Classical Concept of Oxidation and Reduction

The oldest working definition of these reactions comes directly from the two elements chemists met most often in the laboratory: oxygen and hydrogen.

Oxidation, classically, is the addition of oxygen to a substance, or the removal of hydrogen from it. When magnesium ribbon burns in air, $2\text{Mg}(s) + \text{O}_2(g) \to 2\text{MgO}(s)$, magnesium gains oxygen and is said to be oxidized. Similarly, when hydrogen sulphide is treated with chlorine water, $\text{H}_2\text{S} + \text{Cl}_2 \to 2\text{HCl} + \text{S}$, the sulphide loses hydrogen and is oxidized to elemental sulphur.

Reduction is exactly the reverse: the removal of oxygen from a substance, or the addition of hydrogen to it. In $\text{CuO} + \text{H}_2 \to \text{Cu} + \text{H}_2\text{O}$, copper(II) oxide loses its oxygen and is reduced to metallic copper, while hydrogen itself is oxidized (it gains oxygen, becoming water).

The key observation, present in every one of these examples, is that oxidation and reduction always occur together, never alone — one reactant cannot gain oxygen (or lose hydrogen) unless another loses oxygen (or gains hydrogen) at exactly the same time. A reaction in which this paired change happens is called a redox reaction (reduction + oxidation, compressed into one word), and the process as a whole is redox.

For the combination reaction $2\text{Mg} + \text{O}_2 \to 2\text{MgO}$: magnesium starts as the free element and ends up combined with oxygen — by the classical rule this is oxidation, so magnesium is the substance oxidized. The oxygen molecule, in turn, is the species that is consumed by combining with another element; every atom that receives oxygen forces the corresponding reduction to be credited to the oxygen supplied, so $\text{O}_2$ is described as undergoing reduction in this combination even though it was not itself already combined with anything — because the reaction as a whole is classified by tracking which reactant delivers oxygen (the oxidizing agent, here $\text{O}_2$) and which one is oxidized by accepting it (here $\text{Mg}$).

This classical, element-counting definition is intuitive and matches how redox chemistry was first discovered historically, but it has a serious limitation: it says nothing about reactions that involve neither oxygen nor hydrogen at all — for instance, $\text{Zn}(s) + \text{CuSO}_4(aq) \to \text{ZnSO}_4(aq) + \text{Cu}(s)$, where zinc metal displaces copper from solution. No oxygen or hydrogen changes hands anywhere in this equation, yet chemically it is unmistakably the same kind of process as the magnesium-burning reaction — one species is losing something to another. What is actually being transferred, in every one of these cases without exception, is the electron — and recognizing this is what leads to the modern, fully general definition developed in the next section.