Electrolysis

This lesson covers electrolysis as required by the Edexcel GCSE Combined Science specification (1SC0). You need to understand what electrolysis is, the electrolysis of molten compounds and aqueous solutions, how to predict products at each electrode, and — for higher tier — how to write half equations.

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What Is Electrolysis?

Electrolysis is the process of using electricity to decompose (break down) an ionic compound that is either molten (melted) or dissolved in water (aqueous).

Key Terms

Term	Definition
Electrolyte	The ionic compound (molten or in solution) that is decomposed
Electrode	A solid conductor through which electricity enters or leaves the electrolyte
Cathode	The negative electrode (connected to the negative terminal of the power supply)
Anode	The positive electrode (connected to the positive terminal of the power supply)
Cation	A positive ion (e.g. Na⁺, Cu²⁺, H⁺) — attracted to the cathode
Anion	A negative ion (e.g. Cl⁻, O²⁻, OH⁻) — attracted to the anode

How Does Electrolysis Work?

1. The ionic compound must be molten or dissolved so that its ions are free to move.
2. When electricity flows, positive ions (cations) move to the cathode (−) and negative ions (anions) move to the anode (+).
3. At the cathode, cations gain electrons (are reduced) — forming atoms or molecules.
4. At the anode, anions lose electrons (are oxidised) — forming atoms or molecules.

graph TD
    subgraph Electrolysis Cell
        PS["Power Supply<br/>DC"] --- C["Cathode (−)"]
        PS --- An["Anode (+)"]
        C --- E["Electrolyte<br/>(molten or aqueous<br/>ionic compound)"]
        An --- E
    end

    subgraph Ion Movement
        CAT["Cations (+)<br/>move to Cathode (−)<br/>GAIN electrons<br/>(REDUCTION)"]
        ANI["Anions (−)<br/>move to Anode (+)<br/>LOSE electrons<br/>(OXIDATION)"]
    end

    style PS fill:#2c3e50,color:#fff
    style C fill:#2980b9,color:#fff
    style An fill:#c0392b,color:#fff
    style E fill:#27ae60,color:#fff
    style CAT fill:#2980b9,color:#fff
    style ANI fill:#c0392b,color:#fff

Exam Tip: Remember: Cathode = Cations, Anode = Anions. Cations are positive and move to the negative cathode. Anions are negative and move to the positive anode. Also: Reduction at the Cathode, Oxidation at the Anode.

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Electrolysis of Molten Compounds

When a pure molten ionic compound is electrolysed, there are only two ions present — one positive and one negative. This makes predicting the products straightforward.

Example 1: Electrolysis of Molten Lead Bromide (PbBr₂)

Feature	Detail
Electrolyte	Molten lead bromide (PbBr₂)
Ions present	Pb²⁺ and Br⁻
At the cathode (−)	Pb²⁺ + 2e⁻ → Pb (lead metal — silvery liquid)
At the anode (+)	2Br⁻ → Br₂ + 2e⁻ (bromine gas — brown/orange fumes)

Example 2: Electrolysis of Molten Sodium Chloride (NaCl)

Feature	Detail
Electrolyte	Molten sodium chloride
Ions present	Na⁺ and Cl⁻
At the cathode (−)	Na⁺ + e⁻ → Na (sodium metal)
At the anode (+)	2Cl⁻ → Cl₂ + 2e⁻ (chlorine gas)

Example 3: Electrolysis of Molten Aluminium Oxide (Al₂O₃)

Feature	Detail
Electrolyte	Al₂O₃ dissolved in molten cryolite
Ions present	Al³⁺ and O²⁻
At the cathode (−)	Al³⁺ + 3e⁻ → Al (aluminium metal)
At the anode (+)	2O²⁻ → O₂ + 4e⁻ (oxygen gas)

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Electrolysis of Aqueous Solutions

When an ionic compound is dissolved in water, the situation is more complicated because water partially dissociates into H⁺ ions and OH⁻ ions. There are now four ions in the solution.

Example: Aqueous Sodium Chloride (NaCl)

Ions present: Na⁺, H⁺, Cl⁻, OH⁻

Rules for Predicting Products at the Cathode (−)

At the cathode, metal ions or hydrogen ions are reduced.

Rule	Product at Cathode
If the metal is more reactive than hydrogen (above H in reactivity series)	Hydrogen gas is produced (H⁺ ions are reduced instead)
If the metal is less reactive than hydrogen (below H in reactivity series)	The metal is deposited

Rules for Predicting Products at the Anode (+)

At the anode, non-metal ions are oxidised.

Rule	Product at Anode
If a halide ion (Cl⁻, Br⁻, I⁻) is present	The halogen is produced (Cl₂, Br₂ or I₂)
If no halide ion is present (e.g. sulfate, nitrate)	Oxygen gas is produced (from OH⁻ ions)

Exam Tip: For aqueous solutions: at the cathode, if the metal is reactive (above hydrogen), you get hydrogen. At the anode, if a halide is present, you get the halogen; otherwise you get oxygen. These rules cover most exam questions.

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Summary Table: Products of Electrolysis of Aqueous Solutions

Electrolyte	Ions Present	Product at Cathode (−)	Product at Anode (+)
NaCl (aq)	Na⁺, H⁺, Cl⁻, OH⁻	Hydrogen (Na is more reactive than H)	Chlorine (halide present)
CuSO₄ (aq)	Cu²⁺, H⁺, SO₄²⁻, OH⁻	Copper (Cu is less reactive than H)	Oxygen (no halide)
CuCl₂ (aq)	Cu²⁺, H⁺, Cl⁻, OH⁻	Copper (Cu is less reactive than H)	Chlorine (halide present)
H₂SO₄ (aq)	H⁺, SO₄²⁻, OH⁻	Hydrogen	Oxygen (no halide)
NaBr (aq)	Na⁺, H⁺, Br⁻, OH⁻	Hydrogen (Na is more reactive than H)	Bromine (halide present)
KNO₃ (aq)	K⁺, H⁺, NO₃⁻, OH⁻	Hydrogen (K is more reactive than H)	Oxygen (no halide)

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Half Equations (Higher Tier)

For higher tier, you need to be able to write half equations to show what happens at each electrode.

At the Cathode (Reduction — Gain of Electrons)

• Cu²⁺ + 2e⁻ → Cu
• 2H⁺ + 2e⁻ → H₂
• Pb²⁺ + 2e⁻ → Pb
• Al³⁺ + 3e⁻ → Al

At the Anode (Oxidation — Loss of Electrons)

• 2Cl⁻ → Cl₂ + 2e⁻
• 2Br⁻ → Br₂ + 2e⁻
• 4OH⁻ → O₂ + 2H₂O + 4e⁻
• 2O²⁻ → O₂ + 4e⁻