhibitor benazepril. The hydrogenation of the corresponding α-keto ester to ethyl (R)-2-hydroxy-4-phenylbutanoate [HPB ester, Eq. (5), lower part] was developed and scaled up into a production process (10–200-kg scale, chemical yield >98%, ee 79–82%) [113]. Even though this route was competitive at the time, it suffered from problems such as stability of the starting ester or non-crystallinity of the

HPB ester. Recently, an improved new synthetic route to HPB ester [Eq. (5), upper part] which is feasible on a large scale has been developed and piloted [114]. O

1. Pt/Al2O3 - HCd ee 70 - 87%

OH

Ph

COOEt

O

OH

Ph

COOEt

2. crystallization ee >99% H N N

O

Pd/C, EtOH, HCl ee >99%

Ph COOEt

OH Ph

COOEt

CH2COOH

HPBester

Pt/Al2O3 - Cd ee 80−85%

benazepril O Ph

COOEt

(5) References see page 3307

3306

14.10 Hydrogenation Reactions

COOH H OH HO H COOH

O R

Raney nickel

O COOR

O R

X = O, Y= CH2 X = NR, Y = C=O O

Pt/Al2O3

O

R′

R

O

O N

ee 85−98%

R

ee 85−91%

O

H H HO

R

ee 84−98.6%

tartaric acid

O

ee 63−85%

(kinetic resolution) O

O X Y

OR

R′

OR

ee 91−92%

ee 90−97%

O

OMe

R′ R

ee 90−98%

R

CF3

ee 92−93%

COOH OH N

cinchonidine

Pd/Al2O3 Pd/TiO2 O

ee 72%

Fig. 9

O

ee 85%

Best results reported for heterogeneous enantioselective catalysts (modifier, preferred catalyst and substrates) [112].

Chemoselective Hydrogenation of Nitriles The hydrogenation of nitriles is one of the basic methods to obtain primary amines and especially diamines, which are of high industrial importance. Unfortunately, the literature is rather scattered and the most upto-date review was actually written in 1994 [115]. We focus our summary of current results on selectivity to primary amines, catalyst deactivation and functional group tolerance. 14.10.2.9

14.10.2.9.1 Primary Amine Formation In addition to primary amines, secondary and tertiary amines can be formed via condensation of reaction intermediates, and control of this chemoselectivity problem is one of the main issues of nitrile hydrogenation. Addition of ammonia is most widely used to improve the selectivity for primary amines [115], but recently it was reported that less toxic strong bases such as NaOH [116, 117] and LiOH [118] are also effective for Raney Ni and Co catalysts. The OH− ions not only prevent catalyst deactivation by inhibiting polyamine formation on the catalyst surface for dinitrile hydrogenation [116], but also seem to block active sites responsible for by-product formation [117]. Pretreatment of Ni and Co catalysts with CO, CO2 , aldehydes or ketones also gave significantly less secondary amines [119].

Chemoselective Hydrogenation of Unsaturated For fine chemicals applications, functional Nitriles group tolerance is an important issue. Substituents such as aryl, benzylic and C−Hal groups are usually not reduced with skeletal Ni or Co catalysts. More difficult to conserve are heteroaromatic or heteroaryl–halogen 14.10.2.9.2

functions, ketones, aldehydes or a second CN group, but with the proper catalyst, solvent and additives success is often possible [120]. In contrast, the selective hydrogenation of CN groups in the presence of C=C bonds has long been an unsolved problem, particularly if they are conjugated or in close proximity in the molecule [120, 121]. If the C=C bond is sterically hindered, then high selectivity is possible in liquid ammonia, which not only inhibits the formation of secondary amines, but also improves the selectivity to the unsaturated amine, probably by forcing its desorption [122]. Another case of chemoselective nitrile hydrogenation has been described for a fatty acid nitrile where the selective hydrogenation of remote CN functions is possible with high selectivity by applying a Ziegler-type Co−Fe catalyst even in absence of NH3 [123]. Recently, the chemoselective hydrogenation of cinnamonitrile to 3-phenylallylamine was reported to proceed with up to 80% selectivity at conversions >90% with Raney cobalt [124]. The best results were obtained with a doped Raney cobalt catalyst (RaCo/Cr/Ni/Fe 2724) in ammonia-saturated methanol at 100 ◦ C and 80 bar [see Eq. (6)]. Major problems were the formation of hydrocinnamonitrile and of secondary amines and overreduction to 3-phenylpropylamine. The catalytic system tolerates functional groups such as OH, OMe, Cl and C=O, but not aromatic nitro groups. The same catalyst system is also useful for the hydrogenation of other unsaturated nitriles with di- or trisubstituted C=C bonds [125]. The substitution and the position of the double bond relative to the nitrile group are crucial in determining the chemoselectivity for the unsaturated amine. The double bond is not hydrogenated when

