Pros And Cons Of Lead-Acid Electrochemical Cells Lab Report

Determining Standard Reduction Potentials, Equilibrium Constants and Investigating a Lead-Acid Electrolytic Cell Purpose Experimental Methods All procedures were followed according to the lab manual (experiment 10 – Electrochemistry Laboratory). Data and Observations Part A: Grams of FeSO4 used: 0.759 Voltage during voltmeter and battery check: 9.23V Table 1. Electrochemical Cells Data Part B: Table 2. Lead-Acid Battery Results and Calculations In order to make 5ml of 1M FeSO4(aq) (molar mass = 151.92), we calculated the needed grams according to c= n/V; 1 = n/(0.005), n = 0.005. Since n = m/M, 0.05X151.92 = 0.7596g. Part A: Cu2+ (aq) + Zn (s)  Cu (s) + Zn2+ (aq) E°cell = 1.06 V Cu2+ (aq) + 2e-
 Cu(s) E°red = 0.34 V (known standard reduction potential of copper)
0.34 V – 1.06 V E°red (Zn) = -0.72 V E° = RT/nF · lnK; n = 2 as 2 electrons were transferred from Zn to Cu 1.06 = (8.314 · 298)/(2 · 96485) · lnK K = 7.16 · 1035 Zn2+ (aq) + Mg (s)  Zn (s) + Mg2+ (aq), E°cell = 0.66
V K = 2.11 · 1022 Zn (s) + Fe2+ (aq)  Zn2+ (aq) + Fe (s) E°cell = 0.96 V K = 2.97 · 1032 Cu2+ (aq) + Mg (s)  Cu (s) + Mg2+ (aq) E°cell = 1.72 V K = 1.51 · 1058 E°red(Mg)= 0.34 – 1.72 = -1.38V K = 1.72 · 1016 E°red(Fe) = 0.34 – 0.48 = -0.14 Mg (s) + Fe2+ (aq)  Fe (s) + Mg2+ (aq) E°cell = 1.28 V K = 1.98 · 1043 Part B: Cathode: PbO2 (s) + HSO4- (aq) + 3H+ (aq) + 2e-  PbSO4 (s) + 2H2O (l) Anode: Pb(s) + HSO4- (aq)  PbSO4 (s) + H+ (aq) + 2e- Overall: PbO2 (s) + Pb(s) + 2HSO4- (aq) + 2H+ (aq)  2PbSO4 (s) + 2H2O (l) When an external power source applies energy to the cell, the reverse reaction occured, thus regenerating the reactants. 2PbSO4 (s) + 2H2O (l)  PbO2 (s) + Pb(s) + 2HSO4- (aq) + 2H+ (aq) Discussion and Conclusion In part A, the large equilibrium constants that were calculated (K>106) tell us that for each set of reactions, the equilibrium lies far to the right (product side) and that the forward reactions practically went to completion. From this experiment, we were able to determine the standard reduction potentials of other metals (magnesium, iron, and zinc) by starting out only with the value of the standard reduction potential of copper.
This is indeed similar to the case where hydrogen is taken to be the standard for comparison for all the other elements, given a E°red of 0. From our results, E°red(Cu)> E°red(Fe)> E°red(Zn)> E°red(Mg). Although this relation remains true when we look at the electrochemical series, our experimental values differed from the actual values. In magnesium, the experimental value was -1.38V (actual value: -2.37V). Hence, the percentage error was -41.8%. In iron, the experimental value was -0.14V (actual value: -0.44V), and so the percentage error was -68.1%. In zinc, the experimental value was -0.72 (actual value: -0.76), and the percentage error was the lowest: -5.2%. The difference in values could have resulted from the fact that the temperature was not exactly 298°K, thus altering the calculated reduction potential of the
element. In this lab, small voltaic cells were assembled. The anode is the site of oxidation, where the metal loses electrons and gets oxidized, and so acting like a reducing agent. This electrode carries a negative charge due to the abundance of electrons being transferred, which then travel to the cathode. The electrode loses mass as the solid metal goes into solution in the form of aqueous cations. The cathode, on the other hand, is the site of reduction, where the metal ion in solution gains electrons and gets reduced, hence acting like an oxidizing agent. Because of that, the aqueous ions increase the mass of the electrode as more solid metal forms. It would be interesting to weigh each electrode before and after each reaction occurred in order to determine how much mass was gained/lost at each electrode, whether it was a significant amount or not. Electronegativity also explains the direction and spontaneity of the reactions. More electronegative elements have a greater attraction for a shared pair of bond electrons. Hence, they force the less electronegative elements to lose electrons and form cations, while they gain electrons and get reduced. This allows the reaction to occur spontaneously. A positive E°cell value for the reaction means that the reaction occurs spontaneously in that direction; that is why when calculating E°red for each metal we used the positive E°cell value instead of the negative. Performing more trials will reduce the effect of random errors in the experiment. The average of those trials can be used to determine the reduction potentials as well as the equilibrium constants. In part B, the single cell lead-acid battery that we created in this lab was very much like the actual battery used in automobiles where a series of cells are connected to one another. One electrode (anode) is made out of gray Pb (s) and the other electrode (cathode) is made out of PbO2(s), which is brown in color. The cell uses a strong acid electrolyte. Usually a 12-V lead-acid automotive battery consists of six voltaic cells in series, each producing 2V. This type of battery is rechargeable since a current can be applied to the battery to make the non-spontaneous reaction occur and restore the original reactants and the two electrodes. The external battery enabled non-spontaneous reactions to occur. The bubbling produced at the electrodes can be explained with the equation: H2O (l)  O2 (g) + H2 (g) (non-spontaneous) The O2 gas formed at the cathode reacted spontaneously with the lead electrode to form lead oxide (thus giving it a brown coating) according to the equation: Pb (s) + O2 (g)  PbO2 (s) (brown) (spontaneous) Thus, this enabled the lead oxide to react with the lead electrode and form lead sulfate as explained previously: Pb (s) + PbO2 (s) + 2H2SO4 (aq)  2PbSO4 (s) + 2H2O (l) (spontaneous) In this experiment, the forward (voltaic cell) produced a voltage of 8.00V and the reverse reaction (electrolytic cell) produced a voltage of -8.00V as it made the reaction go in the reverse direction The spontaneous reactions are part of the voltaic cells since the reactions do not need an external power source to make the reaction proceed. The non-spontaneous reactions on the other hand are electrolytic reactions since they need an external power source (battery) to occur. In automobiles the energy necessary for recharging the battery is provided by the alternator, and in the experiment, it was the 9.00V battery. The PbSO4 sticks to the electrodes and the external power source forces the electrons from one electrode to another converting the PbSO4 to Pb at one electrode and PbO2 at the other. There is no equilibrium constant for this reaction since solids and pure liquids are not included in the expression. Hence the relative amounts (moles) of the solids involved in the reaction have no effect on the emf of the battery. However, the concentration of the acid does sometimes affect the emf discharge, which is otherwise constant, since it is consumed during the reaction. It would be interesting to perform experiments varying the concentration of sulfuric acid to see to what degree it has an effect on the voltage produced. Post Lab Questions: Mg2+ (aq) + 2e- Mg(s) E°red = -1.38 V Zn2+ (aq) + 2e-  Zn(s) E°red = -0.72 V Fe2+ (aq) + 2e-  Fe(s) E°red = -0.14 V Cu2+ (aq) + 2e-  Cu(s) E°red = 0.34 V Cu2+(aq) is the strongest oxidizing agent Mg(s) is the strongest reducing agent The lead electrodes should not come in contact with one another because the circuit can be shorted and the power supply can be damaged if that were to happen.