The Process of Electrolysis, Electrolysis of Molten Ionic Compounds and Using Electrolysis to Extract Metals

SPEC 4.4.3.1 – 4.4.3.3 • PAPER 1

Electrolysis

Decomposition driven by electricity. Cathode and anode, ions vs electrons, molten compounds, and aluminium extraction. Half-equations on Higher Tier.

What Is Electrolysis?

Electrolysis

The decomposition of an ionic compound using an electric current, when the compound is molten or dissolved in water.

Decomposes — not Separates

Electrolysis is a chemical change: it breaks the compound apart into its elements. New substances form. "Separation" implies physical methods (filtration, evaporation, etc.) — different process, different word, different mark.

Examiner Alert (P1 Nov 2021 Q07.1)

"Many students believe electrolysis is a method of 'separating' compounds, rather than decomposing a compound." Mark scheme accepts "uses electricity to decompose a compound / electrolyte" — never "separates".

Why Ions Must Be Free to Move

For electrolysis to work, the ions in the compound must be free to move to the electrodes. Their state of matter decides whether they can.

StateCan it be electrolysed?Why

Solid ionic compoundNOIons are fixed in the lattice — they cannot move to electrodes

Molten (liquid) ionic compoundYESHeat melts the lattice — ions are free to move

Aqueous (dissolved in water)YESIons are dissociated into solution — free to move (covered in 4.08)

Electrolyte

A molten or aqueous ionic substance — the liquid through which ions move during electrolysis.

The Two Electrodes

Cathode (–)

The negative electrode. Attracts positive ions (cations) — including all metal ions and H⁺.

At the cathode, cations gain electrons → reduction.

Anode (+)

The positive electrode. Attracts negative ions (anions) — including all non-metal ions and OH⁻.

At the anode, anions lose electrons → oxidation.

Memory Hook

Opposites attract. Negative cathode pulls positive cations; positive anode pulls negative anions. A useful mnemonic: OIL RIG — Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons). Anode = oxidation; cathode = reduction.

Discharge

When an ion loses its charge at an electrode, becoming a neutral atom or molecule. The ion is discharged — the resulting element is produced (not "discharged" itself).

Inert Electrodes

Electrodes that don't take part in the reaction — they only supply or accept electrons. Graphite (carbon) and platinum are typical inert electrodes. Spec 4.4.3.2 specifies inert electrodes for molten electrolysis. (Note: in the aluminium process, the carbon anodes do react — they're a special case covered in slide 10.)

Charge Carriers — Ions vs Electrons

This is the single most-confused point in electrolysis. The current flows in one continuous loop, but different particles carry the charge in different parts of the circuit.

WhereCharge carrierWhy

External wires (between battery and electrodes)ElectronsSolid metal conductors have free delocalised electrons

The electrodes themselvesElectronsSolid conductors (graphite or metal) — same as wires

The electrolyte (molten / aqueous)IONSNo free electrons exist in the liquid — only mobile ions

Examiner Alert (P1 Jun 2022 Q06.3)

Q: "Which type of particle passes through the mesh in the electrolysis of molten sodium chloride? (Atom / Electron / Ion / Molecule)" Mark scheme: ion. ER: "Only 60% of students recognised that ions must pass through the mesh. Nearly all other responses were 'electrons'. Many students seem to think that an electric current can only be conducted by electrons moving, even through a solution in which there are no free electrons."

Inside an Electrolysis Cell — Molten PbBr₂

What's Happening

Pb²⁺ ions move to the cathode (–) and gain 2 electrons → molten lead. Br⁻ ions move to the anode (+) and lose electrons → orange-brown bromine vapour. Overall: PbBr₂(l) → Pb(l) + Br₂(g).

Half Equations for Molten PbBr₂ HT Only

A half equation shows what happens at one electrode — gain or loss of electrons. Each must balance in both atoms and charges.

