4.1Electrochemistry

Electrolysis — Cambridge IGCSE Chemistry 0620 Revision Notes & Topical Questions

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📚 Cambridge IGCSE Chemistry 0620 Hub / Topical Revision

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4.1 Electrolysis — Notes — IGCSE Chemistry notes: Explain electrolysis of molten and aqueous ionic compounds, predict the products, write half-equations and describe its industrial applications.Electrolysis: the decomposition of a molten or aqueous ionic compound by an electric c…

Electrolyte
Inert electrode
Electroplating
Electrolysis: the decomposition of a molten or aqueous ionic compound by an electric current
Covalent compounds cannot undergo electrolysis because they have no ions
Active electrodes (for example copper) take part in the reaction

Section 1 of 12Principles of electrolysis

Electrolysis: the decomposition of a molten or aqueous ionic compound by an electric current.

Requirements: a d.c. supply, two electrodes and an electrolyte.

The ions must be free to move to carry charge, so the compound must be molten or dissolved.

Solid ionic compounds do not conduct because their ions are fixed in place.
Covalent compounds cannot undergo electrolysis because they have no ions.

Electrolyte: the molten or aqueous ionic compound that conducts.

Electrodes are rods or plates that carry current into and out of the electrolyte.

Inert electrodes (graphite or platinum) only carry current and do not react.
Active electrodes (for example copper) take part in the reaction.

Section 2 of 12Terms and charge transfer

Term	Meaning
Anode	Positive electrode; anions move to it and lose electrons (oxidation)
Cathode	Negative electrode; cations move to it and gain electrons (reduction)

PANIC: Positive Anode, Negative Is Cathode.

AN OX and RED CAT: oxidation occurs at the anode and reduction at the cathode.

OIL RIG: Oxidation Is Loss of electrons and Reduction Is Gain of electrons.

External circuit: electrons carry the charge, flowing from the negative terminal of the supply to the cathode and from the anode back to the positive terminal.

Electrons do not flow through the electrolyte, where ions carry the charge, with cations moving to the cathode and anions to the anode.

At the electrodes cations gain electrons to become atoms and anions lose electrons.

Charge carriers are electrons in metals and wires, and ions in electrolytes.

A half-equation shows the electron gain or loss at one electrode.

Section 3 of 12Molten binary compounds

A molten binary ionic compound gives its two elements, the metal at the cathode and the non-metal at the anode.

Molten lead(II) bromide is heated until molten with graphite electrodes and a d.c. supply.

Anode2Br⁻ → Br₂ + 2e⁻, and red-brown bromine vapour is given off.
CathodePb²⁺ + 2e⁻ → Pb, and grey lead is deposited.

Molten potassium chloride gives potassium at the cathode and chlorine at the anode.

Molten aluminium oxide: Al³⁺ + 3e⁻ → Al (reduction) and 2O²⁻ → O₂ + 4e⁻ (oxidation).

Section 4 of 12Aqueous solutions: predicting products

Water supplies H⁺ and OH⁻ ions, which compete with the ions of the dissolved compound, so one ion is discharged at each electrode.

Electrode	Rule
Cathode	The metal forms if it is less reactive than hydrogen (Cu, Ag), otherwise hydrogen gas forms
Anode	A halogen forms if a concentrated halide is present, otherwise oxygen forms (from OH⁻ or water)

Concentration matters: concentrated chloride gives chlorine, but a dilute halide gives oxygen from hydroxide ions.

Concentrated BaCl₂ gives chlorine and dilute BaCl₂ gives oxygen, with hydrogen at the cathode in both cases because barium is more reactive than hydrogen.

Method

List all the ions present, including H⁺ and OH⁻ from water.
At the cathode, decide whether the metal is less reactive than hydrogen.
At the anode, decide whether a concentrated halide is present.
Write the half-equation for each electrode.
Identify the ions left in solution, for example Na⁺ and OH⁻ giving NaOH.

Reactive metals stay as ions and hydrogen forms instead, so aqueous electrolysis does not always give the metal.

Section 5 of 12Products and half-equations

Electrolyte (inert electrodes)	Cathode (−)	Anode (+)	Cathode half-equation	Anode half-equation
Molten PbBr₂	Lead	Bromine	Pb²⁺ + 2e⁻ → Pb	2Br⁻ → Br₂ + 2e⁻
Concentrated NaCl(aq)	Hydrogen	Chlorine	2H⁺ + 2e⁻ → H₂	2Cl⁻ → Cl₂ + 2e⁻
Dilute NaCl(aq)	Hydrogen	Oxygen	2H⁺ + 2e⁻ → H₂	4OH⁻ → O₂ + 2H₂O + 4e⁻
Dilute H₂SO₄	Hydrogen	Oxygen	2H⁺ + 2e⁻ → H₂	4OH⁻ → O₂ + 2H₂O + 4e⁻
CuSO₄(aq), graphite	Copper	Oxygen	Cu²⁺ + 2e⁻ → Cu	4OH⁻ → O₂ + 2H₂O + 4e⁻
CuSO₄(aq), copper	Copper deposited	Copper anode dissolves	Cu²⁺ + 2e⁻ → Cu	Cu → Cu²⁺ + 2e⁻

Products and half-equations

Half-equations must balance charge as well as atoms, and the charge on the metal ion gives the electrons gained.

