The Electrochemistry of Corrosion and How It Is Prevented
Rusting is not a slow burning. It is a short-circuited galvanic cell on the surface of the metal, with an anode, a cathode, an electrolyte and a real measurable cell potential — which means the Nernst equation you learned for calculating EMF describes it exactly. Seen that way, every prevention method in the syllabus stops being a list to memorise: each one simply removes something the cell needs. This article computes the driving force for rusting, shows why zinc protects iron and tin does not, and then explains the limitation engineers care about most — potential tells you whether, never how fast.
The formula you already know
Ecell = Ecathode − Eanode ΔG = −nFEcell
The constant 0.0592 is 2.303RT/F at 298.15 K: 2.303 × 8.314 × 298.15 ÷ 96 485 = 0.05916 V. It is temperature-dependent, so a corrosion calculation at 60 °C uses a different number.
The two half-reactions of rusting
Wherever the metal is slightly strained, impure or under a droplet, iron dissolves:
The electrons travel through the metal itself to a place where oxygen is available, and are consumed there. In near-neutral water this is oxygen reduction; in acid the same reduction is written with H⁺:
Cathode (acidic): O₂ + 4H⁺ + 4e⁻ → 2H₂O E° = +1.23 V
The Fe²⁺ produced diffuses away, meets dissolved oxygen and is oxidised further to iron(III), which precipitates as hydrated iron(III) oxide, Fe₂O₃·xH₂O — rust. Note where the rust actually forms: at neither electrode, but between them. That is why rust flakes off instead of sealing the surface, and it is the single reason iron corrodes away completely while aluminium does not.
Worked example 1 — the driving force for rusting
Take an iron surface under a film of aerated neutral water: pH 7.00, p(O₂) = 0.21 bar (atmospheric), and a low dissolved iron concentration, [Fe²⁺] = 1.0 × 10⁻⁶ M.
Cathode — using the acidic form, O₂ + 4H⁺ + 4e⁻ → 2H₂O, n = 4:
E = 1.23 − (0.0592/4) · log[1 ÷ ([H⁺]⁴ · pO₂)]
E = 1.23 + 0.0148 · (4 log[H⁺] + log pO₂)
4 log[H⁺] = 4 × (−7.00) = −28.00 · log 0.21 = −0.678
E = 1.23 + 0.0148 × (−28.678) = 1.23 − 0.424 = +0.806 V
Cross-check with the alkaline form, O₂ + 2H₂O + 4e⁻ → 4OH⁻, E° = +0.401 V,
with [OH⁻] = 1.0 × 10⁻⁷ M:
E = 0.401 − 0.0148 · (4 log[OH⁻] − log pO₂) = 0.401 − 0.0148 × (−28.00 + 0.678)
E = 0.401 + 0.404 = +0.805 V — the two routes agree, as they must.
Anode — Fe²⁺ + 2e⁻ → Fe, n = 2:
E = −0.44 + (0.0592/2) · log[Fe²⁺] = −0.44 + 0.0296 × (−6.00) = −0.44 − 0.178 = −0.618 V
Cell: Ecell = 0.805 − (−0.618) = +1.42 V
ΔG = −nFE = −4 × 96 485 × 1.42 = −5.49 × 10⁵ J = −549 kJ per mole of O₂ reduced
A positive cell potential of this size says rusting is strongly spontaneous under ordinary conditions. It also shows where the levers are: the cathode potential falls by about 59 mV per unit rise in pH, and it falls as oxygen is excluded. Deaerated water and alkaline conditions both reduce the driving force — which is exactly why boiler feedwater is deaerated and kept alkaline.
Worked example 2 — why zinc protects iron and tin does not
Read the electrochemical series and the answer is immediate. Couple iron to a second metal, and whichever half-cell has the more negative potential becomes the anode and corrodes.
| Half-cell | E° (V) | Coupled to iron |
|---|---|---|
| Mg²⁺ + 2e⁻ → Mg | −2.37 | Mg is the anode — protects iron |
| Al³⁺ + 3e⁻ → Al | −1.66 | Al is the anode — protects iron |
| Zn²⁺ + 2e⁻ → Zn | −0.76 | Zn is the anode — protects iron |
| Fe²⁺ + 2e⁻ → Fe | −0.44 | reference |
| Sn²⁺ + 2e⁻ → Sn | −0.14 | Fe is the anode — iron corrodes faster |
| Cu²⁺ + 2e⁻ → Cu | +0.34 | Fe is the anode — iron corrodes faster |
Galvanised iron (zinc coating):
E°cell = E°(Fe²⁺/Fe) − E°(Zn²⁺/Zn) = −0.44 − (−0.76) = +0.32 V
Positive, so the spontaneous reaction is Zn oxidising while Fe²⁺ is reduced. Zinc is consumed
first. A scratch through the coating still leaves the exposed iron protected, because the
surrounding zinc keeps feeding it electrons.
