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CHM 674 ADVANCED ELECTROCHEMISTRY EXPERIMENT 1: GALVANIC CELL

GROUP MEMBER NUR AINA BINTI MOHAMAD THAREEQ HAKEEM BIN SHAJDEE WELLONICA OLGA WILLIE

NAME OF LECTURER

: DR. NYOTIA NYOKAT

GROUP

: AS202 6A

DATE OF EXPERIMENT

: 4th MARCH 2020

STUDENT NUMBER 2017411794 2016709479 2017412048

Objectives a)

To measure the relative reduction potentials for a number of half cell (redox) couples in a galvanic cell.

b)

To develop an understanding of the movement of electrons, anions and cations in a galvanic cell.

c)

To study factors affecting cell potentials.

d)

To estimate the concentration of ions in solution using the Nernst Equation.

Introduction Galvanic cell is a cell that allow us to harness the electron flow in a redox reaction to perform useful work. Such cells find common use as batteries, pH meters, and as fuel cells. The setup of the cell requires that the oxidation and reduction half-reactions are connected by a wire and by a salt bridge. Electrons will flow through that wire creating an electrical current. The salt bridge allows the passage of ions in solution to maintain charge neutrality in each halfcell. Instead of drawing the setup of every cell like, chemists have devised a shorthand line notation discussed in Line Notation. The direction of the current in a cell is determined by the standard reduction potential of each half-cell. For a reaction to be spontaneous, the overall cell potential must be positive. Therefore, the half-reaction with the greater reduction potential will be a reduction and the other half-reaction will be an oxidation. The electrode in the oxidation half-reaction is called the anode. The electrode in the reduction half-reaction is called the cathode.

Figure 1: Galvanic Cell

Methodology A. Galvanic Cell - Reduction Potentials of Several Redox Couples The cell potential for a number of galvanic cells were measured and the redox couples were placed in order of decreasing reduction potentials. Firstly, the electrodes, solutions and equipment were collected. Four small beakers were obtained, and they were filled with 0.1M solutions. These solutions were shared with other groups of students in the laboratory. The strips of copper, zinc and iron metal was polished with sandpaper, then they were rinsed with deionized water. These polished metals will be used as electrodes and they were put in respective beakers. A potentiometer was checked out with two electrical wires attached to crocodile clips. After that, the copper/Zinc cell was set up. A Cu strip (electrode) was placed in a Cu (NO3)2 solution and a Zn strip (electrode) in the Zn (NO 3)2 solution. A piece of filter paper was rolled and flatten, then it was wet with 0.1M KNO 3 solution. The ends of the filter paper were folded and inserted into the solutions in the two beakers; this is the salt bridge. One electrode was connected to the negative terminal of the potentiometer and the other to the positive terminal. Then, the copper/zinc cell potential was determined. If the voltmeter reads a negative potential, the connections has been reversed. The positive cell potential was read and recorded. The metal strips that serve as the cathode (positive terminal) and the anode was identified. An equation for the half-reaction occurring at each electrode was wrote. The two half-reactions were combined to write the equation for the cell reaction. After that, the procedure was repeated for the remaining cells. The cell potentials for all possible galvanic cells that can be constructed from the other redox couples was determined. A ‘new’ salt bridge was prepared for each galvanic cell. Next, the relative reduction potentials were determined. The reduction potential of the Zn+2 (0.1M)/Zn redox couple has been assume with -0.79V. The reduction potentials of all other redox couples was determined. Lastly, the waste zinc, copper, magnesium and iron solutions were disposed in the ‘Waste Metal Solutions’ container. The metals were returned to appropriately market containers.

B. Effect of Concentration Changes on Cell Potential The effects of changes in ion concentrations on cell potentials were observed and analysed. Firstly, the experiment was conducted on the effect of different molar concentrations. The galvanic cell was set up using 1M CuSO 4 and 0.001M CuSO4 solutions. A polished copper electrode was immersed in each solution. A salt bridge was prepared to connect the two redox couples. The cell potential was measured. The anode and the cathode were determined. An equation for the reaction occurring at each electrode was wrote. Lastly, the experiment was conducted on the effect of complex formation. 5ml of 6M NH3 was added to the 0.001M CuSO 4 solution, until any precipitate re-dissolves. Any changes in the half-cell and the cell potential was observed.

C. The Nernst Equation and an Unknown Concentration Firstly, solutions 1 through 4 as outlined was prepared using a 1ml pipet and 100ml volumetric flasks. It has been sure to rinse the pipet with the more concentrated solution before making the transfer. Deionized water was used for dilution ‘to the mark’ in the volumetric flasks. The molar concentration of each solution was calculated. After that, the cell potential for solution 4 was measured and calculated. The experiment was set up using small beakers. The Zn 2+/Zn redox couple is the reference half-cell for this part of the experiment. The two half-cells were connected with a ‘new’ salt bridge. The electrodes were connected to the potentiometer and the potential difference, Ecell was recorded. The theoretical cell potential was calculated. Next, the cell potentials for solutions 3 and 2 was measured and calculated. Part C2 was repeated with solutions 3 and 2 respectively. A freshly prepared salt bridge is required for each cell. Then, Ecell (measured) and Ecell (calculated) versus log [Cu2+] was plotted on a graph paper for the four concentrations of Cu (NO3)2. Lastly, a Cu (NO3)2 solution was obtained with an ‘unknown’ copper ion concentration from lab and a galvanic cell was set up. E cell was determined as in Part C2. Using the graph, the unknown copper ion concentration in the solution was determined.

