Editing Rust (section)

== Chemical reactions ==
[[File:RustyChainEdit1.jpg|thumb|Heavy rust on the links of a chain near the [[Golden Gate Bridge]] in [[San Francisco]]; it was continuously exposed to moisture and [[salt spray]], causing surface breakdown, cracking, and flaking of the metal. The chain was replaced in 2023.]]
[[File:Trabajando una pieza de metal al rojo sobre un yunque.jpg|thumb|Rust scale forming and flaking off from a steel bar heated to its [[forging]] temperature of 1200°C. Rapid oxidation occurs when heated steel is exposed to air.]]

Rust is a general name for a complex of oxides and hydroxides of iron,<ref>{{Cite book|url=https://books.google.com/books?id=lH8GRXstqQEC&q=Rust+is+another+name+for+iron+oxide|title=Accent on science|last1=Sund|first1=Robert B.|last2=Bishop|first2=Jeanne|date=1980|publisher=C.E. Merrill|isbn=9780675075695|language=en|url-status=live|archive-url=https://web.archive.org/web/20171130151505/https://books.google.com/books?id=lH8GRXstqQEC&q=Rust+is+another+name+for+iron+oxide&dq=Rust+is+another+name+for+iron+oxide&hl=en&sa=X&ved=0ahUKEwiEp8Lzj8DUAhVM82MKHUJnBIsQ6AEILTAB|archive-date=2017-11-30}}</ref> which occur when iron or some alloys that contain iron are exposed to oxygen and moisture for a long period of time. Over time, the oxygen combines with the metal, forming new compounds collectively called rust, in a process called rusting.  Rusting is an [[redox|oxidation]] reaction specifically occurring with iron.  Other metals also corrode via similar oxidation, but such corrosion is not called rusting.

The main [[Catalysis|catalyst]] for the rusting process is water. Iron or steel structures might appear to be solid, but water molecules can penetrate the microscopic [[Pitting corrosion|pits]] and cracks in any exposed metal. The hydrogen atoms present in water molecules can combine with other elements to form acids, which will eventually cause more metal to be exposed. If chloride ions are present, as is the case with saltwater, the corrosion is likely to occur more quickly. Meanwhile, the oxygen atoms combine with metallic atoms to form the destructive oxide compound. These iron compounds are brittle and crumbly and replace strong metallic iron, reducing the strength of the object.

===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.

=== Associated reactions ===

The rusting of iron is an electrochemical process that begins with the transfer of [[electron]]s from iron to oxygen.<ref>{{Cite encyclopedia | last1 = Gräfen | first1 = H. | last2 = Horn | first2 = E. M. | last3 = Schlecker | first3 = H. | last4 = Schindler | first4 = H. | year = 2000 | chapter = Corrosion | encyclopedia = Ullmann's Encyclopedia of Industrial Chemistry| publisher= Wiley-VCH | doi = 10.1002/14356007.b01_08 | isbn = 3527306730 }}</ref> The iron is the reducing agent (gives up electrons) while the oxygen is the oxidizing agent (gains electrons). The rate of corrosion is affected by water and accelerated by [[electrolyte]]s, as illustrated by the effects of [[road salt]] on the corrosion of automobiles. The key reaction is the reduction of oxygen:
:O<sub>2</sub> + 4&nbsp;{{e-}} + 2&nbsp;{{H2O}} → 4&nbsp;{{OH-}}
Because it forms [[hydroxide]] [[ion]]s, this process is strongly affected by the presence of acid. Likewise, the corrosion of most metals by oxygen is accelerated at low [[pH]]. Providing the electrons for the above reaction is the oxidation of iron that may be described as follows:
:Fe → Fe<sup>2+</sup> + 2&nbsp;{{e-}}

The following [[redox reaction]] also occurs in the presence of water and is crucial to the formation of rust:
:4&nbsp;Fe<sup>2+</sup> + O<sub>2</sub> → 4&nbsp;Fe<sup>3+</sup> + 2&nbsp;O<sup>2−</sup>

In addition, the following multistep [[acid–base reaction]]s affect the course of rust formation:

:Fe<sup>2+</sup> + 2&nbsp;{{hsp}}H<sub>2</sub>O ⇌ Fe(OH)<sub>2</sub> + 2&nbsp;{{H+}}
:Fe<sup>3+</sup> + 3&nbsp;{{hsp}}H<sub>2</sub>O ⇌ Fe(OH)<sub>3</sub> + 3&nbsp;{{H+}}

as do the following [[Dehydration reaction|dehydration]] equilibria:
:[[Iron|Fe]](OH)<sub>2</sub> ⇌ FeO + {{H2O}}
:[[Iron|Fe]](OH)<sub>3</sub> ⇌ FeO(OH) + {{H2O}}
:2&nbsp;FeO(OH) ⇌ Fe<sub>2</sub>O<sub>3</sub> + {{H2O}}

From the above equations, it is also seen that the corrosion products are dictated by the availability of water and oxygen. With limited dissolved oxygen, iron(II)-containing materials are favoured, including [[iron(II) oxide|FeO]] and black [[lodestone]] or [[magnetite]] (Fe<sub>3</sub>O<sub>4</sub>). High oxygen concentrations favour [[ferric]] materials with the nominal formulae Fe(OH)<sub>3−''x''</sub>O<sub>{{frac|''x''|2}}</sub>. The nature of rust changes with time, reflecting the slow rates of the reactions of solids.<ref name="Bodner"/>

Furthermore, these complex processes are affected by the presence of other ions, such as [[calcium|Ca<sup>2+</sup>]], which serve as electrolytes which accelerate rust formation, or combine with the [[hydroxide]]s and [[oxide]]s of iron to precipitate a variety of Ca, Fe, O, OH species.

The onset of rusting can also be detected in the laboratory with the use of [[ferroxyl indicator solution]]. The solution detects both Fe<sup>2+</sup> ions and hydroxyl ions. Formation of Fe<sup>2+</sup> ions and hydroxyl ions are indicated by blue and pink patches respectively.