Chemistry: Atoms First 2e · Electrochemistry
Corrosion
On this page 9 sections
In 30 seconds
Corrosion Spontaneous oxidation of a metal by its environment Full entry → is the spontaneous, usually unwanted oxidation of a metal by its environment, and it is electrochemistry in action on a grand scale. A piece of iron left outdoors does not simply "wear away" — it becomes the Anode The electrode where oxidation occurs; metal dissolves here Full entry → of an electrochemical cell in which oxygen from the air is reduced at nearby Cathode The electrode where reduction occurs; O2 is consumed here Full entry → sites, with rainwater, soil moisture, or salt spray acting as the Electrolyte A solution containing ions that conducts charge between electrodes Full entry →. Rusting is therefore not a single reaction but a coupled redox process: oxidation of the metal at anodic spots, reduction of oxygen at cathodic spots, electron flow through the metal itself, and ion flow through the surrounding solution. This topic examines the half-reactions behind rusting, why some metals corrode while others resist, and the electrochemical strategies — coatings, Galvanizing Coating steel with zinc to protect it Full entry →, sacrificial anodes, and alloying — used to slow corrosion down.
Why this matters
- Infrastructure and safety: Corrosion degrades bridges, pipelines, ships, and aircraft. Industry studies estimate the global cost of corrosion at roughly 3–4% of world GDP, a large fraction of it preventable.
- Catastrophic failures: Corrosion-weakened gas pipelines, storage tanks, and fasteners can fail without warning; understanding the chemistry helps engineers inspect and protect the right things.
- Medical devices: Metal implants (stainless steel, titanium alloys) must resist corrosion inside the body, where saltwater-like fluids make corrosion more aggressive.
- Prevention economics: Choosing the right coating, alloy, or Sacrificial anode A more active metal attached to corrode in place of the protected metal Full entry → can multiply a structure's lifetime many times over — a direct application of cell-potential logic from this chapter.
- Exams: Corrosion problems test whether you can identify the anode and cathode half-reactions, compute E°cell, and explain why a more active metal protects a less active one.
The college version
Core Concepts
Corrosion is a spontaneous electrochemical cell
For corrosion to occur, four things must be present: a metal that can be oxidized, an oxidizing agent (usually O2), water or moisture to serve as electrolyte, and an electrical path for electrons. In rusting iron, microscopic anodic regions lose electrons while nearby cathodic regions pass them to oxygen. The metal itself conducts electrons, and the water film conducts ions, so the cell is complete without any wires.
The half-reactions of rusting
At anodic sites, iron dissolves:
Fe(s) → Fe2+(aq) + 2e- E°= -0.447 V
At cathodic sites (in neutral or basic water), dissolved oxygen is reduced:
O2(g) + 2H2O(l) + 4e- → 4OH-(aq) E°= +0.401 V
Combining (doubling the iron half-reaction to balance electrons) gives:
2Fe(s) + O2(g) + 2H2O(l) → 2Fe(OH)2(s)
The iron(II) hydroxide then oxidizes further in air to the familiar reddish-brown Rust Hydrated iron(III) oxide, Fe2O3 · xH2O Full entry →, Fe2O3 · xH2O. Notice that both water and oxygen are reactants, which is why iron rusts faster in humid air than in dry air, and faster still where salt (an electrolyte) is present — salt water conducts ions far better than pure water, speeding the ionic half of the circuit.
Why some metals corrode and others do not
Whether a metal corrodes depends on how easily it gives up electrons, quantified by its standard reduction potential. Active metals with very negative E° values (magnesium, zinc, iron) oxidize readily; noble metals with positive E° values (gold, platinum) resist corrosion because oxidation is not energetically favorable. Some metals, notably aluminum and chromium, corrode so quickly at first that they form a thin, tightly adherent oxide film (Passivation Formation of a thin protective oxide film on a metal surface Full entry → layer) that seals the surface and stops further attack — the reason an aluminum ladder does not rust away.
