Chemical changes: electrolysis of molten ionic compounds

What you'll learn

This revision guide covers electrolysis of molten ionic compounds, a key topic in AQA GCSE Chemistry. You'll learn how electricity breaks down molten compounds into their elements, why ions move to oppositely charged electrodes, and how to predict the products formed. Understanding this process is essential for questions on chemical changes and extracting metals.

Key terms and definitions

Electrolysis — the decomposition of an ionic compound when molten or in aqueous solution by the passage of electricity

Electrolyte — a liquid or solution that conducts electricity because it contains free-moving ions

Electrode — a solid conductor through which electric current enters or leaves the electrolyte during electrolysis

Anode — the positive electrode, which attracts negatively charged ions (anions)

Cathode — the negative electrode, which attracts positively charged ions (cations)

Oxidation — the loss of electrons (occurs at the anode)

Reduction — the gain of electrons (occurs at the cathode)

Molten — the liquid state of a substance that is normally solid at room temperature, achieved by heating

Core concepts

Why ionic compounds must be molten or dissolved

Ionic compounds in solid form do not conduct electricity because:

• The ions are held in fixed positions in a giant ionic lattice
• The ions cannot move to carry charge
• Electrolysis requires free-moving ions to conduct electricity

When an ionic compound is melted:

• The lattice structure breaks down
• The ions become mobile and can move freely throughout the liquid
• The molten compound now conducts electricity and can undergo electrolysis
• No water is involved, so the products are simpler to predict than in aqueous electrolysis

For example, solid lead bromide (PbBr₂) is an insulator, but molten lead bromide conducts electricity because the Pb²⁺ and Br⁻ ions can move.

The electrolysis process

During electrolysis of a molten ionic compound:

1. The compound is heated until it melts
2. Electrodes are placed into the molten electrolyte
3. A direct current (DC) is applied
4. Positive ions (cations) move toward the cathode (negative electrode)
5. Negative ions (anions) move toward the anode (positive electrode)
6. Chemical reactions occur at each electrode
7. Elements are produced and can be collected

The key principle: opposites attract. Positive ions are attracted to the negative electrode, and negative ions are attracted to the positive electrode.

What happens at the cathode (negative electrode)

At the cathode, reduction occurs. Positive metal ions gain electrons to form metal atoms.

The general equation: Metal ion + electrons → metal atom

For a metal ion with charge n⁺: M^n⁺ + ne⁻ → M

Specific examples:

Lead bromide electrolysis: Pb²⁺ + 2e⁻ → Pb

Aluminium oxide electrolysis: Al³⁺ + 3e⁻ → Al

Sodium chloride electrolysis: Na⁺ + e⁻ → Na

The metal forms at the cathode. You may observe:

• A silvery coating of metal on the electrode
• Droplets of liquid metal forming (for metals with low melting points)
• The electrode appearing shiny or grey

Remember: reduction is gain of electrons. The mnemonic OIL RIG helps: Oxidation Is Loss, Reduction Is Gain (of electrons).

What happens at the anode (positive electrode)

At the anode, oxidation occurs. Negative non-metal ions lose electrons to form atoms, which then usually form molecules.

The general equation: Non-metal ion → non-metal atom + electrons

For a non-metal ion with charge n⁻: X^n⁻ → X + ne⁻

The atoms often join together to form molecules.

Specific examples:

Lead bromide electrolysis: 2Br⁻ → Br₂ + 2e⁻

(Bromine gas forms)

Aluminium oxide electrolysis: 2O²⁻ → O₂ + 4e⁻

(Oxygen gas forms)

Sodium chloride electrolysis: 2Cl⁻ → Cl₂ + 2e⁻

(Chlorine gas forms)

You may observe:

• Gas bubbles forming at the anode
• Coloured gases (brown for bromine, yellow-green for chlorine)
• Distinctive smells (chlorine has a bleach-like smell)

Remember: oxidation is loss of electrons.

