Electrolysis and Fuel Cells

This guide covers Topic 4, Electrochemistry — subtopics 4.1 Electrolysis and 4.2 Hydrogen–oxygen fuel cells — for Cambridge IGCSE Chemistry 0620 and Cambridge O Level Chemistry 5070, 2026–2028 series.

Where this fits in 0620/5070

Electrolysis connects two ideas from earlier topics — ionic bonding (Topic 2) and redox as electron transfer (Topic 6.4) — into a single practical process. It’s also one of the few IGCSE/O Level topics with a direct, expanded sequel: AS & A Level Chemistry 9701 revisits electrolysis in far more quantitative detail (Faraday’s laws, standard electrode potentials) once you reach Electrochemistry: Electrolysis and Electrode Potentials.

Syllabus coverage

CAMBRIDGE IGCSE CHEMISTRY 0620

Core

• Defining electrolysis as the decomposition of an ionic compound, molten or in aqueous solution, by an electric current (4.1)
• Identifying the anode (positive electrode), cathode (negative electrode) and electrolyte in a simple electrolytic cell (4.1)
• Identifying products and observations for the electrolysis of molten lead(II) bromide, concentrated aqueous sodium chloride, and dilute sulfuric acid, using inert platinum/carbon electrodes (4.1)
• Knowing that metals or hydrogen form at the cathode, and non-metals (other than hydrogen) form at the anode (4.1)
• Predicting the products at each electrode for electrolysis of a molten binary compound (4.1)
• Knowing that metal objects are electroplated for appearance and corrosion resistance, and describing how (4.1)
• Stating that a hydrogen–oxygen fuel cell produces electricity from hydrogen and oxygen, with water as the only chemical product (4.2)

Subtopic 4.2 has almost no Extended-only content — only the advantages/disadvantages comparison with petrol engines below is Supplement.

Supplement / Extended

• Identifying products and observations for the electrolysis of aqueous copper(II) sulfate, both with inert carbon and with copper electrodes (4.1)
• Predicting the products at each electrode for electrolysis of a halide compound in dilute or concentrated aqueous solution (4.1)
• Constructing ionic half-equations for the anode (oxidation) and cathode (reduction) (4.1)
• Describing electrolysis in terms of the transfer of charge — the movement of ions through the solution or melt, and of electrons through the metal conductors and external circuit, to the electrodes (4.1)
• Describing the advantages and disadvantages of hydrogen–oxygen fuel cells compared with petrol engines in vehicles (4.2)

CAMBRIDGE O LEVEL CHEMISTRY 5070

5070 has no Core/Extended split — every outcome above, Core and Supplement alike, is required for every O Level candidate.

What happens during electrolysis

Electrolysis is the decomposition of an ionic compound — molten, or dissolved in water — by passing an electric current through it. Three parts make up an electrolytic cell:

• The electrolyte is the molten or aqueous substance being electrolysed. It must contain freely-moving ions, which is why electrolysis works on molten or dissolved ionic compounds but not on solid ionic compounds (the ions are locked in a lattice) or on covalent substances (no ions at all).
• The anode is the positive electrode.
• The cathode is the negative electrode.

Ions move through the electrolyte toward the electrode of opposite charge: positive ions (cations) move to the cathode, negative ions (anions) move to the anode. At each electrode, ions gain or lose electrons — this is a redox process (Topic 6.4): reduction (gain of electrons) at the cathode, oxidation (loss of electrons) at the anode.

Charge is transferred all the way round the circuit, not just inside the electrolyte. Inside the electrolyte, charge is carried by the movement of ions. Outside the cell, in the wires, electrodes and power supply that make up the external circuit, charge is instead carried by the movement of electrons — electrons flow from the power supply to the cathode (where they are given to cations, reducing them) and from the anode back to the power supply (having been taken from anions there, oxidising them). The complete picture is: ions carry charge through the electrolyte, electrons carry charge through the external circuit, and the two meet at the electrodes.

The general pattern for products:

• Metals or hydrogen form at the cathode.
• Non-metals (other than hydrogen) form at the anode.

Electrolysis of molten compounds

In a molten binary ionic compound, there’s only one possible source for each ion, so the products are predictable directly from the compound’s formula.

Worked example. What forms at each electrode during the electrolysis of molten lead(II) bromide, PbBr₂?

molten PbBr2 contains Pb²⁺ ions and Br⁻ ions

at the cathode (reduction): Pb²⁺ + 2e⁻ → Pb        (lead metal forms)
at the anode (oxidation):   2Br⁻ → Br2 + 2e⁻        (bromine forms)

Electrolysis of aqueous solutions

In aqueous solutions, water itself contributes H⁺ and OH⁻ ions alongside the dissolved compound’s own ions — so predicting the products means knowing which ion is preferentially discharged.

At the cathode: hydrogen forms unless the metal ion present is less reactive than hydrogen (i.e. below hydrogen in the reactivity series), in which case the metal forms instead. This is why electrolysing aqueous copper sulfate deposits copper metal at the cathode, while electrolysing aqueous sodium chloride produces hydrogen gas — sodium is far more reactive than hydrogen and stays in solution.

