Corrosion

A piece of iron exposed to moist air is never electrochemically uniform. Impurities, mechanical stress, and uneven access to oxygen make some tiny regions slightly more reactive than their neighbours. The more reactive spots become anodic, the less reactive spots become cathodic, and the thin film of moisture sitting on the surface — made mildly acidic by dissolved atmospheric $\text{CO}_2$ forming carbonic acid — acts as the electrolyte connecting them. The metal itself carries the electrons from anode to cathode, exactly as a wire would in a lab galvanic cell.

At the anodic region, iron is oxidised:

At the cathodic region, the electrons released are consumed by dissolved oxygen, reduced in the presence of $\text{H}^+$ ions:

where: $E^\circ$ is the standard reduction potential of each half-reaction, measured in volts against the standard hydrogen electrode, and the sign convention is that a more positive $E^\circ$ means a greater tendency to be reduced (to gain electrons). The overall cell potential, cathode minus anode, is

strongly positive, so the process runs on its own without any external push — which is exactly why iron left in humid air rusts so readily and so persistently. The $\text{Fe}^{2+}$ produced does not stop there. Dissolved oxygen oxidises it further to $\text{Fe}^{3+}$, which then combines with water to give hydrated iron(III) oxide, the familiar reddish-brown rust:

The hydration number $x$ is variable, which is why rust is written $\text{Fe}_2\text{O}_3\cdot x\text{H}_2\text{O}$ rather than as a fixed formula. Rust is also loose, flaky and porous, so unlike the tight protective oxide layer some metals form, it never seals off the metal underneath — corrosion just keeps eating inward.

Factors that speed up corrosion: moisture with dissolved electrolytes (near the sea, or in air polluted with $\text{SO}_2$/$\text{CO}_2$); impurities in the metal that sharpen the anode-cathode contrast; and contact between two dissimilar metals, where the more reactive one (lower $E^\circ$) corrodes preferentially.

Prevention, at board level: barrier methods (paint, oil, galvanising) that simply keep moisture and oxygen out; sacrificial (cathodic) protection, where a more reactive metal like zinc or magnesium is attached and corrodes in place of the iron; and alloying, as in stainless steel, which forms a thin, self-healing protective oxide.

The value $E^\circ_{cell} = 1.67\ \text{V}$ is not just an exam number — it is the reason rusting needs no external trigger. Any galvanic couple with $E^\circ_{cell} > 0$ is thermodynamically spontaneous, since $\Delta G^\circ = -nFE^\circ_{cell}$ is then negative. Because the oxygen-reduction half-reaction has such a high standard potential ($+1.23\ \text{V}$), it will pair with almost any common metal's oxidation half-reaction and still come out spontaneous — which is precisely why so many metals corrode in ordinary moist air, not just iron.

Problem: A steel gate is coated first with zinc (galvanised, $E^\circ_{\text{Zn}^{2+}/\text{Zn}} = -0.76\ \text{V}$) and, on a second batch, with tin ($E^\circ_{\text{Sn}^{2+}/\text{Sn}} = -0.14\ \text{V}$). Both coatings get scratched, exposing bare iron ($E^\circ_{\text{Fe}^{2+}/\text{Fe}} = -0.44\ \text{V}$) next to the coating metal. Explain, using the electrode potentials, which coating still protects the iron after the scratch and which one makes corrosion worse.

Step 1. A scratch exposes bare iron right next to the coating metal, with the surface moisture film acting as electrolyte — this is a new, tiny galvanic couple, and whichever metal has the lower (more negative) $E^\circ$ becomes the anode and corrodes preferentially.

Step 2. Compare zinc to iron: $E^\circ_{\text{Zn}^{2+}/\text{Zn}} = -0.76\ \text{V}$ is more negative than $E^\circ_{\text{Fe}^{2+}/\text{Fe}} = -0.44\ \text{V}$, so zinc is more easily oxidised than iron. At the scratch, zinc becomes the anode and keeps oxidising ($\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-$) in place of iron, while the iron sits as the protected cathode. This is sacrificial (cathodic) protection, and it continues to work even where the coating is broken.

Step 3. Compare tin to iron: $E^\circ_{\text{Sn}^{2+}/\text{Sn}} = -0.14\ \text{V}$ is less negative than iron's $-0.44\ \text{V}$, so tin is harder to oxidise than iron. At the scratch, iron itself becomes the anode of the couple, and the tin cathode draws electrons away from it, so the exposed iron corrodes faster than if it had never been coated at all.

The galvanised (zinc) coating still protects the iron after the scratch; the tin-plated coating fails and actively worsens corrosion once breached — because zinc's $E^\circ$ lies below iron's while tin's lies above it.

• Treating corrosion as a plain chemical reaction rather than an electrochemical one — the anodic and cathodic sites, electron flow through the metal, and ion transport through the moisture film are the same three ingredients as any galvanic cell.
• Mixing up galvanising and tin-plating: only a coating metal with a lower $E^\circ$ than iron (like zinc) gives sacrificial protection after a scratch; a coating with a higher $E^\circ$ (like tin) accelerates corrosion once breached.
• Forgetting that pure, dry air or pure water alone rusts iron only very slowly — it is moisture together with dissolved oxygen and electrolytes (like $\text{CO}_2$) that drives the process at a noticeable rate.
• Writing rust simply as $\text{Fe}_2\text{O}_3$ instead of the hydrated, variable-composition $\text{Fe}_2\text{O}_3\cdot x\text{H}_2\text{O}$.

Ask the tutor: want to work through why some metal coatings keep protecting iron after a scratch while others make it worse, or compare electrode potentials for a few more common alloys?