References

it is sterically hindered or far from the nitrile group such as in cyclohex-1-enylacetonitrile [see Eq. (6)] or the double bond at C-6 in geranylnitrile. In contrast, in dimethylacrylonitrile or 2-pentenylnitrile, the activated double bond is hydrogenated preferentially to the saturated nitrile. CN RaneyCo/Cr/Ni/ Fe 2724

CN

NH3 / MeOH 100 °C, 80 bar

NH2 Selectivity up to 80%

NH2

Selectivity up to 87%

(6) References 1. P. N. Rylander, Hydrogenation Methods, Academic Press, New York, 1985, 193 pp. 2. Houben-Weyl, Methoden der organischen Chemie, Vierte Auflage, Reduktionen Teil I, Band IV/1c, Georg Thieme, Stuttgart, 1980, 14–562. 3. M. Freifelder, Practical Catalytic Hydrogenation, WileyInterscience, New York, 1971, 663 pp. 4. F. Zymalkowski, Katalytische Hydrierung, Ferdinand Enke Verlag, Stuttgart, 1965, 360 pp. 5. A. P. G. Kieboom, F. van Rantwijk, Hydrogenation and Hydrogenolysis in Synthetic Organic Chemistry, Delft University Press, Delft, 1977, 157 pp. 6. G. V. Smith, F. Notheisz, Heterogeneous Catalysis in Organic Chemistry, Academic Press, San Diego, CA, 1999, 352 pp. 7. S. Nishimura, Handbook of Heterogeneous Catalytic Hydrogenation for Organic Synthesis, Wiley, New York, 2001, 784 pp. 8. R. L. Augustine, Heterogeneous Catalysis for the Synthetic Chemist, Marcel Dekker, New York, 1995, 640 pp. 9. R. A. Sheldon, Chem. Ind. (London) 1992, 7(12), 90. 10. G. C. Bond, Heterogeneous Catalysis, 2nd Ed., Clarendon Press, Oxford, 1987, 176 pp. 11. I. M. Campbell, Catalysis at Surfaces, Chapman and Hall, London, 1988, 250 pp. 12. O. Levenspiel, Chemical Reaction Engineering, Wiley, New York, 1962, 501 pp. 13. R. J. Farrauto, M. C. Hobson, N. L. Brungard, Catal. Org. React. 1988, 33, 177. 14. R. L. Augustine, P. J. O’Hagan, Catal. Org. React. 1990, 40, 111. 15. P. N. Rylander, Catalysis in Organic Synthesis, Academic Press, New York, 1980, p. 155. 16. R. L. Augustine, P. Techasauvapak, J. Mol. Catal. 1994, 87, 95. 17. C. N. Satterfield, Heterogeneous Catalysis in Practice, McGrawHill, New York, 1981, 416 pp. 18. G. W. Roberts, Catalysis in Organic Synthesis, Academic Press, New York, 1976, p. 1. 19. J. J. Concordia, Chem. Eng. Prog. 1990, 86(3), 50. 20. D. C. Harwell, G. H. Timms, Synth. Commun. 1979, 9, 223. 21. A. G. Caldwell, E. R. H. Jones, J. Chem. Soc. 1946, 597. 22. G. Stork, J. Am. Chem. Soc. 1947, 69, 576. 23. P. Baumeister, H. U. Blaser, U. Siegrist, M. Studer, Catal. Org. React. 1998, 75, 207. 24. M. Studer, U. Siegrist, Ciba-Geigy, unpublished results.