At the anode: a halide ion is discharged in preference to hydroxide only from a concentrated halide solution (so concentrated aqueous sodium chloride gives chlorine gas); from a dilute halide solution, or wherever no halide is present at all, oxygen forms from hydroxide/water instead.

Worked example. Predict the products of electrolysing concentrated aqueous sodium chloride with inert electrodes.

solution contains: Na⁺, Cl⁻, H⁺, OH⁻ (from water)

at the cathode: H⁺ is discharged in preference to Na⁺ → hydrogen gas
at the anode: Cl⁻ is discharged in preference to OH⁻ → chlorine gas
              2Cl⁻ → Cl2 + 2e⁻

The electrolysis of aqueous copper(II) sulfate is a special case worth knowing separately: with inert (carbon/graphite) electrodes, copper deposits at the cathode and oxygen forms at the anode; with copper electrodes, copper still deposits at the cathode, but the copper anode dissolves instead of oxygen being released — this is the basis of electroplating and of purifying copper.

Worked example. Predict the products of electrolysing dilute sulfuric acid with inert electrodes.

dilute H2SO4 contains: H⁺, SO4²⁻, plus H⁺ and OH⁻ from water

at the cathode: H⁺ is discharged → hydrogen gas
                2H⁺ + 2e⁻ → H2

at the anode: no halide is present, so OH⁻/water is discharged → oxygen gas
              4OH⁻ → O2 + 2H2O + 4e⁻

Hydrogen and oxygen are collected in the ratio 2 : 1 by volume (hydrogen : oxygen), matching the formula of water. Confirm each gas with the standard test: hydrogen gives a squeaky pop with a lit splint; oxygen relights a glowing splint. The sulfate ion, SO₄²⁻, is not discharged at all — it, like the sulfuric acid’s hydrogen ion, simply makes the solution a better conductor, and its concentration in solution actually increases slowly as electrolysis proceeds.

Electroplating

Electroplating coats an object in a thin layer of metal, usually to improve appearance or resist corrosion — chromium-plated taps, silver-plated cutlery. The object to be plated is made the cathode, the plating metal is made the anode, and the electrolyte is a solution of the plating metal’s ions — metal ions from the anode dissolve into solution and are deposited onto the cathode.

Hydrogen–oxygen fuel cells

A hydrogen–oxygen fuel cell produces electricity directly from a chemical reaction between hydrogen and oxygen, with water as the only chemical product — no carbon dioxide, unlike burning a fossil fuel.

	Fuel cell	Petrol/gasoline engine
Product	Water only	CO₂, water, and pollutants (CO, NOₓ)
Efficiency	Generally higher	Loses more energy as heat
Fuel storage	Hydrogen is difficult and costly to store and transport	Petrol is easy to store and transport
Infrastructure	Hydrogen refuelling is not widely available	Petrol stations are everywhere

Common mistakes

• Forgetting that solid ionic compounds don’t conduct. Electrolysis needs freely-moving ions — molten or dissolved, never solid.
• Assuming the most reactive ion is always discharged. It’s the opposite: at the cathode, the least reactive species (closest to/below hydrogen) is discharged first; at the anode, halide ions are discharged before hydroxide, but only from a concentrated halide solution — a dilute halide solution gives oxygen instead, just as if no halide were present.
• Mixing up anode and cathode. Anode = positive = oxidation (Anode, Oxidation — both vowels; Cathode, Reduction — both consonants, if that helps you remember).
• Writing half-equations with unbalanced charge. Check that the electrons on each side balance the change in ionic charge, not just the atoms.
• Assuming fuel cells produce zero pollution overall. The cell itself only produces water — but manufacturing and transporting the hydrogen fuel has its own environmental cost, which is why “advantages and disadvantages” questions expect more than just “no CO₂”.

Quick revision checklist

• Definition of electrolysis, and why an electrolyte must contain freely moving ions
• Anode, cathode and electrolyte, correctly identified
• General rule: metals/hydrogen at the cathode, non-metals at the anode
• Predicting products for molten binary compounds and for common aqueous solutions (including copper(II) sulfate with inert vs copper electrodes)
• (0620 Extended, 5070 required) constructing ionic half-equations for anode and cathode reactions
• How and why metal objects are electroplated
• Hydrogen–oxygen fuel cells: what they produce, and how they compare with petrol engines

Related resources

• Ionic, Covalent and Metallic Bonding — the ionic bonding electrolysis depends on
• Redox Reactions — oxidation and reduction as electron transfer, applied here to each electrode
• Electrochemistry: Electrolysis and Electrode Potentials — the AS & A Level continuation, with Faraday’s laws and standard electrode potentials
• Cambridge IGCSE Chemistry hub · Cambridge O Level Chemistry hub

Written against Cambridge IGCSE Chemistry 0620 and Cambridge O Level Chemistry 5070, 2026–2028 series. Always check the current syllabus for your examination year.