Assign oxidation numbers, write and balance half-equations, and use the standard electrode potential series to predict and calculate cell potentials for galvanic and electrolytic cells

What this dot point is asking

Redox reactions involve the transfer of electrons. Oxidation is the loss of electrons; reduction is the gain of electrons (remember OIL RIG: Oxidation Is Loss, Reduction Is Gain). Every redox reaction couples one of each: the species oxidised is the reducing agent, and the species reduced is the oxidising agent.

Oxidation numbers

Oxidation numbers track electron distribution and reveal what is oxidised and reduced. The key rules are: elements in their standard state are 0; oxygen is usually $-2$ (peroxides $-1$); hydrogen is $+1$ with non-metals and $-1$ in metal hydrides; a monatomic ion equals its charge; and the sum of oxidation numbers equals the overall charge of the species. An increase in oxidation number is oxidation; a decrease is reduction.

Writing half-equations

A redox reaction splits into an oxidation half-equation and a reduction half-equation. To balance one (in acidic solution): balance the atoms being oxidised or reduced, balance oxygen with $\text{H}_2\text{O}$, balance hydrogen with $\text{H}^+$, then balance charge with electrons. For example the reduction of dichromate:

To combine two half-equations, multiply each so the electrons cancel, then add them.

Standard electrode potentials

The standard electrode potential ($E^{\circ}$) of a half-reaction measures its tendency to be reduced, relative to the standard hydrogen electrode ($E^{\circ} = 0\ \text{V}$), measured at $25^{\circ}\text{C}$, $100\ \text{kPa}$ and $1\ \text{mol L}^{-1}$ concentrations. The SCSA data booklet lists these as reduction half-equations. A more positive $E^{\circ}$ means a stronger oxidising agent (more readily reduced); a more negative $E^{\circ}$ means a stronger reducing agent (more readily oxidised). The species with the higher (more positive) $E^{\circ}$ is reduced; the other is oxidised.

Galvanic (voltaic) cells

A galvanic cell converts the energy of a spontaneous redox reaction into electrical energy. Two half-cells are connected by an external wire (electron path) and a salt bridge (ion path that keeps each half-cell neutral). Oxidation occurs at the anode (negative electrode in a galvanic cell) and reduction at the cathode (positive electrode); electrons flow through the wire from anode to cathode. The cell potential is

A positive $E^{\circ}_{cell}$ confirms the reaction is spontaneous.

Electrolytic cells

An electrolytic cell uses an external power supply to drive a non-spontaneous redox reaction (for example electrolysis of molten salts or aqueous solutions, electroplating, and electrorefining). Here the anode is positive and the cathode is negative, the reverse polarity of a galvanic cell, but oxidation still occurs at the anode and reduction at the cathode.

When answering redox questions in the WACE examination, read half-equations from the data booklet exactly as written (reductions), identify cathode and anode by comparing $E^{\circ}$ values, and state the electron and ion flow directions explicitly.