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.
Galvanic Cell vs Electrolytic Cell Comparison
| Property | Galvanic Cell | Electrolytic Cell |
|---|---|---|
| Energy conversion | Chemical → Electrical | Electrical → Chemical |
| Reaction type | Spontaneous (ΔG < 0) | Non-spontaneous (ΔG > 0) |
| External power | Not required | Required |
| Anode polarity | Negative | Positive |
| Cathode polarity | Positive | Negative |
| Common use | Batteries | Electroplating, electrolysis |
Interactive Tools
PhET Electrolysis Simulation
Simulate electrolysis of water and molten salts with adjustable voltage
Open ToolKhan Academy – Electrolytic Cells
Detailed explanation of electrolytic cell operation and industrial applications
Open ToolBrilliant – Electrochemistry
Interactive problems and explanations on electrolysis and electrolytic cells
Open Tool
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Related Terms
Galvanic Cell
A galvanic cell (also called a voltaic cell) is an electrochemical device that converts chemical energy into electrical energy through spontaneous redox reactions occurring at two electrodes separated by an electrolyte. The oxidation half-reaction occurs at the anode (negative terminal) and the reduction half-reaction occurs at the cathode (positive terminal), with electrons flowing through an external circuit. Galvanic cells are the basis of all batteries and are fundamental to understanding energy storage and conversion in chemistry.
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.
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.
From Greek "elektron" (amber, source of static electricity) and "lysis" (loosening or decomposition). The prefix "electro-" has been used in chemistry since the early 19th century following Volta's and Faraday's foundational work.