Writing Half Equations for Electrolysis Revision Notes | GCSE Edexcel Higher Chemistry

1. What Are Half Equations and Why Do We Need Them?

During electrolysis, ions move to electrodes and react. But what exactly happens when they get there? Half equations show the precise chemical changes at each electrode — who gains electrons, who loses them, and how many.

Think of electrolysis like a busy train station. Ions are passengers heading to different platforms (electrodes). Half equations tell us what each passenger does when they arrive at their platform.

2. The Golden Rules - Anode and Cathode Reactions

Every half equation follows two simple rules. Master these and you'll never get confused about which electrode does what.

3. Writing Half Equations Step by Step

Writing half equations looks intimidating, but there's a systematic method. Follow these steps and you'll get them right every time.

Worked example: Writing a Half Equation for Copper Ion Reduction

Write the half equation for copper ions being reduced at the cathode during electrolysis of copper sulphate solution.

1. Cu²⁺ ions go to the cathode (negative electrode) — Positive ions are attracted to the negative cathode
2. At cathode = reduction = gain electrons — Cathode always means reduction (gaining electrons)
3. Cu²⁺ → Cu — Ion becomes neutral atom by gaining electrons
4. Cu²⁺ + 2e⁻ → Cu — Need 2 electrons to cancel the 2+ charge

Answer: Cu²⁺ + 2e⁻ → Cu

Notice how the number of electrons equals the charge on the ion. A 2+ ion needs 2 electrons to become neutral. A 3+ ion would need 3 electrons.

4. Oxidation Half Equations at the Anode

At the anode, the opposite happens. Negative ions lose electrons to become neutral atoms or molecules. This is oxidation — remember, oxidation means losing electrons.

Worked example: Writing a Half Equation for Chloride Ion Oxidation

Write the half equation for chloride ions being oxidised at the anode during electrolysis of sodium chloride solution.

1. Cl⁻ ions go to the anode (positive electrode) — Negative ions are attracted to the positive anode
2. At anode = oxidation = lose electrons — Anode always means oxidation (losing electrons)
3. Cl⁻ → Cl — But chlorine exists as Cl₂ molecules, not single atoms
4. 2Cl⁻ → Cl₂ — Need 2 chloride ions to make 1 chlorine molecule
5. 2Cl⁻ → Cl₂ + 2e⁻ — 2 ions each lose 1 electron = 2 electrons total

Answer: 2Cl⁻ → Cl₂ + 2e⁻

Key insight: Some elements form diatomic molecules (H₂, O₂, Cl₂, Br₂, I₂). You need to account for this when writing half equations.

5. Balancing Half Equations - The Electron Count

The trickiest part is getting the electron count right. The number of electrons must balance the charges perfectly. Here's the foolproof method.

Worked example: Balancing a Complex Half Equation

Write the half equation for oxide ions (O²⁻) being oxidised to oxygen gas at the anode.

1. O²⁻ → O₂ — Oxide ions become oxygen molecules
2. 2O²⁻ → O₂ — Need 2 oxide ions to make 1 oxygen molecule
3. Left side: 2 × (2⁻) = 4⁻ Right side: 0 (neutral molecule) — Total charge difference = 4
4. 2O²⁻ → O₂ + 4e⁻ — Need to remove 4 negative charges, so add 4e⁻ on right
5. Left: 4⁻ Right: 4⁻ (from electrons) Balanced! — Charges now balance on both sides

Answer: 2O²⁻ → O₂ + 4e⁻

The key is charge balance. Count all the charges on the left, count all the charges on the right, then add electrons to make them equal.

6. Common Exam Patterns and Mistakes

Examiners have favourite tricks when testing half equations. Spot these patterns and you'll avoid the most common mistakes that cost students marks.

Pattern 1: They give you a partially completed equation and ask you to finish it. Always check the electron count matches the ion charges.

Pattern 2: They ask for equations at both electrodes. Remember — electrons are always on opposite sides. At cathode: ions + electrons → atoms. At anode: ions → atoms + electrons.

The biggest trap: Forgetting that some elements are diatomic. Hydrogen, oxygen, nitrogen, and the halogens (F₂, Cl₂, Br₂, I₂) always exist as molecules, never single atoms.

7. Exam Technique - Getting Full Marks

Half equation questions are worth 2-3 marks each. Here's how to secure every mark by following the examiner's marking scheme exactly.

Mark scheme pattern: [1 mark] for correct species on both sides, [1 mark] for correct electron count, [1 mark] for proper balancing. Miss any part and you lose marks.

Worked example: Exam-Style Half Equation Question

Complete this half equation for the reaction at the anode: 2I⁻ → ___ + ___

1. Anode = oxidation = losing electrons — Iodide ions will lose electrons to form iodine
2. I⁻ ions become I₂ molecules — Iodine is diatomic, so 2I⁻ → I₂
3. 2I⁻ have total charge of 2⁻ I₂ has charge of 0 Difference = 2 electrons lost — Each I⁻ loses 1 electron, so 2I⁻ lose 2e⁻ total
4. 2I⁻ → I₂ + 2e⁻ — [1 mark] for I₂, [1 mark] for 2e⁻

Answer: 2I⁻ → I₂ + 2e⁻

Examiner's secret: They often give you the hardest part (like the number of ions) and ask you to fill in the rest. Use the given information as your starting point.

Exam precision note: In Writing Half Equations for Electrolysis - The process of electrolysis - Electrolysis, always identify the key chemical idea first, then link it to evidence, data, or an equation. A high-scoring GCSE answer names the particles involved, states what changes, and justifies the conclusion using correct scientific language. Where calculations are used, show each step with units and keep significant figures sensible. Finally, check that terminology matches AQA wording so your answer is unambiguous and mark-scheme friendly.

At GCSE, half equations should usually be balanced with whole-number coefficients. For hydroxide ions at the anode, AQA expects 4OH- -> 2H2O + O2 + 4e- rather than a version with fractional coefficients, because the equation must balance charge and atoms using whole numbers.