Electrolysis

Electrolysis uses electrical energy to drive a chemical reaction that would not happen on its own. An external power source pushes electrons through a liquid electrolyte, causing reduction at the cathode and oxidation at the anode. It’s the reverse of what happens in a battery: instead of a spontaneous reaction producing electricity, electricity forces a non-spontaneous reaction to occur. Industrial electrolysis is how aluminium is smelted, how chlorine and sodium hydroxide are made, and how metal objects are electroplated.

The Electrolytic Cell

An electrolytic cell has three essential parts:

• Electrolyte. A liquid (molten salt or aqueous solution) containing mobile ions.
• Cathode. The electrode connected to the negative terminal of the power supply. Cations migrate here and are reduced.
• Anode. The electrode connected to the positive terminal. Anions migrate here and are oxidized.

Note the sign convention is opposite to a galvanic cell: in electrolysis the cathode is negative and the anode is positive, because the power supply pushes electrons into the cathode and out of the anode.

Electrolysis of Water

With a small amount of acid or base added to make it conductive, water splits into hydrogen and oxygen:

Cathode (reduction): \( 2\text{H}_2\text{O} + 2e^- \rightarrow \text{H}_2 + 2\text{OH}^- \)

Anode (oxidation): \( 2\text{H}_2\text{O} \rightarrow \text{O}_2 + 4\text{H}^+ + 4e^- \)

Overall: \( 2\text{H}_2\text{O} \rightarrow 2\text{H}_2 + \text{O}_2 \)

You get twice as many moles of hydrogen as oxygen — visible as the relative volumes collected over the two electrodes. The minimum theoretical voltage is 1.23 V, though in practice 1.7-2.0 V is needed because of overpotential.

Electrolysis of Brine

Electrolysing concentrated aqueous sodium chloride (brine) is the basis of the chlor-alkali industry — among the largest industrial chemical processes in the world:

Cathode: \( 2\text{H}_2\text{O} + 2e^- \rightarrow \text{H}_2 + 2\text{OH}^- \)

Anode: \( 2\text{Cl}^- \rightarrow \text{Cl}_2 + 2e^- \)

Three products are produced simultaneously: hydrogen gas, chlorine gas, and sodium hydroxide solution. All three are essential industrial chemicals used in plastics (PVC), water treatment, paper bleaching, soap manufacture, and chemical synthesis.

Faraday’s Laws of Electrolysis

Michael Faraday’s 1834 quantitative laws relate the amount of substance deposited to the charge passed:

1. First law. The mass of substance deposited at an electrode is proportional to the total electric charge passed through the cell.
2. Second law. For the same quantity of charge, the masses of different substances deposited are proportional to their equivalent weights (molar mass divided by the number of electrons transferred).

Combined, this gives \( m = \frac{QM}{nF} \), where \( m \) is mass deposited, \( Q \) is charge (in coulombs), \( M \) is molar mass, \( n \) is the number of electrons in the half-reaction, and \( F \) is Faraday’s constant (96,485 C/mol).

Industrial Applications

• Aluminium extraction. The Hall-Héroult process electrolyses molten Al₂O₃ dissolved in cryolite. Aluminium is too reactive to extract by traditional reduction; electrolysis is the only economic route. Globally consumes roughly 3% of all electricity.
• Chlor-alkali industry. Brine electrolysis produces chlorine, sodium hydroxide, and hydrogen at enormous scale.
• Electroplating. A thin layer of metal (gold, silver, chromium, nickel) is deposited on an object by making it the cathode in a solution of the plating metal’s salt.
• Copper refining. Impure copper is the anode, pure copper plates out at the cathode, and the impurities settle as anode mud — often containing recoverable silver, gold, and platinum.
• Hydrogen production. Electrolysis of water using renewable electricity (‘green hydrogen’) is a leading candidate for decarbonising steel, ammonia, and long-haul transport.

Related study notes: Redox Reactions, Electronegativity, Ionic vs Covalent Bonds, Molarity.

Frequently Asked Questions