GCSE Chemistry Electrolysis revision

Electrolysis means “splitting with electricity”: a direct current passed through an electrolyte decomposes it at two electrodes. Only ionic compounds can be electrolysed, and only when their ions are free to move — molten, or dissolved in water. A solid crystal locks every ion in place, so it will not conduct; a melted or dissolved one swims with mobile charge carriers. The cathode is the negative electrode and pulls positive ions (cations) onto itself; the anode is positive and pulls negative ions (anions). At the cathode, ions gain electrons (reduction); at the anode they lose electrons (oxidation).

Dissolve an ionic compound and water supplies its own ions: H⁺ (with the cation at the cathode) and OH⁻ (with the anion at the anode). Two rules settle the contests. At the cathode, hydrogen gas forms unless the metal is less reactive than hydrogen — copper and silver are lazy enough to deposit as metal instead. At the anode, if halide ions (Cl⁻, Br⁻, I⁻) are present in reasonable concentration, the halogen forms; if not, oxygen is released from the hydroxide ions.

A half equation shows the electrons explicitly. Cathode (reduction): electrons on the left — Al³⁺ + 3e⁻ → Al; 2H⁺ + 2e⁻ → H₂; Cu²⁺ + 2e⁻ → Cu. Anode (oxidation): electrons on the right — 2Cl⁻ → Cl₂ + 2e⁻; 2O²⁻ → O₂ + 4e⁻; 4OH⁻ → O₂ + 2H₂O + 4e⁻ (aqueous oxygen). Balance the atoms first, then the charge — every half equation must balance both.

Aluminium oxide melts far too high (2072 °C) to electrolyse economically as a pure melt. The cell therefore dissolves the alumina in molten cryolite, a less common aluminium mineral, cutting the working temperature to about 950 °C. Alumina is fed in continuously; Al³⁺ reduces at the carbon-lined cathode (molten aluminium taps out at the bottom) and O²⁻ oxidises at carbon anodes. The oxygen attacks the hot carbon: C + O₂ → CO₂, so the anodes burn away and must be replaced regularly — part of the process’s cost, alongside the enormous electricity demand that puts smelters next to power stations.

1. 1Set-up

   Beaker of solution, two inert electrodes (carbon/graphite or platinum) connected to a low-voltage DC supply.

2. 2Brine (concentrated NaCl)

   Cathode: hydrogen (squeaky pop). Anode: chlorine — pale green, bleaches damp litmus paper. The solution goes alkaline: NaOH forms. Half equations: 2H⁺ + 2e⁻ → H₂ and 2Cl⁻ → Cl₂ + 2e⁻.

3. 3Copper sulfate (inert electrodes)

   Cathode: pink-brown copper coating grows (Cu²⁺ + 2e⁻ → Cu). Anode: oxygen bubbles (4OH⁻ → O₂ + 2H₂O + 4e⁻). The blue colour fades as Cu²⁺ is used up.

4. 4Dilute sulfuric acid

   Hydrogen at the cathode, oxygen at the anode — collected and tested. The volumes come out 2 : 1 H₂ : O₂, because water (H₂O) contains exactly twice as many hydrogen as oxygen atoms.

Variables to control: voltage, electrode separation, concentration and volume of solution. Hazards: chlorine is toxic (ventilate, small quantities); hydrogen is flammable (no flames). Evaluation points: carbon electrodes can crumble and discolour products; platinum is better but pricey.