Overview
What you’ll learn
- Define electrolysis, electrolyte, and the anode/cathode, and explain why ionic compounds conduct only when molten or dissolved.
- Predict the products of electrolysis of a molten binary ionic compound and write the electrode half-equations.
- Apply selective-discharge reasoning to aqueous solutions (cation reactivity, halide concentration, OH⁻ discharge) and predict electrode products.
- Describe the electrolysis of concentrated sodium chloride solution (brine) and of dilute sulfuric acid, with half-equations and observations.
- Describe electroplating and the electrolytic refining of copper, with electrode equations.
- Describe a simple chemical cell and relate the reactivity series to the polarity of its electrodes.
In 30 seconds
- 1Electrolysis is the decomposition of an electrolyte by an electric current; the cathode (−) attracts cations, the anode (+) attracts anions — PANIC: Positive Anode, Negative Is Cathode.
- 2Solid ionic crystals do not conduct: the ions are locked in the lattice; melting or dissolving frees the ions to move and carry charge.
- 3Molten binary compounds split cleanly: metal at the cathode (Mⁿ⁺ + ne⁻ → M), non-metal at the anode (e.g. 2Br⁻ → Br₂ + 2e⁻).
- 4In aqueous electrolytes the least reactive cation discharges at the cathode: Cu²⁺ and Ag⁺ beat H⁺; if only very reactive cations are present, H₂ from water comes off instead.
- 5At the anode: concentrated halide → halogen (Cl₂, Br₂, I₂); dilute halides and sulfate/nitrate solutions give O₂ from OH⁻ (4OH⁻ → O₂ + 2H₂O + 4e⁻).
- 6Brine (concentrated NaCl) gives hydrogen at the cathode, chlorine at the anode, and sodium hydroxide left in solution — three useful products from one cell.
- 7Electrolysis of dilute sulfuric acid (or any inert electrolyte) is electrolysis of water: H₂ at the cathode and O₂ at the anode in a 2 : 1 volume ratio.
- 8In electroplating, the object is the cathode and the plating metal the anode; in a simple cell, the more reactive metal is the negative electrode and pushes electrons around the circuit.
Notes · 1 of 5
Electrolysis set-up and vocabulary
Electrolyte, electrodes, and the PANIC rule for which pole is which.
Electrolysis forces a chemical change with electricity: a battery drives current through an electrolyte — a molten ionic compound or an aqueous solution of ions — between two electrodes. Cations swim to the cathode and gain electrons (reduction); anions swim to the anode and lose electrons (oxidation). Memory hook: PANIC — Positive Anode, Negative Is Cathode.
Why the molten-or-dissolved rule? In a solid ionic crystal the ions vibrate but cannot move, so no charge flows. Melt the crystal or dissolve it and the ions float free: the liquid conducts and decomposes. Covalent substances (and pure water, almost) are non-electrolytes — they own no ions to carry the current.
| Substance | Conducts when solid? | Conducts when molten/aqueous? | Why |
|---|---|---|---|
| Copper metal | Yes | Yes | Delocalised electrons drift (not electrolysis) |
| NaCl crystal | No | Yes | Ions fixed in lattice until molten or dissolved |
| Sucrose solution | No | No | Molecules dissolve whole — no ions |
| Molten PbBr₂ | — | Yes | Mobile Pb²⁺ and Br⁻ carry charge |
Notes · 2 of 5
Electrolysis of molten compounds
No water to compete — every molten binary compound splits into its two elements.
A molten binary compound holds exactly two ions, so prediction is trivial: the metal forms at the cathode, the non-metal at the anode. Molten sodium chloride gives silvery sodium and pale green chlorine; molten aluminium oxide gives aluminium and oxygen (the basis of aluminium extraction).
| Electrode | Ion arriving | Half-equation | Observation |
|---|---|---|---|
| Cathode (−) | Pb²⁺ | Pb²⁺ + 2e⁻ → Pb | Silvery molten beads of lead |
| Anode (+) | Br⁻ | 2Br⁻ → Br₂ + 2e⁻ | Red-brown vapour of bromine |
The half-equations add to PbBr₂ → Pb + Br₂ — decomposition with the battery paying the energy bill. Notice the redox pairing from Chapter 1: reduction at the cathode, oxidation at the anode, every time.
Notes · 3 of 5
Selective discharge in aqueous solutions
Water joins the contest — reactivity and concentration decide the winners.
Dissolve a salt and you double the field: water contributes H⁺ and OH⁻ alongside the salt ions. Who discharges? At the cathode, the least reactive cation wins: Cu²⁺ and Ag⁺ beat H⁺ (copper or silver deposits), while solutions of very reactive cations (Na⁺, K⁺, Ca²⁺) simply evolve hydrogen from water.
