Corrosion and Rusting
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Direct answer
Rust is an electrochemical cell running on a wet iron surface. Tiny anodic patches dissolve — Fe → Fe^2+ + 2e^-, with E° = −0.44 V — while nearby cathodic patches reduce dissolved oxygen — O2 + 4H^+ + 4e^- → 2H2O, E° = +1.23 V — giving an overall cell potential near 1.67 V. The Fe^2+ then oxidises further and precipitates as hydrated ferric oxide, Fe2O3·xH2O, the familiar rust. Both water and oxygen must be present; perfectly dry air or fully deoxygenated water barely rusts iron at all, and salt speeds everything by making the moisture film a better electrolyte.
What you must remember
- Half-reactions with potentials: anode Fe → Fe^2+ + 2e^- (−0.44 V); cathode O2 + 4H^+ + 4e^- → 2H2O (+1.23 V); overall E°cell ≈ 1.67 V.
- The product: rust is Fe2O3·xH2O, not FeO or Fe3O4 — a one-mark distinction examiners reuse every few years.
- Differential aeration: oxygen-starved regions turn anodic and corrode — rust piling at the waterline while the submerged metal pits, pipelines failing under poorly aerated soil lumps; scratches in paint start crevices.
- Accelerators: electrolytes (saline water — why ships corrode faster than bridge girders), acid rain, impurities in the metal, stress.
- Sacrificial protection: blocks of magnesium or zinc bolted to hulls and pipelines oxidise preferentially (more negative E°), sparing the iron; consumed blocks are simply replaced.
- Galvanisation: zinc coating protects physically and, once scratched, electrochemically — zinc becomes the anode.
- Contrast with tin plating: a scratched tin can corrodes faster than bare iron because iron becomes the anode against tin; scratched galvanised iron survives.
How the sacrificial maths runs
A zinc anode of 6.54 kg welded to a ship's steel plate is consumed over its service life. Moles of Zn = 6,540/65.4 = 100 mol, each releasing 2 electrons: 200 mol of electrons. Iron dissolving as Fe^2+ needs 2 electrons per atom, so those electrons protect 100 mol of iron = 100 × 55.85 ≈ 5.6 kg of steel that would otherwise have corroded. Per kilogram of zinc, roughly 0.85 kg of iron is spared — a ratio that turns "sacrificial protection" from a phrase into arithmetic, exactly the kind of quantitative twist JEE Advanced enjoys.
The mechanism picture to hold: electrons flow through the metal from anodic to cathodic patches while ions migrate through the moisture film — a galvanic cell short-circuited by its own body. Anything that breaks either circuit (blocking oxygen, drying the surface, applying an external opposing current — impressed current cathodic protection) stops the rust.
The examiner's angle
Three formats recur. First, identify the anode and cathode reactions from a labelled diagram of a droplet on iron — remember the rust ring forms at the edge where oxygen is plentiful, while pitting happens at the oxygen-starved centre, a direct application of differential aeration. Second, choose the sacrificial metal: magnesium and zinc qualify (E° more negative than iron), tin and copper do not — copper in contact with iron accelerates corrosion the same way a scratched tin can does. Third, explain-coated-item questions: why galvanised iron resists rust even after scratches (zinc still anodic), why stainless steel resists (chromium forms a self-healing Cr2O3 passivating film), why aluminium cookware survives despite reactive metal (tightly adherent oxide layer). JEE Main keeps these single-correct; Advanced frames them as assertion-reason with the electrochemical series as the referee.
Frequently asked questions
What are the half-reactions in the electrochemical theory of rusting?
Anodic: Fe → Fe^2+ + 2e^-; cathodic: O2 + 4H^+ + 4e^- → 2H2O; the ferrous ions later oxidise and hydrate to Fe2O3·xH2O.
Why does iron rust more under a waterline droplet's centre than at its edge?
The oxygen-poor centre becomes anodic by differential aeration, while the oxygen-rich edge stays cathodic — corrosion concentrates where oxygen is least.
How does a magnesium block protect a pipeline?
Magnesium, with a more negative standard potential than iron, oxidises in iron's place; the pipeline becomes a cathode and survives until the block is consumed.
Why does a scratched tin-plated can corrode faster than plain iron?
Tin is nobler than iron, so at the scratch iron becomes the anode of a galvanic couple and corrodes faster than it would alone.
What makes stainless steel corrosion-resistant?
Chromium above about 12% forms an adherent, self-renewing Cr2O3 film that passivates the surface against the anodic dissolution.