Electrolysis
Electrolysis is a process in which an electric current is passed through an electrolyte (an ionic compound in solution or molten state) to drive non-spontaneous chemical reactions, decomposing the electrolyte into its constituent elements or ions at the electrodes. At the cathode, cations are reduced by gaining electrons, while at the anode, anions are oxidised by losing electrons. Electrolysis has wide industrial applications including the production of chlorine and sodium hydroxide (chlor-alkali process), extraction of aluminium (Hall–Héroult process), electroplating, and water splitting to produce hydrogen fuel.
Important Industrial Electrolysis Processes
| Process | Electrolyte | Cathode Product | Anode Product | Application |
|---|---|---|---|---|
| Chlor-alkali | Brine (NaCl solution) | H₂ gas | Cl₂ gas + NaOH | Chemical industry |
| Hall–Héroult | Molten Al₂O₃ in cryolite | Al metal | O₂ / CO₂ | Aluminium production |
| Water splitting | Dilute H₂SO₄ / KOH | H₂ gas | O₂ gas | Hydrogen fuel cells |
| Electroplating | Metal salt solution | Metal deposit | Metal dissolved | Jewellery, corrosion prevention |
| Electrorefining | CuSO₄ solution | Pure Cu | Impure Cu dissolved | Copper purification |
Interactive Tools
PhET Electrolysis Simulation
Interactive simulation to observe electrode reactions and gas production during electrolysis
Open ToolKhan Academy – Electrolysis
Comprehensive lessons on electrolysis reactions, products, and industrial applications
Open ToolBrilliant – Electrolysis
Conceptual and quantitative problems on electrolysis covering Faraday's laws
Open Tool
Wikimedia Commons, CC BY-SA
Related Terms
Electrolytic Cell
An electrolytic cell is an electrochemical device that uses an external electrical energy source to drive non-spontaneous redox reactions, converting electrical energy into chemical energy. Unlike a galvanic cell, the anode is connected to the positive terminal of the power supply and the cathode to the negative terminal, and oxidation still occurs at the anode while reduction occurs at the cathode. Electrolytic cells are used industrially in processes such as electroplating, electro-refining of metals, chlor-alkali production, and the electrolysis of water.
Faraday's Laws of Electrolysis
Faraday's laws of electrolysis, formulated by Michael Faraday in 1833–1834, quantitatively describe the relationship between the amount of substance deposited or dissolved at an electrode and the quantity of electric charge passed through the electrolyte. The first law states that the mass of substance liberated is directly proportional to the charge passed; the second law states that for a given charge, the masses of different substances liberated are proportional to their equivalent masses. Together, these laws enable the precise calculation of the amount of product formed in any electrolytic process.
Cathode
The cathode is the electrode at which reduction occurs in an electrochemical cell, with species gaining electrons from the electrode. In a galvanic (voltaic) cell, the cathode is the positive terminal because reduction of cations (gain of electrons) takes place there and it has a higher electrode potential; in an electrolytic cell, the cathode is connected to the negative terminal of the power supply. The mnemonic "Red Cat" (Reduction at Cathode) helps students remember this consistent rule across all types of electrochemical cells.
From Greek "elektron" (amber, source of electricity) and "lysis" (a loosening or decomposition). The term was introduced by Michael Faraday in 1834, who also coined related terms such as "electrode," "anode," "cathode," and "ion."