At the anode, concentrated halides discharge as Cl₂, Br₂ or I₂ (concentration wins); but dilute halides, sulfates and nitrates let the OH⁻ from water discharge, bubbling oxygen and leaving the solution acidic. The OH⁻ half-equation: 4OH⁻ → O₂ + 2H₂O + 4e⁻.
| Electrolyte | Cathode product | Anode product | Solution change |
|---|---|---|---|
| Concentrated NaCl (brine) | H₂ (Na⁺ too reactive) | Cl₂ (concentrated halide) | NaOH remains — becomes alkaline |
| CuSO₄ (aq) | Cu deposits (Cu²⁺ beats H⁺) | O₂ (SO₄²⁻ never discharges) | Acidic, blue colour fades |
| Dilute H₂SO₄ | H₂ | O₂ | 2 : 1 volume ratio — water is split |
| Molten PbBr₂ | Pb | Br₂ | — (no water present) |
For dilute sodium chloride the anode flips: OH⁻ discharges instead, giving O₂, and the remaining solution of NaOH grows alkaline. Comparing the two is a classic four-marker — quote concentration as the deciding factor.
Notes · 4 of 5
Electroplating and refining
The object goes negative, the plating metal positive — and impure copper cleans itself.
Electroplating coats a cheap object with a precious or protective metal. Set the object as the cathode, a bar of the plating metal as the anode, and an electrolyte of the plating metal’s ions. Silver-plating a spoon: the spoon (cathode) catches silver as Ag⁺ + e⁻ → Ag, while the silver anode dissolves away as Ag → Ag⁺ + e⁻, feeding fresh ions in. The bath thins but never runs dry.
Electrolytic refining of copper uses the same trick industrially: a block of impure copper is the anode, a thin pure-copper sheet the cathode, in copper(II) sulfate solution. Copper leaves the anode as Cu²⁺ (Cu → Cu²⁺ + 2e⁻) and stacks onto the cathode (Cu²⁺ + 2e⁻ → Cu) at 99.99% purity; less reactive impurities (gold, silver) fall off as anode sludge, and more reactive ones (zinc, iron) stay dissolved.
Notes · 5 of 5
The simple chemical cell
Electrolysis in reverse: two metals, an electrolyte, and electron flow you can measure.
A simple cell converts chemical energy to electrical energy: two different metals dip in an electrolyte (dilute acid or salt solution) and a voltmeter reads a potential difference. The more reactive metal becomes the negative electrode: it oxidises, pushing electrons around the external circuit, e.g. Zn → Zn²⁺ + 2e⁻. At the positive electrode the electrons are consumed, e.g. 2H⁺ + 2e⁻ → H₂ on the copper.
Wider metal gaps give bigger voltages (Mg/Cu beats Zn/Cu); identical metals give none. The comparison table between the two directions of energy conversion is a favourite one-marker:
| Feature | Electrolysis | Simple cell |
|---|---|---|
| Energy direction | Electrical → chemical | Chemical → electrical |
| Negative electrode | Cathode (reduction) | More reactive metal (oxidation) |
| Anode behaviour | Oxidation, forced by the battery | Oxidation of the reactive metal |
| Driven by | External battery | Spontaneous redox of the metals |
Revise
Key terms
The words examiners expect you to use precisely.
- Electrolysis
- The decomposition of an electrolyte by an electric current, at two electrodes.
- Electrolyte
- A molten ionic compound or aqueous solution that conducts electricity and decomposes — its ions are mobile.
- Non-electrolyte
- A liquid or solution with no mobile ions (e.g. sucrose solution, pure water) that does not conduct.
- Electrode
- A rod (carbon/graphite or platinum) dipped into the electrolyte to pass current in or out.
- Cathode
- The negative electrode, attracting cations; reduction happens here.
- Anode
- The positive electrode, attracting anions; oxidation happens here.
Show 6 more key termsShow fewer key terms
- Selective discharge
- When several ions compete, the one discharged at an electrode depends on reactivity (cations) and concentration (anions).
- Brine
- Concentrated sodium chloride solution; electrolysed to hydrogen, chlorine and sodium hydroxide.
- Electroplating
- Coating an object (cathode) with a metal from solution, using an anode of that metal.
- Anode sludge
- Impurities (gold, silver) that drop beneath the anode during copper refining.
- Simple cell
- Two different metals in an electrolyte generating a voltage; the more reactive metal is negative.
- Inert electrode
- An electrode (carbon, platinum) that conducts but never reacts — so products come only from the electrolyte.
Revise
Common mistakes
Where students most often lose marks in this chapter — and what to do instead.
Mistake: Calling the cathode positive.
The polarity is set by the battery, and cations are negative-seeking.
Instead: PANIC: Positive Anode, Negative Is Cathode — cathode attracts cations for reduction.
Mistake: Saying solid sodium chloride conducts because it contains ions.
Ions exist but cannot move through the rigid lattice.
