Corrosion Chemistry and Prevention Methods

Corrosion is the gradual destruction of a metal by an electrochemical reaction with its environment, and the rusting of iron is the best-known example: iron is oxidized at anodic areas (Fe → Fe²⁺ + 2e⁻) and oxygen is reduced at cathodic areas, in the presence of water. In brief: iron rusts because the reaction has E° = +0.84 V, and it can be prevented by coating the metal, by protecting it with a more active metal such as zinc, or by making it the cathode of an electric circuit.

Corrosion costs economies a large share of their output: estimates put the global cost at roughly 3–4 % of gross domestic product. It damages bridges, pipelines, ships and cars. Understanding the electrochemistry of corrosion shows how to prevent it. This article explains the mechanism of rusting, the factors that speed it up, the main methods of protection with a comparison table and a calculation for a sacrificial anode.

How does iron rust?

Rusting is an electrochemical process that takes place at the surface of the metal, which acts as a tiny galvanic cell. A drop of water on the iron surface provides the electrolyte. Regions of the surface differ in their composition, stress or oxygen supply, and so they become anodes and cathodes:

• Anode (oxidation): Fe(s) → Fe²⁺(aq) + 2e⁻, E°(reduction) = −0.44 V. This occurs in a pit or at a scratch, often where the oxygen concentration is lowest.
• Cathode (reduction): O₂(g) + 2H₂O(l) + 4e⁻ → 4OH⁻(aq), E° = +0.40 V. This takes place at areas with a better supply of oxygen. In acid solution, O₂ + 4H⁺ + 4e⁻ → 2H₂O (1.23 V) is even more favorable.
• Overall: 2Fe + O₂ + 2H₂O → 2Fe(OH)₂, with E°_cell = 0.40 − (−0.44) = +0.84 V.

The electrons flow through the metal from anode to cathode, and ions move through the water film. The Fe(OH)₂ is further oxidized by oxygen to Fe(OH)₃ and then to hydrated iron(III) oxide, Fe₂O₃·xH₂O, which is rust. Rust is porous and flaky, and so it does not protect the metal below from further attack. The rust forms where the Fe²⁺ meets oxygen, often away from the pit in which the iron dissolved.

What speeds up corrosion?

• Water and oxygen: both are required. Iron in dry air or in boiled, air-free water does not rust.
• Dissolved salts: ions make the water more conductive, so rust forms faster near the sea or on roads salted in winter.
• Acids: a low pH favors the reduction of oxygen and of H⁺.
• Contact with a less active metal: iron in contact with copper or tin corrodes faster, because the other metal acts as a cathode.
• Stress and surface defects: bent or scratched regions tend to become anodes.
• Temperature: as in most reactions, corrosion is faster when it is warm.

How can corrosion be prevented?

1. Barrier coatings: paint, plastic, enamel, oil or grease keep out water and oxygen. They fail if the coating is scratched.
2. Metallic coatings: electroplating or dipping in molten metal. Galvanizing is the coating of steel with zinc; tin plating is used for food cans.
3. Sacrificial protection: a more active metal (zinc, magnesium, aluminum) is connected to the iron and corrodes instead. It is the anode, and the iron is the cathode.
4. Cathodic protection by an impressed current: the structure is connected to the negative terminal of a DC power source, with an inert anode in the ground, which makes the whole structure a cathode. It is used for pipelines and storage tanks.
5. Alloying: stainless steel contains at least 10.5 % chromium, which forms a thin, self-healing layer of chromium oxide.
6. Passivation and inhibitors: chemicals added to the water or surface that slow the reactions, for example chromates, phosphates or nitrites.
7. Design: avoiding crevices where water can collect, and avoiding the contact of dissimilar metals.

Comparison of protection methods

Why is zinc better than tin?

Standard reduction potentials: Zn²⁺/Zn, −0.76 V; Fe²⁺/Fe, −0.44 V; Sn²⁺/Sn, −0.14 V. Zinc is more active than iron, so if the coating is scratched and iron is exposed, zinc becomes the anode and corrodes first, and iron stays protected as the cathode. In addition, zinc forms a protective layer of zinc carbonate. Tin is less active than iron, so a scratch exposes iron as the anode with a large cathodic area of tin around it, and the iron corrodes faster than if it had not been plated. For this reason tin cans must have an intact coating, which is often supplemented with lacquer.

How long does a sacrificial anode last?

Worked example

A pipeline is protected by a 10.0 kg magnesium anode that supplies an average current of 0.050 A. How long does it last, assuming 100 % efficiency?

1. n(Mg) = 10,000 g / 24.305 g/mol = 411.4 mol.
2. Mg → Mg²⁺ + 2e⁻, so n(e⁻) = 2 × 411.4 = 822.8 mol.
3. Q = 822.8 × 96,485 = 7.94 × 10⁷ C.
4. t = Q / I = 7.94 × 10⁷ / 0.050 = 1.59 × 10⁹ s, or about 50 years.

Real anodes are not fully utilized and some of the magnesium corrodes without producing useful current, so their life is shorter, perhaps 25 years or less, and they are replaced periodically.

What about other metals?

Aluminum is much more active than iron (E° = −1.66 V), but it resists corrosion in air because it forms a tightly bound, transparent layer of Al₂O₃ that stops further oxidation. Anodizing makes this layer thicker. The green patina on copper roofs and the tarnish on silver, which is Ag₂S, are other forms of corrosion that form protective or at least slow-growing films. Gold does not corrode, as its reduction potential is so high (+1.50 V for Au³⁺/Au).

How was corrosion explained?

Humphry Davy proposed cathodic protection in 1824 for the copper sheathing of Royal Navy ships, by attaching iron or zinc pieces to it. The idea worked, but the protected copper no longer prevented the growth of marine organisms, which slowed the ships. The electrochemical nature of corrosion was explained in the early 20th century, in the work of Ulick Evans, who showed in the 1920s that rust formation depends on separated anodic and cathodic areas, which he demonstrated with drops of water containing an indicator on iron.

Which mistakes are common?

• Thinking that rust forms only from oxygen: water is also needed.
• Believing that all coatings protect when scratched: tin does not.
• Mixing up the sacrificial metal and the protected metal: the more active metal is the anode.
• Assuming that rust protects the iron below, as aluminum oxide does: rust is porous.
• Forgetting that salt speeds up rusting: it makes the electrolyte more conductive.