What you'll learn
Electrolysis uses electricity to break down ionic compounds into elements, and is used industrially to extract reactive metals such as aluminium and to produce chlorine and other chemicals. In this guide you will learn what electrolysis is, what happens at each electrode, how to predict the products of electrolysing molten compounds and aqueous solutions, why aluminium is extracted by electrolysis, and how to write half equations. These ideas link bonding, the reactivity series and redox reactions.
Key terms and definitions
Electrolysis — the breaking down of an ionic compound (molten or in solution) using electricity.
Electrolyte — the molten or dissolved ionic substance that conducts electricity during electrolysis.
Electrode — a conducting rod (often inert, like graphite or platinum) dipped into the electrolyte.
Cathode — the negative electrode, where positive ions (cations) gain electrons (reduction).
Anode — the positive electrode, where negative ions (anions) lose electrons (oxidation).
Half equation — an equation showing the gain or loss of electrons at one electrode.
Core concepts
What happens during electrolysis
An ionic compound conducts electricity only when molten or dissolved, because the ions are then free to move. When a current passes, positive ions move to the cathode (negative electrode) and negative ions move to the anode (positive electrode). At the cathode, positive ions gain electrons (reduction); at the anode, negative ions lose electrons (oxidation). This decomposes the compound into elements.
Electrolysis of molten compounds
For a molten ionic compound, the products are simply the two elements. The metal forms at the cathode and the non-metal at the anode. For example, molten lead bromide gives lead at the cathode and bromine at the anode.
Electrolysis of aqueous solutions
In solution, water also provides H⁺ and OH⁻ ions, so the rules are more involved:
- At the cathode: if the metal is more reactive than hydrogen, hydrogen gas is produced; if the metal is less reactive than hydrogen (e.g. copper, silver), the metal is deposited.
- At the anode: if a halide (Cl⁻, Br⁻, I⁻) is present, the halogen is produced; otherwise oxygen is produced (from OH⁻).
For example, electrolysing copper sulfate solution with inert electrodes gives copper at the cathode and oxygen at the anode.
Extraction of aluminium
Aluminium is more reactive than carbon, so it cannot be extracted by reduction with carbon — it is extracted by electrolysis of molten aluminium oxide. The aluminium oxide is mixed with cryolite to lower the melting point and save energy. Aluminium forms at the cathode; oxygen forms at the anode, where it reacts with the hot carbon electrodes to form carbon dioxide, so the carbon anodes must be replaced regularly. Electrolysis is expensive because of the large amount of electrical energy needed.
Half equations
A half equation shows electron transfer at one electrode. At the cathode, ions gain electrons: Al³⁺ + 3e⁻ → Al or 2H⁺ + 2e⁻ → H₂. At the anode, ions lose electrons: 2Cl⁻ → Cl₂ + 2e⁻ or 2O²⁻ → O₂ + 4e⁻. Balance the charges by including the right number of electrons.
Worked examples
Example 1: Molten compound
What forms at each electrode when molten potassium chloride is electrolysed?
Potassium (metal) at the cathode; chlorine (non-metal) at the anode.
Example 2: Aqueous solution
Predict the products of electrolysing sodium chloride solution.
At the cathode: sodium is more reactive than hydrogen, so hydrogen is produced. At the anode: a halide is present, so chlorine is produced.
Example 3: Half equation
Write the cathode half equation for the extraction of aluminium.
Al³⁺ + 3e⁻ → Al. Each aluminium ion gains three electrons (reduction) to form an aluminium atom.
Common mistakes and how to avoid them
Mixing up the electrodes. Cathode = negative (attracts positive ions, reduction); anode = positive (attracts negative ions, oxidation).
Forgetting the water rules in solution. Check whether the metal is more or less reactive than hydrogen, and whether a halide is present.
Saying aluminium can be extracted with carbon. It is too reactive, so electrolysis is required.
Unbalanced half equations. Make sure electrons balance the charges on each side.
Forgetting why electrolysis is expensive. It needs large amounts of electrical energy (and, for aluminium, melting the oxide).
Exam technique for Electrolysis
State that ions must be free to move (molten or dissolved) for conduction.
Use cathode/anode rules confidently, including the aqueous solution rules.
Explain aluminium extraction with cryolite, electron cost and anode replacement.
Write balanced half equations, showing electrons gained or lost.
Identify oxidation and reduction at each electrode (OIL RIG: Oxidation Is Loss, Reduction Is Gain of electrons).
Quick revision summary
Electrolysis breaks down an ionic compound using electricity, and works only when the compound is molten or dissolved so ions are free to move. Positive ions go to the cathode (negative electrode) and gain electrons (reduction); negative ions go to the anode (positive electrode) and lose electrons (oxidation). For molten compounds, the metal forms at the cathode and the non-metal at the anode. For aqueous solutions, water's H⁺ and OH⁻ matter: at the cathode you get hydrogen unless the metal is less reactive than hydrogen (then the metal); at the anode you get the halogen if a halide is present, otherwise oxygen. Aluminium is extracted by electrolysing molten aluminium oxide (mixed with cryolite to lower the melting point); it cannot be reduced by carbon because it is too reactive, and the carbon anodes burn away and must be replaced. Write half equations with balanced electrons (Al³⁺ + 3e⁻ → Al; 2Cl⁻ → Cl₂ + 2e⁻), and remember electrolysis is costly because of the electrical energy required.