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66. P. N. Rylander, Catalytic Hydrogenation in Organic Synthesis, Academic Press, New York, 1979, p. 20. 67. E. N. Marvell, T. H. Li, Synthesis 1973, 457. 68. A. V. R. Rao, S. P. Reddy, E. R. Reddy, J. Org. Chem. 1986, 51, 4158 69. R. S. Mann, K. C. Khulbe, Can. J. Chem. 1967, 45, 2755. 70. R. S. Mann, K. C. Khulbe, J. Catal. 1969, 13, 25. 71. R. S. Mann, K. C. Khulbe, J. Catal. 1968, 10, 401. 72. R. S. Mann, K. C. Khulbe, J. Catal. 1970, 17, 46. 73. G. C. Bond, P. B. Wells, J. Catal. 1965, 5, 65. 74. W. Palczewska, in Hydrogen Effects in Catalysis, Z. Paal, P. G. Menon (Eds.), Marcel Dekker, New York, 1988, p. 383. 75. F. Roessler, DSM, unpublished results. 76. T. W. Greene, P. G. M. Wuts, Protective Groups in Organic Synthesis, 2nd Ed., Wiley, New York, 1991. 77. B. P. Czech, R. A. Bartsch, J. Org. Chem. 1984, 49, 4076. 78. Catalytica, Catalyst Modification: Selective Hydrogenation, Catalytica Studies Division, Mountain View, CA, 1992, p. 177. 79. W. M. Pearlman, Tetrahedron Lett. 1967, 17, 1663. 80. S. Hanessian, T. J. Liak, B. Vanasse, Synth. Commun. 1981, 11, 396. 81. R. C. Bernotas, R. V. Cube, Synth. Commun. 1990, 20, 1209. 82. K. Yoshida, S. Nakajima, T. Wakamatsu, Y. Ban, M. Shibasaki, Heterocycles 1988, 27, 1167. 83. V. S. Rao, A. S. Perlin, Can. J. Chem. 1983, 61, 652. 84. M. Bart´ok, Stereochemistry of Heterogeneous Metal Catalysis, Wiley, Chichester, 1985. 85. J. S. Bajwa, Tetrahedron Lett. 1992, 33, 2299. 86. D. Beaup`ere, I. Boutbaia, A. Wadouachi, C. Frechou, G. Demailly, R. Uzan, New J. Chem. 1992, 16, 405. 87. H. Steiner, Ciba-Geigy, unpublished results. 88. Nissan Chemical Industries, Japanese Patent 55 133 341, 1980. 89. K. Hattori, H. Sajiki, K. Hirota, Tetrahedron 2000, 56, 8433; H. Sajiki, K. Hattori, K. Hirota, Chem. Eur. J. 2000, 6, 2200, and references cited therein. 90. H. Sajiki, K. Hirota, Tetrahedron 1998, 54, 13981. 91. H. Sajiki, H. Kuno, K. Hirota, Tetrahedron Lett. 1997, 38, 399. 92. Y. Oikawa, T. Tanaka, K. Horita, O. Yonemitsu, Tetrahedron Lett. 1984, 25, 5397. 93. L. S. Seif, M. Partyaka, J. E. Hengeveld, Catal. Org. React. 1991, 40, 197. 94. Y. Fukayama, Tetrahedron Lett. 1999, 40, 105. 95. F. Cramer, D. Voges, Chem. Ber. 1959, 92, 952. 96. J. Kovacs, L. R. Rodin, J. Org, Chem. 1968, 33, 2418. 97. P. W. Erhardt, Synth. Commun. 1983, 13, 103. 98. L. Velluz, G. Amiard, R. Heymes, Bull. Soc. Chim. Fr. 1954, 1012. 99. W. G. Bornmann, M. E. Kuehne, J. Org. Chem. 1992, 57, 1752. 100. A. Boutros, J.-Y. Legros, J.-C. Fiaud, Tetrahedron 2000, 56, 2239. 101. E. A. Papageorgiou, M. J. Gaunt, J. Yu, J. B. Spencer, Org. Lett. 2000, 2, 1049, and references cited therein. 102. H. Sajiki, H. Kuno, K. Hirota, Tetrahedron Lett. 1997, 38, 399. 103. D. P. Curran, R. Ferritto, H. Ye, Tetrahedron Lett. 1998, 39, 4937. 104. M. A. Clark, B. Ganem, Tetrahedron Lett. 2000, 41, 9523. 105. Y. Sugi, S. Mitsui, Tetrahedron 1973, 29, 2041. 106. M. Besson, F. Delbecq, P. Gallezot, S. Neto, C. Pinel, Chem. Eur. J. 2000, 6, 949. 107. V. Hada, A. Tungler, L. Szepesy, Appl. Catal. A 2001, 210, 165.

