Electrolyzing

Electrolyzing, or electrolysis, is the decomposition of a chemical compound by passing an electric current through an electrolyte. It occurs in an electrochemical cell in which an external power source drives non-spontaneous redox reactions. An electrode pair is placed in a conductive electrolyte; the electrode connected to the positive terminal becomes the anode and the negative terminal the cathode. In an electrolytic cell, oxidation occurs at the anode and reduction at the cathode, with the external circuit enforcing the charge flow.

Ions migrate to the electrodes: cations move toward the cathode where they gain electrons, while anions move toward the anode where they lose electrons (or water is oxidized/reduced, depending on the solution and potentials). In aqueous solutions, the actual half-reactions depend on the relative overpotentials and concentrations of species present.

Common applications include electroplating and metal refining, where a metal is deposited on an object or impurities are removed from a metal. Industrially important processes include the chlor-alkali production of chlorine and caustic soda, and the Hall–Héroult process for aluminum extraction. Water electrolysis is a foundational method for producing hydrogen gas and oxygen gas, often used for fuel applications or chemical synthesis.

Faraday’s laws describe how the amount of material produced at an electrode is proportional to the total electrical charge passed. The deposited or dissolved mass m relates to the charge Q = I t and the substance’s molar mass M and valence n by m = (Q M)/(n F), where F is Faraday’s constant.

Safety and efficiency considerations include controlling overpotentials, electrode materials, electrolyte stability, gas handling, and energy consumption. Electrolysis remains a cornerstone technique in chemistry and industry for driving otherwise non-spontaneous chemical changes with electricity.