OCR GCSE Chemistry A: Electrolysis (J248)

This guide covers sub-topic C3.4 Electrolysis from the J248 specification for OCR GCSE (9–1) Chemistry A (Gateway Science), Version 4.0 (August 2026), where it belongs to Topic C3 Chemical reactions. It teaches learning outcomes C3.4a to C3.4e, the mathematical learning outcome CM3.4i and the practical suggestions tied to PAG C2. No C3.4 statement is Higher-Tier-only, so all of it applies to Foundation and Higher candidates alike. The assessment overview places Topic C3 in Paper 1 (Foundation Tier) and Paper 3 (Higher Tier).

After this guide, the Electrolysis revision notes give a shorter version for the final weeks and the Electrolysis practice questions test it. Every J248 topic is listed on the J248 course hub, and the statement-by-statement checklist has every C3.4 point ready to tick. To find your weakest areas of chemistry, try a free 10-minute diagnostic. Electrolysis builds on ionic bonding, so the OCR bonding guide is worth a look first if ions feel shaky.

The C3.4 outcomes at a glance

Statement	You need to	Tier
C3.4a	Recall the inert-electrode rule (metals or hydrogen at the cathode, non-metals at the anode) and use the terms cation and anion	Both tiers
C3.4b	Predict the products when a molten binary ionic compound, such as NaCl, is electrolysed	Both tiers
C3.4c	Describe competing reactions in aqueous solutions, including aqueous NaCl and CuSO₄ with inert electrodes	Both tiers
C3.4d	Describe electrolysis in terms of the ions present, with equations and half equations for the electrode reactions	Both tiers
C3.4e	Describe the technique of electrolysis with inert and non-inert electrodes	Both tiers
CM3.4i	Use arithmetic and ratio for empirical formulae and for balancing equations	Both tiers
PAG C2	Electrolyse aqueous sodium chloride or copper sulfate and test the gases produced	Both tiers

Why an ionic compound must melt or dissolve

Electrolysis is the decomposition of a liquid by passing a direct current through it. The liquid is called the electrolyte. The two rods that dip into it are the electrodes.

An ionic solid contains positive and negative ions held in a lattice by strong electrostatic forces. The ions vibrate about fixed positions; they cannot drift through the solid, so no current flows. When the compound melts, or dissolves in water, the ions are free to move and the liquid conducts.

The specification names two misconceptions to avoid:

• Ionic solutions do not conduct because electrons move through them. The charge is carried by ions moving through the liquid. Electrons flow only in the external wires and are transferred at the electrode surfaces.
• An ionic solid does not fail to conduct because “electrons cannot move”. It fails because its ions are fixed in place.

C3.4a: which electrode, which product

The cathode is the electrode joined to the negative terminal of the power supply. The anode is joined to the positive terminal.

• Cations are positive ions. Opposite charges attract, so cations move to the cathode.
• Anions are negative ions. They move to the anode.

Metal ions and hydrogen ions are positive, so with inert electrodes metals (or hydrogen) form at the cathode. Non-metal ions such as chloride, bromide, iodide and oxide are negative, so non-metals form at the anode.

At the cathode, cations gain electrons. At the anode, anions lose electrons. Describing these changes as reduction and oxidation in terms of electrons is statement C3.3b, which is Higher Tier only.

C3.4b: molten binary ionic compounds

A binary compound contains just two elements. When it is molten, the only ions present are the metal cation and the non-metal anion, so the products follow directly:

• cathode: the metal
• anode: the non-metal (halogens and oxygen form as diatomic molecules: Cl₂, Br₂, I₂, O₂)

For molten sodium chloride, the specification’s own example, sodium metal forms at the cathode and chlorine gas at the anode.

Worked example 1: predicting products

Predict the products when each molten compound is electrolysed with graphite electrodes.

