Redox Reactions

The word redox is a contraction of reduction and oxidation — two halves of a single event that always occur together. The classical picture talks of oxygen and hydrogen: oxidation is gain of oxygen or loss of hydrogen; reduction is loss of oxygen or gain of hydrogen. The modern and far more powerful picture is purely about electrons.

Oxidation is loss of electrons; reduction is gain of electrons. The two are inseparable — electrons lost by one species are gained by another, so a redox reaction is simply an electron transfer. The mnemonics OIL RIG (Oxidation Is Loss, Reduction Is Gain) and LEO the lion says GER (Lose Electrons = Oxidation, Gain Electrons = Reduction) capture it.

To track electrons without drawing every bond, chemists use the oxidation number (or oxidation state) — the charge an atom would carry if every bond were treated as fully ionic. A rise in oxidation number is oxidation; a fall is reduction. The rules below are applied in order of priority.

The species that causes oxidation is the oxidising agent — and to do so it must accept electrons, so it is itself reduced. The species that causes reduction is the reducing agent, which donates electrons and is itself oxidised. The roles always cross.

Most familiar reactions sort into a handful of redox families. Recognising the family often tells you what to expect of the products.

In a disproportionation, a single element in one intermediate oxidation state is simultaneously oxidised and reduced — part of it goes up, part goes down. The element must have an accessible higher and lower state for this to occur.

The reverse process, comproportionation, takes two different oxidation states of the same element and merges them into one intermediate state — for example \(\ce{2H2S + SO2 -> 3S + 2H2O}\), where sulphur at \(-2\) and \(+4\) both meet at \(0\).

The first balancing technique bookkeeps the change in oxidation number. The total increase (oxidation) must equal the total decrease (reduction), because the electrons lost equal the electrons gained.

The ion-electron method splits the reaction into an oxidation half and a reduction half, balances each completely, then recombines them. It is the cleaner route for ionic equations and the one that mirrors what happens in a cell.

Redox stoichiometry is often handled with equivalents rather than moles, because in a titration the equivalents of oxidant and reductant are always equal. The bridge is the n-factor — the number of electrons exchanged per formula unit.

A redox titration uses an oxidant–reductant pair to find an unknown concentration. Permanganometry uses \(\ce{KMnO4}\), which is self-indicating — the first faint pink persists once the reductant is used up. Iodometry/iodimetry uses the \(\ce{I2}/\ce{I-}\) couple with starch as indicator. Dichromatometry uses \(\ce{K2Cr2O7}\) with a redox indicator such as diphenylamine.

Why does zinc reduce \(\ce{Cu^2+}\) but not the reverse? Each half-reaction has a tendency to gain electrons measured by its standard electrode potential \(E^\circ\), referenced to the standard hydrogen electrode at \(0\ \text{V}\). Arranged by \(E^\circ\), these form the electrochemical series.

Assign oxidation numbers first, then decide which species moves up and which moves down. For titration questions, fix every n-factor from the medium before calculating. Difficulty rises down the list.