Oxidation states (rules for assigning them), oxidation as loss of electrons and reduction as gain, oxidising and reducing agents, and writing and balancing half-equations and overall redox equations.

What this dot point is asking

AQA wants you to assign oxidation states using the standard rules, define oxidation and reduction in terms of electron transfer, identify oxidising and reducing agents, and write and combine half-equations into balanced overall redox equations.

Oxidation states

The oxidation state (oxidation number) is the charge an atom would have if every bond were fully ionic. It is a bookkeeping device for tracking electron transfer, and roman numerals in names such as iron(III) or manganate(VII) state it directly. Assigning oxidation states lets you spot redox, identify what is oxidised and reduced, and balance equations that would be hard to balance by inspection.

Oxidation and reduction

• An oxidising agent accepts electrons from another species and is itself reduced (e.g. $\text{MnO}_4^-$, which is a powerful oxidising agent in acidic solution).
• A reducing agent donates electrons to another species and is itself oxidised (e.g. $\text{Fe}^{2+}$, which is oxidised to $\text{Fe}^{3+}$).

A change in oxidation state is the quickest way to spot redox: if any element's oxidation state changes during a reaction, it is a redox reaction. The element whose oxidation state rises has been oxidised; the element whose oxidation state falls has been reduced. Disproportionation is a special case where the same element is both oxidised and reduced in one reaction, for example chlorine in $\text{Cl}_2 + 2\text{NaOH} \rightarrow \text{NaCl} + \text{NaClO} + \text{H}_2\text{O}$, where chlorine goes from 0 to both $-1$ and $+1$.

Half-equations and overall redox equations

A half-equation shows the electrons transferred for one species, written with the electrons on the side that balances the charge. To balance a half-equation in acidic solution: balance the main atoms, then balance oxygen by adding $\text{H}_2\text{O}$, then balance hydrogen by adding $\text{H}^+$, and finally balance the charge by adding electrons. Combine an oxidation half-equation with a reduction half-equation so the electrons cancel (scale each up to the lowest common multiple of electrons), then add them to give the overall equation. This is the basis of redox titrations, where the volume of (for example) manganate(VII) needed to react with a known reducing agent lets you calculate an unknown concentration.

Try this

Q1. Give the oxidation state of sulfur in $SO_4^{2-}$. [1 mark]

• Cue. $x + 4(-2) = -2$, so $x = +6$.

Q2. State whether a reducing agent is oxidised or reduced. [1 mark]

• Cue. Oxidised (it loses electrons to the other species).