Cathode (–) — Reduction

Pb²⁺ + 2e⁻ → Pb

Pb²⁺ gains 2 electrons; charge balances (+2 + −2 = 0).

Anode (+) — Oxidation

2Br⁻ → Br₂ + 2e⁻

Two Br⁻ lose 2 electrons total; charge balances (−2 = 0 + −2).

Examiner Alert (P1 Jun 2022 Q06.1)

Q: "Which is the correct half equation for the production of sodium?" The mark-scheme answer was Na⁺ + e⁻ → Na. ER: "A significant number chose a distractor in which Na was present at the start rather than formed in the reaction." Watch out: the ion is the reactant; the element is the product. Reverse direction = no mark.

Chemical Cell vs Electrolytic Cell

These two devices use the same components — electrodes in an electrolyte — but with opposite directions of energy flow. The examiner tests this distinction directly.

Electrolytic Cell

USES electricity to drive a chemical reaction.

Electrical energy → chemical change (decomposition).

Examples: molten NaCl, aluminium extraction, extraction of reactive metals (Mg, Na, K).

Chemical Cell (Battery)

PRODUCES electricity from a chemical reaction.

Chemical change → electrical energy.

Examples: AA battery, hydrogen fuel cell, lead-acid car battery.

Examiner Alert (P1 Nov 2021 Q07.1)

Mark scheme: "electrolysis uses electricity to produce a chemical reaction" + "cells use a chemical reaction to produce electricity". ER: "Very few realised that a chemical cell uses a reaction to produce electricity, with many making statements about a cell using electricity." Reversing the direction loses both marks.

Why Electrolysis to Extract Metals?

Most metals can be extracted by reduction with carbon — cheap and energy-efficient (just heat the metal oxide with coke). Electrolysis is reserved for metals that can't be reduced by carbon.

The Reactivity Rule

Electrolysis is used in two cases (Spec 4.4.3.3): (1) the metal is too reactive to be extracted by reduction with carbon — examples: aluminium, magnesium, calcium, sodium, potassium; or (2) the metal reacts with carbon — for example, titanium reacts with carbon to form titanium carbide, which makes the metal brittle. Carbon can only reduce metals less reactive than itself (zinc, iron, copper, etc.) without forming unwanted products.

Energy Cost

Electrolysis requires large amounts of electrical energy — to melt the compound AND to drive the decomposition. This makes electrolytically-extracted metals more expensive than carbon-reduced metals.

Mark-Scheme Phrase (P1 Jun 2022 Q07.2)

"More energy is needed (to obtain aluminium); (because) aluminium is obtained (from aluminium oxide) by electrolysis." Note the comparison: explain WHY more energy — name the process and (where contrast applies) compare to the carbon-reduction alternative.

Aluminium Extraction from Al₂O₃

Aluminium is more reactive than carbon, so it cannot be reduced by carbon. The industrial process electrolyses molten aluminium oxide.

The Process

Aluminium oxide (Al₂O₃) is dissolved in molten cryolite. Pure Al₂O₃ melts at ~2050 °C; in cryolite the mix melts at ~950 °C.

Cryolite lowers the melting point — saves enormous amounts of energy.

Carbon (graphite) electrodes are used.

Aluminium forms at the cathode (Al³⁺ gains electrons → Al). Molten aluminium is denser than the cryolite mix and collects at the bottom.

Oxygen forms at the anode (2 O²⁻ → O₂ + 4e⁻).

Oxygen reacts with the carbon anodes at high temperatures → CO₂. The anodes burn away and must be replaced regularly (a major operating cost).

Half Equations (HT)

Cathode: Al³⁺ + 3e⁻ → Al | Anode: 2 O²⁻ → O₂ + 4e⁻

Charges balance: cathode (+3 + −3 = 0); anode (−4 = 0 + −4). Atoms balance.

Inside the Aluminium Cell

Industrial aluminium extraction uses a unique cell design — the cell itself acts as the cathode, and multiple carbon anodes hang into the molten mixture from above.