Molten compound	Ions present	Cathode	Anode
calcium iodide, CaI₂	Ca²⁺, I⁻	calcium	iodine
magnesium bromide, MgBr₂	Mg²⁺, Br⁻	magnesium	bromine
lithium oxide, Li₂O	Li⁺, O²⁻	lithium	oxygen

Step 1: split the compound into its ions. Step 2: send the cation to the cathode and the anion to the anode. Step 3: write each product as the element, with halogens and oxygen as diatomic molecules.

Worked example 2: empirical formula from electrolysis (CM3.4i)

A molten metal iodide is electrolysed until it is used up. The cathode collects 0.80 g of calcium and 5.08 g of iodine is formed at the anode. Use Ar values Ca = 40 and I = 127 to find the empirical formula.

	Ca	I
mass (g)	0.80	5.08
mass ÷ Ar	0.80 ÷ 40 = 0.02	5.08 ÷ 127 = 0.04
divide by smaller	1	2

The empirical formula is CaI₂. This matches the charges: one Ca²⁺ balances two I⁻.

C3.4c: competing reactions in aqueous solutions

When an ionic compound dissolves in water, the solution contains the compound’s ions and a small number of hydrogen ions (H⁺) and hydroxide ions (OH⁻) from water. Two cations compete at the cathode and two anions compete at the anode. Only one product usually forms at each electrode.

At the cathode (with inert electrodes):

• Metal below hydrogen in reactivity (copper, silver): the metal is deposited.
• Metal above hydrogen in reactivity (sodium, potassium, magnesium): hydrogen is given off and the metal ions remain in solution.

At the anode (with inert electrodes):

• If chloride, bromide or iodide ions are present, the halogen forms.
• If no halide is present (for example sulfate or nitrate solutions), oxygen forms from hydroxide ions.

Aqueous sodium chloride

Ions present: Na⁺, Cl⁻, H⁺, OH⁻.

• Cathode: sodium is more reactive than hydrogen, so hydrogen gas forms.
• Anode: chloride is a halide, so chlorine gas forms. In a very dilute solution some oxygen can form as well.
• Left in solution: Na⁺ and OH⁻, which make sodium hydroxide. The solution near the cathode turns alkaline.

Aqueous copper(II) sulfate

Ions present: Cu²⁺, SO₄²⁻, H⁺, OH⁻.

• Cathode: copper sits below hydrogen in the reactivity series, so it plates out as a pink-brown layer of copper.
• Anode: there is no halide, so bubbles of oxygen form.
• Left in solution: H⁺ and SO₄²⁻ (sulfuric acid). The blue colour fades as Cu²⁺ ions are removed, and the solution becomes acidic.

Worked example 3: aqueous predictions

Solution (inert electrodes)	Cathode	Reason	Anode	Reason
potassium bromide	hydrogen	K more reactive than H	bromine	halide present
copper(II) bromide	copper	Cu less reactive than H	bromine	halide present
sodium nitrate	hydrogen	Na more reactive than H	oxygen	no halide
lithium iodide	hydrogen	Li more reactive than H	iodine	halide present

Bromine shows as an orange-brown colour around the anode and iodine as a brown colour; neither bubbles off as a gas from cold solution.

C3.4d: half equations

Each electrode has its own half equation, which includes the electrons (e⁻) gained or lost. To build one:

1. Write the ion as the reactant and the element as the product.
2. Make the numbers of each atom match.
3. Add e⁻ to whichever side needs it until the total charges are equal.

Product	Where	Half equation
chlorine	anode	2Cl⁻ → Cl₂ + 2e⁻
sodium, from molten NaCl	cathode	Na⁺ + e⁻ → Na
oxygen, from aqueous solutions	anode	4OH⁻ → O₂ + 2H₂O + 4e⁻
copper	cathode	Cu²⁺ + 2e⁻ → Cu
hydrogen	cathode	2H⁺ + 2e⁻ → H₂

Worked example 4: balancing and ratio (CM3.4i)

Write the half equations for molten calcium iodide and the overall equation.

