Corrosion is the gradual destruction of a metal by chemical reactions with its environment, and for iron or steel it often appears as rust. It matters because bridges, pipelines, ships, cars, and buried tanks can lose strength when metal atoms are converted into ions and oxides. In most real settings, corrosion is electrochemical, meaning oxidation and reduction happen at different spots on the metal surface.
Water, dissolved oxygen, salts, and differences in stress or composition can turn one region into an anode and another into a cathode.
In rusting iron, Fe atoms at the anode lose electrons and become Fe2+ ions, while oxygen dissolved in water is reduced at the cathode. The ions and products then react further to form hydrated iron(III) oxide, the reddish-brown material commonly called rust. Cathodic protection slows corrosion by forcing the iron or steel to act as the cathode, either by attaching a more reactive sacrificial metal or by using an external power supply.
Coatings and galvanizing help by blocking water and oxygen, and zinc galvanizing also protects exposed scratches because zinc oxidizes more readily than iron.
Key Facts
- Anodic oxidation of iron: Fe(s) -> Fe2+(aq) + 2e-
- Cathodic reduction in neutral water: O2(g) + 2H2O(l) + 4e- -> 4OH-(aq)
- Overall early corrosion reaction: 2Fe(s) + O2(g) + 2H2O(l) -> 2Fe(OH)2(s)
- Rust is mainly hydrated iron(III) oxide, often written as Fe2O3·xH2O.
- A sacrificial anode protects steel when E° of the anode metal is more negative than E° of Fe2+/Fe.
- Faraday relationship for metal loss or plating: n = Q/(zF), where Q = It and F = 96485 C/mol e-.
Vocabulary
- Corrosion
- Corrosion is the chemical or electrochemical breakdown of a material, usually a metal, due to reactions with its surroundings.
- Anode
- The anode is the electrode or metal region where oxidation occurs and electrons are produced.
- Cathode
- The cathode is the electrode or metal region where reduction occurs and electrons are consumed.
- Sacrificial anode
- A sacrificial anode is a more reactive metal attached to a structure so it corrodes instead of the protected metal.
- Galvanizing
- Galvanizing is the process of coating iron or steel with zinc to protect it from corrosion.
Common Mistakes to Avoid
- Calling rust pure Fe2O3, which is wrong because common rust is a mixture that includes hydrated iron(III) oxides and hydroxides.
- Thinking corrosion only happens where the metal touches air, which is wrong because water or moist soil can carry ions and allow electrochemical cells to form below the surface.
- Labeling the protected steel as the anode in cathodic protection, which is wrong because protection works by making the steel the cathode so it does not oxidize.
- Assuming paint always stops corrosion permanently, which is wrong because scratches or pores can let water and oxygen reach the metal and create small intense corrosion sites.
Practice Questions
- 1 A steel tank is protected by a magnesium sacrificial anode. If a corrosion current of 0.50 A flows for 24.0 h and Mg oxidizes as Mg -> Mg2+ + 2e-, how many moles of Mg are consumed? Use F = 96485 C/mol e-.
- 2 For the reaction Fe(s) -> Fe2+(aq) + 2e-, how many grams of iron would be lost if 19300 C of charge passed through an unprotected corrosion cell? Use molar mass Fe = 55.85 g/mol and F = 96485 C/mol e-.
- 3 A galvanized steel sheet is scratched so that both zinc and iron are exposed to rainwater. Explain which metal acts as the anode and why the iron can still be protected.