Editing Iodine (section)

== Chemistry and compounds ==
{{Main|Iodine compounds}}
{| class="wikitable" style="float:right; width:25%;"
|+ style="margin-bottom: 5px;" | Halogen bond energies (kJ/mol)<ref name="Greenwood804" />
|-
! X
! XX
! HX
! BX<sub>3</sub>
! AlX<sub>3</sub>
! CX<sub>4</sub>
|-
! F
| 159
| 574
| 645
| 582
| 456
|-
! Cl
|243
|428
|444
|427
|327
|-
! Br
|193
|363
|368
|360
|272
|-
! I
|151
|294
|272
|285
|239
|}
Iodine is quite reactive, but it is less so than the lighter halogens, and it is a weaker oxidant. For example, it does not [[Halogenation|halogenate]] [[carbon monoxide]], [[nitric oxide]], and [[sulfur dioxide]], which [[chlorine]] does. Many metals react with iodine.<ref name="Greenwood800" /> By the same token, however, since iodine has the lowest ionisation energy among the halogens and is the most easily oxidised of them, it has a more significant cationic chemistry and its higher oxidation states are rather more stable than those of bromine and chlorine, for example in [[iodine heptafluoride]].<ref name="Greenwood804" />

===Charge-transfer complexes ===
[[File:Iodine-triphenylphosphine charge-transfer complex in dichloromethane.jpg|thumb|upright=1.8|right|I<sub>2</sub>•[[triphenylphosphine|PPh<sub>3</sub>]] charge-transfer complexes in [[dichloromethane|CH<sub>2</sub>Cl<sub>2</sub>]]. From left to right: (1) I<sub>2</sub> dissolved in dichloromethane – no CT complex. (2) A few seconds after excess PPh<sub>3</sub> was added – CT complex is forming. (3) One minute later after excess PPh<sub>3</sub> was added, the CT complex [Ph<sub>3</sub>PI]<sup>+</sup>I<sup>−</sup> has been formed. (4) Immediately after excess I<sub>2</sub> was added, which contains [Ph<sub>3</sub>PI]<sup>+</sup>[I<sub>3</sub>]<sup>−</sup>.<ref name="InorgChem">{{Housecroft3rd|page=541}}</ref>]]

The iodine molecule, I<sub>2</sub>, dissolves in CCl<sub>4</sub> and aliphatic hydrocarbons to give bright violet solutions. In these solvents the absorption band maximum occurs in the 520 &ndash; 540&nbsp;nm region and is assigned to a {{pi}}<sup>*</sup> to ''σ''<sup>*</sup> transition.   When I<sub>2</sub> reacts with Lewis bases in these solvents a blue shift in I<sub>2</sub> peak is seen and the new peak (230 &ndash; 330&nbsp;nm) arises that is due to the formation of adducts, which are referred to as charge-transfer complexes.<ref name="Greenwood806">Greenwood and Earnshaw, pp. 806–07</ref>

===Hydrogen iodide===
The simplest compound of iodine is [[hydrogen iodide]], HI. It is a colourless gas that reacts with oxygen to give water and iodine. Although it is useful in [[Halogenation|iodination]] reactions in the laboratory, it does not have large-scale industrial uses, unlike the other hydrogen halides. Commercially, it is usually made by reacting iodine with [[hydrogen sulfide]] or [[hydrazine]]:<ref name="Greenwood809">Greenwood and Earnshaw, pp. 809–812</ref>

:2 I<sub>2</sub> + N<sub>2</sub>H<sub>4</sub> {{overset|H<sub>2</sub>O|⟶}} 4 HI + N<sub>2</sub>

At room temperature, it is a colourless gas, like all of the hydrogen halides except [[hydrogen fluoride]], since hydrogen cannot form strong [[hydrogen bond]]s to the large and only mildly electronegative iodine atom. It melts at {{convert|−51.0|°C}} and boils at {{convert|−35.1|°C}}. It is an [[Endothermic process|endothermic]] compound that can exothermically dissociate at room temperature, although the process is very slow unless a [[Catalysis|catalyst]] is present: the reaction between hydrogen and iodine at room temperature to give hydrogen iodide does not proceed to completion. The H–I [[Bond-dissociation energy|bond dissociation energy]] is likewise the smallest of the hydrogen halides, at 295&nbsp;kJ/mol.<ref name="Greenwood812">Greenwood and Earnshaw, pp. 812–819</ref>

