Corrosion
Corrosion is the gradual electrochemical degradation of metals through reactions with their environment, most commonly with oxygen and water. The rusting of iron is the most economically significant example, costing the global economy hundreds of billions of dollars annually in infrastructure damage, vehicle deterioration, and industrial equipment failure. Far from being a simple surface reaction, rust formation is a sophisticated electrochemical process involving spatially separated oxidation and reduction half-reactions linked by electron flow through the metal itself. Understanding corrosion requires applying the principles of redox chemistry and electrochemistry. A corroding iron surface acts like a short-circuited galvanic cell: anodic regions where iron is oxidised lose metal ions into solution, while cathodic regions nearby support the reduction of dissolved oxygen. The presence of electrolytes — particularly chloride ions from sea salt or road de-icers — dramatically accelerates the process by increasing solution conductivity and enabling faster ion transport. The final rust product, hydrated iron(III) oxide (Fe₂O₃·xH₂O), forms away from the original metal surface, which means that rust cannot act as a protective barrier the way that aluminium oxide does on aluminium. Because corrosion is electrochemical in nature, prevention strategies exploit electrochemical principles. Galvanisation applies a sacrificial zinc coating whose lower reduction potential ensures zinc corrodes preferentially, protecting the underlying iron even when the coating is scratched. Cathodic protection supplies electrons to the structure directly — either through an impressed current or a sacrificial anode — making the entire metal surface a cathode where reduction, not oxidation, occurs. This article examines the electrochemical mechanism of rusting in detail and explains how each major prevention method disrupts it.
- Formation of Anodic and Cathodic Sites
- Anodic Half-Reaction: Oxidation of Iron
- Cathodic Half-Reaction: Reduction of Oxygen
- Formation of Rust
- Role of Electrolytes and Chloride Ions
- Galvanisation as Prevention
- Cathodic Protection
Formation of Anodic and Cathodic Sites
Real iron surfaces are heterogeneous: differences in grain boundaries, surface stress, impurities, and local oxygen concentration create regions of slightly different electrical potential. Areas with lower potential (less noble) become anodes, while areas of higher potential become cathodes. Oxygen concentration cells form where iron in contact with low-O₂ water (e.g. under a water droplet centre) becomes anodic, while iron at the droplet edge — in contact with air-saturated water — becomes cathodic. This spatial separation is key: electrons flow through the metal from anode to cathode, just as in a galvanic cell.
Anodic Half-Reaction: Oxidation of Iron
At anodic sites, iron atoms lose electrons and enter solution as iron(II) ions: Fe(s) → Fe²⁺(aq) + 2e⁻. The standard reduction potential for this process is E° = −0.44 V (so the oxidation potential is +0.44 V). These electrons migrate through the metallic lattice toward cathodic sites. Iron(II) ions migrate through the electrolyte solution toward cathodic regions driven by the electric field. The rate of metal dissolution depends on the anodic current density, which increases with electrolyte conductivity and temperature.
Cathodic Half-Reaction: Reduction of Oxygen
In neutral to slightly basic conditions with dissolved oxygen present, the dominant cathodic reaction is the reduction of oxygen: O₂(g) + 2H₂O(l) + 4e⁻ → 4OH⁻(aq), with E° = +0.40 V. In acidic conditions (e.g. near industrial pollution or acid rain), hydrogen evolution can also occur: 2H⁺(aq) + 2e⁻ → H₂(g). The overall cell potential for the iron/oxygen cell is E°cell = E°cathode − E°anode = +0.40 − (−0.44) = +0.84 V, confirming the reaction is thermodynamically spontaneous (ΔG = −nFE°cell < 0).
Formation of Rust
Fe²⁺ ions migrating from anodic regions and OH⁻ ions from cathodic regions meet in solution to precipitate iron(II) hydroxide: Fe²⁺(aq) + 2OH⁻(aq) → Fe(OH)₂(s). This white-green solid is further oxidised by dissolved oxygen: 4Fe(OH)₂(s) + O₂(g) + 2H₂O(l) → 4Fe(OH)₃(s). Iron(III) hydroxide dehydrates over time to give the familiar reddish-brown rust: Fe(OH)₃ → FeO(OH) + H₂O, and ultimately Fe₂O₃·xH₂O. Unlike Al₂O₃ on aluminium, rust is porous, non-adherent, and non-passivating — it flakes off, exposing fresh iron to further attack.
