The potential difference between the two half-cells drove electrons through the conducting wire from the negative copper wire to the positive nichrome wire. Thus, the chemical energy was converted into electrical energy. A filter paper soaked with saturated KNO3 was dipped into the two solutions acting as the salt bridge. It could enable electrical conduction and prevent direct mixing of two solutions at the same time. If the two solutions were directly mixed, spontaneous reactions occurred in one cell and most energy was released as the less useful heat.
With changes in concentration of the chemical species, the equilibrium position for
the above half reactions would change. Therefore, the electrode potential of the half-cell would also change. The change in electrode potential could then be predicted by the Nernst equation.
For a Cu(s)/ Cu 2+(aq) half-cell,
E = E0 + {0.059 log [Cu 2+(aq)]}/n E: electrode potential, E0: standard electrode potential
n: number of electrons involved
For a Fe3+ (aq)/ Fe2+ (aq) half-cell,
E = E0 + {0.059 log [Fe3+(aq)]/ [Fe2+ (aq)] /n }
For the overall cell reaction with concentration of CuSO4 kept at 1M
Ecell = E0cell + {0.059 log [Fe3+(aq)]2/ [Fe2+ (aq)] 2 } / n
Given E0 Fe(II)/ Fe(III) = +0.77 V and E0 Cu/Cu(II) = +0.34 V,
Ecell = (0.77-0.34) + {0.059 log [Fe3+(aq)]2/ [Fe2+ (aq)] 2 } / n
= 0.43 + {0.059 x2 x log [Fe3+(aq)]/ [Fe2+ (aq)] } / n
When a graph of Ecell was plotted against log [Fe3+(aq)]/ [Fe2+ (aq)], the trend of change of cell e.m.f with the variation of concentration ratio [Fe3+(aq)]/ [Fe2+ (aq)] could be investigated. As the slope of the curve equaled to 0.059/ n, the number of electrons involved in the Fe2+ (aq)/ Fe3+ (aq) equilibrium could be found out.
Results:
Interpretation:
The above graph was a straight line with negative slope. This showed that with decreasing concentration ratio of Fe2+ (aq)/ Fe3+ (aq), the e.m.f measured decreased.
This agreed with the Nerst equation, Ecell = 0.43 + {0.059 x2 x log [Fe3+(aq)]/ [Fe2+ (aq)] } / n ,
when [Fe3+(aq)]/ [Fe2+ (aq)] decreased, log [Fe3+(aq)]/ [Fe2+ (aq)] decreased and Ecell decreased.
This could also be proved from the equilibrium equation.
2Fe3+ + Cu(s) 2Fe2+ (aq) + Cu2+ (aq)
When the ratio [Fe3+(aq)]/ [Fe2+ (aq)] decreased and concentration of [Cu2+ (aq)] being unchanged, the proportion of [Fe3+(aq)] to [Fe2+ (aq)] decreased. Thus, the equilibrium position would shift to the left. The standard electrode potential of the system to be measured refer to reduction and a more positive value of electrode potential meant the reaction tend to proceed in the direction of reduction. Therefore, when the equilibrium position shifted to the left, the electrode potential of cathode and hence the cell e.m.f became less negative.
From the graph, the slope = 0.104
∵slope = 0.0592 x 2 / n
0.104 = 0.0592 x 2 / n
n = 1.14
~ 1
∴The number of electrons involved in the Fe2+ (aq)/ Fe3+ (aq) equilibrium was 1.
This agreed with the equation Fe3+ + e- Fe2+ (aq).
Discussion:
Criterion for predicting polarities of electrodes
In the beginning of the experiment, it was assumed that the copper wire was the negative pole. As the e.m.f measured in all cases showed a positive value, the assumption was correct and copper was
indeed the negative pole and the nichrome wire was the positive pole. For the copper wire, the electrode was negative and had a higher tendency to lose electron to undergo oxidation. As the electrode potential refer to reduction, the electrode potential should be lower. However, if a negative value of cell e.m.f was recorded, the assumption was not true. The copper should be the positive pole and the nichrome wire should be the negative pole instead. The other electrode had a higher tendency to lose electron to undergo oxidation. Then, the electrode potential of copper should be higher than the other.
Discrepancy between literature and experimental value of cell e.m.f
From the graph, y-intercept = 0.3686 V
E0cell = 0.3686 V
Given E0 Cu/Cu(II) = +0.34 V,
∵E0cell = E0 Fe(II)/ Fe(III) - E0 Cu/Cu(II)
E0 Fe(II)/ Fe(III) = E0cell + E0 Cu/Cu(II)
∴E0 Fe(II)/ Fe(III) = 0.3686 + 0.34
= 0.7083 V
Theoretical value = 0.77
Percentage error = (0.77-0.7083)/ 0.77 X 100% = 8.01%
The percentage error may be caused by the following sources and should be improved.
- The theoretical e.m.f could only be measured in a condition of zero current passing through the electrochemical cell, and hence the cell had zero resistance. However, it was impossible in real case. Resistance present in the electrochemical cell and therefore the cell e.m.f in real case would be smaller.
Yet, the internal resistance could be minimized by using electrodes without much impurities coated on the surface. Also, by moving closer the two electrodes, the internal resistance could also be reduced. Moreover, to have a minimum current passing through the voltmeter, a high impedance voltmeter or a potentiometer could be used.
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As iron (II) ion was a quite strong reducing agent, it may be oxidized to iron (III) ion even in the air. Therefore, if the iron (II) sulphate solution was placed in room condition for a long time without proper storage, it may be oxidized. Then, even the solutions were added according to the proportion given in the table, [Fe3+ (aq)] was larger and [Fe2+ (aq)] was smaller due to reduction in air. Thus, the ratio of [Fe3+(aq)]/ [Fe2+ (aq)] would be larger than the expected ratio. In this way, the cell e.m.f measured would be larger. To improve, only small volume of iron (II) sulphate solution was poured out into a beaker each time. So that, even it was oxidized in air, only a small amount was oxidized. Moreover, the time for the solution to be exposed to the air could be reduced as all of the solution in the beaker was added in one trial. To further prevent oxidation, the excess iron (II) sulphate solution in the beaker could be discarded each time and use the iron (II) sulphate solution which was freshly poured out.
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As the determination of cell e.m.f depended on the concentration of ions inside. Contamination may varied the concentration ratio and hence affect the accurate determination of cell e.m.f. To minimize this error, different plastic containers were used for different trials. Or the plastic containers should be rinsed with distilled water thoroughly after each trial. Otherwise, the ions from the previous trial may remained and affect the concentration ratio. After rinsing, the containers should be dried such that no water remained and diluted the solution. Then, the concentration ratio of [Fe3+(aq)]/ [Fe2+ (aq)] only depended on the specific proportion of Fe(NO3)3 and FeSO4 added.
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The volume of Fe(NO3)3 and FeSO4 added were expressed in term of the number of drop added. Yet, the size of each drop was not necessarily the same. Therefore, the actual volume of solutions added may not be the same as stated in the table. To minimize this error, the force when squeezing the dropper should be approximately constant and the dropper should be held vertically during addition of solution.
Conclusion
From the experiment, it was found out that a decrease in ionic concentration would decreased the cell emf and vice versa. Moreover, the number of electrons involved in the Fe2+ (aq)/ Fe3+ (aq) equilibrium was 1.
