Question

A student conducted an experiment to calculate the Faraday constant using a silver-copper voltaic cell. The masses of the copper and silver electrodes were 4.023g and 4.721g respectively. The cell was run for 782 seconds at a constant current of 0.250A. Calculate the value of the Faraday constant for this experiment.

Answer 1

To calculate the Faraday constant, we first need to determine the moles of electrons transferred in the cell. We can do this by finding the moles of copper and silver deposited and then using the stoichiometry of the reaction.First, let's find the moles of copper and silver deposited:Copper:Molar mass of Cu = 63.55 g/molMass of Cu deposited = 4.023 gMoles of Cu = mass / molar massMoles of Cu = 4.023 g / 63.55 g/mol = 0.0633 molSilver:Molar mass of Ag = 107.87 g/molMass of Ag deposited = 4.721 gMoles of Ag = mass / molar massMoles of Ag = 4.721 g / 107.87 g/mol = 0.0437 molNow, we need to find the moles of electrons transferred. The balanced redox reaction for the silver-copper voltaic cell is:Cu s  + 2Ag aq   Cu aq  + 2Ag s From the stoichiometry of the reaction, we can see that 1 mole of Cu reacts with 2 moles of Ag, so the moles of electrons transferred will be twice the moles of Cu.Moles of electrons = 2 * moles of Cu = 2 * 0.0633 mol = 0.1266 molNow, we can use the current and time to find the total charge transferred:Current  I  = 0.250 ATime  t  = 782 sCharge  Q  = current * time = 0.250 A * 782 s = 195.5 C  Coulombs Finally, we can calculate the Faraday constant  F  using the charge and moles of electrons:Faraday constant  F  = charge / moles of electrons = 195.5 C / 0.1266 mol = 1543.6 C/molThe value of the Faraday constant for this experiment is approximately 1543.6 C/mol.