Editing Iodine (section)

===Iodine oxides and oxoacids===
[[File:Iodine-pentoxide-3D-balls.png|thumb|right|upright=0.7|Structure of iodine pentoxide]]
[[Iodine oxide]]s are the most stable of all the halogen oxides, because of the strong I–O bonds resulting from the large electronegativity difference between iodine and oxygen, and they have been known for the longest time.<ref name="King" /> The stable, white, [[Hygroscopy|hygroscopic]] [[iodine pentoxide]] (I<sub>2</sub>O<sub>5</sub>) has been known since its formation in 1813 by Gay-Lussac and Davy. It is most easily made by the dehydration of [[iodic acid]] (HIO<sub>3</sub>), of which it is the anhydride. It will quickly oxidise carbon monoxide completely to [[carbon dioxide]] at room temperature, and is thus a useful reagent in determining carbon monoxide concentration. It also oxidises [[nitrogen oxide]], [[ethylene]], and [[hydrogen sulfide]]. It reacts with [[sulfur trioxide]] and peroxydisulfuryl difluoride (S<sub>2</sub>O<sub>6</sub>F<sub>2</sub>) to form salts of the iodyl cation, [IO<sub>2</sub>]<sup>+</sup>, and is reduced by concentrated [[sulfuric acid]] to iodosyl salts involving [IO]<sup>+</sup>. It may be fluorinated by [[fluorine]], [[bromine trifluoride]], [[sulfur tetrafluoride]], or [[chloryl fluoride]], resulting [[iodine pentafluoride]], which also reacts with [[iodine pentoxide]], giving iodine(V) oxyfluoride, IOF<sub>3</sub>. A few other less stable oxides are known, notably I<sub>4</sub>O<sub>9</sub> and I<sub>2</sub>O<sub>4</sub>; their structures have not been determined, but reasonable guesses are I<sup>III</sup>(I<sup>V</sup>O<sub>3</sub>)<sub>3</sub> and [IO]<sup>+</sup>[IO<sub>3</sub>]<sup>−</sup> respectively.<ref name="Greenwood851">Greenwood and Earnshaw, pp. 851–853</ref>

{| class="wikitable" style="float:right; width:25%;"
|+ Standard reduction potentials for aqueous I species<ref name="Greenwood853" />
! {{nowrap|E°(couple)}}!!{{nowrap|''a''(H<sup>+</sup>) {{=}} 1}}<br>(acid)!!{{nowrap|E°(couple)}}!!{{nowrap|''a''(OH<sup>−</sup>) {{=}} 1}}<br>(base)
|-
|I<sub>2</sub>/I<sup>−</sup>||+0.535|||I<sub>2</sub>/I<sup>−</sup>||+0.535
|-
|HOI/I<sup>−</sup>||+0.987||IO<sup>−</sup>/I<sup>−</sup>||+0.48
|-
|0||0||{{chem|IO|3|-}}/I<sup>−</sup>||+0.26
|-
|HOI/I<sub>2</sub>||+1.439||IO<sup>−</sup>/I<sub>2</sub>||+0.42
|-
|{{chem|IO|3|-}}/I<sub>2</sub>||+1.195||0||0
|-
|{{chem|IO|3|-}}/HOI||+1.134||{{chem|IO|3|-}}/IO<sup>−</sup>||+0.15
|-
|{{chem|IO|4|-}}/{{chem|IO|3|-}}||+1.653||0||0
|-
|H<sub>5</sub>IO<sub>6</sub>/{{chem|IO|3|-}}||+1.601||{{chem|H|3|IO|6|2-}}/{{chem|IO|3|-}}||+0.65
|}
More important are the four oxoacids: [[hypoiodous acid]] (HIO), [[Iodite|iodous acid]] (HIO<sub>2</sub>), [[iodic acid]] (HIO<sub>3</sub>), and [[periodic acid]] (HIO<sub>4</sub> or H<sub>5</sub>IO<sub>6</sub>). When iodine dissolves in aqueous solution, the following reactions occur:<ref name="Greenwood853">Greenwood and Earnshaw, pp. 853–9</ref>