Result and Observation A. Galvanic Cell – Reduction Potentials of Several Half Cell Couples 1. Observations and interpretations from the galvanic cells. Galvanic Measured Anode Equation for Cell Ecell Anode Reaction Cu-Mg Cu-Fe Mg-Fe

0.3 V 0.1 V 0.15 V

Mg Fe Mg

−¿ ¿

Mg → Mg2 +¿+2e Fe → Fe 2+¿+2 e ¿ Mg → Mg2 +¿+2e

¿

−¿¿

−¿ ¿

Cathode Equation for Cathode Reaction ¿ Cu Cu2+¿+2 e ¿ Cu Cu2+¿+2 e ¿ Fe Fe2+ ¿+2 e −¿→Cu ¿

−¿→Cu ¿

¿

−¿→ Fe ¿

Overall equations for the three cell reactions. Cu-Mg: Mg+ Cu2 +¿→ Mg Cu-Fe: Fe+Cu 2+¿ → Fe Mg-Fe: Mg+ Fe 2+¿→ Mg

2+ ¿+Cu ¿

2+ ¿+Cu ¿

¿

¿

2 +¿+ Fe ¿

¿

2. Arrange the three half cells in order of decreasing reduction potentials. List the reduction potential for each half cell couple relative to that of the Zn2+¿¿(0.1M)/Zn couple, which is – 0.79 V. Use a table of standard reduction potentials and the Nernst equation to calculate the reduction potentials for each of these half-cell couples. Redox Couple Cu-Mg Mg-Fe Cu-Fe

Reduction Potential (Measured) 0.3 V 0.15 V 0.1 V

Calculation for Reduction Potential:

Reduction Potential (Calculated) 2.71 V 1.93 V 0.78 V

% Error 88.93% 92.23% 87.18%

1. Cu-Mg Ecell = Ecation - Eanion = 0.34 - (-2.37) = 2.71 V 2. Mg-Fe Ecell = Ecation - Eanion = -0.44 – (-2.37) = 1.93 V 3. Cu-Fe Ecell = Ecation - Eanion = 0.34 - (-0.44) = 0.78 V Calculation for % Error: |2.71−0.3| × 100 % 1. % Error = 2.71 = 88.93% |1.93−0.15| × 100 % 2. % Error = 1.93 = 92.23% |0.78−0.1| × 100 % 3. % Error = 0.78 = 87.18%

B. Effect of Concentration Changes on Cell Potential 1. Cell potential of concentration cell: 0.05 V Anode cell: Cu→ Cu2+¿+2 e ¿ ¿ Cathode cell: Cu2+¿+2 e −¿¿

−¿→Cu ¿

Potential is recorded in order to show there is increase in concentration that cause the equilibrium to make the electrode potential more positive. And a decrease in concentration must have the opposite effect which is make it more negative. 2. Cell potential from complex formation: 0.5 V Observation of solution in half-cell: The colour of CuSO4 solution changes from colourless to light blue colour. The potential changes with the addition of NH 3 cause any change in the concentration of Cu2+ ions will alter the electrode potential. More Cu 2+ ions flow from cathode and the concentration of Cu2+ ions increase. 3. How would the cell potential have been affected if the NH 3 had been added to the 1M CuSO4 solution instead of the 0.0010 M CuSO4 solution of the cell? Explain. The cell potential has been affected or decreases after adding the NH3 into 1M CuSO4 solution instead of the 0.0010M CuSO 4 solution because the mole concentration of 1M CuSO4 are higher than 0.001M CuSO4.

C. The Nernst Equation and an Unknown Concentration 1. Complete the following table with the concentrations of the Cu (NO3)2 solutions and the measured cell potentials. Use a table of standard reduction potentials and the Nernst equation to calculate the Ecell. Solutio n Number 1 2 3 4

Concentration of Cu(NO3)2

log ¿

0.1 mol/L 0.001 mol/L 0.00001 mol/L 0.000001 mol/L

Ecell (measured) -1 -3 -5 -7

Calculation of concentration of Cu (NO3)2: 1. Solution 2 M1V1=M2V2 (0.1) (0.001) = M2(0.1) M2 = 0.001 mol/L 2. Solution 3 M1V1=M2V2 (0.001) (0.001) = M2(0.1) M2 = 0.00001 mol/L 3. Solution 4 M1V1=M2V2 (0.00001) (0.001) = M2(0.1) M2 = 0.0000001 mol/L Calculation of Ecell: Cu – Zn Anode: Zn→ Zn2 +¿+2 e ¿ ¿ Cathode: Cu2+¿ 2e Overall equation: Zn+ Cu2 +¿→ Zn −¿ ¿

