How to balance redox and half equations

Redox questions reward a clear method more than almost any other topic in redox equilibrium. Once you can write a half equation and combine two of them so the electrons cancel, a whole family of exam questions becomes routine. This guide teaches that method step by step, with worked examples you can copy.

The two words behind every redox reaction

Remember OIL RIG: Oxidation Is Loss of electrons, Reduction Is Gain of electrons. In every redox reaction one species is oxidised (loses electrons) and another is reduced (gains them) at the same time, the electrons lost by one are exactly the electrons gained by the other. That balance of electrons is the heart of the whole method.

A second way to spot redox is the oxidation number: oxidation is an increase in oxidation number, reduction is a decrease. If you are rusty on assigning oxidation numbers, work through the oxidation number method first, because half equations lean on it.

What a half equation is

A half equation shows just the oxidation or just the reduction, with electrons written in. For example, when zinc is oxidised:

Zn → Zn²⁺ + 2e⁻

Zinc loses two electrons, so two electrons appear on the right (the product side) of an oxidation half equation. When copper(II) ions are reduced:

Cu²⁺ + 2e⁻ → Cu

The electrons appear on the left (they are gained) of a reduction half equation. A quick rule: electrons on the right means oxidation; electrons on the left means reduction.

The step-by-step method

1. Write the oxidation half equation. Put the reduced form on the left, the oxidised form on the right.
2. Balance the atoms of the element changing.
3. Balance the charge by adding electrons to the more positive side.
4. Repeat for the reduction half equation.
5. Multiply one or both half equations so the number of electrons is the same in each.
6. Add the two half equations and cancel the electrons, plus anything identical on both sides.

Worked example 1: zinc and copper(II)

Balance the reaction between zinc metal and copper(II) ions.

• Oxidation: Zn → Zn²⁺ + 2e⁻
• Reduction: Cu²⁺ + 2e⁻ → Cu
• Both have 2 electrons, so add directly and cancel the electrons: Zn + Cu²⁺ → Zn²⁺ + Cu

Here zinc is the reducing agent (it gives electrons away) and Cu²⁺ is the oxidising agent. You will meet this exact pattern in the displacement of metals.

Worked example 2: halogen displacement

Balance chlorine reacting with bromide ions.

• Reduction: Cl₂ + 2e⁻ → 2Cl⁻
• Oxidation: 2Br⁻ → Br₂ + 2e⁻
• Electrons already match at 2 each, so add and cancel: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂

Chlorine is more electronegative than bromine, so it takes the electrons, chlorine is reduced and bromide is oxidised.

Worked example 3: scaling the electrons

This is the step that trips people up. Balance iron(II) ions being oxidised to iron(III) by chlorine.

• Oxidation: Fe²⁺ → Fe³⁺ + e⁻ (one electron)
• Reduction: Cl₂ + 2e⁻ → 2Cl⁻ (two electrons)

The electrons do not match, one versus two. Multiply the oxidation half equation by 2 so both involve two electrons:

• 2Fe²⁺ → 2Fe³⁺ + 2e⁻

Now add to the reduction half and cancel the two electrons:

2Fe²⁺ + Cl₂ → 2Fe³⁺ + 2Cl⁻

Check the charge: left side is 2 × (+2) + 0 = +4; right side is 2 × (+3) + 2 × (−1) = +6 − 2 = +4. The charges balance, so the equation is correct.

Half equations at electrodes

The same skill writes the reactions at electrodes during electrolysis, which is heavily examined. At the cathode (reduction), positive ions gain electrons:

2H⁺ + 2e⁻ → H₂ or Cu²⁺ + 2e⁻ → Cu

At the anode (oxidation), negative ions lose electrons:

2Cl⁻ → Cl₂ + 2e⁻ or 4OH⁻ → O₂ + 2H₂O + 4e⁻

Learning these standard half equations means you can write electrode reactions straight down in the exam. Which ion is discharged depends on the electrochemical series and the conditions.

Checks that catch mistakes

• Atoms balance on both sides of each half equation.
• Charge balances, the total charge is the same on the left and right.
• Electrons cancel completely when the halves are added; if any remain, you forgot to scale.
• Oxidation numbers move sensibly, up for the species oxidised, down for the one reduced.

Practise until it is automatic

Redox is a method, not a memory test, so drill it. Write out the two half equations, scale the electrons, add and cancel, the same four moves every time. Start with displacement and electrode reactions, then build to oxidising and reducing agents. If scaling the electrons or assigning oxidation numbers keeps tripping you up, that is a quick fix with a teacher, our online one-to-one lessons run in English from RM50 an hour, with a paid one-hour trial; see how it works if you would like redox drilled against real Paper 2 questions.