corrosion
/ kuh-ROH-zhun /
A bright new nail left in the rain turns flaky and brown; a shiny car fender bubbles with rust along its edges. Corrosion is the slow, unwanted eating-away of a metal as it reacts with its surroundings — most often a metal returning, atom by atom, to the dull oxide it came from in the ore.
Beneath the surface it is electrochemistry happening without any wires. Tiny patches of the metal act as anodes, where metal atoms lose electrons and dissolve; other patches act as cathodes, where oxygen and water take those electrons up. A thin film of moisture is the electrolyte, and the metal itself carries the electrons between the patches. The metal is, in effect, running a galvanic cell against itself and slowly destroying itself in the process.
Corrosion matters because it quietly costs the world enormous sums — eaten bridges, leaking pipes, failing machinery. Understanding it electrochemically points straight to the cures: keep out moisture and oxygen with paint or grease, or attach a more reactive "sacrificial" metal (like zinc on a ship's hull) that corrodes first and spares the metal you want to protect.
Zinc blocks bolted to a steel ship's hull corrode away while the steel stays sound. The zinc, being more reactive, becomes the anode and sacrifices itself — a deliberate use of corrosion to prevent corrosion.
A sacrificial zinc block corrodes first so the steel hull does not.
Not all corrosion eats deep. On aluminium the oxide forms a tight, invisible skin that seals the metal and stops further attack; on iron the rust is loose and flaky, so it keeps falling away and exposing fresh metal.