Question

(1) Most redox reactions are performed under non-standard conditions. Consider the
permanganate to manganese (II) ion half-cell, represented by the reaction:
MnO4
-
(aq) + 8H+
(aq) + 5e- ----> Mn2+(aq) + 4H2O(l) Eθ
= +1.51V Suppose the pH is increased to 4.0, but the concentrations of MnO4
- and Mn2+ are kept at 1.0
M (standard conditions). Using the Nernst equation, determine the half-cell potential under these
new conditions.
Is the permanganate ion a stronger or weaker oxidant at these less acidic conditions?
Discuss the sensitivity of this reaction (or lack thereof) to pH. Is it large or small?

Answer 1

E= E0 -0.059/n LnQ

E= 1.51V - 0.059/5 Ln [Mn+2]/[H+]⁸[MnO4-]

pH= -log [H+] =4 ----> [H+]= 1x10^-4M

E= 1.51 -0.059/5 log (1)/(1x10^-4)⁸(1)

E= 1.132V

An ion is a better oxidant when it reduction potential is higher. In this case the reduction potential for the permanganate ion goes from 1.51V to 1.132V, so, the strength as oxidant is lower.

I think this reaction is very sensitive to pH, another way to write the Nerst equation above is:

E= E0 -0.472/n x pH

The potential of the cell is proportional to pH, so, it is very sensitive.