Preventing corrosion electrochemically
- Barrier coatings (paint, oil, plastic, enamel) keep water and oxygen away from the metal surface. Any scratch or chip, however, reopens the cell.
- Galvanizing coats steel with zinc. Zinc protects in two ways: it is a physical barrier, and if scratched, it becomes the sacrificial anode because Zn2+/Zn (-0.762 V) is more negative than Fe2+/Fe (-0.447 V). The zinc corrodes instead of the steel.
- Cathodic protection Forcing a structure to act as cathode so it cannot oxidize Full entry → deliberately attaches a more active metal (zinc or magnesium) to pipelines, ship hulls, and underground tanks; the attached metal corrodes sacrificially while the protected structure is forced to act as cathode. Impressed-current systems use a rectifier to drive the same effect.
- Alloying produces stainless steel: adding chromium (and nickel) creates a self-healing passivation layer of Cr2O3.
- Environmental control — removing oxygen (deaerated boilers), removing electrolytes, or adding corrosion inhibitors — attacks the conditions that sustain the cell.
Common Confusions
| Do Not Confuse | With | Difference |
|---|---|---|
| Rust (Fe2O3 · xH2O) | Fe(OH)2 | The hydroxide forms first and is then oxidized in air to rust; they are different compounds at different stages |
| Corrosion requiring two metals | Corrosion of a single metal | One metal alone corrodes if it can act as both anode and cathode — rusting iron needs no second metal |
| "Rusting is just oxidation" | Rusting as a redox couple | Both oxidation (Fe) and reduction (O2) occur; a corrosion explanation must name both half-reactions |
| A more positive E° metal protecting iron | A more negative E° metal protecting iron | The more negative potential metal (Zn, Mg) is the sacrificial anode; noble metals like copper accelerate iron's corrosion when in contact |
| Standard E° values | Real-world corrosion conditions | E° assumes 1 M solutions, 1 atm, 25 °C; real corrosion rates also depend on kinetics, oxygen access, and electrolyte |
| Aluminum "doesn't corrode" | Aluminum passivates | Aluminum oxidizes readily but the oxide film seals the surface; the corrosion is happening, just stopped at the surface |
| Passivation layer | Paint coating | Passivation is a chemical oxide film that self-heals; paint is an applied physical barrier that does not |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Rusting is like a tiny battery draining on the metal's surface. One spot on the iron gives up electrons and slowly dissolves away, while another spot hands those electrons to oxygen in the air. Water acts like the wet wire that completes the battery's circuit. If you cover the iron with paint, the battery cannot form; if you bolt a piece of zinc onto it, the zinc becomes the part that "battery-drains" instead of the iron.
Worked example
Example 1: Is rusting spontaneous? Compute E°cell
Use standard potentials to show that the iron–oxygen couple that drives rusting is spontaneous.
Step 1 — Write the half-reactions with potentials:
Fe2+ + 2e- → Fe E°= -0.447 V O2 + 2H2O + 4e- → 4OH- E°= +0.401 V
Step 2 — Identify anode and cathode. Iron is oxidized, so the Fe2+/Fe couple is the anode (oxidation reverses it); oxygen is reduced at the cathode.
Step 3 — Apply the cell-potential formula:
E°cell = E°cathode - E°anode = 0.401 V - (-0.447 V) = +0.848 V
A positive E°cell means ΔG°= -nFE°cell < 0, so the reaction is spontaneous — corrosion happens on its own, which is exactly why unprotected iron rusts without any external power source.
Example 2: Why zinc protects iron (galvanizing)
Compare the standard reduction potentials to show why a scratch in a galvanized coating corrodes the zinc, not the steel underneath.
Step 1 — Write the relevant potentials:
Fe2+ + 2e- → Fe E°= -0.447 V Zn2+ + 2e- → Zn E°= -0.762 V
Step 2 — Identify which metal is the stronger reducing agent. The more negative potential (Zn, -0.762 V) means zinc loses electrons more readily, so in a zinc–iron couple zinc is the anode and iron the cathode.