108. L. Hegedus, V. H´ada, A. Tungler, T. M´ath´e, L. Szepesy, Appl. Catal. A 2000, 201, 107 (the originally claimed de value of 98% had to be corrected; A. Tungler, personal communication). 109. F. Glorius, N. Spielkamp, S. Holle, R. Goddard, C. W. Lehmann, Angew. Chem. Int. Ed. 2004, 43, 2850. 110. H. Steiner, P. Giannousis, A. Pische-Jacques, H. U. Blaser, Top. Catal. 2000, 13, 191, and references cited therein. 111. R. Burmeister, A. Freund, P. Panster, T. Tacke, S. Wieland, Stud. Surf. Sci. Catal. 1995, 92, 343. 112. F. M. Studer, H. U. Blaser, C. Exner, Adv. Synth. Catal. 2003, 345, 45. 113. G. H. Sedelmeier, H. U. Blaser, H. P. Jalett; European Patent 206 993, assigned to Ciba-Geigy, 1986. 114. P. Herold, A. F. Indolese, M. Studer, H. P. Jalett, H. U. Blaser, Tetrahedron 2000, 56, 6497; M. Studer, S. Burkhardt, A. F. Indolese, H. U. Blaser, Chem. Commun. 2000, 1327. 115. C. De Bellefon, P. Fouilloux, Catal. Rev. 1994, 36, 459. 116. A. M. Allgeier, M. W. Duch, Catal. Org. React. 2001, 82, 229. 117. S. N. Thomas-Pryor, T. A. Manz, Z. Liu, T. A. Koch, S. K. Sengupta, W. N. Delgass, Catal. Org. React. 1998, 75, 195. 118. T. A. Johnson, D. P. Freyberger, Catal. Org. React. 2001, 82, 201. 119. O. G. Degischer, F. Roessler, European Patent 1 108 469, assigned to Hoffmann-La Roche, 2001. 120. G. D. Yadav, M. R. Kharkara, Appl. Catal. A 1995, 123, 115. 121. P. Tinapp, in Houben-Weyl, Methoden der organischen Chemie, Vierte Auflage, Reduktionen Teil 1, Band IV/1c, Georg Thieme, Stuttgart, 1980, p. 111. 122. W. Poepel, J. Gaube, Dechema Monogr. 1991, 122, 189. 123. B. Fell, J. Sojka, Fett Wiss. Technol. 1991, 93, 79. 124. P. Kukula, M. Studer, H. U. Blaser, Adv. Synth. Catal. 2004, 245, 1487. 125. P. Kukula, K. Koprivova, J. Catal. 2005, 234, 161.

14.10.3

Regioselective Hydrogenations .. .. Peter Claus∗ and Yucel Onal

Generalities in Regioselectivity with Heterogeneous Catalysts

14.10.3.1

14.10.3.1.1 Definitions The most fascinating characteristic of a catalyst is its ability to favor selectively one among various competitive or consecutive reactions. This ability of a catalyst plays a decisive role in the chemical industry, where selective hydrogenation reactions of more or less complicated organic reactants are carried out in order to produce bulk and fine chemicals. Selectivity in chemical reactions can be classified in terms of chemo-, regioand stereoselectivity. Whereas stereoselectivity reflects the production of one stereoisomer of the product in preference to another, chemoselectivity is the ability of a catalyst to discriminate between different functional groups (Scheme 1). ∗

Corresponding author.

14.10.3 Regioselective Hydrogenations

Thereby, one of the functional groups is preferably hydrogenated in a competitive reaction, e.g. a carbonyl group in the presence of another reactive group such as an olefinic bond. A chemical reaction is called regioselective when a reactant which exhibits two (or more) identical functional groups at different positions is converted preferentially (regioselectively) at one of the identical functional groups. However, the term regioselectivity also applies to conjugated systems of double bonds belonging to different functional groups, as in the case of α, βunsaturated aldehydes: their hydrogenation at either the olefinic or the carbonyl group is regioselective because the addition of hydrogen takes place at a given position of the conjugated molecule. The aim of this chapter is to review some recent and important developments in the field of regioselective hydrogenations with heterogeneous catalysts, which are usually supported metal catalysts and, to a much lesser extent, oxides (e.g. in transfer hydrogenations). This excludes regioselective reductions by stoichiometric agents (metal hydrides) and the application of soluble metal complexes in homogeneously catalyzed hydrogenations and such reactions which are carried out in a multiphase gas/liquid/liquid (G/L/L) mode. Selectivity in heterogeneous catalysis is always (i) the result of the concentration of the active sites and their (electronic and geometric) structure, i.e. a function of the catalyst design, and (ii) the result of the concentration of the reactants at the active sites, i.e. a function of catalytic reaction engineering aspects (process parameters, reaction kinetics, mass and heat transport, reactor design). Therefore, the latter will also be considered in this chapter. Some of the factors controlling the intramolecular selectivity, i.e. the preferred hydrogenation of one functional group compared with another of a molecule with conjugated double bonds, are well understood and will be presented in the following sections exemplified by certain reactions.

