Corrosion
Quick Notes
- Corrosion is the gradual destruction of metals by chemical or electrochemical reaction with their environment.
Example: Rusting of iron is the most common form. - Corrosion typically occurs when a metal reacts with air, moisture, or acids.
- Rust is a hydrated form of ferric oxide: Fe₂O₃·xH₂O.
- Rusting is an electrochemical process involving anodic oxidation of Fe and cathodic reduction of O₂.
- At the anode: Fe → Fe²⁺ + 2e⁻
- At the cathode: O₂ + 4H⁺ + 4e⁻ → 2H₂O
- Overall reaction (in acidic medium): 2Fe + O₂ + 4H⁺ → 2Fe²⁺ + 2H₂O
- Fe²⁺ gets further oxidised to Fe³⁺ and forms rust.
- Moisture (H₂O), O₂, CO₂, and acidic conditions accelerate corrosion.
- Corrosion is prevented by:
- Barrier protection (painting, greasing)
- Alloying (e.g., stainless steel)
- Galvanisation (coating with zinc)
- Cathodic protection
Full Notes
Definition of Corrosion
Corrosion is a natural process where metals deteriorate due to chemical reactions with their environment. It typically involves the oxidation of metals by air (oxygen) and moisture (water), leading to the formation of oxides and other compounds.
A common example is the rusting of iron, which is the formation of iron oxides (like Fe2O3·xH2O). Other forms include the tarnishing of silver and the green coating on copper.
Chemical Basis of Corrosion
Corrosion can be considered an electrochemical process where the metal loses electrons and gets oxidised. It involves the following:
- Anode (oxidation site): The metal atoms lose electrons and form metal ions.
For iron: Fe(s) → Fe2+(aq) + 2e− - Cathode (reduction site): Oxygen is reduced by gaining electrons in the presence of hydrogen ions.
Reaction: O2(g) + 4H+(aq) + 4e− → 2H2O(l)
Overall redox reaction: 2Fe(s) + O2(g) + 4H+(aq) → 2Fe2+(aq) + 2H2O(l)
Standard electrode potentials: E°(Fe2+/Fe) = −0.44 V, E°(O2/H2O) = +1.23 V, E°(cell) = 1.67 V
Mechanism of Rust Formation
At a localised anode on the iron surface: Fe(s) → Fe2+ + 2e−
Electrons flow to the cathodic site, often another area of the same surface.
At the cathode, oxygen reacts with H+ (from H2CO3 or other acids in water) and the electrons: O2 + 4H+ + 4e− → 2H2O
The Fe2+ ions formed at the anode are further oxidised to Fe3+ by atmospheric oxygen, forming rust, mainly hydrated iron(III) oxide: Fe2O3·xH2O.
Prevention of Corrosion
Corrosion causes severe economic losses and structural damage, so its prevention is essential. Methods include:
- Barrier Protection: Coating the metal surface with paint, oil, grease, or chemicals to prevent exposure to air and moisture.
- Alloying: Using metals that do not corrode easily (e.g., stainless steel).
- Sacrificial Protection: Attaching a more reactive metal (like zinc or magnesium) that corrodes in place of the protected metal.
- Electrochemical Protection: Using an external power source or a sacrificial anode to force the protected metal to stay in its reduced state.
Summary
- Corrosion is an electrochemical process where metals oxidise in their environment.
- Rusting of iron involves anodic Fe oxidation and cathodic O2 reduction.
- Acidic conditions and moisture accelerate rusting of iron.
- Protection methods include barrier coatings, alloying, galvanisation, and cathodic protection.