Electrolysis: GCSE Chemistry Explained with Equations & Examples

Electrolysis is a core topic in GCSE Chemistry that brings together ideas about ions, bonding, and chemical reactions. It regularly appears in exams as both short-answer and extended-response questions. This guide covers everything you need to know, from the basic definitions through to predicting products and writing half equations.

What Is Electrolysis?

Electrolysis is the process of using an electric current to decompose (break down) an ionic compound. The compound must be either molten (melted) or dissolved in water so that its ions are free to move and carry the charge. A solid ionic compound cannot undergo electrolysis because its ions are locked in a fixed lattice and cannot move to the electrodes.

Key Definitions

Before diving into the detail, make sure you know these terms precisely:

Electrolyte: the ionic compound that is being decomposed, either molten or in aqueous solution
Electrode: a solid conductor through which the electric current enters or leaves the electrolyte (usually made of graphite or platinum, which are inert)
Anode: the positive electrode (connected to the positive terminal of the power supply)
Cathode: the negative electrode (connected to the negative terminal of the power supply)

A helpful mnemonic: the anode is positive (vowels: a and p go together), or remember PANIC — Positive Anode, Negative Is Cathode.

What Happens at Each Electrode?

During electrolysis, positive ions (cations) are attracted to the negative cathode, and negative ions (anions) are attracted to the positive anode.

At the Cathode (Negative Electrode)

Positive ions gain electrons. This is reduction. Metal ions or hydrogen ions are deposited or released here.

At the Anode (Positive Electrode)

Negative ions lose electrons. This is oxidation. Non-metal atoms are released here, often as a gas.

OILRIG

The mnemonic OILRIG helps you remember which process happens where:

Oxidation Is Loss (of electrons) — happens at the anode
Reduction Is Gain (of electrons) — happens at the cathode

Electrolysis of Molten Compounds

When an ionic compound is molten, only two types of ion are present. The products are straightforward to predict:

The metal forms at the cathode
The non-metal forms at the anode

Example: Molten Lead Bromide (PbBr₂)

At the cathode: Pb²⁺ + 2e⁻ —> Pb (lead metal is deposited)

At the anode: 2Br⁻ —> Br₂ + 2e⁻ (bromine gas is released)

The half equations must balance for both atoms and charge.

Example: Molten Sodium Chloride (NaCl)

At the cathode: Na⁺ + e⁻ —> Na (sodium metal)

At the anode: 2Cl⁻ —> Cl₂ + 2e⁻ (chlorine gas)

Electrolysis of Aqueous Solutions

Aqueous solutions are more complex because water also provides ions: H⁺ and OH⁻ (from the partial dissociation of water). This means there are competing ions at each electrode, and you need rules to predict which one reacts.

At the Cathode: Which Positive Ion Is Discharged?

Use the reactivity series:

If the metal is more reactive than hydrogen (e.g., sodium, potassium, calcium, magnesium, aluminium), hydrogen gas is produced instead of the metal. The metal ions stay in solution.
If the metal is less reactive than hydrogen (e.g., copper, silver, gold), the metal is deposited at the cathode.

At the Anode: Which Negative Ion Is Discharged?

If a halide ion is present (Cl⁻, Br⁻, or I⁻), the halogen is produced (chlorine, bromine, or iodine).
If no halide is present (e.g., sulfate or nitrate solutions), oxygen is produced from the discharge of OH⁻ ions.

Example: Aqueous Sodium Chloride (Brine)

At the cathode: Sodium is more reactive than hydrogen, so hydrogen is produced.

2H⁺ + 2e⁻ —> H₂

At the anode: Chloride ions are present (halide), so chlorine is produced.

2Cl⁻ —> Cl₂ + 2e⁻

The remaining ions (Na⁺ and OH⁻) stay in solution, forming sodium hydroxide. This is the basis of the chlor-alkali industry.

Example: Aqueous Copper Sulfate

At the cathode: Copper is less reactive than hydrogen, so copper is deposited.

Cu²⁺ + 2e⁻ —> Cu

At the anode: No halide is present, so oxygen is produced.

