View source for Chemistry:Iodine

{{Infobox iodine|engvar=en-GB}}

'''Iodine''' is a [[Chemistry:Chemical element|chemical element]]; it has [[Chemistry:Symbol|symbol]] '''I''' and [[Physics:Atomic number|atomic number]] 53. The heaviest of the stable [[Chemistry:Halogen|halogen]]s, it exists  at standard conditions as a semi-lustrous, non-metallic solid that melts to form a deep violet liquid at {{convert|114|C}}, and boils to a violet gas at {{convert|184|C}}. The element was discovered by the French chemist [[Biography:Bernard Courtois|Bernard Courtois]] in 1811 and was named two years later by [[Biography:Joseph Louis Gay-Lussac|Joseph Louis Gay-Lussac]], after the [[Social:Ancient Greek|Ancient Greek]] {{lang|grc|Ιώδης}} 'violet-coloured'.

Iodine occurs in many oxidation states, including [[Chemistry:Iodide|iodide]] (I<sup>−</sup>), [[Chemistry:Iodate|iodate]] ({{chem|IO|3|-}}), and the various periodate anions. It is the least abundant of the stable [[Chemistry:Halogen|halogen]]s, being the sixty-first most abundant element. As the heaviest essential mineral nutrient, iodine is required for the synthesis of thyroid hormones.<ref name="lpi">{{cite web|url=http://lpi.oregonstate.edu/mic/minerals/iodine|title=Iodine|publisher=Micronutrient Information Center, Linus Pauling Institute, [[Organization:Oregon State University|Oregon State University]], Corvallis|date=2015|access-date=20 November 2017}}</ref> [[Medicine:Iodine deficiency|Iodine deficiency]] affects about two billion people and is the leading preventable cause of intellectual disabilities.<ref>{{cite news|url= https://query.nytimes.com/gst/fullpage.html?res=9E05E3D81231F935A25751C1A9609C8B63|work=The New York Times|title=In Raising the World's I.Q., the Secret's in the Salt| vauthors = McNeil Jr DG |date=2006-12-16|access-date=2009-07-21|url-status=live|archive-url= https://web.archive.org/web/20100712011551/http://query.nytimes.com/gst/fullpage.html?res=9E05E3D81231F935A25751C1A9609C8B63|archive-date=2010-07-12}}</ref>

The dominant producers of iodine today are {{wipe|Chile}} and {{wipe|Japan}}. Due to its high atomic number and ease of attachment to [[Chemistry:Organic compound|organic compound]]s, it has also found favour as a non-toxic radiocontrast material. Because of the specificity of its uptake by the human body, radioactive isotopes of iodine can also be used to treat [[Medicine:Thyroid cancer|thyroid cancer]]. Iodine is also used as a [[Chemistry:Catalyst|catalyst]] in the industrial production of [[Chemistry:Acetic acid|acetic acid]] and some [[Chemistry:Polymer|polymer]]s.

It is on the [[Medicine:WHO Model List of Essential Medicines|World Health Organization's List of Essential Medicines]].<ref name="WHO22nd">{{cite book | vauthors = ((World Health Organization)) | title = World Health Organization model list of essential medicines: 22nd list (2021) | year = 2021 | hdl = 10665/345533 |  publisher = World Health Organization | location = Geneva | id = WHO/MHP/HPS/EML/2021.02 | hdl-access=free }}</ref>

