The use of reduction potentials to predict the course of reactions is limited to thermodynamic considerations only: if the overall reduction potential for the reaction, which is the difference of the reduction potentials of the individual couples, is positive, then there is a negative Gibbs free reaction energy and the reaction is thermodynamically spontaneous.
Whether the reaction occurs, however, and the rate at which it does occur depend on the kinetic factors which influence the process. The most important kinetic factor controlling the speed and ease of redox reactions is the overpotential.
It is not easy to predict when a reaction is likely to be fast, even when it is thermodynamically favourable. One observation that is useful, however, is that redox couples with potentials that are more negative than -0.6 V cause the reduction of H+ to H2 to take place at a rapid rate. Similarly, redox couples with potentials more positive than +1.83 V cause O2 to be evolved from water (the oxidation of the O2- ion to O) at a rapid rate. The reduction potential for the O2,H+/H2O couple is +1.23 V, and so the potential of +1.83 V represents an additional potential of +0.6 V.
The extra +0.6 V above the +1.23 V potential for oxidation of O2-, or the extra -0.6 V below the 0 V potential for reduction of H+, is known as the overpotential.
The overpotential is the potential beyond the standard, zero-current, reduction potential which must exist before the reaction proceeds at a significant rate.
In neutral solution, the reduction potential for the H+/H2 couple is -0.41 V (remember that pH = 7 is not standard conditions, where E(H+,H2) = 0 V). When the pH of the solution is lowered, and the solution becomes more acidic, the reduction potential for the H+/H2 couple becomes more positive. Some metals, such as iron and zinc, have negative reduction potentials, and so might be expected to reduce water, but they do not in neutral solution, whilst they do in acidic solution. This can be explained in terms of the overpotential.
Taking zinc, where the Zn2+/Zn couple has E = -0.76, the potential for the overall reaction in neutral solution is (-0.41) – (-0.76) = +0.35 V, which is thermodynamically favourable, and yet the reaction does not proceed. When the pH is reduced, E(H+,H2) becomes more positive, and so the overall reaction potential becomes more positive, and when Eoverall > +0.6 V, the reaction proceeds. This is the value required for the overpotential to be exceeded, and explains why zinc will reduce an acid but not water itself.
In cases of outer-sphere electron transfer in complexes, there is little change in geometry of the complexes as the reaction proceeds, and it is found that the rate of electron transfer is proportional to the exponential of the difference in the standard potentials of the two individual couples. This electron transfer, which can be very fast, is therefore favoured by a large potential difference, and hence a large overpotential.
Even though the inner-sphere electron transfer has a more complicated mechanism, involving the transfer of a ligand from one complex to the other, it is also found that the greater the overpotential, or the difference in the redox couples, then the faster the rate of the electron transfer reaction.