Electrolysis
What Is Electrolysis?
Electrolysis uses electricity to decompose an ionic compound into its elements. The compound being broken down is called the electrolyte. For electrolysis to work, the ions must be free to move — so the electrolyte must either be molten or dissolved in water (aqueous).
The electrolyte is placed between two electrodes connected to a direct current supply. The cathode is the negative electrode; the anode is the positive electrode.
Cathode = negative electrode → attracts positive ions (cations) → cations gain electrons here (reduction). Anode = positive electrode → attracts negative ions (anions) → anions lose electrons here (oxidation). Memory aid: PANIC — Positive Anode, Negative Is Cathode.
Electrolysis of Molten Ionic Compounds
When a pure ionic compound is melted, its ions become free to move. Passing a current through molten lead bromide (\(PbBr_2\)), for example, produces lead at the cathode and bromine at the anode. In general:
- At the cathode: metal cations are attracted, gain electrons and are deposited as atoms of the metal.
- At the anode: non-metal anions are attracted, lose electrons and are released as atoms or molecules of the non-metal.
A half-equation shows what happens at one electrode. Electrons are included, represented as \(e^-\).
For molten lead bromide:
Cathode: \(Pb^{2+} + 2e^- \rightarrow Pb\) (reduction — gain of electrons)
Anode: \(2Br^- \rightarrow Br_2 + 2e^-\) (oxidation — loss of electrons)
Key rule: the number of electrons in each half-equation must balance when you add the two half-equations together. Scale one equation by a multiplier if the charges do not match.
Electrolysis of Aqueous Solutions
When an ionic compound is dissolved in water, there are additional \(H^+\) and \(OH^-\) ions from the water itself. The electrode product therefore depends on which ions are preferentially discharged:
At the cathode: if the metal is more reactive than hydrogen, hydrogen gas is produced instead of the metal (water's \(H^+\) ions are discharged). If the metal is less reactive than hydrogen (e.g. copper), the metal is deposited.
At the anode: if a halide ion (\(Cl^-\), \(Br^-\)) is present at significant concentration, the halogen is produced (e.g. chlorine gas from sodium chloride solution). If no halide is present, oxygen is produced from the discharge of hydroxide ions (\(OH^-\)).
Question: predict the products at each electrode when copper sulfate (\(CuSO_4\)) solution is electrolysed with inert carbon electrodes.
Step 1: Ions present — \(Cu^{2+}\), \(SO_4^{2-}\), \(H^+\) (from water), \(OH^-\) (from water).
Step 2: At cathode — \(Cu^{2+}\) or \(H^+\). Copper is less reactive than hydrogen, so \(Cu^{2+}\) is preferentially discharged. Cathode product: copper metal.
Step 3: At anode — \(SO_4^{2-}\) or \(OH^-\). No halide present, so \(OH^-\) is discharged. Anode product: oxygen gas.
Observations: a reddish-brown coating of copper on the cathode; colourless gas (oxygen) at the anode; the blue colour of the solution fades as \(Cu^{2+}\) ions are removed.
Extraction of Aluminium by Electrolysis
Aluminium is the most abundant metal in the Earth's crust, yet it is expensive to produce because it sits above carbon in the reactivity series. It cannot be extracted by reduction with carbon, so electrolysis of molten aluminium oxide (\(Al_2O_3\)) is used instead.
Aluminium oxide has a very high melting point, so dissolving it in molten cryolite (sodium aluminium fluoride) lowers the operating temperature dramatically, reducing energy costs. The electrolysis is carried out in large steel cells lined with graphite, which acts as the cathode. Carbon anodes dip into the molten mixture from above.
Aluminium extraction: Al₂O₃ dissolved in molten cryolite. Cathode: Al³⁺ + 3e⁻ → Al (aluminium deposited). Anode: 2O²⁻ → O₂ + 4e⁻ (oxygen released). Carbon anodes burn away because oxygen reacts with them at high temperature — they must be replaced regularly, adding to costs.
Aluminium is above carbon in the reactivity series, which means carbon is not reactive enough to reduce aluminium oxide. Even at very high temperatures, the reaction \(2Al_2O_3 + 3C \rightarrow 4Al + 3CO_2\) does not proceed as written, because the stronger bond in aluminium oxide is not broken by carbon.
Additionally, there is a common misconception that aluminium 'rusts'. In fact, aluminium reacts quickly with oxygen in the air to form a thin, tightly bonded aluminium oxide layer. Unlike iron rust, this layer is impermeable and protective — it prevents any further oxidation. This self-sealing property is why aluminium appears unreactive in everyday life despite being quite high in the reactivity series.