==History==
In 1811, iodine was discovered by French chemist [[Biography:Bernard Courtois|Bernard Courtois]],<ref name="court">{{cite journal| vauthors = Courtois B  |title=Découverte d'une substance nouvelle dans le Vareck |trans-title=Discovery of a new substance in seaweed |journal=Annales de chimie |volume=88 |pages=304–310 |date=1813 |url=https://books.google.com/books?id=YGwri-w7sMAC&pg=RA2-PA304|language=French}} In French, seaweed that had been washed onto the shore was called "varec", "varech", or "vareck", whence the English word "wrack". Later, "varec" also referred to the ashes of such seaweed: the ashes were used as a source of iodine and salts of sodium and potassium.</ref><ref>{{cite journal | vauthors = Swain PA |title=Bernard Courtois (1777–1838) famed for discovering iodine (1811), and his life in Paris from 1798 |journal=Bulletin for the History of Chemistry |volume=30 |issue=2 |page=103 |date=2005 |url=http://www.scs.uiuc.edu/~mainzv/HIST/awards/OPA%20Papers/2007-Swain.pdf |access-date=2 April 2009 |archive-url=https://web.archive.org/web/20100714110757/http://www.scs.uiuc.edu/~mainzv/HIST/awards/OPA%20Papers/2007-Swain.pdf |archive-date=14 July 2010 |url-status=dead }}</ref> who was born to a manufacturer of [[Chemistry:Potassium nitrate|saltpetre]] (an essential component of [[Chemistry:Gunpowder|gunpowder]]). At the time of the Napoleonic Wars, saltpetre was in great demand in France. Saltpetre produced from French nitre beds required [[Chemistry:Sodium carbonate|sodium carbonate]], which could be isolated from [[Biology:Seaweed|seaweed]] collected on the coasts of [[Place:Normandy|Normandy]] and Brittany. To isolate the sodium carbonate, seaweed was burned and the ash washed with water. The remaining waste was destroyed by adding [[Chemistry:Sulfuric acid|sulfuric acid]]. Courtois once added excessive sulfuric acid and a cloud of purple vapour rose. He noted that the vapour crystallised on cold surfaces, making dark crystals.<ref name="Greenwood794">Greenwood and Earnshaw, p. 794</ref> Courtois suspected that this material was a new element but lacked funding to pursue it further.<ref name="vdK">{{cite web|url=http://elements.vanderkrogt.net/element.php?sym=i |title=53 Iodine |publisher=Elements.vanderkrogt.net |access-date=23 October 2016}}</ref>

Courtois gave samples to his friends, [[Biography:Charles Bernard Desormes|Charles Bernard Desormes]] (1777–1838) and [[Biography:Nicolas Clément|Nicolas Clément]] (1779–1841), to continue research. He also gave some of the substance to [[Chemist|chemist]] [[Biography:Joseph Louis Gay-Lussac|Joseph Louis Gay-Lussac]] (1778–1850), and to [[Physicist|physicist]] [[Biography:André-Marie Ampère|André-Marie Ampère]] (1775–1836). On 29 November 1813, Desormes and Clément made Courtois' discovery public. They described the substance to a meeting of the Imperial Institute of France.<ref>Desormes and Clément made their announcement at the Institut impérial de France on 29 November 1813; a summary of their announcement appeared in the ''Gazette nationale ou Le Moniteur Universel'' of 2 December 1813. See: 
*  {{cite journal |last1=(Staff) |title=Institut Imperial de France |journal=Le Moniteur Universel |date=2 December 1813 |issue=336 |page=1344 |url=https://www.retronews.fr/journal/gazette-nationale-ou-le-moniteur-universel/02-decembre-1813/149/1332251/2 |language=French}}
*  {{cite journal | vauthors = Chattaway FD |title=The discovery of iodine |journal=Chemical News and Journal of Industrial Science |date=23 April 1909 |volume=99 |issue=2578 |pages=193–195 |url=https://books.google.com/books?id=Rco_AQAAIAAJ&pg=PA193}}</ref> On 6 December, Gay-Lussac announced that the new substance was either an element or a compound of [[Chemistry:Oxygen|oxygen]].<ref name="Gay-Lussac">{{cite journal | vauthors = Gay-Lussac J |title=Sur un nouvel acide formé avec la substance décourverte par M. Courtois |trans-title=On a new acid formed by the substance discovered by Mr. Courtois |journal=Annales de Chimie |volume=88|pages=311–318|date=1813 |url=https://books.google.com/books?id=YGwri-w7sMAC&pg=PA311|language=French}}</ref><ref>{{cite journal| vauthors = Gay-Lussac J |title=Sur la combination de l'iode avec d'oxigène |trans-title=On the combination of iodine with oxygen |journal=Annales de Chimie |volume=88 |pages=319–321|date=1813 |url=https://books.google.com/books?id=YGwri-w7sMAC&pg=PA319|language=French}}</ref><ref>{{cite journal| vauthors = Gay-Lussac J |title=Mémoire sur l'iode |trans-title=Memoir on iodine |journal=Annales de Chimie |volume=91 |pages=5–160|date=1814 |url=https://books.google.com/books?id=Efms0Fri1CQC&pg=PA5|language=French}}</ref> Gay-Lussac suggested the name ''"iode"'', from the Ancient Greek {{lang|grc|ἰοειδής}} ({{transliteration|grc|ioeidēs}}, "violet"), because of the colour of iodine vapor.<ref name="court" /><ref name="Gay-Lussac" /> Ampère had given some of his sample to English chemist [[Biography:Humphry Davy|Humphry Davy]] (1778–1829), who experimented on the substance and noted its similarity to [[Chemistry:Chlorine|chlorine]].<ref>{{cite journal| vauthors = Davy H |title=Sur la nouvelle substance découverte par M. Courtois, dans le sel de Vareck |trans-title=On the new substance discovered by Mr. Courtois in the salt of seaweed |journal=Annales de Chimie |volume=88|pages=322–329 |date=1813 |url=https://books.google.com/books?id=YGwri-w7sMAC&pg=PA322|language=French}}</ref> Davy sent a letter dated 10 December to the Royal Society of London stating that he had identified a new element.<ref>{{cite journal| vauthors = Davy H |author-link=Humphry Davy |title=Some experiments and observations on a new substance which becomes a violet coloured gas by heat |journal=Philosophical Transactions of the Royal Society of London |volume=104 |pages=74–93 |date=1 January 1814 |doi=10.1098/rstl.1814.0007 |doi-access=free }}</ref> Arguments erupted between Davy and Gay-Lussac over who identified iodine first, but both scientists acknowledged Courtois as the first to isolate the element.<ref name="vdK" />