Tin-plated iron (a tin can):
E°cell = E°(Sn²⁺/Sn) − E°(Fe²⁺/Fe) = −0.14 − (−0.44) = +0.30 V
Now iron is the anode. While the tin layer is intact it is a perfectly good barrier, but the
moment it is scratched the exposed iron becomes a small anode connected to a large tin
cathode — and corrodes faster than bare iron would.
Same magnitude of potential, opposite engineering outcome. This contrast is a standard exam question and it is worth being able to derive rather than recall.
How corrosion is actually prevented
| Method | Which part of the cell it attacks |
|---|---|
| Paint, grease, polymer and enamel coatings | Removes the electrolyte and the oxygen supply |
| Galvanising (zinc) | Barrier plus sacrificial anode — works even when scratched |
| Sacrificial anodes (Mg, Zn, Al blocks on ships, tanks, pipelines) | Makes the structure the cathode; the block is replaced periodically |
| Impressed-current cathodic protection | An external supply holds the structure at a potential where iron does not dissolve |
| Alloying — chromium in stainless steel | Forms a thin, adherent, self-repairing passive oxide film |
| Corrosion inhibitors in closed systems | Adsorb on the surface, or shift the local chemistry, to block one half-reaction |
| Deaeration and pH control of process water | Lowers the cathode potential and removes the oxidising agent |
| Design — drainage, avoiding crevices and dissimilar-metal joints | Prevents the trapped, stagnant, oxygen-starved conditions that start pits |
The honest limit — potential does not give you a rate
- Thermodynamics says "yes", kinetics decides "how fast". The +1.42 V calculated above tells you rusting can happen, not that it will happen quickly. The actual rate depends on the exchange current densities of both half-reactions, the overpotential at each, the conductivity of the electrolyte and, very often, the rate at which oxygen can diffuse to the surface. A corrosion rate in millimetres per year is measured, never predicted from E° alone.
- Aluminium is the proof. E°(Al³⁺/Al) = −1.66 V makes aluminium thermodynamically far more reactive than iron, yet aluminium window frames survive decades. The reason is entirely kinetic: a compact, adherent Al₂O₃ film forms in seconds and stops further attack. Iron oxide is porous and flakes, so iron gets no such protection. The electrochemical series ranks tendency, not durability.
- Passivity can be broken. Chloride ions locally destroy passive films on stainless steel and aluminium, producing pitting — small, deep, fast-growing holes in a surface that looks clean. This is why coastal and marine service is a different design problem from inland service, and why a passive alloy can fail suddenly rather than thin gradually.
- Real metals do not have a single equilibrium potential. An alloy surface adopts a mixed (corrosion) potential where the total oxidation and reduction currents balance. Nernst calculations for pure metals in standard states give the direction, not a description of the surface.
- Area ratio matters more than the potential difference. A small anode joined to a large cathode concentrates all the dissolution into a tiny area — a steel bolt in a copper plate fails quickly, while a copper rivet in a steel plate barely matters. Nothing in E°cell tells you this.
- Differential aeration corrodes uniform metal. No second metal is needed. The area under a droplet, gasket or layer of dirt is oxygen-poor and becomes the anode relative to the aerated metal beside it — which is why rust appears under seals and inside crevices rather than on open surfaces.
- Common exam slip: forgetting that the electrochemical series is tabulated as reduction potentials. Reverse the sign only when you write the half-reaction as an oxidation, and never reverse it inside Ecell = Ecathode − Eanode, which already accounts for it.
Why this matters for JAM, GATE, NET and CUET-PG
| Exam area | What is typically asked |
|---|---|
| Electrochemistry | Nernst equation with pH and partial pressure in Q; Ecell and ΔG = −nFE |
| Electrochemical series | Predicting which metal corrodes in a given couple |
| Applied/industrial chemistry | Galvanising, cathodic protection, inhibitors, passivation |
| Inorganic chemistry | Oxide film structure and why some metals passivate |
| Conceptual/assertion questions | Thermodynamic tendency versus observed kinetic stability |
Run the numbers instead of guessing the sign. The Nernst calculator takes E°, n and the reaction quotient and returns the cell potential under real conditions — ideal for pH-dependent and pressure-dependent half-cells like the ones above.
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