Mechanism: corrosion is a galvanic cell

• Very pure iron does not rust easily.
• Industrial iron and steel contain impurities (for example carbon) and are alloys that are never perfectly uniform.
• Wherever a more active region (iron) touches a less active region (carbon or another metal) in the presence of moisture, a tiny local galvanic cell forms.
• The more active metal is the anode and is corroded; the less active one is the cathode.

For iron in damp air:

• Anode (the iron itself): 2Fe(s) → 2Fe²⁺(aq) + 4e⁻. The Fe²⁺ passes into the water film, which, containing dissolved ions, acts as the electrolyte; the electrons travel through the metal to the cathode regions (carbon impurities).
• Cathode (carbon/impurity regions): O₂(g) + 2H₂O(l) + 4e⁻ → 4OH⁻(aq)
• Ions meet: 2Fe²⁺ + 4OH⁻ → 2Fe(OH)₂(s) (iron(II) hydroxide)
• Further oxidation by dissolved oxygen: 2Fe(OH)₂ + ½O₂ + H₂O → 2Fe(OH)₃ (iron(III) hydroxide; after losing water it becomes the brown rust, Fe₂O₃·xH₂O)
• Overall: 2Fe + 3/2 O₂ + 3H₂O → 2Fe(OH)₃

Rusting is slow in pure water because very few ions are present, but it is much faster in sea water or salty water because the dissolved salts make the electrolyte a much better conductor.