Role of Electrolytes and Chloride Ions
Pure water corrodes iron slowly because its low ionic conductivity limits current flow between anodic and cathodic sites. Dissolved salts — especially NaCl from marine environments or road salt — provide ions that dramatically increase conductivity. Chloride ions are particularly aggressive: Cl⁻ can penetrate the thin oxide layer on iron, displacing oxygen ions in the lattice and forming soluble FeCl₂ complexes that prevent any passive film from forming. This is why coastal infrastructure and vehicles in cold climates corrode far faster than those in dry, clean environments.
Galvanisation as Prevention
Hot-dip galvanisation coats iron or steel with a layer of zinc (typically 45–85 μm thick). Zinc has a more negative standard reduction potential (E° = −0.76 V) than iron (E° = −0.44 V), making it the anode whenever both metals are in electrical contact in an electrolyte. Even if the zinc coating is scratched, the exposed iron remains cathodically protected: zinc preferentially oxidises (Zn → Zn²⁺ + 2e⁻) and electrons flow to iron, suppressing its oxidation. The zinc corrosion product, Zn(OH)₂, further converts to ZnCO₃ (zinc patina), which is relatively adherent and slows further zinc loss.
Cathodic Protection
Cathodic protection (CP) forces the entire structure to act as a cathode, eliminating anodic dissolution. Two methods exist. (1) Sacrificial anode CP: a block of a more active metal — zinc, magnesium (E° = −2.37 V), or aluminium alloy — is bolted directly to the structure. This metal becomes the anode and corrodes instead of the iron. Used on ship hulls, buried pipelines, and offshore rigs; anodes must be periodically replaced. (2) Impressed current CP (ICCP): a DC power supply drives current from an inert anode (e.g. platinum-coated titanium or graphite) through the soil or water to the structure, making it cathodic. ICCP is used for long pipelines and large structures where sacrificial anodes would be impractical. Both methods rely on the same electrochemical principle: supply sufficient electrons to the structure so that its potential is shifted below the equilibrium potential for iron oxidation.
The Key Equation
- \ce{Fe} — Iron (—): The metal being oxidised from 0 to +3 oxidation state; acts as the reducing agent.
- \ce{O2} — Oxygen (—): The oxidising agent; reduced from 0 to −2 oxidation state at cathodic sites.
- \ce{H2O} — Water (—): Acts as the electrolyte medium enabling ion transport between anode and cathode; participates in the cathodic reduction reaction.
- \ce{Fe2O3 * xH2O} — Hydrated Iron(III) Oxide (Rust) (—): The final corrosion product; x is variable (typically 1–3), giving rust its variable composition. It is porous and non-protective.
Worked Example
Problem. A steel pipeline buried in moist, saline soil is to be protected using magnesium sacrificial anodes. (a) Write the half-reactions and the overall cell reaction. (b) Calculate E°cell. (c) Explain why the steel is protected even where it is not in direct contact with the magnesium block. Given: E°(Fe²⁺/Fe) = −0.44 V; E°(Mg²⁺/Mg) = −2.37 V.
Answer. E°cell = +2.77 V; the pipeline is cathodically protected because electrons supplied by Mg oxidation flow to the steel, making it the cathode where only reduction can occur.
Origins · 1 minSir Humphry Davy and how this was discovered
Davy was one of the first scientists to apply electrochemical principles to the problem of corrosion. In 1824, commissioned by the Royal Navy, he demonstrated that attaching blocks of zinc or iron to the copper sheeting on warship hulls made the copper cathodic, preventing its dissolution into seawater — the first deliberate application of cathodic protection. His experiments established that the driving force for corrosion was electrochemical in nature, laying the groundwork for all modern corrosion prevention strategies. Davy's cathodic protection on navy ships worked so well at stopping copper corrosion that it also prevented the natural release of toxic copper ions that had kept ships' hulls free of barnacles and algae — leading to heavily fouled hulls and slower vessels. The navy ultimately abandoned his scheme because a clean hull was less strategically useful than a fast one!
Related Concepts
- Galvanic Cells — Corrosion is fundamentally a short-circuited galvanic cell; understanding how galvanic cells generate EMF from spontaneous redox reactions explains why anodic and cathodic sites form on a metal surface.
- Standard Electrode Potentials — The relative E° values of metals determine which acts as anode and which as cathode in a corrosion cell, and predict whether galvanisation or sacrificial anodes will provide protection.
- Electrolysis — Impressed current cathodic protection (ICCP) is essentially forced electrolysis: an external power supply drives current to make the structure cathodic, directly applying electrolytic principles.
- Oxidation States and Redox — Tracking changes in oxidation state of iron (0 → +2 → +3) and oxygen (0 → −2) through the rusting process is essential for writing and balancing the half-equations involved.