Predicting products of electrolysis

To predict the products when any molten ionic compound undergoes electrolysis:

Step 1: Identify the ions present in the compound

For example, molten zinc chloride (ZnCl₂) contains Zn²⁺ and Cl⁻ ions.

Step 2: Determine which ions go to which electrode

• Positive ions (cations) → cathode (negative electrode)
• Negative ions (anions) → anode (positive electrode)

For zinc chloride:

• Zn²⁺ ions move to the cathode
• Cl⁻ ions move to the anode

Step 3: Write the electrode reactions

At the cathode (reduction): Zn²⁺ + 2e⁻ → Zn

At the anode (oxidation): 2Cl⁻ → Cl₂ + 2e⁻

Step 4: State the products

• Cathode product: zinc metal (Zn)
• Anode product: chlorine gas (Cl₂)

This method works for any binary molten ionic compound (containing two elements).

Industrial applications

Electrolysis of molten ionic compounds is crucial for extracting reactive metals that cannot be extracted by reduction with carbon.

Extraction of aluminium:

• Aluminium is extracted from aluminium oxide (Al₂O₃) by electrolysis
• Aluminium oxide is dissolved in molten cryolite to lower the melting point (from 2050°C to about 950°C), saving energy and costs
• Carbon electrodes are used
• Aluminium forms at the cathode: Al³⁺ + 3e⁻ → Al
• Oxygen forms at the anode: 2O²⁻ → O₂ + 4e⁻
• The oxygen reacts with the carbon anodes, producing carbon dioxide and wearing away the electrodes
• Anodes must be replaced regularly

Why electrolysis is used:

Aluminium is more reactive than carbon, so it cannot be extracted by reduction with carbon. Electrolysis is the only viable method, though it requires large amounts of electricity, making it expensive.

Other metals extracted by electrolysis of molten compounds include sodium, calcium, magnesium, and potassium.

Half equations

You must be able to write half equations for the reactions at each electrode. A half equation shows what happens at one electrode only.

Rules for writing half equations:

1. Identify the ion reacting
2. Show the ion becoming an atom or molecule
3. Add electrons (e⁻) to balance the charges
4. Balance atoms if needed

Cathode half equations (reduction — electrons on the left):

Li⁺ + e⁻ → Li Ca²⁺ + 2e⁻ → Ca Al³⁺ + 3e⁻ → Al

Anode half equations (oxidation — electrons on the right):

2Cl⁻ → Cl₂ + 2e⁻ 2Br⁻ → Br₂ + 2e⁻ 2O²⁻ → O₂ + 4e⁻

Notice that:

• The number of electrons equals the charge on the ion
• Non-metal atoms pair up to form diatomic molecules (Cl₂, Br₂, O₂)
• The charges must balance on both sides

Worked examples

Example 1: Electrolysis of molten lithium chloride

Question: Lithium chloride is an ionic compound with the formula LiCl.

a) Name the ions present in molten lithium chloride. [2 marks]

b) During electrolysis of molten lithium chloride, lithium forms at one electrode.

i) At which electrode does lithium form? [1 mark]

ii) Write a half equation for the formation of lithium. [1 mark]

c) Name the product formed at the other electrode. [1 mark]

Answers:

a) Lithium ions / Li⁺ [1 mark] Chloride ions / Cl⁻ [1 mark]

b) i) Cathode / negative electrode [1 mark]

ii) Li⁺ + e⁻ → Li [1 mark]

c) Chlorine / Cl₂ [1 mark]

Mark scheme guidance: In part (a), both the name and formula are acceptable. In part (b)(i), either term is acceptable. The half equation must be balanced with state symbols not required at GCSE level.

