Corrosion Thermodynamics

Introduction

Every year, corrosion destroys a large fraction of the steel ever made and costs economies several per cent of their gross domestic product. Rusting is not simply "reacting with air". It is electrochemistry: a tiny galvanic cell, often invisible, operating on a single piece of metal. This page asks the thermodynamic question: under what conditions can a metal corrode at all? The next page asks how fast it happens.

Core explanation

Two coupled half-reactions. Corrosion needs an anodic (oxidation) reaction in which the metal loses electrons, for example

Fe → Fe²⁺ + 2e⁻ (E° = −0.44 V)

and a cathodic (reduction) reaction that consumes those electrons. The metal itself conducts electrons from anode sites to cathode sites, and a water film acts as the electrolyte. The two common cathodic reactions are:

- hydrogen evolution: 2H⁺ + 2e⁻ → H₂ (E° = 0.00 V), important in acids; - oxygen reduction: O₂ + 4H⁺ + 4e⁻ → 2H₂O (E° = +1.23 V), or O₂ + 2H₂O + 4e⁻ → 4OH⁻ in neutral and alkaline water.

The thermodynamic test. Corrosion can occur only if the overall reaction has ΔG < 0, that is if the equilibrium potential of the cathodic reaction lies above that of the metal's oxidation, under the actual conditions. Since oxygen reduction has such a high potential, almost every common metal is thermodynamically unstable in moist air. Gold (E°(Au³⁺/Au) ≈ +1.50 V) is a rare exception; that is why it is found native and does not tarnish.

Why conditions matter. The Nernst equation adjusts these potentials. For the hydrogen reaction, E = −0.059 pH V at 25 °C and 1 bar H₂, so hydrogen evolution becomes less favourable as pH rises. For the metal, E rises as the metal-ion activity increases; corrosion analyses often assume a small ion activity such as 10⁻⁶ to define when the metal "counts" as corroding. Oxygen concentration also matters: less dissolved oxygen means a lower cathodic potential.

Pourbaix diagrams. Plotting these equilibria on axes of potential (E) and pH produces a Pourbaix diagram. For iron in water it shows three kinds of regions:

- Immunity: at low potentials Fe metal is stable and cannot corrode. - Corrosion: where soluble ions such as Fe²⁺ or Fe³⁺ are stable (acidic solutions) or, at very high pH, soluble ferrate(II) species. - Passivity: where solid oxides or hydroxides, such as Fe₂O₃ or Fe₃O₄, are stable, as in mildly alkaline conditions.

Two sloping dashed lines mark the stability limits of water itself. Where the metal's corrosion region overlaps the region where water or oxygen can act as oxidant, corrosion is thermodynamically possible.

Rust. In neutral water, Fe²⁺ from the anode meets OH⁻ from oxygen reduction to precipitate Fe(OH)₂, which is further oxidised by dissolved oxygen to hydrated iron(III) oxide, Fe₂O₃·xH₂O. Rust is porous and flakes off, so it does not protect the underlying steel.

Step-by-step reasoning

To decide whether a metal can corrode:

1. Identify the possible cathodic reactions in the environment (H⁺, O₂ or other oxidants). 2. Find their equilibrium potentials, correcting for pH and concentration. 3. Find the equilibrium potential of the metal's oxidation. 4. If the cathodic potential is higher, ΔG is negative and corrosion is possible. 5. Check a Pourbaix diagram to see whether the product is soluble or a protective solid.

Visual explanation

Picture a water droplet on a steel plate. Near the edge, where oxygen diffuses in easily, the steel acts as a cathode and OH⁻ forms. Under the centre, where oxygen is scarce, the steel acts as an anode and dissolves. Rust forms in a ring between them, where the ions meet.

Real-world analogy

Thermodynamics is like a map showing which slopes lead downhill. A metal in moist air sits on a hillside above its oxide. The map says it will roll down eventually, but it says nothing about how steep the path is or whether a fence (a protective film) blocks the way.

Real-world example

Aluminium has a very negative electrode potential (about −1.66 V), so it is far more reactive than iron. Yet aluminium window frames last for decades, because in near-neutral conditions its Pourbaix diagram shows a stable, adherent Al₂O₃ layer. In strongly acidic or alkaline solutions, where the oxide dissolves, aluminium corrodes quickly.

Why?

Why does iron rust faster in seawater than in pure water? The thermodynamic driving force is similar, but dissolved salts greatly increase the conductivity of the electrolyte film, and chloride ions break down protective films, so the electrochemical cell operates much more effectively.

Common misconception

"Corrosion requires the metal to be directly attacked by oxygen at the same spot." In electrochemical corrosion, oxygen is reduced at one site while the metal dissolves at another, often millimetres away; electrons travel through the metal between them.

Worked example

Question: Can copper corrode in oxygen-free acid at pH 0? (E°(Cu²⁺/Cu) = +0.34 V)

Reasoning: Without oxygen, the only oxidant is H⁺, with E = 0.00 V at pH 0. Since 0.00 V is below +0.34 V, hydrogen ions cannot oxidise copper; ΔG for Cu + 2H⁺ → Cu²⁺ + H₂ is positive.

Answer: No. Copper is stable in non-oxidising acid without dissolved oxygen, but it can corrode if oxygen is present, because +1.23 V exceeds +0.34 V.

Quick check

1. Which cathodic reaction drives the rusting of steel in neutral, aerated water? Answer: The reduction of dissolved oxygen, O₂ + 2H₂O + 4e⁻ → 4OH⁻.

Exam focus

State that corrosion requires an anode, a cathode, an electrolyte and an electronic path. Use electrode potentials to decide feasibility and remember that thermodynamics predicts possibility, not rate. Be able to label immunity, corrosion and passivity regions on a Pourbaix diagram.

Advanced insight

Pourbaix diagrams are equilibrium maps: they ignore kinetics, assume particular ion activities and do not include complexing agents. Chloride, for example, forms soluble complexes and can shrink passive regions, while metastable oxides not on the diagram often control real behaviour. They are powerful starting points, not final answers.

Summary

Corrosion couples metal oxidation with a cathodic reduction, usually of oxygen or hydrogen ions, on the same surface. It is thermodynamically possible when the cathodic potential exceeds the metal's oxidation potential under the real conditions, which depend on pH and concentration via the Nernst equation. Pourbaix diagrams map immunity, corrosion and passivity against potential and pH.

Practice questions

1. Write the anodic and cathodic half-equations for iron corroding in acidic, oxygen-free water. Answer: Anode: Fe → Fe²⁺ + 2e⁻; cathode: 2H⁺ + 2e⁻ → H₂. 2. Calculate the equilibrium potential of the hydrogen electrode at pH 7 and 25 °C. Answer: E = −0.059 × 7 ≈ −0.41 V. 3. Why is gold found native in nature while iron is not? Answer: Gold's oxidation potential is higher than that of oxygen reduction under natural conditions, so no common oxidant can oxidise it; iron is easily oxidised. 4. In which region of a Pourbaix diagram would a metal be protected by a solid film? Answer: The passivity region, where an insoluble oxide or hydroxide is the stable species.