# Corrosion and Rusting

> Corrosion for JEE Chemistry: electrochemical theory of rusting, anode and cathode half-reactions, differential aeration, sacrificial protection, galvanisation.

- Canonical URL: https://prepelephant.com/topics/jee/chemistry/corrosion-jee
- Exam / course: JEE · Subject: Chemistry
- Publisher: PrepElephant (https://prepelephant.com) — Prepared and reviewed by the PrepElephant Academic Review Team
- First published: 2026-10-02
- Last updated: 2026-10-02
- How to cite: "Corrosion and Rusting", PrepElephant, https://prepelephant.com/topics/jee/chemistry/corrosion-jee

## 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.