• Cathode: Ca²⁺ + 2e⁻ → Ca
• Anode: 2I⁻ → I₂ + 2e⁻

Both half equations involve 2 electrons, so they add directly: CaI₂ → Ca + I₂.

When electron numbers differ, multiply first. In aqueous sodium nitrate, hydrogen forms with 2e⁻ per H₂ and oxygen with 4e⁻ per O₂. Doubling the hydrogen half equation gives 4H⁺ + 4e⁻ → 2H₂, so two molecules of hydrogen form for every molecule of oxygen. Gas volumes follow the same 2 : 1 ratio. If 46 cm³ of hydrogen is collected, you would expect 23 cm³ of oxygen. A slightly smaller reading is common because oxygen is a little more soluble in water than hydrogen.

C3.4e and PAG C2: the technique

Inert electrodes

Graphite and platinum are the usual inert electrodes: they conduct electricity without being changed themselves. A typical set-up for an aqueous solution:

• the solution in a beaker
• a pair of graphite rods, kept apart, wired to a low-voltage d.c. supply
• a small test tube over each rod, full of solution and upside down, so any gas pushes the liquid out

For a molten compound, the solid is heated in a crucible until it melts, with the electrodes dipped in. Chlorine and bromine are toxic, so work with them in a fume cupboard or in a well-ventilated room with small quantities.

Testing the gases (PAG C2)

Gas	Test	Positive result
hydrogen	hold a burning splint at the tube’s mouth	goes out with a squeaky pop
oxygen	put a glowing splint into the gas	bursts back into flame
chlorine	hold moist blue litmus paper in the gas	goes red, then white as it is bleached

Non-inert electrodes

A non-inert electrode takes part in the reaction. The standard case is copper electrodes in copper(II) sulfate solution:

• Anode: the copper electrode dissolves. Cu → Cu²⁺ + 2e⁻. No oxygen forms.
• Cathode: copper is deposited. Cu²⁺ + 2e⁻ → Cu.

Copper ions enter the solution at the anode as fast as they leave at the cathode, so the blue colour stays the same. The anode loses mass and the cathode gains the same mass.

Worked example 5: copper electrode masses

Electrode	Mass before (g)	Mass after (g)
cathode	24.36	25.04
anode	26.10	25.42

Cathode change: 25.04 − 24.36 = +0.68 g. Anode change: 25.42 − 26.10 = −0.68 g. The two changes are equal in size, which is what the half equations predict: every copper atom that leaves the anode as Cu²⁺ is replaced by one deposited at the cathode. As a percentage, the cathode gained 0.68 ÷ 24.36 × 100 = 2.79% of its starting mass.

For reliable masses, rinse each electrode gently with distilled water, dry it fully before reweighing, and avoid knocking off loose copper from the cathode.

Common errors

• Saying electrons travel through the solution. Ions carry the charge in the electrolyte.
• Mixing up the electrodes. Cathode is negative and attracts cations.
• Writing “sodium” at the cathode for aqueous sodium chloride. Hydrogen forms because sodium is more reactive than hydrogen.
• Writing “sulfur” or “sulfur dioxide” at the anode for copper sulfate. Sulfate ions stay in solution; oxygen forms.
• Leaving the 2 off Cl₂. At the anode, the electrons belong on the right-hand side as products.
• Forgetting that copper electrodes change the anode reaction: the anode dissolves instead of giving oxygen.
• Describing the chlorine test with dry litmus. The paper must be damp.

Next steps

• Condense the content with the Electrolysis revision notes.
• Test yourself with the Electrolysis practice questions.
• Revisit formulae, ions and balancing in the Introducing chemical reactions guide.
• Once C3.4 feels secure, tick it off on your checklist, then check yourself with a 10-minute diagnostic.

Official syllabus

Source: OCR’s specification for GCSE (9–1) Chemistry A (Gateway Science), J248, Version 4.0 (August 2026), for first assessment in 2018. Content taught here: C3.4 Electrolysis (statements C3.4a–C3.4e), mathematical learning outcome CM3.4i and the PAG C2 practical suggestions.