Chemistry 2e · Electrochemistry
Corrosion
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In 30 seconds
Corrosion Unwanted oxidation of a metal by its environment Full entry → is the unwanted oxidation of a metal by its environment — the reason a steel railing rusts, a silver spoon tarnishes, and a copper roof turns green. Far from being a simple surface reaction, corrosion is electrochemistry in the open: the metal acts as an anode, oxygen (or another oxidant) acts as the cathode, and the metal itself conducts the electrons between them. This topic uses the half-reaction framework to explain why iron rusts, why Rust Hydrated iron(III) oxide, Fe₂O₃·xH₂O Full entry → keeps corroding instead of protecting the metal, and how coatings, alloying, and Cathodic protection Making the structure the cathode by supplying electrons Full entry → interrupt the electrochemical circuit.
Why this matters
Corrosion is one of the most expensive chemical reactions on Earth, with direct and indirect costs estimated in the hundreds of billions of dollars per year worldwide: crumbling bridges, leaking pipelines, failing aircraft fasteners, and corroded rebar in concrete. In healthcare, corrosion explains why surgical implants must be carefully chosen metals and why sterilization and body chemistry interact with them. Understanding the electrochemistry lets engineers prevent failure — and lets you explain why galvanized nails outlast plain steel ones, why salt accelerates rusting on cars, and why a magnesium block strapped to a water heater protects the steel tank.
The college version
Core Concepts
Rusting is an electrochemical cell
Iron rusts only when both water and oxygen are present. The metal surface splits into microscopic electrodes: at anodic regions, iron is oxidized; at cathodic regions, oxygen is reduced:
Anode: Fe(s) → Fe2+(aq) + 2e- Cathode: O2(g) + 2H2O(l) + 4e- → 4OH-(aq)
The electrons travel through the iron itself, and the ions travel through the water film, completing the circuit. The Fe²⁺ then oxidizes further and combines with hydroxide to form hydrated iron(III) oxide — the familiar reddish-brown rust, Fe2O3 · xH2O. Because both half-reactions are spontaneous, no external energy is needed; the cell runs itself as long as water and oxygen are supplied.
Why rust is worse than a surface film
Some metals protect themselves: aluminum instantly forms a thin, tightly adherent oxide layer that seals the surface and stops further attack. Rust does not. Rust is porous, flaky, and larger in volume than the iron it replaced, so it cracks away and exposes fresh metal to water and oxygen — the corrosion continues until the object is consumed. This is why a scratched aluminum ladder survives while a scratched steel beam keeps rusting.
The role of the environment
Anything that helps the electrochemical circuit speeds corrosion. Salts (road salt, sea spray) increase the conductivity of the water film, so the cell carries more current. Acids supply H⁺ that can be reduced directly and dissolve protective oxide films. Even ordinary moisture condensing on metal — a "water line" or a droplet — creates aeration differences, with oxygen-poor regions becoming anodes. Galvanic corrosion is the same idea on purpose: when two dissimilar metals touch in an electrolyte (for example, copper pipe joined to steel), the less noble metal corrodes faster than it would alone.
Prevention: break the circuit
Because corrosion is an electrochemical cell, it is stopped by removing any one of its parts. Coatings (paint, oil, plastic) exclude water and oxygen. Galvanizing Coating iron with zinc Full entry → coats iron with zinc, which both seals the surface and protects sacrificially: zinc is more easily oxidized than iron, so if the coating is scratched, the zinc corrodes instead of the steel. Alloying can change the metal's behavior — stainless steel adds chromium, which forms a protective oxide film much like aluminum's. Cathodic protection forces the metal to be the cathode: either a Sacrificial anode A more active metal (Mg, Zn) attached to the protected metal Full entry → (a block of magnesium or zinc, more active than the protected metal, that corrodes in its place) or an impressed current from an external power supply, used for pipelines, ship hulls, and buried tanks.