What's Unique About This Cell

Unlike a simple electrolysis cell (one cathode rod, one anode rod), this cell has multiple carbon anodes suspended into the mixture from above, and the cell itself is the cathode (steel container lined with carbon). Molten aluminium collects at the bottom because it is denser than the cryolite mix.

Why Aluminium Is Expensive

The Three Cost Drivers

High electrical energy demand — both to melt the cryolite-Al₂O₃ mix and to drive the electrolysis itself.

Anode replacement — carbon anodes react with the oxygen produced and burn away as CO₂. They must be replaced regularly.

Aluminium is more reactive than carbon — so cannot be cheaply reduced by carbon (the standard, low-cost route used for iron, copper, etc.).

Examiner Alert (P1 Jun 2022 Q07.2)

ER: "Although many mentioned the high energy cost of electrolysis, few compared electrolysis to the lower energy cost of heating coal to obtain carbon. Other students did state that extracting aluminium requires more energy than forming carbon, but did not explain why." Lesson: when explaining cost, name electrolysis AND name the cheaper alternative method for the comparison.

Environmental Note

Because aluminium production is so energy-intensive, recycling aluminium uses only ~5% of the energy needed to extract it from ore — a strong environmental and economic argument for recycling.

Keeping Products Apart

In molten electrolysis cells, the two products form on opposite electrodes — but in the same liquid. If they meet, they can react and reform the original compound.

"Most students realised the need to prevent the products from reacting together and reforming sodium chloride. However, a significant minority failed to appreciate that the reverse reaction would occur and instead focused on preventing the products from mixing or ensuring product purity." — AQA Examiner Report, P1 Jun 2022 Q06.2

Mark-Scheme Answer (Jun 2022 Q06.2)

"So the products do not react (to reform sodium chloride)." The mesh keeps Na (cathode) and Cl₂ (anode) apart so they cannot recombine. Note: the reactivity makes the reverse reaction spontaneous if products meet — so industrial cells design in physical barriers (mesh, separator).

Recall the Charge Carrier

Q06.3 (same paper) asked which particle passes through the mesh. Answer: ion. The mesh lets ions through (so the circuit completes) but blocks the gaseous/liquid products from mixing.

Common Errors → Mark-Scheme Phrases

Common errorCorrect mark-scheme phrase

"Electrolysis separates compounds""Electrolysis decomposes compounds"

"Chemical cells use electricity""Chemical cells produce electricity"

"Electrons flow through the electrolyte"" Ions move through the electrolyte"

"Oxygen is discharged""O²⁻ ions are discharged; oxygen is produced "

"Sodium is discharged""Na⁺ ions are discharged; sodium is produced "

"Chlorine ions"" Chloride ions" (Cl⁻ is the ion of the element chlorine)

"Cryolite makes it more conductive""Cryolite lowers the melting point "

Half equation: Na as reactantNa⁺ + e⁻ → Na (ion is reactant; element is product)

Pattern

The vocabulary precision matters: "discharge" applies to ions, not products. "Decompose" describes a chemical change; "separate" describes a physical method. "Chloride" names the ion; "chlorine" names the element. Get the right word and the mark follows.

Practice 1 — Foundation

Molten lead bromide (PbBr₂) is electrolysed using inert electrodes. State the product formed at each electrode and explain why each forms there.3 MARKS

MP1

Cathode product: lead (Pb) — molten lead.

MP2

Anode product: bromine (Br₂) — orange-brown vapour.

MP3

Reasoning: Pb²⁺ ions (positive) are attracted to the negative cathode where they gain electrons; Br⁻ ions (negative) are attracted to the positive anode where they lose electrons.

⚠️

Watch out: use the names "lead" and "bromine" (not "Pb" and "Br" or "PbBr"). And it's "bromide ions" (Br⁻) — not "bromine ions".

