Calculate the minimum concentration of Ag⁺ that would remain unreduced by a standard Fe²⁺/Fe³⁺ electrodes at equilibrium?

Given Fe³⁺ + e^- → Fe²⁺ + 0.671 v Ag⁺ + e^- → Ag + 0.699 v

The cell reaction is Fe²⁺ (aq) + Ag⁺ (aq) → Fe³⁺ (aq) + Ag (s)E^o cell = E^o Fe²⁺ → Fe³⁺ + E Ag⁺ → Ag (s)E^o cell = (- 671) + (0.699)E^o cell = + 0.028 vE^o cell = 0.059/n log (Fe³⁺ x Ag)/(Fe²⁺ x [Ag⁺]) Now E^o cell = 0.028 vn = 1, [Fe³⁺] = [ Fe²⁺] = 1 M, [Ag⁺] = ? 0.028 = 0.059/1 log (1 x 1)/(1 x [Ag + ]) 0.028/0.059 = log 1/([Ag⁺])log 1/([Ag⁺]) = 0.4745log 1/([Ag⁺]) = 10^0.4745 = 2.98 [Ag+] = 1/2.98 = 0.335 M.