H2C C CH2

H2C CH CH3

H2C CH CH CH2

H2C CH CH CH3

H2C CH C

O H

H2C CH CH2OH

NO2

NO2 Scheme 1

NH2

NO2

3309

14.10.3.1.2 General Factors Determining Selectivity Heterogeneously catalyzed hydrogenations of olefinic double bonds and other unsaturated functionalities on a catalytically active surface involve a reaction mechanism according to Horiuti and Polanyi [1]. First, the reactant is adsorbed on the catalyst surface via π- or σ -bonding, which is followed by the addition of chemisorbed hydrogen. The half-hydrogenated species can now be either (i) hydrogenated into the saturated product or (ii) dehydrogenated (depending on the molecular structure), a process that can be accompanied by double bond migration or cis/trans isomerization. One of the main factors determining the course of the mechanism is the degree of hydrogen coverage. Parameters influencing the hydrogen coverage are – in addition to the metal character (see Section 14.10.3.3.2) – the partial pressure of hydrogen, the agitation speed and the texture of the support. At high surface coverage of hydrogen, the reactant undergoes preferably hydrogenation, whereas at low hydrogen coverage isomerization or even hydrogenolysis of the reactant can be observed. The influence of hydrogen coverage on the selectivity was clearly shown for the interplay between hydrogenation and selective isomerization of linoleic acid to conjugated linoleic acids (CLAs) over Ru/C and Ag/SiO2 catalysts [2]. Thereby, the hydrogen coverage of the catalyst surface was controlled either by a special pretreatment of the catalyst prior to reaction (Ru) or by the nature of the catalyst metal (Ru versus Ag). In experiments described in the literature with heterogeneous ruthenium catalysts supported on Al2 O3 and carbon [3], the catalyst was first covered with hydrogen, followed by the conversion of linoleic acid to CLA under nitrogen. In order to produce CLA at all, this two-step reaction is necessary because of the catalyst properties and the subsequent reactions competing with the desired isomerization. On the one hand, the excellent hydrogenation qualities of the ruthenium catalysts result in the fact that if the reaction is carried out under hydrogen, linoleic acid is hydrogenated completely and very quickly into oleic acid and consecutively to stearic acid. On the other hand, obviously the presence of only a small quantity of hydrogen is required for the isomerization of linoleic acid to cis−9, trans−11and trans−10, cis−12-CLA. In contrast, since hydrogen chemisorption on Ag is very weak, the catalyst must not be precovered with a small amount of hydrogen as in the case of Ru and the reaction could be carried out in presence of hydrogen with the same high selectivity to CLA. Scheme 2 illustrates how weakly bonded hydrogen (on Ag) causes the desired CLA formation via an addition/elimination mechanism, whereas strongly bonded hydrogen (on Ru) causes the formation of oleic acid via consecutive H addition. The adsorption properties of hydrogen also play References see page 3327

3310

14.10 Hydrogenation Reactions

R2

R1

cis 9-C18:1 Oleic Acid + H

*

1:

R H3C(CH2)4 R2: (CH2)7COOH R1

1

R H

R2

H *

+ H

R2

*

*

H H R1

cis9, cis12-C18:2 Linoleic Acid

H

R2

* − H

* R2 R1

cis9, trans11-CLA

H H R1

R2 H * *

Scheme 2

R

R Cl

NO2

Cl

H2

NH2

Scheme 3

Relative activity 74

1.3

1

0.002

Scheme 4

a crucial role in controlling the intramolecular selectivity of the regioselective hydrogenation of α, β-unsaturated aldehydes (see Section 14.10.3.3.2). Competitive reaction pathways, namely hydrogenation versus hydrogenolysis as a function of the degree of

