Corrosion is an everyday chemical process that connects directly to IB Chemistry. Whether it’s rust on iron, tarnish on silver, or the gradual deterioration of metals in the environment, corrosion is fundamentally a redox reaction. Understanding how corrosion occurs helps explain oxidation, electrochemical cells, and methods used to protect metals.
What Is Corrosion?
Corrosion is the gradual destruction of a metal through chemical reactions with its environment, usually involving oxidation.
In most cases:
- The metal loses electrons
- It forms metal ions
- New compounds form on the surface (such as oxides or hydroxides)
Corrosion weakens structures, damages equipment, and leads to economic loss. It is a natural, spontaneous process.
Rusting: The Most Common Form of Corrosion
Rusting is the corrosion of iron. It requires:
- Oxygen
- Water
- An electrolyte (such as salt) to speed up the process
The overall rusting reaction produces hydrated iron(III) oxide—commonly known as rust.
Key reactions involved:
Oxidation (at anode region):
Fe → Fe²⁺ + 2e⁻
Reduction (at cathode region):
O₂ + 2H₂O + 4e⁻ → 4OH⁻
These ions then combine to form rust.
Rust is flaky, porous, and does not protect the metal underneath, so corrosion continues.
Corrosion Is a Redox Process
Corrosion is fundamentally a redox reaction:
- The metal is oxidized (loses electrons).
- A substance in the environment—oxygen or hydrogen ions—is reduced.
