Editing Iron (section)

===Solution chemistry===
[[File:Ferrate_and_permanganate_solution.jpg|thumb|upright=0.7|right|Comparison of colors of solutions of ferrate (left) and [[permanganate]] (right)]]
The [[standard reduction potential]]s in acidic aqueous solution for some common iron ions are given below:{{sfn|Greenwood|Earnshaw|1997|pp=1075–79}}
:{|
|-
| [Fe(H<sub>2</sub>O)<sub>6</sub>]<sup>2+</sup> + 2 e<sup>−</sup>|| {{eqm}} Fe || E<sup>0</sup> = −0.447 V
|-
| [Fe(H<sub>2</sub>O)<sub>6</sub>]<sup>3+</sup> + e<sup>−</sup>|| {{eqm}} [Fe(H<sub>2</sub>O)<sub>6</sub>]<sup>2+</sup> || E<sup>0</sup> = +0.77 V
|-
| {{chem|FeO|4|2-}} + 8 H<sub>3</sub>O<sup>+</sup> + 3 e<sup>−</sup>|| {{eqm}} [Fe(H<sub>2</sub>O)<sub>6</sub>]<sup>3+</sup> + 6 H<sub>2</sub>O || E<sup>0</sup> = +2.20 V
|}

The red-purple tetrahedral [[ferrate]](VI) anion is such a strong oxidizing agent that it oxidizes ammonia to nitrogen (N<sub>2</sub>) and water to oxygen:{{sfn|Greenwood|Earnshaw|1997|pp=1082–84}}
:4 {{chem|FeO|4|2-}} + 34 {{chem|H|2|O}} → 4 {{chem2|[Fe(H2O)6](3+)}} + 20 {{chem|OH|-}} + 3 O<sub>2</sub>

The pale-violet hex[[aquo complex]] {{chem2|[Fe(H2O)6](3+)}} is an acid such that above pH 0 it is fully hydrolyzed:{{sfn|Greenwood|Earnshaw|1997|pp=1088–91}}

:{|
|-
| {{chem2|[Fe(H2O)6](3+)}} || {{eqm}} {{chem2|[Fe(H2O)5(OH)](2+) + H(+)}} || ''[[equilibrium constant|K]]'' = 10<sup>−3.05</sup> mol dm<sup>−3</sup>
|-
| {{chem2|[Fe(H2O)5(OH)](2+)}} || {{eqm}} {{chem2|[Fe(H2O)4(OH)2](+) + H(+)}} || ''K'' = 10<sup>−3.26</sup> mol dm<sup>−3</sup>
|-
| {{chem2|2[Fe(H2O)6](3+)}} || {{eqm}} {{chem2|[Fe(H2O)4(OH)]2(4+) + 2H(+) + 2H2O}} || ''K'' = 10<sup>−2.91</sup> mol dm<sup>−3</sup>
|}

[[File:Iron(II)-sulfate-heptahydrate-sample.jpg|thumb|right|upright=0.7|Blue-green [[iron(II) sulfate]] heptahydrate]]

As pH rises above 0 the above yellow hydrolyzed species form and as it rises above 2–3, reddish-brown hydrous [[iron(III) oxide]] precipitates out of solution. Although Fe<sup>3+</sup> has a d<sup>5</sup> configuration, its absorption spectrum is not like that of Mn<sup>2+</sup> with its weak, spin-forbidden d–d bands, because Fe<sup>3+</sup> has higher positive charge and is more polarizing, lowering the energy of its ligand-to-metal [[charge-transfer complex|charge transfer]] absorptions. Thus, all the above complexes are rather strongly colored, with the single exception of the hexaquo ion – and even that has a spectrum dominated by charge transfer in the near ultraviolet region.{{sfn|Greenwood|Earnshaw|1997|pp=1088–91}} On the other hand, the pale green iron(II) hexaquo ion {{chem2|[Fe(H2O)6](2+)}} does not undergo appreciable hydrolysis. Carbon dioxide is not evolved when [[carbonate]] anions are added, which instead results in white [[iron(II) carbonate]] being precipitated out. In excess carbon dioxide this forms the slightly soluble bicarbonate, which occurs commonly in groundwater, but it oxidises quickly in air to form [[iron(III) oxide]] that accounts for the brown deposits present in a sizeable number of streams.{{sfn|Greenwood|Earnshaw|1997|pp=1091–97}}