electrolytic cell
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Picture plugging a charger into a tank of liquid to force a reaction that nature refuses to do on its own. That tank is an electrolytic cell: a setup where an outside power source drives an electric current through an electrolyte to make a non-spontaneous chemical change happen — splitting water, refining a metal, or plating one metal onto another.
It has the same parts as any electrochemical cell — two electrodes in an electrolyte — but here an external battery or power supply pushes the electrons. The supply pulls electrons off the anode (forcing oxidation there) and shoves them onto the cathode (forcing reduction there). Because the reaction is being driven uphill, electrical energy is consumed and stored as chemical change in the products.
Electrolytic cells are the backbone of electrochemical industry: aluminium and chlorine production, copper refining, electroplating, and the charging side of every rechargeable battery. The key contrast to remember is with the galvanic cell. There a spontaneous reaction makes electricity; here electricity is spent to drive a reaction. The hardware can be identical — only the direction of energy flow differs.
In an aluminium smelter, a huge electrolytic cell passes thousands of amperes through molten aluminium oxide. The current rips the metal free from the oxygen — a reaction that, left alone, would never run.
An aluminium cell: brute electrical force tears metal out of its ore.
The signs flip between cell types. In an electrolytic cell the anode is the positive terminal and the cathode negative — the opposite of a galvanic cell — yet oxidation still happens at the anode in both.