hydrogen coverage, was also observed in the hydrogenation of halonitrobenzenes (Scheme 3). Selectivity to the amine is high if hydrogenolysis of the C−halogen bond, which is favored at low hydrogen coverage on the surface, can be avoided. Therefore, regioselective hydrogenations with these organic molecules must be carried out at high hydrogen pressures [4]. The chemical structure of the molecule plays an important role in regioselective hydrogenations. In the case of alkene hydrogenation, the double bond which exhibits the lowest degree of substitution and which is therefore least hindered is preferably hydrogenated. Hence, external C=C double bonds are hydrogenated faster than internal olefinic bonds. This was shown for the hydrogenation of 1-hexene in comparison with 2-methyl-1-pentene and 2,3-dimethyl-2-butene [5], exhibiting relative reactivities in the range 74–0.002 depending on the degree of substitution (Scheme 4). The impact of steric hindrance on the selectivity in the hydrogenation of alkenes was also shown for methyl 2,4-hexadienate on carbon supported Pt, Pd, Rh and Ru (Scheme 5) [6]. Because of the lower degree of substitution, the double bond between C4 and C5 was hydrogenated with higher selectivity. Furthermore, exocyclic double bonds are reduced more easily than endocyclic bonds. This applies to the hydrogenation of limonene on Raney nickel, exhibiting 96% selectivity to carvomenthene (Scheme 6) [7].

COOCH3 COOCH3

Pt, Pd, Rh, Ru/C COOCH3 COOCH3

Scheme 5

14.10.3 Regioselective Hydrogenations

The degree of substitution at the olefinic double bond also plays a decisive role in the selective hydrogenation of α, β-unsaturated aldehydes to the corresponding unsaturated alcohols on several carbon- or metal oxidesupported Group VIII metals. It turned out that the selectivity to the unsaturated alcohol increases with increasing steric hindrance at the C=C double bond. Thus, selectivities obtained with heterogeneous catalysts are worst for the hydrogenation of acrolein on conventional hydrogenation catalysts, whereas hydrogenation of crotonaldehyde, prenal, cinnamaldehyde and citral can be achieved with reasonably high selectivities [8–10]. Therefore, it must be kept in mind that in the case of α, β-unsaturated aldehydes higher than acrolein the selectivities reported do not represent the intrinsic selectivity of the metal (see Section 14.10.3.3.2) because the substituent on the C=C double bond contributes to the overall selectivity by steric repulsion of the olefinic group changing the adsorption configuration of the aldehyde, as shown by a number of experimental data and theoretical calculations [9–14]. Many heterogeneously catalyzed hydrogenation reactions are carried out in multiphase G/L/S reaction systems comprising a liquid solution with the reactant. Since hydrogenation reactions are exothermic, the dilution is useful to obtain better control of the reaction temperature. Thereby the solvent characteristics can have an important impact on the observed selectivity [15]. One of the main factors controlling the rate of hydrogenation and the selectivity is the solubility of hydrogen in the liquid phase determining the hydrogen coverage on the catalyst surface. Polarity and acid–base features of the solvent can also influence the adsorption equilibrium of the educt and product on the catalyst surface, thereby influencing the selectivity of the reaction. For example, in the presence of pyridine, hydrogenation of nitroaromatics can stop at the hydroxylamine stage, which is easily desorbed [16]. Another example is the hydrogenation of isoquinoline (Scheme 7). In hydrochloric acid solution, isoquinoline is transformed into the cyclohexyl derivative with a selectivity S of 97%, since the chemisorption of the protonated pyridine ring on the metal surface is suppressed, whereas in methanol the selectivity to the phenyl derivative is 87% [17]. Recently, the selective hydrogenation of geraniol over a Cu/Al2 O3 catalyst under mild hydrogenation conditions