Aqueous hydrogen iodide is known as [[hydroiodic acid]], which is a strong acid. Hydrogen iodide is exceptionally soluble in water: one litre of water will dissolve 425 litres of hydrogen iodide, and the saturated solution has only four water molecules per molecule of hydrogen iodide.<ref>{{Cite book |last1=Holleman |first1=A. F. |title=Inorganic Chemistry |last2=Wiberg |first2=E. |publisher=Academic Press |year=2001 |isbn=0-12-352651-5 |location=San Diego}}</ref> Commercial so-called "concentrated" hydroiodic acid usually contains 48–57% HI by mass; the solution forms an [[azeotrope]] with boiling point {{convert|126.7|°C}} at 56.7&nbsp;g HI per 100&nbsp;g solution. Hence hydroiodic acid cannot be concentrated past this point by evaporation of water.<ref name="Greenwood812" /> Unlike gaseous hydrogen iodide, hydroiodic acid has major industrial use in the manufacture of [[acetic acid]] by the [[Cativa process]].<ref name="Cativa">{{Cite journal |last=Jones |first=J. H. |year=2000 |title=The Cativa Process for the Manufacture of Acetic Acid |url=http://www.platinummetalsreview.com/pdf/pmr-v44-i3-094-105.pdf |url-status=live |journal=[[Platinum Metals Review]] |volume=44 |issue=3 |pages=94–105 |doi=10.1595/003214000X44394105 |archive-url=https://web.archive.org/web/20150924074441/http://www.platinummetalsreview.com/pdf/pmr-v44-i3-094-105.pdf |archive-date=24 September 2015 |access-date=26 August 2023}}</ref><ref>{{Cite journal |last1=Sunley |first1=G. J. |last2=Watsonv |first2=D. J. |year=2000 |title=High productivity methanol carbonylation catalysis using iridium – The Cativa process for the manufacture of acetic acid |journal=Catalysis Today |volume=58 |issue=4 |pages=293–307 |doi=10.1016/S0920-5861(00)00263-7}}</ref>

===Other binary iodine compounds===
With the exception of the [[Noble gas|noble gases]], nearly all elements on the periodic table up to einsteinium ([[Einsteinium(III) iodide|EsI<sub>3</sub>]] is known) are known to form binary compounds with iodine. Until 1990, [[nitrogen triiodide]]<ref>The ammonia adduct NI<sub>3</sub>•NH<sub>3</sub> is more stable and can be isolated at room temperature as a notoriously shock-sensitive black solid.</ref> was only known as an ammonia adduct. Ammonia-free NI<sub>3</sub> was found to be isolable at –196 °C but spontaneously decomposes at 0 °C.<ref>{{cite journal |last1=Tornieporth-Oetting |first1=Inis |last2=Klapötke |first2=Thomas |date=June 1990 |title=Nitrogen Triiodide |url=https://onlinelibrary.wiley.com/doi/10.1002/anie.199006771 |journal=Angewandte Chemie |edition=international |language=en |volume=29 |issue=6 |pages=677–679 |doi=10.1002/anie.199006771 |issn=0570-0833 |access-date=5 March 2023 |archive-date=5 March 2023 |archive-url=https://web.archive.org/web/20230305194218/https://onlinelibrary.wiley.com/doi/10.1002/anie.199006771 |url-status=live }}</ref> For thermodynamic reasons related to electronegativity of the elements, neutral sulfur and selenium iodides that are stable at room temperature are also nonexistent, although S<sub>2</sub>I<sub>2</sub> and SI<sub>2</sub> are stable up to 183 and 9 K, respectively.  As of 2022, no neutral binary selenium iodide has been unambiguously identified (at any temperature).<ref>{{cite journal |last=Vilarrubias |first=Pere |date=17 November 2022 |title=The elusive diiodosulphanes and diiodoselenanes |url=https://doi.org/10.1080/00268976.2022.2129106 |journal=Molecular Physics |volume=120 |issue=22 |pages=e2129106 |doi=10.1080/00268976.2022.2129106 |bibcode=2022MolPh.12029106V |s2cid=252744393 |issn=0026-8976 |access-date=5 March 2023 |archive-date=19 March 2024 |archive-url=https://web.archive.org/web/20240319070247/https://www.tandfonline.com/pb/css/t1709911000430-v1707891316000/head_4_698_en.css |url-status=live }}</ref> Sulfur-iodine and selenium-iodine polyatomic cations (e.g., [S<sub>2</sub>I<sub>4</sub><sup>2+</sup>][AsF<sub>6</sub><sup>–</sup>]<sub>2</sub> and [Se<sub>2</sub>I<sub>4</sub><sup>2+</sup>][Sb<sub>2</sub>F<sub>11</sub><sup>–</sup>]<sub>2</sub>) have been prepared and characterised crystallographically.<ref>{{cite journal |last1=Klapoetke |first1=T. |last2=Passmore |first2=J. |date=1 July 1989 |title=Sulfur and selenium iodine compounds: from non-existence to significance |url=https://pubs.acs.org/doi/abs/10.1021/ar00163a002 |journal=Accounts of Chemical Research |language=en |volume=22 |issue=7 |pages=234–240 |doi=10.1021/ar00163a002 |issn=0001-4842 |access-date=15 January 2023 |archive-date=15 January 2023 |archive-url=https://web.archive.org/web/20230115160630/https://pubs.acs.org/doi/abs/10.1021/ar00163a002 |url-status=live }}</ref>

