Electrolysis: KS3 Chemistry Guide

Electrolysis is the process of using electricity to break down an ionic compound — an electrolyte — into its elements. When a direct current passes through a molten or dissolved ionic compound, positive ions move to the negative electrode (cathode) and negative ions move to the positive electrode (anode).

What is an ionic compound and what is an electrolyte?

Ionic compounds form when a metal and a non-metal react, producing positive ions (cations) and negative ions (anions) held together in a lattice by electrostatic attraction. Common examples include sodium chloride (NaCl → Na⁺ + Cl⁻), copper sulfate (CuSO₄ → Cu²⁺ + SO₄²⁻), and aluminium oxide (Al₂O₃ → Al³⁺ + O²⁻).

In solid form, those ions sit in fixed positions and cannot move — so a solid ionic compound does not conduct electricity. When the compound is melted (molten) or dissolved in water, the lattice breaks apart and the ions become free to travel through the liquid. That liquid is now called an electrolyte.

Before writing any equation, picture the ions as charged spheres floating in a liquid, free to drift. The moment you connect a battery, the positive spheres begin moving one way and the negative spheres the other. That movement of charge is the electric current — and it is what makes electrolysis possible.

What is the setup for electrolysis?

Electrolysis requires four components working together:

Power supply — must provide direct current (DC), not alternating current. DC keeps a consistent direction so ions always travel the same way.

Electrodes — conductors dipped into the electrolyte. Graphite or platinum electrodes are described as inert because they do not react with the electrolyte. Copper electrodes are used when copper transfer is the goal.

Cathode — the electrode connected to the negative terminal. Because opposite charges attract, positive ions (cations) migrate towards it.

Anode — the electrode connected to the positive terminal. Negative ions (anions) migrate towards it.

A useful memory trick: CAtions go to the CAthode; ANions go to the ANode — think "CAT" and "ANna". At the electrodes, ions gain or lose electrons: reduction (gain of electrons) occurs at the cathode; oxidation (loss of electrons) occurs at the anode.

What happens at each electrode?

At the cathode (negative electrode): positive metal ions arrive and gain electrons, being reduced to neutral atoms. For example, Cu²⁺ + 2e⁻ → Cu deposits copper metal onto the cathode. In aqueous solutions, hydrogen ions are also reduced: 2H⁺ + 2e⁻ → H₂, producing hydrogen gas. The OIL RIG mnemonic helps here: Oxidation Is Loss, Reduction Is Gain.

At the anode (positive electrode): negative ions arrive and lose electrons, being oxidised. Chloride ions are discharged as chlorine gas: 2Cl⁻ → Cl₂ + 2e⁻. In dilute or non-halide solutions, hydroxide ions are oxidised instead: 4OH⁻ → 2H₂O + O₂ + 4e⁻, releasing oxygen gas.

What are worked examples of electrolysis?

Example 1: Electrolysis of molten lead bromide (PbBr₂)

The electrolyte is molten PbBr₂ → Pb²⁺ + 2Br⁻ (ions free to move once melted).

At the cathode: Pb²⁺ + 2e⁻ → Pb — liquid lead metal forms and sinks to the bottom.

At the anode: 2Br⁻ → Br₂ + 2e⁻ — orange-brown bromine vapour is produced.

Overall: PbBr₂ → Pb + Br₂ — the compound is split into its two elements.

Example 2: Electrolysis of copper sulfate solution with copper electrodes

The electrolyte contains Cu²⁺ and SO₄²⁻ ions.

At the cathode: Cu²⁺ + 2e⁻ → Cu — copper deposits and the cathode grows in mass.

At the anode (copper): Cu → Cu²⁺ + 2e⁻ — the copper anode dissolves.

Net result: copper is transferred from anode to cathode; the Cu²⁺ concentration stays roughly constant.

Applications: electroplating and copper purification (electrorefining).

Example 3: Electrolysis of brine (sodium chloride solution)

At the anode: 2Cl⁻ → Cl₂ + 2e⁻ — chlorine gas is produced.

At the cathode: 2H⁺ + 2e⁻ → H₂ — hydrogen gas is produced.

Sodium hydroxide (NaOH) solution remains.

