Oxidation–Reduction (Redox) Reactions

Learning Objective 4.9.A Represent a balanced redox reaction equation using half-reactions.

Quick Notes

Redox reactions involve the transfer of electrons between species.

Oxidation = loss of electrons
Reduction = gain of electrons

Half-reactions represent each process separately.
Electrons lost in oxidation must equal electrons gained in reduction.
Half-reactions can be combined to form a balanced overall redox equation.

Full Notes

What is a Redox Reaction?

A redox (reduction–oxidation) reaction involves the transfer of electrons between two chemical species.

Oxidation: a species loses electrons
Reduction: a species gains electrons

The species that gets oxidized donates electrons, and the one that gets reduced accepts electrons.

You can remember:
OIL RIG = Oxidation Is Loss, Reduction Is Gain

Example Reaction between Zn(s) + Cu²⁺(aq)

AP Chemistry Electron transfer between Zn and Cu2+ showing oxidation and reduction.

Zinc loses two electrons → becomes Zn²⁺ → oxidation
Copper ion gains two electrons → becomes Cu → reduction
Electrons move from Zn to Cu²⁺

Half-Reactions

Redox reactions can be split into two parts called half-reactions:

The oxidation half-reaction shows the species losing electrons.
The reduction half-reaction shows the species gaining electrons.

Each half-reaction must be balanced for atoms and charge.

Example Reaction between Zn(s) + Cu²⁺(aq)

AP Chemistry Displacement reaction between zinc and copper showing electron transfer.

Step 1 – Write oxidation and reduction half-reactions:
Oxidation: Zn (s) → Zn²⁺ (aq) + 2e⁻
Reduction: Cu²⁺ (aq) + 2e⁻ → Cu (s)

Step 2 – Combine the half-reactions:
Make sure the number of electrons lost = electrons gained.
Zn (s) + Cu²⁺ (aq) → Zn²⁺ (aq) + Cu (s) ✔

This is now the balanced redox equation.

When to Use H⁺ and H₂O in Half-Equations

For more complex redox reactions, especially in aqueous solution, additional steps may be needed to help balance O and H in reactants (and products):

Balance atoms other than H and O.
Balance O atoms by adding H₂O.
Balance H atoms by adding H⁺ (acidic) or OH⁻ (basic).
Balance charges by adding electrons.
Multiply half-reactions to equalize electrons, then add.

This method ensures mass and charge are both conserved.

Worked Example: Acidic Reduction of Manganate(VII)

Balance the half-reaction in acidic solution: MnO₄⁻ → Mn²⁺

Balance atoms other than H and O: Mn is already 1 on each side.
Balance O with H₂O: MnO₄⁻ has 4 O, so add 4 H₂O to the products.
MnO₄⁻ → Mn²⁺ + 4H₂O
Balance H with H⁺: 4 H₂O contains 8 H, so add 8 H⁺ to the reactants.
8H⁺ + MnO₄⁻ → Mn²⁺ + 4H₂O
Balance charge with e⁻: Left: +8 − 1 = +7; right: +2. Add 5 e⁻ to the left to reduce the charge to +2.
8H⁺ + MnO₄⁻ + 5e⁻ → Mn²⁺ + 4H₂O

Balanced half-reaction (acidic): 8H⁺ + MnO₄⁻ + 5e⁻ → Mn²⁺ + 4H₂O

Matt’s exam tip

If a redox reaction seems complicated, start with half-reactions. They break the problem into manageable parts and make it much easier to follow the flow of electrons. Don't forget to check that electrons cancel when the half-reactions are combined.

Summary

Redox reactions involve the transfer of electrons, with one species being oxidized and another reduced.
These reactions can be systematically balanced by separating them into oxidation and reduction half-reactions.
Balance each for mass and charge, then combine them to form the complete balanced equation.