Given the large size of the iodide anion and iodine's weak oxidising power, high oxidation states are difficult to achieve in binary iodides, the maximum known being in the pentaiodides of [[niobium]], [[tantalum]], and [[protactinium]]. Iodides can be made by reaction of an element or its oxide, hydroxide, or carbonate with hydroiodic acid, and then dehydrated by mildly high temperatures combined with either low pressure or anhydrous hydrogen iodide gas. These methods work best when the iodide product is stable to hydrolysis. Other syntheses include high-temperature oxidative iodination of the element with iodine or hydrogen iodide, high-temperature iodination of a metal oxide or other halide by iodine, a volatile metal halide, [[carbon tetraiodide]], or an organic iodide. For example, [[Molybdenum dioxide|molybdenum(IV) oxide]] reacts with [[Aluminium iodide|aluminium(III) iodide]] at 230&nbsp;°C to give [[molybdenum(II) iodide]]. An example involving halogen exchange is given below, involving the reaction of [[tantalum(V) chloride]] with excess aluminium(III) iodide at 400&nbsp;°C to give [[tantalum(V) iodide]]:<ref name="Greenwood821">Greenwood and Earnshaw, pp. 821–4</ref>

<chem display="block">3TaCl5 + \underset{(excess)}{5AlI3} -> 3TaI5 + 5AlCl3</chem>

Lower iodides may be produced either through thermal decomposition or disproportionation, or by reducing the higher iodide with hydrogen or a metal, for example:<ref name="Greenwood821" />

<chem display="block">TaI5{} + Ta ->[\text{thermal gradient}] [\ce{630^\circ C\ ->\ 575^\circ C}] Ta6I14</chem>

Most metal iodides with the metal in low oxidation states (+1 to +3) are ionic. Nonmetals tend to form covalent molecular iodides, as do metals in high oxidation states from +3 and above. Both ionic and covalent iodides are known for metals in oxidation state +3 (e.g. [[Scandium triiodide|scandium iodide]] is mostly ionic, but [[aluminium iodide]] is not). Ionic iodides MI<sub>''n''</sub> tend to have the lowest melting and boiling points among the halides MX<sub>''n''</sub> of the same element, because the electrostatic forces of attraction between the cations and anions are weakest for the large iodide anion. In contrast, covalent iodides tend to instead have the highest melting and boiling points among the halides of the same element, since iodine is the most polarisable of the halogens and, having the most electrons among them, can contribute the most to van der Waals forces. Naturally, exceptions abound in intermediate iodides where one trend gives way to the other. Similarly, solubilities in water of predominantly ionic iodides (e.g. [[potassium]] and [[calcium]]) are the greatest among ionic halides of that element, while those of covalent iodides (e.g. [[silver]]) are the lowest of that element. In particular, [[silver iodide]] is very insoluble in water and its formation is often used as a qualitative test for iodine.<ref name="Greenwood821" />

