Electrolysis: GCSE Chemistry Guide

Picture an ionic compound melted or dissolved in water: the ions are now free to move. Apply a voltage — positive ions drift to the negative electrode, negative ions to the positive electrode — and at each electrode electrons transfer, converting electrical energy into chemical change. That particle picture is the whole of electrolysis.

What is electrolysis?

Electrolysis is the process of using electrical energy to decompose an ionic compound. It requires:

An electrolyte — a substance that conducts electricity when molten or dissolved in water, because it contains free-moving ions (e.g. molten sodium chloride, or sodium chloride solution)

Two electrodes — conductors (usually graphite or metal) dipped into the electrolyte and connected to a direct current (d.c.) power supply

A d.c. power supply — which maintains a potential difference across the electrodes

When the circuit is complete, ions in the electrolyte are attracted to the electrode of opposite charge. At each electrode, ions either lose or gain electrons, converting them back to neutral atoms or molecules. Electrolysis is a redox process — both oxidation and reduction occur simultaneously.

What happens at the cathode and anode?

Electrode | Charge | Ions attracted | Process | OIL RIG
Cathode | Negative (−) | Positive ions (cations) e.g. Na⁺, Cu²⁺, H⁺ | Cations gain electrons → neutral atoms/molecules | Reduction — RIG (gain of electrons)
Anode | Positive (+) | Negative ions (anions) e.g. Cl⁻, OH⁻, O²⁻ | Anions lose electrons → neutral atoms/molecules | Oxidation — OIL (loss of electrons)

Memory hook: CAThode = CATions (opposite charges attract); OIL RIG — Oxidation Is Loss, Reduction Is Gain.

A reliable way to remember which electrode is which: the cathode is negative, and cats always negative-ly scratch — CAT NEG.

What is the electrolysis of brine?

Brine is concentrated sodium chloride solution (NaCl dissolved in water). Electrolysis of brine is one of the most industrially important chemical processes, producing three useful products.

The ions present are: Na⁺, Cl⁻, H⁺ (from water), OH⁻ (from water).

At the cathode (negative electrode): H⁺ ions are preferentially discharged (rather than Na⁺, because hydrogen is lower in the reactivity series and more easily reduced):

2H⁺ + 2e⁻ → H₂

Hydrogen gas is produced.

At the anode (positive electrode): Cl⁻ ions are preferentially discharged (rather than OH⁻, because of the high concentration of Cl⁻):

2Cl⁻ → Cl₂ + 2e⁻

Chlorine gas is produced.

Left in solution: Na⁺ and OH⁻ ions remain, forming sodium hydroxide (NaOH) solution.

Summary of products:

Product | Where produced | Industrial use
Hydrogen (H₂) | Cathode | Making ammonia (Haber process), fuel cells
Chlorine (Cl₂) | Anode | Making bleach, disinfectants, PVC plastic
Sodium hydroxide (NaOH) | Remaining solution | Soap, paper, cleaning products

How do you write half equations for electrolysis?

Half equations for electrolysis follow the same rules as for any redox half equation: balance atoms, then balance charge using electrons.

Cathode half equation (reduction):

Cation gains electrons

Electrons appear on the LEFT-hand side

Example — copper deposition from copper sulfate solution:

Cu²⁺ + 2e⁻ → Cu

Example — hydrogen produced at cathode:

2H⁺ + 2e⁻ → H₂

Anode half equation (oxidation):

Anion loses electrons

Electrons appear on the RIGHT-hand side

Example — chlorine produced at anode:

2Cl⁻ → Cl₂ + 2e⁻

Example — oxygen produced when no halide present (OH⁻ is discharged):

4OH⁻ → 2H₂O + O₂ + 4e⁻

Steps for writing any electrolysis half equation:

Identify the ion being discharged at that electrode

Write the ion on the left (cathode) or right (anode)

Write the product (atom or molecule) on the other side

Balance atoms, then add electrons to balance the charge

What is the electrolysis of aluminium oxide and why is it used?

