Corrosion

Corrosion is the gradual, spontaneous deterioration of a metal caused by chemical or electrochemical attack from its environment — oxygen, moisture, acidic gases, dissolved salts — which converts the metal into an oxide, hydroxide, sulphide, or carbonate. The most important example, and the one the syllabus focuses on, is the rusting of iron to hydrated iron(III) oxide, $\text{Fe}_2\text{O}_3\cdot x\text{H}_2\text{O}$.

The mechanism is electrochemical, not a plain surface reaction. A droplet of water sitting on iron, with dissolved $\text{O}_2$ and often $\text{CO}_2$ making it mildly acidic and conducting, sets up a miniature galvanic cell across the metal's own surface:

• Anode — typically the centre of the droplet or an impure/strained region, where oxygen access is lower: $\text{Fe(s)} \rightarrow \text{Fe}^{2+}(aq) + 2e^-$.
• Cathode — typically the edge of the droplet, where dissolved $\text{O}_2$ is more available: $\text{O}_2(g) + 4\text{H}^+(aq) + 4e^- \rightarrow 2\text{H}_2\text{O(l)}$.

Electrons released at the anode travel through the metal lattice itself to the cathodic region — the iron is simultaneously anode, cathode, and internal wire, with the water film supplying the electrolyte path for ions. The $\text{Fe}^{2+}$ produced at the anode does not stay put: dissolved oxygen oxidises it further to $\text{Fe}^{3+}$, which precipitates out as rust,

where: $x$ is the number of water molecules of hydration in the rust, and the $\text{H}^+$ released feeds back into the cathodic reaction, sustaining the cycle. Rust does not form where iron dissolved (the anode) but out in the surrounding solution wherever $\text{Fe}^{2+}$ meets enough oxygen — so rust often appears as a ring around, not exactly on, the pitted spot.

What speeds corrosion up. Dissolved electrolytes (salty air, de-icing salt) raise the conductivity of the moisture film and let the corrosion current flow faster. Acidic gases ($\text{CO}_2$, $\text{SO}_2$, $\text{NO}_x$) lower the pH of the film, supplying more $\text{H}^+$ for the cathodic step. Impure or strained metal (a bent nail, a weld line) becomes locally anodic, concentrating the attack. Contact with a less reactive metal creates a larger, more efficient galvanic couple. Higher temperature and humidity speed up both the chemistry and the ion mobility that carries the current.

Every prevention method either keeps the electrolyte film away from the metal (a barrier) or rigs the galvanic couple so the metal to be protected acts as the cathode (sacrificial protection). Barrier protection — paint, oil, grease, or plating with a less reactive metal such as chromium or tin — excludes moisture and oxygen, but works only while the barrier stays intact. Galvanisation coats iron with zinc; since zinc is more reactive, a scratch that exposes bare iron leaves zinc as the anode, so zinc corrodes preferentially and the iron stays protected — a barrier that keeps working after damage. Cathodic protection connects the structure to be protected (a pipeline, a ship's hull) to a still more reactive metal, the sacrificial anode (magnesium or zinc), consumed in place of the structure and periodically replaced. Alloying, as in chromium-bearing stainless steel, forms a thin, adherent, self-healing oxide layer that passivates the surface.

Problem. An underground steel pipeline is protected by connecting it to a buried magnesium block. Explain how this achieves cathodic protection, write the two half-reactions, and state what limits how long the protection lasts.

Step 1. The magnesium block and the steel pipeline are wired together and both sit in moist soil, which acts as the electrolyte — this is a deliberately built galvanic couple, following the same $E^\circ$ logic as the derivation above.

Step 2. Compare potentials: $E^\circ_{\text{Mg}^{2+}/\text{Mg}} = -2.37\ \text{V}$ versus $E^\circ_{\text{Fe}^{2+}/\text{Fe}} = -0.44\ \text{V}$. Magnesium is far more negative, so it is oxidised far more readily than iron and takes the role of anode.

Step 3. Write the half-reactions: anode (magnesium, sacrificed), $\text{Mg(s)} \rightarrow \text{Mg}^{2+}(aq) + 2e^-$; cathode (pipeline, protected), $\text{O}_2(g) + 2\text{H}_2\text{O(l)} + 4e^- \rightarrow 4\text{OH}^-(aq)$ — a reduction, so no iron is lost here.

Step 4. The magnesium block is steadily consumed as it keeps donating electrons to the pipeline, so it must be inspected and replaced periodically; the moment it is fully consumed or disconnected, the pipeline itself becomes the most reactive metal in contact with the soil and starts corroding. Answer: the pipeline stays protected because magnesium, being far more negative in $E^\circ$, is preferentially oxidised in its place; protection lasts only as long as the sacrificial magnesium block remains connected and unconsumed.

• Treating rusting as one uniform chemical reaction rather than two spatially separated half-reactions (anode and cathode) linked by electron flow through the metal and ion flow through the moisture film.
• Assuming any metal coating protects equally — a coating below the base metal in the electrochemical series (tin on iron) accelerates corrosion once scratched, because it forces the base metal to be the anode.
• Forgetting that a sacrificial anode must have a more negative $E^\circ$ than the metal it protects, and that it is consumed over time, not a permanent one-time fix.
• Assuming dry air alone rusts iron — the moisture film is the electrolyte that lets the corrosion current flow; iron in genuinely dry air corrodes negligibly.

Ask the tutor: want to work through more electrochemical-series predictions for other metal-coating pairs, or see how real ships and pipelines choose and size their sacrificial anodes?