===Iodine halides===
The halogens form many binary, [[Diamagnetism|diamagnetic]] [[interhalogen]] compounds with stoichiometries XY, XY<sub>3</sub>, XY<sub>5</sub>, and XY<sub>7</sub> (where X is heavier than Y), and iodine is no exception. Iodine forms all three possible diatomic interhalogens, a trifluoride and trichloride, as well as a pentafluoride and, exceptionally among the halogens, a heptafluoride. Numerous cationic and anionic derivatives are also characterised, such as the wine-red or bright orange compounds of {{chem|ICl|2|+}} and the dark brown or purplish black compounds of I<sub>2</sub>Cl<sup>+</sup>. Apart from these, some [[pseudohalogen|pseudohalides]] are also known, such as [[cyanogen iodide]] (ICN), iodine [[thiocyanate]] (ISCN), and iodine [[azide]] (IN<sub>3</sub>).<ref name="Greenwood824">Greenwood and Earnshaw, pp. 824–828</ref>

[[File:Iodine monochloride1.jpg|thumb|right|Iodine monochloride]]
[[Iodine monofluoride]] (IF) is unstable at room temperature and disproportionates very readily and irreversibly to iodine and [[iodine pentafluoride]], and thus cannot be obtained pure. It can be synthesised from the reaction of iodine with fluorine gas in [[trichlorofluoromethane]] at −45&nbsp;°C, with [[iodine trifluoride]] in trichlorofluoromethane at −78&nbsp;°C, or with [[silver(I) fluoride]] at 0&nbsp;°C.<ref name="Greenwood824" /> [[Iodine monochloride]] (ICl) and [[iodine monobromide]] (IBr), on the other hand, are moderately stable. The former, a volatile red-brown compound, was discovered independently by [[Joseph Louis Gay-Lussac]] and [[Humphry Davy]] in 1813–1814 not long after the discoveries of chlorine and iodine, and it mimics the intermediate halogen bromine so well that [[Justus von Liebig]] was misled into mistaking bromine (which he had found) for iodine monochloride. Iodine monochloride and iodine monobromide may be prepared simply by reacting iodine with chlorine or bromine at room temperature and purified by [[fractional crystallization (chemistry)|fractional crystallisation]]. Both are quite reactive and attack even [[platinum]] and [[gold]], though not [[boron]], [[carbon]], [[cadmium]], [[lead]], [[zirconium]], [[niobium]], [[molybdenum]], and [[tungsten]]. Their reaction with organic compounds depends on conditions. Iodine chloride vapour tends to chlorinate [[phenol]] and [[salicylic acid]], since when iodine chloride undergoes [[Homolysis (chemistry)|homolytic fission]], chlorine and iodine are produced and the former is more reactive. However, iodine chloride in [[carbon tetrachloride]] solution results in iodination being the main reaction, since now [[Heterolysis (chemistry)|heterolytic fission]] of the I–Cl bond occurs and I<sup>+</sup> attacks phenol as an electrophile. However, iodine monobromide tends to brominate phenol even in carbon tetrachloride solution because it tends to dissociate into its elements in solution, and bromine is more reactive than iodine.<ref name="Greenwood824" /> When liquid, iodine monochloride and iodine monobromide dissociate into {{chem|I|2|X|+}} and {{chem|IX|2|-}} ions (X = Cl, Br); thus they are significant conductors of electricity and can be used as ionising solvents.<ref name="Greenwood824" />

[[Iodine trifluoride]] (IF<sub>3</sub>) is an unstable yellow solid that decomposes above −28&nbsp;°C. It is thus little-known. It is difficult to produce because fluorine gas would tend to oxidise iodine all the way to the pentafluoride; reaction at low temperature with [[xenon difluoride]] is necessary. [[Iodine trichloride]], which exists in the solid state as the planar dimer I<sub>2</sub>Cl<sub>6</sub>, is a bright yellow solid, synthesised by reacting iodine with liquid chlorine at −80&nbsp;°C; caution is necessary during purification because it easily dissociates to iodine monochloride and chlorine and hence can act as a strong chlorinating agent. Liquid iodine trichloride conducts electricity, possibly indicating dissociation to {{chem|ICl|2|+}} and {{chem|ICl|4|-}} ions.<ref name="Greenwood828">Greenwood and Earnshaw, pp. 828–831</ref>