Uses: Cl₂ → PVC plastic, bleach, disinfectants; H₂ → margarine, rocket fuel; NaOH → paper-making, soap, bleach.

What are the industrial uses of electrolysis?

Process | Electrolyte | Cathode product | Anode product | Use
Aluminium extraction | Molten aluminium oxide (with cryolite) | Aluminium metal | Oxygen gas | Aircraft, packaging
Copper purification | Copper sulfate solution | Pure copper | Impure copper dissolves | Electrical wiring
Chlorine production (chlor-alkali) | Brine (NaCl solution) | Hydrogen gas + NaOH | Chlorine gas | PVC, bleach, soap
Electroplating | Solution of the coating metal | Object being plated (e.g. steel spoon) | Coating metal anode | Cutlery, jewellery, car parts
Electrorefining | Solution of the metal | Pure metal deposited | Impure metal dissolves | Purifying copper, silver

Why is aluminium extracted by electrolysis rather than by reduction with carbon?

Carbon can only displace metals that sit below it in the reactivity series — iron, copper, and tin, for example. Aluminium sits above carbon, so carbon cannot reduce aluminium oxide into aluminium metal. Electrolysis is therefore the only practical extraction route.

The process dissolves aluminium oxide (Al₂O₃, called alumina) in molten cryolite. Cryolite lowers the melting point from roughly 2 000 °C to around 1 000 °C, dramatically reducing energy costs. At the graphite cathode lining, Al³⁺ + 3e⁻ → Al produces molten aluminium that is tapped from the base of the cell. At the graphite anodes, O²⁻ is oxidised to oxygen gas, which reacts with the hot carbon — so anodes must be replaced regularly, adding to running costs.

Before the invention of electrolysis in the 1880s, aluminium was rarer and more expensive than gold. Electrolysis made it the abundant, affordable metal we rely on today.

Frequently asked questions

Why must the ionic compound be molten or dissolved for electrolysis to work?

In a solid ionic compound, ions are locked into a rigid lattice and cannot move. Without mobile ions, no electric current can pass through the substance and no electrolysis occurs. When the compound is melted or dissolved in water, the lattice breaks apart and the ions become free to move. A voltage then drives positive ions to the cathode and negative ions to the anode — that movement of charges constitutes an electric current, allowing ions to be discharged at the electrodes.

What is the difference between electrolysis of a molten compound versus an aqueous solution?

In a pure molten ionic compound, only the ions from that compound are present. Molten lead bromide produces only lead at the cathode and bromine at the anode. In an aqueous solution, water molecules contribute H⁺ and OH⁻ ions, creating competition at each electrode. At the cathode, either metal cations or H⁺ might be reduced; at the anode, either the anion or OH⁻ might be oxidised. The key KS3 rule: hydrogen is produced at the cathode from dilute aqueous solutions, and a high-concentration halide (such as Cl⁻) is preferentially discharged at the anode; otherwise oxygen is produced.

What is electroplating and how does it work?

Electroplating uses electrolysis to coat a thin layer of one metal onto the surface of another object. The object to be plated is made the cathode; an electrode of the coating metal is the anode; and the electrolyte is a solution of a salt of the coating metal. To silver-plate a spoon: the spoon is the cathode, a piece of silver is the anode, and silver nitrate solution is the electrolyte. Silver ions (Ag⁺ + e⁻ → Ag) are reduced at the spoon's surface, depositing a thin layer, while the silver anode dissolves to replenish Ag⁺ ions in solution.

How is electrolysis used to extract aluminium from its ore?

Bauxite ore is first purified to aluminium oxide (alumina, Al₂O₃). Because aluminium is too reactive for carbon reduction, the alumina is dissolved in molten cryolite at around 1 000 °C to form a conducting melt. A large direct current is passed through: at the graphite-lined cathode, Al³⁺ + 3e⁻ → Al produces liquid aluminium that sinks and is tapped off. At the graphite anodes, O²⁻ → O₂ + electrons — the oxygen reacts with the hot carbon anodes, which must therefore be replaced regularly. The process consumes enormous quantities of electricity, which is why recycling aluminium uses only about 5 % of the energy needed to produce new metal by electrolysis.

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