Gas tests

Hydrogen gives a squeaky pop with a lighted splint.
Oxygen relights a glowing splint.
Chlorine turns damp litmus paper red and then bleaches it white.

Section 6 of 12Aqueous sodium chloride and dilute sulfuric acid

Concentrated brine contains Na⁺, Cl⁻, H⁺ and OH⁻ ions.

H⁺ is discharged at the cathode as it is less reactive than Na⁺, and Cl⁻ is discharged at the anode.
Na⁺ and OH⁻ remain, forming sodium hydroxide solution.

Electrolysis of brine is done on a very large scale,

and in the industrial cell a membrane divides the cell so that chlorine and sodium hydroxide cannot react together.

Uses of the products

Chlorinewater treatment, bleach, PVC and solvents.
Hydrogenmaking ammonia and margarine, and as a fuel.
Sodium hydroxidesoap and detergents, paper, bleach and aluminium purification.

Dilute sulfuric acid: effectively water is decomposed, giving hydrogen and oxygen in a 2 : 1 volume ratio.

A Hoffman voltameter shows that twice as much hydrogen forms as oxygen.

Section 7 of 12Copper(II) sulfate solution and refining copper

Copper(II) sulfate solution and refining copper

Graphite electrodes: the cathode is coated with pink-brown copper and oxygen bubbles form at the anode.

The blue colour fades as Cu²⁺ is removed, and the solution becomes acidic.
Copper is less reactive than hydrogen, so Cu²⁺ is discharged at the cathode.
OH⁻ loses electrons more readily than SO₄²⁻, so oxygen forms at the anode.

Copper electrodes: the blue colour stays constant, the cathode gains mass and the anode loses the same mass, so the Cu²⁺ concentration stays constant.

Refining copper

The anode is impure copper, which shrinks as copper atoms lose electrons: Cu → Cu²⁺ + 2e⁻.
The cathode is a thin sheet of pure copper, which grows as copper ions gain electrons: Cu²⁺ + 2e⁻ → Cu.
The electrolyte is copper(II) sulfate solution.
Insoluble impurities fall to the bottom of the cell as anode sludge.
Copper for electrical wiring must be very pure because impurities greatly increase resistance.

Section 8 of 12Electroplating

Electroplating: coating the surface of one metal with a thin layer of a different metal.

Cell set-up: the object to be plated is the cathode, the anode is a bar of the plating metal,

and the electrolyte is a solution of a soluble salt of that metal.

Purposes: to protect against corrosion or damage, and to improve appearance.

Chromium and nickel plate protect metals, and zinc plating (galvanising) protects iron and steel.
Silver plating improves the appearance of cutlery and jewellery, and tin plates food cans.

Method

Clean the object so that the coating sticks.
Connect the object to the negative terminal and place it in the electrolyte.
Connect a bar of the plating metal to the positive terminal.
Pass a current so that metal ions gain electrons and deposit on the object.

At the anode the plating metal atoms lose electrons (oxidation) and dissolve to replace the ions used up,

so the concentration of the solution stays constant.

Silver plating a spoon: cathode Ag⁺ + e⁻ → Ag and anode Ag → Ag⁺ + e⁻.

Tin plating iron uses tin(II) chloride solution: anode Sn → Sn²⁺ + 2e⁻ and cathode Sn²⁺ + 2e⁻ → Sn.

Section 9 of 12Extraction of aluminium

Aluminium is more reactive than carbon, so it cannot be extracted by heating with carbon and electrolysis is used.

Electrolyte: purified aluminium oxide (alumina, from bauxite) dissolved in molten cryolite.

Aluminium oxide melts above 2000 °C, so cryolite lowers the operating temperature to about 950 °C and saves energy.

Cathode: molten aluminium forms at the graphite-lined cathode, Al³⁺ + 3e⁻ → Al, and is tapped off.

Anode: oxygen forms at the graphite anodes, 2O²⁻ → O₂ + 4e⁻.

The hot oxygen burns the carbon anodes to carbon dioxide, so the anodes must be replaced regularly.

The process uses huge amounts of electricity, so plants are built near cheap power such as hydroelectric stations.

Recycling aluminium needs only about 5% of the energy used to extract it.

Section 10 of 12More from the specification

Reactivity series
- A list of metals in order of how readily they react.

Section 11 of 12Key terms

Term	Meaning
Electrolysis	the decomposition of a molten or aqueous ionic compound by an electric current.
Electrolyte	the molten or dissolved ionic compound that conducts.
Anode	the positive electrode, where oxidation occurs.
Cathode	the negative electrode, where reduction occurs.
Inert electrode	an electrode, such as graphite or platinum, that does not react.
Half-equation	an equation showing electron gain or loss at one electrode.
Electroplating	coating one metal with a thin layer of another using electrolysis.
Cryolite	a molten solvent that lowers the temperature needed to electrolyse aluminium oxide.

Section 12 of 12Exam technique

Saying the anode is negativethe anode is the positive electrode.

Saying electrons flow through the electrolyteions move through the electrolyte, and electrons flow only in the external circuit.

Saying sodium forms at the cathode in aqueous NaCl: hydrogen forms because sodium is more reactive than hydrogen.

Saying solid NaCl conducts when electrolysedit must be molten or aqueous.

Using the impure metal as the cathode in refining: the impure metal is the anode and the pure metal is the cathode.

Last updated: 14 October 2026