Rusting of Iron

In-depth explanation of Rusting of Iron

Rusting of iron is an electrochemical process in which iron oxidizes at anodic sites and oxygen is reduced at cathodic sites in the presence of water. At the anode, Fe -> Fe2+ + 2e-. At the cathode, O2 + 2H2O + 4e- -> 4OH-. The iron(II) ions are further oxidized by oxygen to iron(III) and form hydrated iron(III) oxide, the reddish rust. Salt water speeds the process by increasing the conductivity of the electrolyte. Rust is porous and does not protect the metal, so corrosion continues.

Going deeper: Rusting of Iron

Key ideas: Anode: Fe giving Fe2+ + 2e- (at pits). Cathode: O2 + 2H2O + 4e- giving 4OH- on the surface. Fe2+ + 2OH- gives Fe(OH)2, then oxygen oxidizes it to Fe(OH)3, which loses water to form rust Fe2O3 times xH2O. Needs both water and O2; salt water speeds it by raising conductivity.

Common mistakes and connections: Mistakes: stating rust forms by direct reaction of iron with O2 gas; it is electrochemical. Rust forms away from the pit where Fe2+ originates since ions diffuse. Rust is porous and flakes off, unlike aluminum oxide, so protection fails. Hence painting and galvanizing are used.

Visualization of Rusting of Iron

Rust forms when Fe2+ is further oxidized by oxygen to hydrated Fe2O3; salt in the water raises conductivity and speeds the cell up. The interactive lesson shows this as a live animation or chart.

Worked example and use case

Cars corrode faster near the sea because salt spray raises the conductivity of the surface water film.

Where Rusting of Iron fits in Acids, Bases & Electrochemistry

Rusting of Iron is topic 76 of 105 in Acids, Bases & Electrochemistry.

Related topics

• Fuel CellsIntermediate
• Solid-State BatteriesAdvanced
• CorrosionBeginner
• Prevention of CorrosionIntermediate
• Sacrificial AnodeAdvanced
• GalvanizationIntermediate

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