corrosion and passivation

Left to itself, a refined metal wants to go home. We spend energy reducing ores to win shiny metals, and corrosion is the universe slowly clawing that energy back, re-oxidizing the metal toward the compound it came from. Rust on an iron gate, green tarnish on copper, the white bloom on old aluminium — all are the spontaneous reverse of metal extraction, electrons leaking out of the metal back into the environment.

Rusting is a tiny electrochemical cell set up on the metal's own surface. At anodic patches iron is oxidized, Fe -> Fe2+ + 2 e-; at cathodic patches oxygen and water consume those electrons, O2 + 2 H2O + 4 e- -> 4 OH-; the iron ions and hydroxide then react and are further oxidized by air into hydrated iron(III) oxide, the flaky brown rust. Crucially, rust is porous and flakes off, so fresh metal is always exposed and the rot continues. Passivation is the lucky opposite: some metals (aluminium, chromium, titanium, stainless steel) form an oxide that is thin, dense, and tightly bonded — an invisible, self-healing skin that seals the surface and stops further attack, which is why aluminium foil does not crumble and why stainless steel stays bright.

Understanding the cell lets us fight it. Coating (paint, galvanizing) blocks oxygen and water. Sacrificial protection wires the iron to a more reactive metal such as zinc or magnesium, which sits lower in the electrochemical series and so corrodes preferentially, feeding electrons to the iron and keeping it cathodic and safe — the basis of galvanized steel and the anodes bolted to ship hulls and pipelines. The honest nuance is that corrosion is electrochemistry plus kinetics: thermodynamics says nearly every structural metal 'should' corrode, yet a good passive film or a slow cathodic reaction can make the difference between a bridge that lasts a century and one that rusts away in a decade.