4OH⁻ —> O₂ + 2H₂O + 4e⁻

The solution changes from blue to colourless as copper ions are removed.

The Required Practical

The required practical involves the electrolysis of copper sulfate solution using inert (graphite or carbon) electrodes.

What You Observe

Cathode: a pink/brown coating of copper metal appears on the electrode
Anode: bubbles of oxygen gas are produced
Solution: the blue colour fades over time as Cu²⁺ ions are removed

With Copper Electrodes

If you use copper electrodes instead of inert ones, something different happens. The copper anode dissolves, replacing the Cu²⁺ ions in solution. The cathode gains copper. The solution stays blue. This is called purification of copper and is used industrially to produce pure copper for electrical wiring.

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

At the cathode: Cu²⁺ + 2e⁻ —> Cu (pure copper deposited)

Writing Half Equations

Half equations show what happens at each electrode separately. To write one:

Write the ion involved
Add electrons to balance the charge
Check atoms and charges balance on both sides

Reduction (cathode): electrons appear on the left (gained by the ion)

Oxidation (anode): electrons appear on the right (lost by the ion)

Example Half Equations

Pb²⁺ + 2e⁻ —> Pb (reduction)
2Cl⁻ —> Cl₂ + 2e⁻ (oxidation)
2H⁺ + 2e⁻ —> H₂ (reduction)
4OH⁻ —> O₂ + 2H₂O + 4e⁻ (oxidation)

The OH⁻ equation is the hardest to remember. Notice that water is produced as well as oxygen, and four electrons are transferred.

Industrial Uses of Electrolysis

Aluminium Extraction

Aluminium is too reactive to extract by heating with carbon, so electrolysis is used. Aluminium oxide (Al₂O₃) is dissolved in molten cryolite (to lower the melting point from over 2000 degrees Celsius to about 950 degrees Celsius) and electrolysed.

At the cathode: Al³⁺ + 3e⁻ —> Al (molten aluminium)

At the anode: 2O²⁻ —> O₂ + 4e⁻ (oxygen gas)

The carbon anodes react with the oxygen and gradually burn away, so they must be replaced regularly. This is a frequent exam question.

Electroplating

Objects can be coated with a thin layer of metal using electrolysis. The object to be plated is made the cathode, and the plating metal is the anode. The electrolyte is a solution of the plating metal’s salt.

Common Mistakes to Avoid

Confusing Anode and Cathode

Remember: the anode is positive, the cathode is negative. Positive ions go to the negative electrode (cathode), and negative ions go to the positive electrode (anode). Getting this the wrong way round will cause every subsequent answer to be wrong.

Forgetting the Rules for Aqueous Solutions

In molten electrolysis, the products are simply the metal and the non-metal. In aqueous solutions, you must apply the reactivity series (cathode) and the halide test (anode). Many students apply molten rules to aqueous solutions and lose marks.

Unbalanced Half Equations

Every half equation must balance for both atoms and charge. Check that the number of electrons matches the total charge change. For example, Al³⁺ needs 3 electrons, not 2.

Saying Electrolysis Works on Solids

Electrolysis requires ions that are free to move. Solid ionic compounds have ions fixed in a lattice, so they cannot conduct electricity or undergo electrolysis. This is a key point that examiners test regularly.

Exam Technique Tips

Define terms precisely. If asked “what is electrolysis?”, include all the key parts: decomposition, ionic compound, molten or dissolved, using electricity.

State the electrode. When predicting products, always say which electrode the product forms at. “Copper is deposited at the cathode” earns more marks than just “copper is produced.”

Use OILRIG. When writing half equations, explicitly state whether oxidation or reduction is occurring. This shows the examiner you understand the electron transfer.

Summary

Electrolysis decomposes ionic compounds using electricity. Positive ions move to the cathode (reduction), negative ions move to the anode (oxidation). For molten compounds, products are the constituent metal and non-metal. For aqueous solutions, use the reactivity series and halide rules to predict products. Master the half equations and required practical, and you will be well prepared for any electrolysis question.

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