Instead: Conduction needs mobile charge: melt or dissolve the salt first.
Show 10 more mistakesShow fewer mistakes
Mistake: Predicting sodium metal from brine electrolysis.
Na⁺ is far too reactive to discharge from water.
Instead: H⁺ from water discharges instead: hydrogen gas at the cathode; NaOH stays in solution.
Mistake: Expecting oxygen at the anode of concentrated brine.
High halide concentration beats OH⁻ in the discharge contest.
Instead: Concentrated halide → halogen; only dilute solutions (or sulfate/nitrate) give O₂.
Mistake: Swapping the electrodes when plating.
Metal ions are positive; they must be reduced onto the negative object.
Instead: Object = cathode; plating metal = anode; electrolyte = ions of the plating metal.
Mistake: Claiming the sulfate ion discharges to make SO₂.
Sulfate (and nitrate) ions are too stable to discharge at O-Level.
Instead: OH⁻ from water discharges: oxygen bubbles form.
Mistake: Forgetting that Cu²⁺ beats H⁺ at the cathode.
The reactivity series, not concentration, rules the cathode at O-Level.
Instead: Less reactive cations (Cu²⁺, Ag⁺) discharge first; hydrogen comes last.
Mistake: Writing electrons on the wrong side of a half-equation.
Reduction gains electrons; oxidation loses them.
Instead: Cathode half-equations show e⁻ as reactants; anode half-equations as products.
Mistake: Saying electrons flow through the solution.
Ions carry the current inside the electrolyte.
Instead: Electrons move only through the wires; ions move through the electrolyte.
Mistake: Making the reactive metal positive in a simple cell.
The reactive metal pushes electrons out — that makes it negative.
Instead: More reactive metal = negative electrode = site of oxidation.
Mistake: Expecting a 1 : 1 gas ratio from dilute sulfuric acid electrolysis.
Water is H₂O — two hydrogens per oxygen.
Instead: H₂ : O₂ = 2 : 1 by volume (twice as much hydrogen).
Mistake: Using copper electrodes for brine electrolysis.
Copper is not inert — it would enter the reaction.
Instead: Use carbon (graphite) or platinum for inert electrodes.
Check yourself
Quick quiz
20 questions with instant explanations. About 5 minutes.
Question 1 of 3
- Question 1: current question
- Question 2: not answered
- Question 3: not answered
Recall electrolysis terms
During electrolysis, cations move towards the
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Practise
Topical questions
Exam-style practice in three tiers, from recall to data questions.
- 4Foundation
- 5Standard
- 3Challenging
12 questions worth 53 marks, each with a model answer, mark scheme and examiner tips.
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Practice paper
A 30-mark paper set out like a school paper, with a separate mark scheme. Give yourself 40 min.
Question paper
Chapter 2 Practice Paper
30 marks · 40 min · 4 pages
5 multiple-choice questions and structured questions with answer lines. Chapter 2 (molten and aqueous electrolysis, selective discharge, plating, refining, simple cells).
Mark scheme
PROChapter 2 Practice Paper — mark scheme
Answers · Marking points · 3 pages
Model answers with a mark for each creditworthy point — for marking at home or in class.
Where this chapter is tested
Each school sets its own scope — check your assessment outline.
Questions
Frequently asked questions
What does Sec 4 Chemistry Chapter 2 cover?
It covers sub-topic 3.4 of O-Level Pure Chemistry 6092: electrolysis of molten and aqueous ionic compounds, the selective-discharge rules (cation reactivity, halide concentration, hydroxide discharge), brine and dilute sulfuric acid electrolysis, electroplating, electrolytic refining of copper, and the simple chemical cell.
What is the fastest way to remember electrode polarity?
PANIC: Positive Anode, Negative Is Cathode. The cathode attracts cations (reduction), the anode attracts anions (oxidation) — in electrolysis the battery sets the polarity; in a simple cell the more reactive metal becomes the negative terminal.
Why do concentrated and dilute sodium chloride give different anode gases?
At the anode, discharge is a concentration contest: a high chloride concentration out-competes hydroxide, giving chlorine; in dilute solution the OH⁻ ions from water win, giving oxygen. Cathode products, in contrast, are decided by cation reactivity.
Why does copper(II) sulfate solution lose its blue colour during electrolysis?
The blue colour belongs to Cu²⁺ ions. They are less reactive than hydrogen, so they discharge at the cathode as copper metal (Cu²⁺ + 2e⁻ → Cu); as the blue ions are removed, the colour fades and oxygen bubbles at the anode.
When is this chapter usually tested in school?
Electrolysis lands early in Sec 4 and reappears in every weighted assessment, the prelim, and both O-Level chemistry papers — often fused with redox and qualitative analysis. Scope varies by school, so check your school’s assessment outline.
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Chapter 3: Chemical Energetics