Aluminium is too reactive to be extracted by reduction with carbon — it sits above carbon in the reactivity series, so carbon cannot displace it from its oxide. Instead, electrolysis of molten aluminium oxide (Al₂O₃) is used, with the oxide dissolved in molten cryolite (Na₃AlF₆) to lower the melting point from ~2,050 °C to ~950 °C (saving enormous amounts of energy).

At the cathode:

Al³⁺ + 3e⁻ → Al (liquid aluminium sinks to the bottom and is tapped off)

At the anode:

2O²⁻ → O₂ + 4e⁻ (oxygen gas is produced)

The oxygen produced at the anode reacts with the graphite (carbon) anode at high temperatures, forming CO₂ and gradually burning away the anode. This means graphite anodes must be replaced regularly — adding to the operating cost.

Aluminium extraction by electrolysis is energy-intensive and therefore expensive, which is why recycling aluminium (which requires only ~5% of the energy of primary extraction) is economically and environmentally important.

What factors affect what is produced at each electrode?

When the electrolyte contains more than two types of ion, several factors determine which ion is discharged:

Factor | Effect
Position in reactivity series (at cathode) | The less reactive metal/ion is preferentially discharged. Cu²⁺ is discharged before H⁺ (copper is less reactive than hydrogen); H⁺ is discharged before Na⁺ (hydrogen is less reactive than sodium)
Concentration of ions (at anode) | At high Cl⁻ concentration, Cl⁻ is discharged in preference to OH⁻, producing Cl₂. At low Cl⁻ concentration, OH⁻ is discharged, producing O₂
State of electrolyte (molten vs aqueous) | In molten ionic compounds, only the ions of that compound are present. In aqueous solutions, H⁺ and OH⁻ from water also compete

Key rule: In aqueous electrolysis, if the metal ion is below hydrogen in the reactivity series (e.g. Cu²⁺, Ag⁺), the metal is deposited at the cathode. If the metal ion is above hydrogen (e.g. Na⁺, Ca²⁺), hydrogen gas is produced instead.

Frequently asked questions

What is electrolysis in simple terms?

Electrolysis is using electricity to split an ionic compound into its elements. You pass direct current through a liquid that contains freely moving ions (the electrolyte). Positive ions move to the negative electrode (cathode) and gain electrons — they are reduced. Negative ions move to the positive electrode (anode) and lose electrons — they are oxidised. The result is that the compound is broken apart and its elements are deposited or released as gases at the electrodes.

What is the difference between the cathode and the anode in electrolysis?

The cathode is the negative electrode, where positive ions (cations) are attracted and gain electrons — this is reduction. The anode is the positive electrode, where negative ions (anions) are attracted and lose electrons — this is oxidation. A reliable memory aid: the cathode is negative (CAT NEG), and OIL RIG — Oxidation Is Loss (anode), Reduction Is Gain (cathode).

Why is aluminium extracted by electrolysis and not by reduction with carbon?

Aluminium is more reactive than carbon, so carbon cannot reduce aluminium oxide — aluminium clings to oxygen more strongly than carbon does. The only way to supply enough energy to force the reduction of Al³⁺ ions to Al metal is to use electrical energy directly. Electrolysis of molten aluminium oxide (dissolved in cryolite to lower the operating temperature) provides the electrical energy needed. This makes aluminium extraction far more energy-intensive and expensive than iron extraction, which uses carbon (coke) in a blast furnace.

What products are made by electrolysis of brine and what are they used for?

Electrolysis of concentrated sodium chloride solution (brine) produces three products: hydrogen gas at the cathode (used in making ammonia via the Haber process and as a clean fuel), chlorine gas at the anode (used in making bleach, disinfectants, and PVC plastic), and sodium hydroxide solution remaining in the electrolyte (used in soap manufacture, paper production, and cleaning products). These three products make the chloralkali industry one of the most economically important applications of electrolysis globally.

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