{{block indent|{{wikitable|
|-
| I<sub>2</sub> + H<sub>2</sub>O || {{eqm}} HIO + H<sup>+</sup> + I<sup>−</sup> || ''K''<sub>ac</sub> = 2.0 × 10<sup>−13</sup> mol<sup>2</sup> L<sup>−2</sup>
|-
| I<sub>2</sub> + 2 OH<sup>−</sup> || {{eqm}} IO<sup>−</sup> + H<sub>2</sub>O + I<sup>−</sup> || ''K''<sub>alk</sub> {{=}} 30&nbsp;mol<sup>2</sup> L<sup>−2</sup>
}}}}

Hypoiodous acid is unstable to disproportionation. The hypoiodite ions thus formed disproportionate immediately to give iodide and iodate:<ref name="Greenwood853" />

{{block indent| 3 IO<sup>−</sup> {{eqm}} 2 I<sup>−</sup> + {{chem|IO|3|-}} ''K'' {{=}} 10<sup>20</sup>}}

Iodous acid and iodite are even less stable and exist only as a fleeting intermediate in the oxidation of iodide to iodate, if at all.<ref name="Greenwood853" /> Iodates are by far the most important of these compounds, which can be made by oxidising [[alkali metal]] iodides with oxygen at 600&nbsp;°C and high pressure, or by oxidising iodine with [[chlorate]]s. Unlike chlorates, which disproportionate very slowly to form chloride and perchlorate, iodates are stable to disproportionation in both acidic and alkaline solutions. From these, salts of most metals can be obtained. Iodic acid is most easily made by oxidation of an aqueous iodine suspension by [[electrolysis]] or fuming [[nitric acid]]. Iodate has the weakest oxidising power of the halates, but reacts the quickest.<ref name="Greenwood863">Greenwood and Earnshaw, pp. 863–4</ref>

Many periodates are known, including not only the expected tetrahedral {{chem|IO|4|-}}, but also square-pyramidal {{chem|IO|5|3-}}, octahedral orthoperiodate {{chem|IO|6|5-}}, [IO<sub>3</sub>(OH)<sub>3</sub>]<sup>2−</sup>, [I<sub>2</sub>O<sub>8</sub>(OH<sub>2</sub>)]<sup>4−</sup>, and {{chem|I|2|O|9|4-}}. They are usually made by oxidising alkaline [[sodium iodate]] electrochemically (with [[Lead dioxide|lead(IV) oxide]] as the anode) or by chlorine gas:<ref name="Greenwood872">Greenwood and Earnshaw, pp. 872–5</ref>

{{block indent|{{chem|IO|3|-}} + 6 OH<sup>−</sup> → {{chem|IO|6|5-}} + 3 H<sub>2</sub>O + 2 e<sup>−</sup>}}
{{block indent|{{chem|IO|3|-}} + 6 OH<sup>−</sup> + Cl<sub>2</sub> → {{chem|IO|6|5-}} + 2 Cl<sup>−</sup> + 3 H<sub>2</sub>O}}

They are thermodymically and kinetically powerful oxidising agents, quickly oxidising Mn<sup>2+</sup> to [[permanganate|{{chem|MnO|4|-}}]], and cleaving [[Diol|glycols]], α-[[Dicarbonyl|diketones]], α-[[Hydroxy ketone|ketols]], α-[[Alkanolamine|aminoalcohols]], and α-[[diamine]]s.<ref name="Greenwood872" /> Orthoperiodate especially stabilises high oxidation states among metals because of its very high negative charge of −5. [[Periodic acid|Orthoperiodic acid]], H<sub>5</sub>IO<sub>6</sub>, is stable, and dehydrates at 100&nbsp;°C in a vacuum to [[Periodic acid|Metaperiodic acid]], HIO<sub>4</sub>. Attempting to go further does not result in the nonexistent iodine heptoxide (I<sub>2</sub>O<sub>7</sub>), but rather iodine pentoxide and oxygen. Periodic acid may be protonated by [[sulfuric acid]] to give the {{chem|I(OH)|6|+}} cation, isoelectronic to Te(OH)<sub>6</sub> and {{chem|Sb(OH)|6|-}}, and giving salts with bisulfate and sulfate.<ref name="King" />