Example 2: Predicting products and writing equations

Question: Molten calcium bromide undergoes electrolysis.

a) Predict the products at: i) the cathode [1 mark] ii) the anode [1 mark]

b) Write a half equation for the reaction at the anode. [2 marks]

c) Explain why calcium bromide must be molten for electrolysis to occur. [2 marks]

Answers:

a) i) Calcium / Ca [1 mark]

ii) Bromine / Br₂ [1 mark]

b) 2Br⁻ → Br₂ + 2e⁻ [2 marks]

(Award 1 mark for correct species, 1 mark for balanced equation with electrons)

c) The ions must be free to move [1 mark] so that they can carry charge / conduct electricity [1 mark]

Alternative: In solid form, ions are fixed in position [1 mark], but when molten the ions can move [1 mark]

Mark scheme guidance: In part (b), the equation must be balanced to gain full marks. In part (c), both points about movement and charge/conduction are needed.

Example 3: Understanding oxidation and reduction

Question: Magnesium oxide is electrolysed when molten.

a) Complete the half equation for the reaction at the cathode: Mg²⁺ + _____ → _____ [2 marks]

b) State whether oxidation or reduction occurs at the cathode. Explain your answer. [2 marks]

c) Write the half equation for the reaction at the anode. [2 marks]

Answers:

a) Mg²⁺ + 2e⁻ → Mg [2 marks]

(Award 1 mark for 2e⁻, 1 mark for Mg)

b) Reduction [1 mark]

Because electrons are gained / magnesium ions gain electrons [1 mark]

c) 2O²⁻ → O₂ + 4e⁻ [2 marks]

(Award 1 mark for correct species, 1 mark for correctly balanced equation)

Mark scheme guidance: Explanations of reduction must reference electrons, not just "oxygen is removed" (which applies to different contexts).

Common mistakes and how to avoid them

• Confusing which electrode is which: Remember that the cathode is the negative electrode (both contain 't' when written). The anode is positive. Alternatively, use "a-node" = "a-dding electrons" = positive.

• Getting products at the wrong electrode: Metals always form at the cathode (negative), non-metals at the anode (positive). Think "opposite charges attract" — positive metal ions go to the negative electrode.

• Writing unbalanced half equations: Always check that charges balance. If an ion has a 2+ charge, you need 2 electrons in the half equation. For example, Cu²⁺ + 2e⁻ → Cu, not Cu²⁺ + e⁻ → Cu.

• Forgetting that non-metals form molecules: Elements like chlorine, bromine, and oxygen exist as diatomic molecules (Cl₂, Br₂, O₂), not as single atoms. Your half equations must show this: 2Cl⁻ → Cl₂ + 2e⁻.

• Confusing oxidation and reduction: Use OIL RIG — Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons). Oxidation always occurs at the anode, reduction at the cathode.

• Stating that molten compounds conduct because they contain electrons: Molten ionic compounds conduct electricity because they contain mobile ions, not free electrons. Metals conduct using electrons; ionic compounds use ions.

Exam technique for "Chemical changes: electrolysis of molten ionic compounds"

• Command word "Explain": You must give reasons or mechanisms, not just state facts. For example, "Explain why the compound must be molten" requires you to mention that ions must be free to move AND that they carry charge/conduct electricity.

• Half equations are worth 2 marks typically: One mark for correct species, one mark for balancing (atoms and charges). Always show electrons with the symbol e⁻ and include their number.

• Predicting products: Make sure you clearly identify which product forms at which electrode. Writing "chlorine and copper" without specifying locations loses marks if the question asks where each forms.

• Link structure to function: Questions often ask why ionic compounds conduct when molten but not when solid. Your answer must link the solid lattice structure (fixed ions) to lack of conduction, and the molten state (mobile ions) to conduction.

Quick revision summary

Electrolysis decomposes molten ionic compounds using electricity. The compound must be molten so ions can move freely. Positive metal ions move to the cathode (negative electrode) where reduction occurs — they gain electrons to form metal atoms. Negative non-metal ions move to the anode (positive electrode) where oxidation occurs — they lose electrons to form atoms that usually pair up as molecules. This process extracts reactive metals like aluminium from their ores industrially.