TNPSC Daily Notes: Electrochemistry — Electrolysis, Faraday’s Laws, Electroplating, Galvanic Cells & Corrosion | மின்வேதியியல் — மின்னாற்பகுப்பு, பாரடே விதிகள், மின்முலாம், மின்கலன்கள் — 12 October 2026

1. Conductors, Electrolytes and Ionisation

All substances that conduct electricity fall into two families. Metallic conductors such as copper, silver and aluminium carry current through a drift of free electrons; the conductor itself is chemically unchanged, and its resistance rises with temperature. Electrolytic conductors such as molten salts and aqueous acids, bases and salts carry current through the migration of ions; the electrolyte is chemically decomposed, and its resistance falls with temperature because ionisation increases. Substances that ionise almost completely in water (NaCl, HCl, H₂SO₄, NaOH) are strong electrolytes; those that ionise only partially (acetic acid, ammonium hydroxide) are weak electrolytes; and covalent molecular solids such as sugar, glucose, urea and benzene do not ionise at all and are non-electrolytes. A useful exam point: solid sodium chloride does not conduct, because its ions are locked in the lattice — it conducts only when molten or dissolved, when the ions become free to move.

2. The Electrolytic Cell and the Direction of Change

An electrolytic cell has two electrodes dipped in an electrolyte and connected to a source of direct current. The electrode joined to the positive terminal of the battery is the anode; the one joined to the negative terminal is the cathode. Oxidation (loss of electrons) always happens at the anode and reduction (gain of electrons) always happens at the cathode — this is true in every cell, galvanic or electrolytic, and is worth memorising as An Ox, Red Cat. What flips between the two kinds of cell is only the sign: the anode is positive in an electrolytic cell but negative in a galvanic cell. Because negative anions are attracted to the positive anode, they are discharged there; positive cations travel to the cathode and are discharged there. When acidified water is electrolysed, hydrogen collects at the cathode and oxygen at the anode in a 2:1 volume ratio, exactly matching the formula H₂O.

3. Faraday’s Two Laws — the Quantitative Core

Michael Faraday, who also gave us the words ion, anode, cathode and electrolyte, reduced electrolysis to two simple laws. The first law states that the mass liberated at an electrode is directly proportional to the charge passed: m = Z I t, where I is the current in amperes, t the time in seconds, and Z the electrochemical equivalent — the mass deposited by one coulomb. The second law states that when the same charge is sent through several electrolytes in series, the masses deposited are in the ratio of their chemical equivalent weights. The bridge between the two laws is the faraday: one faraday, equal to 96,485 coulombs (taken as 96,500 C in examinations), is the charge on one mole of electrons, and it deposits precisely one gram-equivalent of any substance. So 1 F liberates 1.008 g of hydrogen, 8 g of oxygen, 108 g of silver, but only 31.75 g of copper from Cu²⁺ — because copper needs two electrons per atom. A standard numerical runs: a current of 9.65 A for 1000 s passes 9650 C, that is 0.1 F, and therefore deposits 10.8 g of silver.

4. Industrial Applications

Electroplating puts a thin protective or decorative film of one metal on another: the article is made the cathode, the plating metal the anode, and a soluble salt of that metal the electrolyte. Chromium plating on taps, nickel on cycle parts, zinc on buckets and silver on cutlery all follow this scheme. For silver and gold the electrolyte is a complex cyanide salt rather than the simple nitrate, because slow ionisation gives a fine, even, well-adhering coat. Electrolytic refining purifies copper to about 99.95%, with silver, gold and platinum collecting as anode mud. Electrometallurgy wins the highly reactive metals that carbon cannot reduce: sodium from fused NaCl (Down’s process), and aluminium by the Hall–Héroult process, in which alumina is dissolved in molten cryolite at roughly 950 °C and reduced at a carbon-lined cathode while the graphite anode is slowly consumed as carbon dioxide. The chlor-alkali industry electrolyses concentrated brine to produce chlorine, hydrogen and caustic soda simultaneously — the raw materials of bleaching powder, PVC, soap and water treatment. Anodising thickens the natural oxide film on aluminium electrolytically, which is why aluminium vessels resist corrosion so well.

5. Cells, Batteries and Corrosion

A galvanic cell runs the process backwards, turning spontaneous chemistry into current. The Daniell cell gives about 1.1 V; the familiar dry cell gives 1.5 V from a zinc can, a graphite rod packed with manganese dioxide, and a paste of ammonium chloride and zinc chloride — it is a primary cell and cannot be recharged. The lead-acid accumulator is a secondary cell: lead anode, lead dioxide cathode, about 35% sulphuric acid, 2 V per cell, so a 12 V car battery contains six cells in series. Lithium-ion cells (about 3.7 V) now dominate phones and electric vehicles, and fuel cells convert hydrogen and oxygen directly into electricity with water as the only product. Corrosion is the unwanted mirror image of all this: iron rusts to hydrated ferric oxide only when oxygen and moisture are both present, with one patch of the metal acting as an anode and another as a cathode. Barrier methods (paint, grease, tin, anodising) and alloying (stainless steel) help, but the most elegant answer is sacrificial protection — attaching a more reactive metal such as zinc or magnesium, which corrodes in place of the iron. This is why galvanised iron survives a scratch while tin-plated iron rusts faster at one.