Oxidation does not necessarily require the outright removal of electrons, as we have said. Oxidation-reduction reactions in which electrons are actually moved from one substance to another are especially useful, since if the donor and recipient can be isolated, and the electrons made to flow through an external wire or circuit, some of the free energy of the oxidation-reduction process can be harnessed to do useful work. As an example, zinc metal has less of an affinity for its 4 outer electrons than metallic copper does. In a competition between Cu²⁺ and Zn²⁺ ions for electrons, copper ions will win. The reaction

is highly spontaneous, with a standard free energy change of -50.7 kcal per mol. If we dip a zinc strip into a copper sulfate solution, as shown opposite, the zinc will be eaten away, a spongy layer of metallic copper will plate out on the zinc strip, and the deep blue color of copper sulfate will gradually fade. (Zinc sulfate, which is formed, is colorless.) In contrast, if we immerse a copper strip in a zinc sulfate solution, no reaction will occur because the reverse reaction is highly nonspontaneous, with a +50.7 kcal per mol free energy barrier to surmount.

This spontaneous transfer of electrons from zinc to copper is not useful because the free energy released is dissipated as heat. It is analogous to burning a spoonful of sugar with a match instead of eating it and converting the free energy of oxidation into useful muscle work. If some means could be found to separate the removal of electrons from zinc (oxidation) from the donation of electrons to copper ions (reduction), then the electrons might be made to do something useful along the way.