Practice 2 — Application

Electrolysis and chemical cells both involve chemical reactions and electricity. Explain the difference between the processes in electrolysis and in a chemical cell.2 MARKS

Citation: P1 Jun 2021 (Nov21) Q07.1 — verbatim AQA stem.

MP1

Electrolysis uses electricity to produce a chemical reaction (or to decompose a compound / electrolyte).

MP2

(But) cells use a chemical reaction to produce electricity.

★

Mark-scheme answer (verbatim): "electrolysis uses electricity to produce a chemical reaction (but) cells use a chemical reaction to produce electricity". Both clauses required; reversed direction loses both.

!

Allowed alternatives: "voltage" / "potential difference" / "(electrical) current" all accepted in place of "electricity".

Practice 3 — HT Half Equation HT Only

A teacher demonstrates the electrolysis of molten lead bromide. Bromine is produced at the positive electrode. Complete the half equation for the production of bromine. You should balance the half equation.2 MARKS

Br⁻ → ___ + ___Citation: P1 Jun 2021 (Nov21) Q07.2 — verbatim AQA stem.

1

Identify the product: bromine = Br₂ (a diatomic molecule — never just "Br").

2

Balance atoms: need 2 Br on the left to make 1 Br₂. So: 2Br⁻ → Br₂ + ...

3

Balance charges: left side is 2 × (−1) = −2; right side is 0 so far. Add 2e⁻ to the right: 2Br⁻ → Br₂ + 2e⁻.

★

Mark-scheme answer: 2Br⁻ → Br₂ + 2e⁻ (allow multiples; allow 1 mark for Br₂ AND e⁻ even if not balanced correctly).

⚠️

Examiner caution: ER reports many students forgot Br is diatomic (Br₂) or didn't balance the electrons. Always check both atoms AND charge balance.

Lesson Checkpoint

Foundation Outcomes

Define electrolysis: decomposition of an ionic compound using electricity (NOT separation)

State requirement: ions must be free to move — molten or aqueous

Identify cathode (–) and anode (+) and what each attracts

Recognise that ions carry charge in the electrolyte; electrons carry charge in the wires

State products of molten electrolysis (e.g. PbBr₂ → Pb at cathode, Br₂ at anode)

Distinguish chemical cell (produces electricity) from electrolytic cell (uses electricity)

Explain aluminium extraction: cryolite lowers mp; carbon anodes burn away

Higher Tier Outcomes

Write half-equations for molten electrolysis (e.g. Pb²⁺ + 2e⁻ → Pb; 2Br⁻ → Br₂ + 2e⁻)

Identify reduction (cathode) and oxidation (anode); balance both atoms and charges

Next Lesson

4.08: Electrolysis of Aqueous Solutions and Half Equations (RP3). When water is present, four ions exist in solution (compound's two ions + H⁺ and OH⁻ from water). New decision rules for which forms at each electrode — covered in detail next.

Exam-Ready Summary

The Five Essentials

Electrolysis = decomposition of an ionic compound using electricity.

Cathode (–) attracts cations → reduction; Anode (+) attracts anions → oxidation.

Ions in electrolyte; electrons in wires.

Aluminium = electrolysis of Al₂O₃ in molten cryolite (cryolite lowers mp; C anodes burn away).

Chemical cell produces electricity; electrolytic cell uses electricity.

Top Lost Marks

"Separates" instead of "decomposes" (Nov21 Q07.1).

"Electrons flow through the electrolyte" — should be "ions" (Jun22 Q06.3).

Half-equation reversed: Na as reactant (should be Na⁺) (Jun22 Q06.1).

"Cells use electricity" — should be "produce electricity" (Nov21 Q07.1).

Cost answer without comparison to carbon-reduction alternative (Jun22 Q07.2).

Exam Technique

Half equations: ion is the reactant, element is the product. Always check both atoms and charges balance. For "explain why aluminium is expensive": name electrolysis AND name the cheap alternative (carbon reduction) for comparison.