Step 3 — Compute the cell potential for the protecting couple:
E°cell = E°cathode - E°anode = (-0.447 V) - (-0.762 V) = +0.315 V
Because this couple is spontaneous with zinc as anode, the zinc dissolves preferentially; the iron (cathode) is spared. The same logic explains why magnesium anodes (-2.372 V) protect buried pipelines.
Example 3: Tracing corrosion at a scratched surface
A galvanized steel gate is scratched down to bare steel and left in humid coastal air. Walk through what happens: the scratch exposes steel while zinc surrounds it; seawater spray provides an electrolyte; zinc, having the more negative E°, becomes the anode and oxidizes to Zn2+, releasing electrons that flow through the steel to cathodic sites where oxygen is reduced. The steel stays intact (cathodically protected) while the zinc coating slowly sacrifices itself — the reason galvanized steel lasts far longer than painted steel once both are scratched.
Key takeaways
- Corrosion is a spontaneous redox cell: metal oxidation at the anode, O2 reduction at the cathode, with moisture as electrolyte.
- Rusting half-reactions: Fe → Fe2+ + 2e- (-0.447 V) and O2 + 2H2O + 4e- → 4OH- (+0.401 V).
- E°cell = E°cathode - E°anode; a positive value means the corrosion reaction is spontaneous under standard conditions.
- More negative E° = more active metal = stronger tendency to corrode = better sacrificial anode (Zn, Mg protect Fe).
- Galvanizing and sacrificial anodes work because zinc or magnesium oxidizes instead of iron.
- Aluminum and chromium "self-protect" via passivating oxide films; this is why they resist corrosion despite active E° values.
- Salt speeds corrosion by making the electrolyte more conductive; dry air and pure water slow it.
Check yourself
6 review questions from the chapter. Try each one, then open the answer.
Write the two half-reactions of iron rusting and label the anode and cathode.
Show answer
Anode: Fe → Fe2+ + 2e-; cathode: O2 + 2H2O + 4e- → 4OH-.
Why does iron corrode faster in salt spray than in pure water?
Show answer
Salt dissolves into the water film, making the electrolyte more conductive, so the ionic half of the corrosion circuit carries charge faster.
Why does galvanized steel protect iron even after the zinc coating is scratched?
Show answer
Zinc has a more negative reduction potential (-0.762 V) than iron (-0.447 V), so zinc oxidizes preferentially at the scratch, keeping the exposed steel cathodic and protected.
Aluminum has a very negative E° yet aluminum structures resist corrosion. Explain.
Show answer
Aluminum oxidizes to form a thin, tightly adherent Al2O3 passivation film that blocks further contact between the metal and oxygen/water.
Compute E°cell for the zinc–iron couple and state which metal corrodes.
Show answer
E°cell = E°cathode - E°anode = (-0.447) - (-0.762) = +0.315 V; zinc is the anode and corrodes.
Name two electrochemical methods used to protect buried pipelines.
Show answer
Sacrificial (galvanic) anodes such as magnesium, and impressed-current cathodic protection.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- Corrosion
- Spontaneous oxidation of a metal by its environment
- Anode
- The electrode where oxidation occurs; metal dissolves here
- Cathode
- The electrode where reduction occurs; O2 is consumed here
- Electrolyte
- A solution containing ions that conducts charge between electrodes
- Rust
- Hydrated iron(III) oxide, Fe2O3 · xH2O
- Passivation
- Formation of a thin protective oxide film on a metal surface
- Sacrificial anode
- A more active metal attached to corrode in place of the protected metal
- Galvanizing
- Coating steel with zinc to protect it
- Standard reduction potential E°
- Voltage of a half-reaction measured against the standard hydrogen electrode
- Cathodic protection
- Forcing a structure to act as cathode so it cannot oxidize
Sources & references
This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.
Educational content only. It is not medical, legal or professional advice. Found an error? Tell us.