Editing Rust (section)

=== 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.