In 1873 the French medical researcher Casimir Joseph Davaine (1812–1882) discovered the antiseptic action of iodine.<ref>{{cite journal | vauthors = Davaine C |title=Recherches relatives à l'action des substances dites ''antiseptiques'' sur le virus charbonneux |journal=Comptes rendus hebdomadaires des séances de l'Académie des Sciences |date=1873 |volume=77 |pages=821–825 |url=https://babel.hathitrust.org/cgi/pt?id=uiug.30112025711521&view=1up&seq=829 |trans-title=Investigations regarding the action of so-called ''antiseptic'' substances on the anthrax bacterium |language=French}}</ref> Antonio Grossich (1849–1926), an Istrian-born surgeon, was among the first to use [[Biology:Sterilization (microbiology)|sterilisation]] of the operative field. In 1908, he introduced tincture of iodine as a way to rapidly sterilise the human skin in the surgical field.<ref>{{cite journal | vauthors = Grossich A |title=Eine neue Sterilisierungsmethode der Haut bei Operationen |journal=Zentralblatt für Chirurgie |date=31 October 1908 |volume=35 |issue=44 |pages=1289–1292 |url=https://babel.hathitrust.org/cgi/pt?id=uc1.b4150494&view=1up&seq=1305 |trans-title=A new method of sterilization of the skin for operations |language=German}}</ref>

In early [[Chemistry:Periodic table|periodic table]]s, iodine was often given the symbol ''J'', for ''Jod'', its name in [[Social:German language|German]].<ref>{{cite web |title=Mendeleev's First Periodic Table |url=https://web.lemoyne.edu/giunta/EA/MENDELEEVann.HTML |website=web.lemoyne.edu}}</ref>

==Properties==
[[File:IodoAtomico.JPG|thumb|left|upright=0.7|alt=Round bottom flask filled with violet iodine vapour|Iodine vapour in a flask.]]
Iodine is the fourth [[Chemistry:Halogen|halogen]], being a member of group 17 in the periodic table, below [[Chemistry:Fluorine|fluorine]], [[Chemistry:Chlorine|chlorine]], and [[Chemistry:Bromine|bromine]]; it is the heaviest stable member of its group. (The fifth and sixth halogens, the radioactive [[Chemistry:Astatine|astatine]] and [[Chemistry:Tennessine|tennessine]], are not well-studied due to their expense and inaccessibility in large quantities, but appear to show various unusual properties for the group due to relativistic effects.) Iodine has an electron configuration of [Kr]4d<sup>10</sup>5s<sup>2</sup>5p<sup>5</sup>, with the seven electrons in the fifth and outermost shell being its [[Physics:Valence electron|valence electron]]s. Like the other halogens, it is one electron short of a full octet and is hence an oxidising agent, reacting with many elements in order to complete its outer shell, although in keeping with periodic trends, it is the weakest oxidising agent among the stable halogens: it has the lowest [[Chemistry:Electronegativity|electronegativity]] among them, just 2.66 on the Pauling scale (compare fluorine, chlorine, and bromine at 3.98, 3.16, and 2.96 respectively; astatine continues the trend with an electronegativity of 2.2). Elemental iodine hence forms [[Chemistry:Diatomic molecule|diatomic molecule]]s with chemical formula I<sub>2</sub>, where two iodine atoms share a pair of electrons in order to each achieve a stable octet for themselves; at high temperatures, these diatomic molecules reversibly dissociate a pair of iodine atoms. Similarly, the iodide anion, I<sup>−</sup>, is the strongest reducing agent among the stable halogens, being the most easily oxidised back to diatomic I<sub>2</sub>.<ref name="Greenwood800">Greenwood and Earnshaw, pp. 800–4</ref> (Astatine goes further, being indeed unstable as At<sup>−</sup> and readily oxidised to At<sup>0</sup> or At<sup>+</sup>.)<ref>{{cite book | series = Gmelin Handbook of Inorganic and Organometallic Chemistry | title = 'At, Astatine', System No. 8a | edition=8th | year = 1985 | publisher = Springer-Verlag | isbn = 978-3-540-93516-2 | vauthors = Kugler HK, Keller C | volume = 8 }}</ref>