Common Confusions
| Do Not Confuse | With | Difference |
|---|---|---|
| Rusting | Tarnishing or patina | Rust is porous and self-sustaining; silver tarnish and copper patina form protective films |
| Corrosion of iron | Corrosion of aluminum | Aluminum's oxide film stops attack; iron's rust does not |
| Anode in corrosion | Anode in a battery | Same definition (oxidation site) — corrosion is just an unwanted, uncontrolled battery |
| Galvanized = coated with zinc | Galvanized = coated with tin | Tin protects by sealing only (and is less active than iron); zinc protects sacrificially |
| Sacrificial anode | Cathode of the cell | The sacrificial anode is the anode — it corrodes so the structure can be the cathode |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Rust is like a tiny, invisible battery spreading across the surface of iron. The iron is one side of the battery, the air and water are the other, and the metal itself is the wire between them — the battery runs until the iron is used up. A scratch lets water in, and rust is crumbly, so it keeps exposing new iron to keep the "battery" running. To stop it, you cover the metal, or attach a metal that would rather be the battery — like zinc — so the zinc gets eaten instead of the iron.
Worked example
Worked Example 1 — Is rusting spontaneous? Combine the iron and oxygen half-reactions with standard potentials E°(Fe2+/Fe) = -0.45 V and E°(O2/OH-) = +0.40 V:
E°cell = E°cathode - E°anode = 0.40 V - (-0.45 V) = 0.85 V
The positive potential means the reaction is spontaneous, exactly as everyday experience suggests. The corresponding free-energy change for n = 4 electrons per mole of O₂:
ΔG°= -nFE°cell = -(4 mol e-)(96,485 Cmol e-)(0.85 V) = -3.28 × 105 J = -328 kJ
per mole of O₂ — a strongly downhill reaction that explains why unprotected iron does not stay shiny for long.
Worked Example 2 — How fast does iron disappear? Faraday's law relates the mass of metal oxidized to the charge passed: m = MItnF, where M is the molar mass, I the current, and t the time. Suppose a corroding steel structure behaves like a cell carrying an average corrosion current of 0.50 A (as Fe → Fe²⁺ + 2e⁻, so n = 2). Over 30 days (t = 30 × 24 × 3600 = 2.59 × 106 s):
m = (55.85 g/mol)(0.50 A)(2.59 × 106 s)(2)(96,485 C/mol) = 375 g
Unit check: gmol × A × s × molC = gmol × Cs × s × molC = g. Even a modest "leakage current" of half an amp quietly removes about 0.38 kg of iron a month — a powerful argument for protective coatings.
Key takeaways
- Rusting requires water and oxygen together; iron is the anode, oxygen is reduced at the cathode.
- Rust is porous and flakes off, exposing fresh metal — corrosion is self-sustaining.
- Salt accelerates corrosion by increasing electrolyte conductivity; acids accelerate it by supplying H⁺.
- Aluminum and stainless steel (chromium) form adherent oxide films that stop corrosion; iron's rust does not.
- Galvanizing = zinc coating: seals the surface and protects sacrificially if scratched.
- Cathodic protection: sacrificial anodes (Mg, Zn) or impressed current keep the structure as the cathode.
- E° comparisons decide which metal corrodes: the more negative (less noble) metal is oxidized first.
Check yourself
5 review questions from the chapter. Try each one, then open the answer.
What two environmental components must be present for iron to rust, and why?
Show answer
Water and oxygen. Water carries ions that complete the electrochemical circuit, and oxygen is the oxidant reduced at the cathodic regions.
Why does rust keep corroding while aluminum oxide stops further attack?
Show answer
Rust is porous and flakes off, exposing fresh iron; aluminum oxide is a thin, adherent, impermeable film that seals the metal.
Why does road salt make cars rust faster?
Show answer
Salt dissolves in the water film and increases its ionic conductivity, so the corrosion cell carries more current and corrodes faster.
How does galvanizing protect a scratch in the zinc coating?
Show answer
Zinc is more easily oxidized than iron, so at the scratch the zinc corrodes sacrificially, feeding electrons to the iron and keeping it as the cathode.
Explain cathodic protection in one sentence, and name one everyday example.
Show answer
Supplying electrons to the structure keeps it as the cathode so it cannot be oxidized — for example, the sacrificial magnesium anode rod in a water heater (or zinc blocks on ship hulls).
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- Corrosion
- Unwanted oxidation of a metal by its environment
- Rust
- Hydrated iron(III) oxide, Fe₂O₃·xH₂O
- Anodic region
- Spot where the metal is oxidized, losing electrons
- Cathodic region
- Spot where oxygen (or H⁺) is reduced
- Galvanizing
- Coating iron with zinc
- Cathodic protection
- Making the structure the cathode by supplying electrons
- Sacrificial anode
- A more active metal (Mg, Zn) attached to the protected metal
- Passivation
- Formation of a thin, adherent protective oxide film
Sources & references
This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.
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