Assign oxidation numbers, identify oxidation and reduction, and balance redox half-equations and overall equations

What this dot point is asking

A redox reaction involves the transfer of electrons. To track that transfer we use oxidation numbers (oxidation states), a set of rules that assign a charge-like number to each atom.

Rules for assigning oxidation numbers

Apply these in order:

1. An element in its standard state has oxidation number 0 (for example $\text{O}_2$, $\text{Na}$, $\text{Cl}_2$).
2. A monatomic ion has an oxidation number equal to its charge ($\text{Na}^+$ is $+1$, $\text{S}^{2-}$ is $-2$).
3. Oxygen is usually $-2$ (except in peroxides, where it is $-1$).
4. Hydrogen is usually $+1$ (except in metal hydrides, where it is $-1$).
5. The sum of oxidation numbers equals the overall charge of the species (0 for a neutral molecule).

Writing half-equations

A half-equation shows just the oxidation or just the reduction, including the electrons. For example, the oxidation of iron(II) to iron(III):

and the reduction of chlorine:

Balancing half-equations in acidic solution

For more complex half-equations involving oxygen, use this sequence:

1. Balance the atom being oxidised or reduced.
2. Balance oxygen by adding $\text{H}_2\text{O}$.
3. Balance hydrogen by adding $\text{H}^+$.
4. Balance the charge by adding electrons.

Combining half-equations

To get the overall equation, multiply each half-equation so the number of electrons lost equals the number gained, then add them and cancel the electrons. The electrons must cancel exactly; if they do not, the multiplication is wrong.

Why this matters

Half-equations are the language of all of electrochemistry. You need them to construct galvanic and electrolytic cell reactions, to use the standard electrode potential table, and to perform redox titration and electrolysis calculations later in Unit 3.