Everything on the GCSE Chemistry Electrolysis poster is written out below, section by section. Use it to search the sheet, copy parts into your own notes, or check a fact quickly.
What electrolysis is
Electrolysis uses an electric current to split an ionic compound (the electrolyte). The compound must be molten or dissolved so that ions can move freely.
- Cations (+) move to the cathode (-ve electrode) and gain electrons (reduction).
- Anions (-) move to the anode (+ve electrode) and lose electrons (oxidation).
- Electrodes are usually inert (graphite or platinum).
A D.C. power supply drives the process. The electrolyte is molten or in solution between the two electrodes.
Molten example - aluminium oxide
Aluminium is more reactive than carbon, so it cannot be extracted by reduction with carbon - it has to be electrolysed.
- Aluminium oxide (Al₂O₃) comes from the ore bauxite.
- Its melting point is very high, so it is dissolved in molten cryolite to lower the melting point and save energy.
- Carbon electrodes are used.
- O₂ gas bubbles off at the anode; molten Al₂O₃ dissolves in cryolite.
Half equations:
- Cathode (reduction): Al³⁺ + 3e⁻ → Al
- Anode (oxidation): 2O²⁻ → O₂ + 4e⁻
Rules for aqueous examples
Water adds H⁺ and OH⁻ to the mix, so four ions compete: the metal cation, the anion, H⁺, and OH⁻.
Cathode (reduction) - is the metal less reactive than hydrogen?
- Yes → metal forms (Cu, Ag, Au)
- No → hydrogen forms (H₂)
Anode (oxidation) - is a halide (Cl⁻, Br⁻, I⁻) present?
- Yes → halogen forms (Cl₂, Br₂, I₂)
- No → oxygen forms (from OH⁻, producing O₂)
Example 1 - CuSO₄(aq) (Cu less reactive than H)
- Cathode: Cu²⁺ + 2e⁻ → Cu (no halide)
- Anode: 4OH⁻ → O₂ + 2H₂O + 4e⁻
Example 2 - NaCl(aq) (Na more reactive than H)
- Cathode: 2H₂O + 2e⁻ → H₂ + 2OH⁻
- Anode (halide present): 2Cl⁻ → Cl₂ + 2e⁻