[[Iodine pentafluoride]] (IF<sub>5</sub>), a colourless, volatile liquid, is the most thermodynamically stable iodine fluoride, and can be made by reacting iodine with fluorine gas at room temperature. It is a fluorinating agent, but is mild enough to store in glass apparatus. Again, slight electrical conductivity is present in the liquid state because of dissociation to {{chem|IF|4|+}} and {{chem|IF|6|-}}. The [[pentagonal bipyramidal molecular geometry|pentagonal bipyramidal]] [[iodine heptafluoride]] (IF<sub>7</sub>) is an extremely powerful fluorinating agent, behind only [[chlorine trifluoride]], [[chlorine pentafluoride]], and [[bromine pentafluoride]] among the interhalogens: it reacts with almost all the elements even at low temperatures, fluorinates [[Pyrex]] glass to form iodine(VII) oxyfluoride (IOF<sub>5</sub>), and sets [[carbon monoxide]] on fire.<ref name="Greenwood832">Greenwood and Earnshaw, pp. 832–835</ref>

===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" />

===Polyiodine compounds===
When iodine dissolves in strong acids, such as fuming sulfuric acid, a bright blue [[Paramagnetism|paramagnetic]] solution including {{chem|I|2|+}} cations is formed. A solid salt of the diiodine cation may be obtained by oxidising iodine with [[antimony pentafluoride]]:<ref name="King" />

{{block indent|2 I<sub>2</sub> + 5 SbF<sub>5</sub> {{overunderset|{{big|⟶}}|SO<sub>2</sub>|20&nbsp;°C}} 2 I<sub>2</sub>Sb<sub>2</sub>F<sub>11</sub> + SbF<sub>3</sub>}}

The salt I<sub>2</sub>Sb<sub>2</sub>F<sub>11</sub> is dark blue, and the blue [[tantalum]] analogue I<sub>2</sub>Ta<sub>2</sub>F<sub>11</sub> is also known. Whereas the I–I bond length in I<sub>2</sub> is 267&nbsp;pm, that in {{chem|I|2|+}} is only 256&nbsp;pm as the missing electron in the latter has been removed from an antibonding orbital, making the bond stronger and hence shorter. In [[fluorosulfuric acid]] solution, deep-blue {{chem|I|2|+}} reversibly dimerises below −60&nbsp;°C, forming red rectangular diamagnetic {{chem|I|4|2+}}. Other polyiodine cations are not as well-characterised, including bent dark-brown or black {{chem|I|3|+}} and centrosymmetric ''C''<sub>2''h''</sub> green or black {{chem|I|5|+}}, known in the {{chem|AsF|6|-}} and {{chem|AlCl|4|-}} salts among others.<ref name="King" /><ref name="Greenwood842">Greenwood and Earnshaw, pp. 842–4</ref>

The only important polyiodide anion in aqueous solution is linear [[triiodide]], {{chem|I|3|-}}. Its formation explains why the solubility of iodine in water may be increased by the addition of potassium iodide solution:<ref name="King" />

{{block indent|I<sub>2</sub> + I<sup>−</sup> {{eqm}} {{chem|I|3|-}} (''K''<sub>eq</sub> {{=}} c. 700 at 20&nbsp;°C)}}

Many other polyiodides may be found when solutions containing iodine and iodide crystallise, such as {{chem|I|5|-}}, {{chem|I|9|-}}, {{chem|I|4|2-}}, and {{chem|I|8|2-}}, whose salts with large, weakly polarising cations such as [[caesium|Cs<sup>+</sup>]] may be isolated.<ref name="King" /><ref name="Greenwood835">Greenwood and Earnshaw, pp. 835–9</ref>