Raney-Ni H2 0.1 MPa

Limonene Scheme 6

Carvomenthene

Pt, HCl H2

N

3311

Pt, MeOH H2

N

NH

Scheme 7

was reported (Scheme 8) [18]. It was demonstrated how the selectivity of the reaction with respect to citronellol and menthol could be switched based on the basicity of the solvent used, which is due to the complex reaction mechanism in the liquid phase. Whereas geraniol could be converted in n-heptane to a mixture of citronellol (S = 56%) and menthol (S = 30%), in the presence of 2-propanol citronellol was formed almost quantitatively (S > 98%). However, the reaction rate was slower when 2-propanol was used as solvent. The same selectivity with respect to citronellol also applied to nerol and citronellal hydrogenation. In order to achieve selectivity at a maximum level, besides optimization of the catalyst, the reaction conditions, including reaction temperature, reactant concentration and hydrogen partial pressure, have to be thoroughly tuned. If the reaction is carried out at very high temperatures, all functional groups will be hydrogenated regardless of steric hindrance. On the other hand, an appropriate reaction temperature has to be maintained in order to run the reaction at an acceptable rate. Therefore, for each reaction the influence of reaction temperature on the activity and selectivity has to be evaluated and optimized. Another point which has to be taken into account is catalyst deactivation with increasing reaction temperature. Thus, in the hydrogenation of α, β-unsaturated aldehydes, decarbonylation of the unsaturated alcohol can induce significant deactivation of the catalyst, since active sites of the catalyst surface become reversibly blocked by CO. For example, an activity minimum was observed at 100 ◦ C in citral hydrogenation on Pt/TiO2 −LTR and Pt/SiO2 [15]. With increasing reaction temperature, more CO is generated, which adsorbs on the catalyst surface. Increasing the temperature further results in a faster desorption rate of CO than the decarbonylation rate, thus leading again to an increased hydrogenation rate. Hydrogen partial pressure determines not only the rate of reaction in G/L/S systems according to Henry’s law, but in some cases also the selectivity of a heterogeneously catalyzed gas-phase reaction. For example, regioselective hydrogenation of acrolein to allyl alcohol was examined in the gas phase on an Ag/SiO2 catalyst in the pressure range 50 mbar–20 bar [19–21]. The selectivity to allyl alcohol increased with increasing acrolein and hydrogen partial pressure, which was explained in terms of a pressure-dependent adsorption geometry of acrolein References see page 3327

3312

14.10 Hydrogenation Reactions

OH

Cu/Al2O3

+

H2, 90 °C

Geraniol

OH

Citronellol

OH

Menthol

Scheme 8

determining the intramolecular selectivity of the reaction. In the case of crotonaldehyde hydrogenation at 413 K over an Rh−Sn/SiO2 catalyst, a marked shift in selectivity to the allylic alcohol from 12 to 66% was observed on increasing the total pressure from 0.1 to 2 MPa and, thus, the corresponding partial pressures in the hydrogenation [9]. At 2 MPa, the turnover frequency (TOF) for the formation of crotyl alcohol was increased by a factor of 5.3, whereas that for the formation of butyraldehyde was decreased to about one-third of the specific activity at 0.1 MPa. In a similar way, the selectivity to unsaturated alcohols can be influenced by varying the concentration of the reactant in the liquid phase. This behavior was observed in selective cinnamaldehyde hydrogenation to cinnamyl alcohol on Pt/SiO2 in ethanol at 50 ◦ C and 29 bar hydrogen partial pressure and on Pt/C/monolith in toluene at 30 ◦ C and 50 bar hydrogen partial pressure [22, 23], showing an increase in selectivity with increase in the initial reactant concentration. Again, the concentration-dependent adsorption structure of cinnamaldehyde might be responsible for the change in selectivity. Thus, at higher reactant concentrations the self-assembly of the aromatic ring is preferred, resulting in an end-on mode of adsorption of cinnamaldehyde on the catalyst surface. Reaction Engineering and Selectivity One of the important factors which is often neglected in the design and preparation of catalysts in order to improve productivity and selectivity is an appropriate reactor concept for carrying out the reaction. Hydrogenation reactions in the liquid phase are often performed batchwise in G/L/S multiphase slurry reactors. Although the reactor concept is very facile for fast and efficient testing of catalytic reactions, it does not allow precise control of the contact time of the reactant with the catalyst. Hence the products remain in the reaction medium and can easily readsorb on the catalyst, decreasing the selectivity of the catalyst. Therefore, when selectivity to a certain product is emphasized in competitive or consecutive reactions, innovative reactor concepts which are often driven in continuous mode should be considered in addition to catalyst design. In general, the selectivity to an intermediate 14.10.3.2

