Editing Rust (section)

===Oxidation of iron===

When iron is in contact with water and oxygen, it rusts.<ref name="Bodner">{{cite web|url=https://chemed.chem.purdue.edu/genchem/topicreview/bp/ch19/oxred_1.php|publisher=Bodner Research Web|access-date=28 April 2020|title=Oxidation Reduction Reactions}}</ref> If [[salt]] is present, for example in [[seawater]] or [[salt spray]], the iron tends to rust more quickly, as a result of chemical reactions. Iron metal is relatively unaffected by pure water or by dry oxygen. As with other metals, like aluminium, a tightly adhering oxide coating, a [[Passivation (chemistry)|passivation layer]], protects the bulk iron from further oxidation. The conversion of the passivating [[ferrous oxide]] layer to rust results from the combined action of two agents, usually oxygen and water.

Other degrading solutions are [[sulfur dioxide]] in water and [[carbon dioxide]] in water. Under these corrosive conditions, [[iron hydroxide]] species are formed. Unlike ferrous oxides, the hydroxides do not adhere to the bulk metal. As they form and flake off from the surface, fresh iron is exposed, and the corrosion process continues until either all of the iron is consumed or all of the oxygen, water, carbon dioxide or sulfur dioxide in the system are removed or consumed.<ref>{{cite book|last1=Holleman|first1=A. F.|last2=Wiberg|first2=E.|title=Inorganic Chemistry|publisher=Academic Press|location=San Diego|date=2001|isbn=0-12-352651-5}}</ref>

When iron rusts, the oxides take up more volume than the original metal; this expansion can generate enormous forces, damaging structures made with iron.  See ''[[#Economic effect|economic effect]]'' for more details.