In a redox reaction, one species loses electrons and another gains them at the same time. You identify which is which by comparing oxidation numbers before and after the reaction.
This lesson is part of SPM Chemistry redox equilibrium. It uses the skill from assigning oxidation numbers.
What four questions label an equation?
Ask these in order for each redox equation.
- Which atom’s oxidation number increases? That species is oxidised.
- Which atom’s oxidation number decreases? That species is reduced.
- The substance that is oxidised is the reducing agent.
- The substance that is reduced is the oxidising agent.
Worked example 1: iron and copper(II) sulfate
Fe + CuSO₄ → FeSO₄ + Cu
Iron goes from 0 to +2, so iron is oxidised. Copper goes from +2 to 0, so copper(II) ions are reduced.
The reducing agent is iron, because it is oxidised. The oxidising agent is the copper(II) ion in CuSO₄, because it is reduced.
Worked example 2: chlorine and potassium bromide
Cl₂ + 2KBr → 2KCl + Br₂
Chlorine goes from 0 to −1, so it is reduced and is the oxidising agent. Bromide goes from −1 to 0, so it is oxidised and is the reducing agent.
Is every reaction redox?
Compare with AgNO₃ + NaCl → AgCl + NaNO₃. Silver stays +1, chlorine stays −1, sodium stays +1 and nitrogen stays +5. Nothing changes, so this is a precipitation and not redox.
The mistake that costs marks
The common slip is to call the substance that is oxidised “the oxidising agent”. In worked example 1, iron is oxidised but it is the reducing agent.
| Question | Wrong answer | Right answer |
|---|---|---|
| Which is oxidised? | Iron | Iron |
| Oxidising agent? | Iron | Copper(II) ions |
| Reducing agent? | Copper(II) ions | Iron |
A useful habit: write the two agent labels only after you have written what is oxidised and reduced, and always flip them.
Check yourself
In 2Fe³⁺ + 2I⁻ → 2Fe²⁺ + I₂, state what is oxidised, what is reduced, and identify the oxidising agent and reducing agent.
Answer
Iron goes from +3 to +2, which is a decrease, so Fe³⁺ is reduced and is the oxidising agent.
Iodine goes from −1 to 0, which is an increase, so I⁻ is oxidised and is the reducing agent.
Check the charge: left is 2(+3) + 2(−1) = +4, and right is 2(+2) + 0 = +4. The equation balances.
What to study next
Continue with writing half-equations, which shows the electron transfer for each change above. Then test yourself with the redox practice set.
The mistake log and paper-error review tool helps you keep track of which labels you swap. If you want a teacher to check your labels on your own equations, see online one-to-one Chemistry tuition.