product in consecutive hydrogenation reactions can be enhanced by applying accurately determined residence times. Simultaneously, the continuous mode allows an easier way to record and model catalyst stability and deactivation. Furthermore, catalyst separation from the products is often redundant. Trickle-bed reactors were used in the liquid-phase hydrogenation of citral [24], crotonaldehyde [25] and benzaldehyde [26, 27]. In citral hydrogenation selectivities towards nerol and geraniol were about 97% at a conversion of 97% on a Pt−Sn/MgO bimetallic catalyst at 100 ◦ C and 20 bar hydrogen partial pressure. In order to achieve better wetting of the catalyst and hence better effectiveness, crotonaldehyde hydrogenation was carried out with periodic flow interruptions, giving rise to enhanced productivities on a Pd/γ -Al2 O3 catalyst at 25 ◦ C and 11 bar hydrogen partial pressure. Benzaldehyde hydrogenation was examined as a model reaction system in order to evaluate the potential of a trickle bed reactor in comparison with a monolith reactor [27]. Ni/γ -Al2 O3 was embedded in the former whereas for the latter a washcoated Ni−cordierite catalyst was applied. The monolith reactor provided beneficial characteristics with respect to both productivity and selectivity to benzyl alcohol at 100 ◦ C and 10 bar hydrogen partial pressure due to the sharper residence time of the fluid within the monolith channels. Thereby the monolith with the higher channel density (600 cpsi) gave rise to an improved selectivity with respect to benzyl alcohol compared with monoliths having 400 cpsi. The effect can be explained in terms of a shorter diffusion path of the reactant to the active sites on the monolith, thus enhancing the selectivity to the intermediate hydrogenation product. Additionally, monolith reactors exhibit a lower pressure drop, reduced catalyst attrition and easier scale-up approach. The advantageous effects in using monolithic type catalysts were further demonstrated in a screw-impeller tank reactor with regard to citral hydrogenation on Ni monoliths [28]. The selective hydrogenation to citronellal was examined under different reaction conditions (temperature and hydrogen partial pressure). The best results (S = 96%, X = 80%) were obtained at 40 ◦ C and 5 bar hydrogen partial pressure. By increasing the partial pressure further, enhanced

14.10.3 Regioselective Hydrogenations

FBR AC H2

PBMR dosing of AC AC AC H2 AyOH H2

3313

PBMR dosing of H2 H2 AC H2 AyOH AC

AC

PA

PA

AC

H2

AC H2 AyOH PA

H2

Scheme 9

selectivity to the consecutive hydrogenation product 3,7dimethyloctanal was observed. Micro-structured reactors exhibit beneficial characteristics with respect to heat and mass transport. Accurate adjustment of reaction conditions and narrow residence time distributions provide safe operation, simultaneously giving rise to enhanced selectivities in complex reaction systems. A micro-structured, carbon-coated reactor was used for the gas-phase hydrogenation of acrolein to allyl alcohol [29]. Ru deposition was performed via ion exchange with [Ru(NH3 )4 Cl]Cl2 . The product distribution obtained was comparable to that measured in a fixed-bed reactor (FBR) on a Ru/C catalyst. The beneficial potential of packed-bed membrane reactors (PBMRs) in contrast to conventional co-feed fixedbed reactors (FBRs) was recently demonstrated for the selective gas-phase hydrogenation of acrolein to allyl alcohol based on theoretical calculations and kinetic data from simple power rate laws [30]. Thereby the conventional co-feed mode of the reactants into the catalyst bed is abandoned and one of the reactants, acrolein or hydrogen, is added to the reactor in a distributed manner over the tubular membrane (Scheme 9). The decision as to which reactant is fed directly into the reactor and which one through the membrane is made by first analyzing the kinetics of the reaction network. Thus, Ag/SiO2 was chosen as model catalyst and the kinetics of the reaction with regard to allyl alcohol and propionaldehyde formation were first determined in the FBR reactor according to a simple power law. An

expression for the differential selectivity to allyl alcohol then revealed that the largest impact on selectivity can be achieved by increasing the local acrolein partial pressure. Then, according to the calculated kinetic parameters and using a 1D isothermal and isobaric tubular reactor model, the reaction system was investigated theoretically analyzing different dosing strategies. Both single- and multi-stage PBMR reactors provided enhanced yields with respect to allyl alcohol when hydrogen was dosed in excess through the reactor wall. The effect was attributed to a beneficial change of the local acrolein concentration and to increased residence time profiles. After establishing an appropriate reactor design for the hydrogenation reaction considered, in the following sections the catalytic properties will be discussed in te