===Organoiodine compounds===
{{main|Organoiodine compound}}
[[File:IBXAcid.png|thumb|right|Structure of the oxidising agent [[2-Iodoxybenzoic acid|2-iodoxybenzoic acid]]]]
Organoiodine compounds have been fundamental in the development of organic synthesis, such as in the [[Hofmann elimination]] of [[amine]]s,<ref>{{cite journal | title = Beiträge zur Kenntniss der flüchtigen organischen Basen | journal = [[Annalen der Chemie und Pharmacie]] | volume = 78 | issue = 3 | year = 1851 | pages = 253–286 | vauthors = Hofmann AW | doi = 10.1002/jlac.18510780302 | url = https://zenodo.org/record/1427040 | access-date = 30 June 2019 | archive-date = 1 December 2022 | archive-url = https://web.archive.org/web/20221201072415/https://zenodo.org/record/1427040 | url-status = live }}</ref> the [[Williamson ether synthesis]],<ref>{{cite journal | title = Theory of Aetherification | journal = Philosophical Magazine | volume = 37 | issue = 251 | pages = 350–356 | year = 1850 | doi = 10.1080/14786445008646627 | vauthors = Williamson A | url = https://zenodo.org/record/1431121 | access-date = 29 September 2020 | archive-date = 9 November 2022 | archive-url = https://web.archive.org/web/20221109194527/https://zenodo.org/record/1431121 | url-status = live }} ([http://web.lemoyne.edu/~giunta/williamson.html Link to excerpt]. {{Webarchive|url=https://web.archive.org/web/20190423075534/http://web.lemoyne.edu/~giunta/williamson.html |date=23 April 2019 }})</ref> the [[Wurtz reaction|Wurtz coupling reaction]],<ref>{{cite journal | title = Ueber eine neue Klasse organischer Radicale | vauthors = Wurtz A | journal = [[Annalen der Chemie und Pharmacie]] | volume = 96 | issue = 3 | pages = 364–375 | year = 1855 | url = https://zenodo.org/record/1427074 | doi = 10.1002/jlac.18550960310 | access-date = 30 June 2019 | archive-date = 3 February 2023 | archive-url = https://web.archive.org/web/20230203205851/https://zenodo.org/record/1427074 | url-status = live }}</ref> and in [[Grignard reagent]]s.<ref>{{cite journal | vauthors = Grignard V | title = Sur quelques nouvelles combinaisons organométaliques du magnésium et leur application à des synthèses d'alcools et d'hydrocabures | journal = Comptes rendus de l'Académie des Sciences | year = 1900 | volume = 130 | pages = 1322–25 | url = http://gallica.bnf.fr/ark:/12148/bpt6k3086n/f1322.table | author-link = Victor Grignard | access-date = 2 October 2016 | archive-date = 8 August 2019 | archive-url = https://web.archive.org/web/20190808225609/https://gallica.bnf.fr/ark:/12148/bpt6k3086n/f1322.table | url-status = live }}</ref>

The [[carbon]]–iodine bond is a common functional group that forms part of core [[organic chemistry]]; formally, these compounds may be thought of as organic derivatives of the [[Iodide|iodide anion]]. The simplest [[Organoiodine chemistry|organoiodine compounds]], [[Organoiodine chemistry|alkyl iodides]], may be synthesised by the reaction of [[Alcohol (chemistry)|alcohol]]s with [[phosphorus triiodide]]; these may then be used in [[nucleophilic substitution]] reactions, or for preparing [[Grignard reagent]]s. The C–I bond is the weakest of all the carbon–halogen bonds due to the minuscule difference in electronegativity between carbon (2.55) and iodine (2.66). As such, iodide is the best [[leaving group]] among the halogens, to such an extent that many organoiodine compounds turn yellow when stored over time due to decomposition into elemental iodine; as such, they are commonly used in [[organic synthesis]], because of the easy formation and cleavage of the C–I bond.<ref>{{Ullmann | vauthors = Lyday PA | title = Iodine and Iodine Compounds | doi = 10.1002/14356007.a14_381}}</ref> They are also significantly denser than the other organohalogen compounds thanks to the high atomic weight of iodine.<ref name="blanksby">{{cite journal | vauthors = Blanksby SJ, Ellison GB | title = Bond dissociation energies of organic molecules | journal = Accounts of Chemical Research | volume = 36 | issue = 4 | pages = 255–263 | date = April 2003 | pmid = 12693923 | doi = 10.1021/ar020230d | url = http://www.colorado.edu/chem/ellison/papers/Blanksby_Acct_Chem_Res_2003.pdf | access-date = 25 October 2017 | url-status = dead | citeseerx = 10.1.1.616.3043 | archive-url = https://web.archive.org/web/20090206144739/http://colorado.edu/chem/ellison/papers/Blanksby_Acct_Chem_Res_2003.pdf | archive-date = 6 February 2009 }}</ref> A few organic oxidising agents like the [[Hypervalent organoiodine compounds|iodanes]] contain iodine in a higher oxidation state than −1, such as [[2-Iodoxybenzoic acid|2-iodoxybenzoic acid]], a common reagent for the oxidation of alcohols to [[aldehyde]]s,<ref>{{ OrgSynth | title = Dess–Martin periodinane: 1,1,1-Triacetoxy-1,1-dihydro-1,2-benziodoxol-3(1''H'')-one | vauthors = Boeckman Jr RK, Shao P, Mullins JJ | year = 2000 | volume = 77 | pages = 141 | collvol = 10 | collvolpages = 696 | prep = v77p0141 }}</ref> and [[iodobenzene dichloride]] (PhICl<sub>2</sub>), used for the selective chlorination of [[alkene]]s and [[alkyne]]s.<ref>{{cite journal | vauthors = Jung ME, Parker MH | title = Synthesis of Several Naturally Occurring Polyhalogenated Monoterpenes of the Halomon Class(1) | journal = The Journal of Organic Chemistry | volume = 62 | issue = 21 | pages = 7094–7095 | date = October 1997 | pmid = 11671809 | doi = 10.1021/jo971371 }}</ref> One of the more well-known uses of organoiodine compounds is the so-called [[Haloform reaction|iodoform test]], where [[iodoform]] (CHI<sub>3</sub>) is produced by the exhaustive iodination of a [[Ketone|methyl ketone]] (or another compound capable of being oxidised to a methyl ketone), as follows:<ref name="March">{{March6th}}</ref>

