Redox Reactions Revisited: The Basis of Electrochemistry

Electrochemistry is the branch of chemistry concerned with the interconversion of chemical energy and electrical energy, and it rests entirely on the redox (reduction-oxidation) reactions already familiar from earlier study: reactions in which one species loses electrons (is oxidized) while another gains those same electrons (is reduced).

In an ordinary redox reaction carried out in a single beaker — for example, dropping a strip of zinc metal directly into a copper sulfate solution — the electron transfer from zinc atoms to copper ions happens by direct contact, at the surface of the zinc strip, and the chemical energy released is dissipated entirely as heat. The reaction $\text{Zn}(s) + \text{Cu}^{2+}(aq) \to \text{Zn}^{2+}(aq) + \text{Cu}(s)$ still occurs — zinc is oxidized, losing two electrons per atom, and copper is reduced, gaining those same two electrons — but because both half-processes happen at the same physical location, none of the electrical energy in this electron transfer can be captured or put to use.

The central idea of electrochemistry is to physically separate the two half-reactions — the oxidation and the reduction — into two different compartments, each containing its own electrode, so that the electrons can only pass from the reducing agent to the oxidizing agent by travelling through an external wire connecting the two compartments. Forcing the electron transfer through an external circuit in this way is what allows the energy of the reaction to be captured as usable electrical work, measured as a voltage (EMF) and a current, rather than being wasted entirely as heat.

This physical separation, and the devices built around it, are the subject of the rest of this chapter: how a spontaneous redox reaction, split this way, becomes a galvanic cell that generates electricity; how the reverse idea — supplying external electrical energy to force a non-spontaneous redox reaction — gives an electrolytic cell; and how measuring the electrical properties of ionic solutions themselves (their conductance) gives further insight into how well those solutions carry current.