Corrosion and Rusting

Introduction to Corrosion

Definition: Corrosion refers to the process by which metals break down by reacting with substances in their environment.
Example: Rusting of iron involves the reaction with oxygen and water to form hydrated iron oxide (rust).

Atomic Level Reaction:
- Iron atoms lose three electrons → become iron ions (oxidation).
- Oxygen molecules gain electrons → become oxygen ions (reduction).
Type of Reaction: The rusting process is a redox reaction (oxidation + reduction).

Essential Elements: Both oxygen and water must be present for rusting to occur.
Experiment with Test Tubes:
1. Test Tube 1: Contains water (boiled to remove oxygen) → No rusting.
2. Test Tube 2: Contains air (oxygen) but no water (calcium chloride used) → No rusting.
3. Test Tube 3: Contains both air and water → Rusting occurs.

Surface Corrosion: Corrosion affects only the surface of metals due to exposure to environmental substances.
Rust Flaking: Rust formed at the surface flakes off, exposing new metal underneath, leading to further corrosion.
Comparison with Aluminium: In aluminium, a protective layer (aluminium oxide) forms, preventing further corrosion of the underlying metal.

Barrier Methods: Prevent oxygen and water from contacting iron.
- Painting: Effective for large iron structures (e.g., ships).
- Oiling/Greasings: Necessary for moving parts (e.g., bike chains).
- Electroplating: Uses electrolysis to coat iron with another metal, providing protection.
Sacrificial Methods: Involves using a more reactive metal to protect iron.
- Example: Adding aluminium or zinc to iron so they are oxidized instead of iron.
- Galvanising: Coating iron with zinc acts as both barrier and sacrificial protection; if scratched, zinc protects iron from rusting.

Understanding corrosion and rusting is essential for protecting metal structures and components.
Methods for prevention are key to extending the life of iron objects.