Consider the following oxidative half reaction:

Question

Mn_(s) --> Mn²⁺ + 2e^-

A) Couple the oxidation reaction with a reduction half reaction that makes the overall redox reaction non-spontaneous (electrolytic cell). Write the balanced redox reaction and do not worry about adding spectator ions but pick a suitable reduction reaction which does not require water, H+, or OH-.

B) Calculate the standard potential for the electrochemical cell using half cell potentials. Show the equation and show your work.

C) How much voltage must the current supplied to the electrolytic cell have in order to drive the reaction expressed in part A and B?

Answer 1

A). Since Mn(s) => Mn2+ is the oxidation reaction, it will be the anode. To select a cathode such that the reaction will be non spontaneous, we want a metal with a lower standard reduction potential than Mn. A reasonable choice would be Mg. From a table of reduction potentials, I found the following:

Mn²⁺ + 2e- ==> Mn(s) Eº = -1.18

Mg²⁺ + 2e- ==> Mg(s) Eº = -2.38

Eºcell = -2.38 - (-1.18) = -1.2 V (non spontaneous since the Eº is negative)

Balanced redox equation:

Mn(s) + Mg²⁺ ==> Mg(s) + Mn²⁺

B) reduction potential @ cathode - reduction potential @ anode = -2.38 - (-1.18) = -1.2 V

C) >1.2 V