Products at each electrode, half equations, molten vs aqueous electrolytes — the complete guide to electrolysis at GCSE.
Electrolysis is one of the topics where GCSE Chemistry questions are very predictable — the same types of question appear every year, just with different electrolytes. If you understand the principles clearly and practise the half equation writing, you can pick up almost every mark available on any electrolysis question. This guide covers everything from the basics through to Higher tier half equations and industrial applications.
Electrolysis is the decomposition of a substance using electrical energy. It requires an electrolyte — a substance that conducts electricity when molten or dissolved in water. Electrolytes conduct electricity because they contain ions that are free to move.
In solid ionic compounds, the ions are held in fixed positions in a lattice and cannot move — so solid sodium chloride does not conduct electricity. But when melted or dissolved, the ions become free to move toward electrodes and carry charge through the solution.
The setup always involves two electrodes (usually carbon or platinum, which are inert) connected to a power supply, placed into the electrolyte. The electrode connected to the positive terminal of the power supply is the anode. The electrode connected to the negative terminal is the cathode.
At the cathode (negative electrode): positive ions (cations) are attracted and gain electrons — they are reduced. At the anode (positive electrode): negative ions (anions) are attracted and lose electrons — they are oxidised. Memory aid: OILRIG — Oxidation Is Loss, Reduction Is Gain (of electrons).
When a simple ionic compound is melted and electrolysed, the products are straightforward — the metal is produced at the cathode and the non-metal at the anode.
Example: Molten lead bromide (PbBr₂)
The key point about molten electrolytes: there are only two types of ion present (from the single ionic compound), so there is no competition and the products are entirely predictable.
When an ionic compound is dissolved in water, there are additional ions present from the water itself: H⁺ and OH⁻ ions. This creates competition at the electrodes, and the products depend on which ions are preferentially discharged.
Both metal cations and H⁺ ions are attracted to the cathode. Which one is discharged depends on the reactivity of the metal:
Both anions from the compound and OH⁻ ions from water are attracted to the anode. The discharge depends on concentration:
Cathode: Metal more reactive than H → get H₂. Metal less reactive than H → get the metal deposited.
Anode: Concentrated halide → get the halogen gas. Dilute or non-halide → get O₂.
These two rules cover the vast majority of GCSE aqueous electrolysis questions.
Ions present: H⁺ and SO₄²⁻ from the acid, plus H⁺ and OH⁻ from water.
Ions present: Cu²⁺ and SO₄²⁻ from copper sulfate, plus H⁺ and OH⁻ from water.
This is the chlor-alkali process — one of the most important industrial applications of electrolysis.
Products and their uses: hydrogen (fuel cells, margarine production), chlorine (bleach, PVC, water treatment), sodium hydroxide (soap, paper, bleach production).
Half equations show what happens at a single electrode. Writing them correctly requires balancing both atoms and charge.
❌ Most common error: forgetting to balance the number of atoms before balancing charge. Writing Cl⁻ → Cl + e⁻ is wrong — chlorine is diatomic, so it should be 2Cl⁻ → Cl₂ + 2e⁻. Always check that non-metals like Cl₂, Br₂, O₂ and H₂ are written as diatomic molecules.
Aluminium is extracted from its ore (bauxite, which contains aluminium oxide Al₂O₃) by electrolysis. Aluminium is too reactive to be extracted by reduction with carbon.
The aluminium oxide is dissolved in molten cryolite (Na₃AlF₆) to lower the melting point from 2000°C to around 850°C — making the process far less energy-intensive. The molten mixture is then electrolysed using carbon electrodes.
The full AQA Chemistry electrolysis specification is at the AQA GCSE Chemistry specification page.
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