KarinaM_Exp6

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San Diego State University *

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200

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Chemistry

Date

Apr 3, 2024

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pdf

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Uploaded by BarristerMetalRaccoon37

Martina Karina Lab partner: Emma Smith 11/19/22 Chem 201: Experiment 6 Electrochemical Cells Procedures: The procedures listed in my lab notebook were followed, with the exception of diving the class into groups of two to complete the experiment. The procedures were followed with an individual set of probes and concentrated solutions, shared between my lab partner and I. Other than this, no changes were made. The solutions were of ZnSO 4 (0.1 and 0.01 M) and CuSO 4 (0.1, 0.01, and 0.001 M). The salt bridge consisted of a 0.50 M Na 2 SO 4 solution. Experimental & Calculated: Trial Red Lead: Solution (M) Black Lead: Solution (M) Measured Cell Potential (V) Calculated Cell Potential (V) Percentage Error A 0.10 CuSO 4 0.10 ZnSO 4 1.102 1.100 0.1818% B 0.0010 CuSO 4 0.10 ZnSO 4 1.049 1.041 0.7685% C 0.10 CuSO 4 0.0010 ZnSO 4 1.093 1.159 5.695% D 0.010 CuSO 4 0.10 CuSO 4 -0.035 -0.030 17% E 0.0010 CuSO 4 0.10 CuSO 4 -0.054 -0.059 8.5% Discussion: The calculated percent errors between the experimental and theoretical cell potentials provides the best insight into the accuracy of our experiment. Experiments A, B, C, and E all had percent errors of less than 10%, with A and B even having percent errors of less than 1%. Trial D had a much higher percent error of 17%. It is likely that the salt bridge may have been touching the electrodes, which would result in a faulty voltmeter reading. An error like this is more probable than an issue with concentration or the dilution process because all other trials produced voltages that were much more accurate to the theoretical values. Overall though, it would have been impossible to obtain the exact voltage values. The theoretical cell potential values are calculated at standard conditions, where temperature is 25 degrees Celsius. The classroom

temperature was less than that (around 22 degrees or so), meaning our values could never be the exact same as they were theoretically calculated in those conditions. When comparing the Daniell cell experiments A and B, it is clear that a decrease in copper solution concentration leads to a decrease in cell potential. But in experiment C, when zinc concentration is decreased, the cell potential goes up. In the Nernst equation for the copper and zinc half reactions, plugging in a lower zinc concentration leads to a higher value. This higher value half reaction potential indicates a higher willingness to give electrons, as zinc is the ion being oxidized. Zinc being more oxidized means a greater flow of electrons will occur between the voltage wires, which would provide a higher voltage value. The opposite is true for copper, explaining the decrease in trial B versus the increase in trial C. In experiments D and E, we measured the effect of concentration differences between only copper cells. D had a higher cell potential than E. Because of this, it can be concluded that when concentration is higher, the cell potential is also higher. Concentration and cell potential have a direct positive relationship.

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