{{block indent|[[Image:Iodoform synthesis.svg|450px]]}}

Some drawbacks of using organoiodine compounds as compared to organochlorine or organobromine compounds is the greater expense and toxicity of the iodine derivatives, since iodine is expensive and organoiodine compounds are stronger alkylating agents.<ref>{{cite web|publisher = Oxford University|title = Safety data for iodomethane|url = http://msds.chem.ox.ac.uk/IO/iodomethane.html|access-date = 12 December 2008|archive-date = 10 August 2010|archive-url = https://web.archive.org/web/20100810211004/http://msds.chem.ox.ac.uk/IO/iodomethane.html|url-status = dead}}</ref> For example, [[iodoacetamide]] and [[iodoacetic acid]] denature proteins by irreversibly alkylating [[cysteine]] residues and preventing the reformation of [[disulfide]] linkages.<ref>{{cite journal | vauthors = Polgár L | title = Deuterium isotope effects on papain acylation. Evidence for lack of general base catalysis and for enzyme–leaving-group interaction | journal = European Journal of Biochemistry | volume = 98 | issue = 2 | pages = 369–374 | date = August 1979 | pmid = 488108 | doi = 10.1111/j.1432-1033.1979.tb13196.x | doi-access = free }}</ref>

Halogen exchange to produce iodoalkanes by the [[Finkelstein reaction]] is slightly complicated by the fact that iodide is a better leaving group than chloride or bromide. The difference is nevertheless small enough that the reaction can be driven to completion by exploiting the differential solubility of halide salts, or by using a large excess of the halide salt.<ref name="March" /> In the classic Finkelstein reaction, an [[Organochlorine chemistry|alkyl chloride]] or an [[Organobromine chemistry|alkyl bromide]] is converted to an [[Organoiodine chemistry|alkyl iodide]] by treatment with a solution of [[sodium iodide]] in [[acetone]]. Sodium iodide is soluble in acetone and [[sodium chloride]] and [[sodium bromide]] are not.<ref>{{cite journal | vauthors = Ervithayasuporn V, Ervithayasuporn V, Pornsamutsin N, Pornsamutsin N, Prangyoo P, Prangyoo P, Sammawutthichai K, Sammawutthichai K, Jaroentomeechai T, Jaroentomeechai T, Phurat C, Phurat C, Teerawatananond T, Teerawatananond T | title = One-pot synthesis of halogen exchanged silsesquioxanes: octakis(3-bromopropyl)octasilsesquioxane and octakis(3-iodopropyl)octasilsesquioxane | journal = Dalton Transactions | volume = 42 | issue = 37 | pages = 13747–13753 | date = October 2013 | pmid = 23907310 | doi = 10.1039/C3DT51373D | s2cid = 41232118 }}</ref> The reaction is driven toward products by [[Law of mass action|mass action]] due to the precipitation of the insoluble salt.<ref>{{cite journal | vauthors = Streitwieser A | year = 1956 | title = Solvolytic Displacement Reactions at Saturated Carbon Atoms | journal = [[Chemical Reviews]] | volume = 56 | pages = 571–752 | doi = 10.1021/cr50010a001 | issue = 4}}</ref><ref>{{cite journal | title = The Effect of the Carbonyl and Related Groups on the Reactivity of Halides in S<sub>N</sub>2 Reactions | vauthors = Bordwell FG, Brannen WT | journal = [[Journal of the American Chemical Society]] | year = 1964 | volume = 86 | pages = 4645–4650 | doi = 10.1021/ja01075a025 | issue = 21}}</ref>