This mini-lesson walks you through the whole of Cambridge IGCSE (0620) Topic 4 — Electrochemistry: electrolysis of molten and aqueous compounds, the electrode rules, the half equations, electroplating and hydrogen–oxygen fuel cells.
Supplement Some statements are extended-only (Paper 4/6). Wherever you see the Supplement flag, that content is needed only by extended candidates — core candidates can read it for interest.
Work through each screen, answer the questions as you go (predicting products, electrode reactions, half equations) and collect ⭐ stars. Press Start when you're ready.
The key words
What electrolysis means
Electrolysis is the breakdown of an ionic compound, when molten or in aqueous solution, by the passage of electricity.
Electrolyte — the molten or aqueous ionic compound that is broken down. It conducts because its ions are free to move.
Electrode — a rod (often graphite or a metal) that carries the current into and out of the electrolyte.
Anode — the positive electrode.
Cathode — the negative electrode.
The anode is +, the cathode is –. A useful memory aid: An Ox (anode oxidation) and Red Cat (reduction at cathode).
Watch out: the cathode is NEGATIVE. It attracts the positive ions (cations), which is the opposite of what many students first guess.
Quick check
Anode or cathode?
?In any electrolysis cell, which statement is correct about the cathode?
Electrolysis of molten compounds
Splitting molten lead(II) bromide
Heat an ionic solid until it is molten and its ions become free to move. The classic example is molten lead(II) bromide, PbBr₂:
Negative Br⁻ ions move to the anode; positive Pb²⁺ ions move to the cathode.
For any molten binary (two-element) compound the rule is simple:
The metal forms at the cathode — here, molten lead.
The non-metal forms at the anode — here, bromine (brown vapour).
Sort it
Predict the products
Each molten compound is electrolysed with inert electrodes. Tap the cathode product.
Supplement
How charge moves Supplement
During electrolysis charge is carried in two different ways — and it is a common error to mix them up:
In the wires and electrodes: charge is carried by a flow of electrons.
In the electrolyte: charge is carried by moving ions — cations to the cathode, anions to the anode. Electrons do NOT travel through the solution.
At the electrodes, ions are discharged by losing or gaining electrons:
At the cathode positive ions gain electrons — this is reduction.
At the anode negative ions lose electrons — this is oxidation.
An Ox · Red CatANode = OXidation | REDuction = CAThode
Watch out: electrons flow in the external circuit (the wires), never through the electrolyte. Inside the liquid it is the ions that move.
Supplement · half equations
Half equations for molten PbBr₂ Supplement
A half equation shows the electrons (e⁻) gained or lost at one electrode. Both atoms and charges must balance.
Cathode: Pb²⁺ + 2e⁻ → Pbreduction — the Pb²⁺ ion GAINS 2 electrons
Anode: 2Br⁻ → Br₂ + 2e⁻oxidation — each Br⁻ LOSES an electron; two combine to make Br₂
Balancing check
Cathode: charge on left = (+2) + (−2) = 0 = charge on right ✓
Anode: 2 Br on each side; charge left = 2×(−1) = −2, right = 0 + 2×(−1) = −2 ✓
Supplement check
Complete the half equation
?Supplement. Molten aluminium oxide is electrolysed. Which is the correct cathode half equation for the aluminium ion Al³⁺?
Electrolysis of aqueous solutions
The two electrode rules
In a water-based solution there is a competition: the water itself can give H₂ or O₂. Cambridge gives two rules:
CATHODE ruleHydrogen (H₂) is given off — UNLESS the metal is BELOW hydrogen in the reactivity series, in which case the metal is deposited.
ANODE ruleOxygen (O₂) is given off — UNLESS a CONCENTRATED halide (Cl⁻, Br⁻, I⁻) solution is used, in which case the halogen is given off.
Decide the cathode product from the reactivity series; decide the anode product from whether a concentrated halide is present.
Watch out: the anode halogen exception needs the halide to be concentrated. A dilute halide solution mostly gives oxygen instead.
Worked examples
Four solutions to know
Dilute sulfuric acid (H₂SO₄): cathode → H₂; anode → O₂ (no halide present). Overall, water is split 2:1 H₂:O₂.
Concentrated sodium chloride (brine): cathode → H₂ (Na is above H); anode → Cl₂ (concentrated halide). NaOH is left in solution.
Dilute sodium chloride: cathode → H₂; anode → O₂ (halide too dilute, so oxygen wins).
Aqueous copper(II) sulfate, inert (carbon/platinum) electrodes: cathode → copper (Cu is below H); anode → O₂ (sulfate, no halide).
With copper electrodes instead of inert ones, copper(II) sulfate behaves differently — the anode dissolves and copper transfers across:
Cathode: copper is deposited (the cathode gains mass).
Anode: the copper electrode dissolves into the solution (the anode loses mass). This is the basis of purifying copper.
Key contrast: inert electrodes give oxygen at the anode; a copper anode dissolves instead, so the blue colour of the solution stays the same.
Supplement · half equations
Aqueous half equations Supplement
The half equations you should be able to write for the aqueous examples:
4 O and 4 H on each side ✓ · charge left = 4×(−1) = −4; right = 0 + 0 + 4×(−1) = −4 ✓
Predict
Your turn — concentrated brine
?Concentrated aqueous sodium chloride (NaCl) is electrolysed with inert electrodes. What is produced at the anode?
Electroplating
Coating one metal with another
Electroplating uses electrolysis to put a thin layer of metal onto an object — to make it look better or to protect it from corrosion. To electroplate an object with copper:
The object is the cathode; the plating metal (copper) is the anode; the electrolyte contains ions of that metal (copper(II) sulfate).
Make the object the cathode — copper deposits onto it (Cu²⁺ + 2e⁻ → Cu).
Make a bar of pure copper the anode — it dissolves to top up the Cu²⁺ ions.
Use a solution of a soluble salt of the plating metal — here copper(II) sulfate.
Quick check
Why electroplate?
?A steel tap is electroplated with chromium. To plate the object, where should it be placed, and why is plating used?
Supplement
Hydrogen–oxygen fuel cells Supplement
A hydrogen–oxygen fuel cell uses hydrogen and oxygen to produce electricity, with water as the only chemical product. The overall reaction is simply the formation of water:
2H₂ + O₂ → 2H₂Ohydrogen + oxygen → water (this releases energy as electricity)
Hydrogen and oxygen flow in; the cell drives a current through the external load and produces only water.
Advantages (vs petrol/diesel engines or normal cells):
The only product is water — no carbon dioxide or other pollutants at the point of use.
More efficient than a petrol/diesel engine and runs continuously while fuel is supplied (unlike a battery that runs flat).
Disadvantages:
Hydrogen is flammable and hard and expensive to store (high-pressure tanks) and transport.
Making the hydrogen often uses energy from fossil fuels; few refuelling stations exist and the cells are costly.
Supplement check
Reading the fuel cell
?Supplement. What is the only chemical product of a hydrogen–oxygen fuel cell, and what is its main environmental advantage?
Sort it
Anode or cathode?
Tap each process, then tap the electrode where it happens.