Yurani Muratalla Experiment 11 Oxidation Reduction Electrochemistry In this experiment the relationship between reduction potentials and reaction spontaneity was explored through the measuring of relative metal reduction potentials In addition Faraday s constant and Avogadro s number were calculated by the method of electrolysis and two additional reactions were observed Standard Reduction Potentials at 25 C E volts referenced to Cu E volts referenced to Ag Half reactions Measured Actual Measured Actual Ag e Ag s 0 443 0 460 0 0 0 0 2 Cu 2e Cu s 0 0 0 0 0 437 0 460 2H 2e H2 s Not measured Not measured Not measured Not measured 2 Fe 2e Fe s 0 579 0 750 1 012 1 210 2 Zn 2e Zn s 0 457 1 100 1 255 1 560 Compared to the actual reduction potentials the measured values in part one the electro chemical series of Metals were always smaller in magnitude The measurements were acceptable as they were close to the actual reductions potentials Overall the only measurement that strayed a little from the actual value was the reduction potential for zinc referenced to the copper electrode This could have been due to either a defect in the equipment such as not being able to handle a great current causing it to report a lower number or due to the solutions of the respective metals mixing together and thus complicating the reading of the voltmeter Part 2 2Al s 3Cu2 2Al3 3Cu s E 2 04 V Part 3 Cu2 2Ag Cu2 s 2Ag s E 0 46 V In part two and three two reactions were observed For both the value of E was positive making them spontaneous reactions This spontaneity prediction was provided true in part 2 when solid copper formed on the aluminum foil as the aluminum dissolved and for part 3 solid silver formed on the copper wire as the copper dissolved It was also evident that copper dissolved since the once clear solution turned into a clear blue color Part 4 Anode Cu s Cu2 aq 2e aq Cathode Overall 2H aq 2e aq H2 g Cu s 2H aq Cu2 aq H2 g In part 4 there were two methods to calculate F Faraday s Constant and NA Avogadro s number Through the method of collecting H2 g F was calculated to be 96485C mol e and NA was calculated to be 6 02 x 1023 molecules mol Through the method of using the moles of copper reacted however F was calculated to be 90353 C mol e and NA was calculated to be 5 64 x 1023 molecules mol Being that the actual values for F and NA are 96485 C mol e and 6 03 x 1023 molecules mol respectively It is noticeable that the results from method one were much closer to the actual values than the results from method two Throughout the course of part four there were a few possible sources of error If any extra volume leaked from the burette used for electrolysis the total volume measured for H2 g produced would be much less than actuality decreasing the moles of H2 calculated and thus a lesser value for F and NA also for method one calculations In addition if the copper anode wasn t properly dried when it was weighed the value for copper consumer would be less making both values for F and NA greater for method two When using the value for volume of H2 g produced the percent errors for F and NA were 0 and 0 respectively Basically the values measured were pretty close to the actual accepted ones and if there was any error it was really minute However when using the value for moles of copper consumed there were percent errors of 6 34 and 6 31 This shows that measuring the volume of H2 g produced is a more reliable method