−¿ →Cu¿

2+ ¿+Cu ¿

E° cell = E° cathode - E° anode E° cell = 0.34 V- (-0.76 V) E° cell = 1.1 V Calculation of calculated E° cell: 0.0592 ln Q Ecell = E ° cell− n Solution 1

¿

0.35 V 0.30 V 0.25 V 0.15 V

Ecell (calculated) 1.1 V 1.04 V 0.98 V 0.92 V

0.0592 0.1 ln 2 0.1 Ecell = 1.1 V −0.0296 log 1 Ecell = 1.1 V Solution 2 0.0592 0.1 ln Ecell = 1.1 V − 2 0.001 Ecell = 1.1 V −0.0296 log 100 Ecell = 1.04 V Ecell = 1.1 V −

Solution 3 0.0592 0.1 ln 2 0.00001 Ecell = 1.1 V −0.0296 log 10000 Ecell = 0.98 V Ecell = 1.1 V −

Solution 1 0.0592 0.1 ln 2 0.0000001 Ecell = 1.1 V −0.0296 log 1000000 Ecell = 0.92 V Ecell = 1.1 V −

2. From Plot of Ecell versus log ¿ - Account for any significant difference between the measured and calculated Ecell values. Calculated value is theoretical value while measured value is practical value. 3. Ecell for the solution of unknown concentration: Molar concentration of Cu2+ in the unknown: 4. How would you adjust the concentrations of Cu 2+ and Zn2+ for the Cu-Zn cell to maximize the cell potential? Explain. Adjust the concentrations of Cu2+ and Zn2+ for the Cu-Zn cell by increases [Zn2+] and decreases [Cu2+] to maximize the cell potential using spontaneous reaction. If the Zn half-cell is diluted, lower [Zn2+], the reaction is moving away from equilibrium and so the cell potential goes up. If Cu half-cell is diluted, lower [Cu2+], the reaction moving towards equilibrium, and so the cell potential goes down.

Discussion In this experiment, a galvanic cell is conducted in order to measure the relative reduction potentials for a number of half-cell couples in a galvanic cells, to develop an understanding of the movement of electrons, anions and cations in a galvanic cell, to study factors affecting cell potentials and to estimate the concentration of ions in solution using the Nernst Equation. There are three parts of experimental procedure, which are part A that is related to the reduction potentials of several redox couples, part B which are related to the effect of concentration changes on cell potential and part C that is related to the Nernst Equation and an unknown concentration. Based on the result and observation that is recorded from the experiment in part A, it can be concluded that there is huge difference between measured and calculated reduction potential. For redox couple between Cu-Fe, the reduction potential differences between measured and calculated is not as big as Cu-Mg and Mg-Fe. The percentage error for Cu-Fe redox couple is 87.18%. However, Mg-Fe redox couples shows huge percentage error with the value 92.23%. Meanwhile the percentage error for redox couple Cu-Mg is 88.93%. These percentage error are quite high due to several factor that cause error while conducting the experiments. Error has occurred due to rusty electrodes and not well-polished electrodes. In experiment procedure part B, the cell potential of ‘concentration cell’ is 0.05 V and the cell potential from complex formation gives 0.5 V value. The colour change of copper sulphate solution in half-cell is recorded. The colour change from colourless to light blue colour solution after 5 mL of 6 M of ammonia solution is added into 0.001 M copper sulphate solution. Furthermore, the cell potential is changed after the addition of 6 M of ammonia solution. Ammonia solution with 6 M of concentration is really high enough to raise the cell potential value.

In experiment procedure part C, Nernst equation can be calculated when the condition of concentration of the solution used is vary. The highest concentration value produced higher value of measured cell potential as well as the calculated cell potential, whereas increasing of concentration will increase the cell potential. From the result of this experiment, the result is different from the cell potential value, which is due to some factors. The wire used is to rusty and can create more resistance towards the electron flow from the outer circuit. During dilution process, the position of eye is not perpendicular to the meniscus that cause concentration obtained is not accurate. Next, in order to reduce the error, occur, make sure the apparatus and material used is in a good condition so that the result obtained will not be affected.

Conclusion By the end of this experiment, we are able to measure the relative reduction potentials for a number of half cell couples in galvanic cell. Besides, we are able to develop our understanding on the movement of electrons, anions and cations in a galvanic cell. Moreover, we are able to study the factors affecting cell potential and finally able to estimate the concentration of ions in solution using the Nernst equation.

References 1. https://www.sparknotes.com/chemistry/electrochemistry/galvanic/summary/ 2. https://www.uccs.edu/Documents/chemistry/nsf/106%20Expt9V-GalvanicCell.pdf 3. https://www.academia.edu/32311201/galvanic_and_electrolytic_cell_exp_1_cmt555.docx 4. Schauer, F., Gerhatova, Z., Ozvoldova, M., Cernansky, P., & Tkac, L. (2012). Electrochemistry remote experiment - galvanic cell - II. 2012 9th International Conference on Remote Engineering and Virtual Instrumentation (REV), 1–6. https://doi.org/10.1109/rev.2012.6293184