The halogens darken in colour as the group is descended: fluorine is a very pale yellow, chlorine is greenish-yellow, bromine is reddish-brown, and iodine is violet.

Elemental iodine is slightly soluble in water, with one gram dissolving in 3450&nbsp;mL at 20&nbsp;°C and 1280&nbsp;mL at 50&nbsp;°C; [[Biology:Potassium iodide|potassium iodide]] may be added to increase solubility via formation of [[Chemistry:Triiodide|triiodide]] ions, among other polyiodides.<ref name="Greenwood804">Greenwood and Earnshaw, pp. 804–9</ref> Nonpolar solvents such as [[Chemistry:Hexane|hexane]] and [[Chemistry:Carbon tetrachloride|carbon tetrachloride]] provide a higher solubility.<ref>{{cite book| title = Merck Index of Chemicals and Drugs| edition = 9th| date = 1976| isbn=978-0-911910-26-1| editor = Windholz, Martha| editor2 = Budavari, Susan| editor3 = Stroumtsos, Lorraine Y.| editor4 = Fertig, Margaret Noether| publisher = J A Majors Company}}</ref> Polar solutions, such as aqueous solutions, are brown, reflecting the role of these solvents as Lewis bases; on the other hand, nonpolar solutions are violet, the color of iodine vapour.<ref name="Greenwood804" /> [[Chemistry:Charge-transfer complex|Charge-transfer complex]]es form when iodine is dissolved in polar solvents, hence changing the colour. Iodine is violet when dissolved in carbon tetrachloride and saturated hydrocarbons but deep brown in [[Chemistry:Alcohol|alcohol]]s and [[Chemistry:Amine|amine]]s, solvents that form charge-transfer adducts.<ref name="King">{{cite book | vauthors = King RB |date=1995 |title=Inorganic Chemistry of Main Group Elements |publisher=Wiley-VCH |pages=173–98|isbn=978-0-471-18602-1}}</ref>

[[File:Iodine-triphenylphosphine charge-transfer complex in dichloromethane.jpg|thumb|upright=1.8|right|I<sub>2</sub>•[[Chemistry:Triphenylphosphine|PPh<sub>3</sub>]] charge-transfer complexes in 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 melting and boiling points of iodine are the highest among the halogens, conforming to the increasing trend down the group, since iodine has the largest electron cloud among them that is the most easily polarised, resulting in its molecules having the strongest van der Waals interactions among the halogens. Similarly, iodine is the least volatile of the halogens, though the solid still can be observed to give off purple vapor.<ref name="Greenwood800" /> Due to this property iodine is commonly used to demonstrate [[Sublimation (phase transition)|sublimation]] directly from [[Physics:Solid|solid]] to [[Physics:Gas|gas]], which gives rise to a misconception that it does not [[Physics:Melting|melt]] in [[Physics:Atmospheric pressure|atmospheric pressure]].<ref>{{cite journal |title=The concept of sublimation – iodine as an example |journal=Educación Química |date=1 March 2012 |volume=23 |pages=171–175 |doi=10.1016/S0187-893X(17)30149-0 |language=en |issn=0187-893X|doi-access=free | vauthors = Stojanovska M, Petruševski VM, Šoptrajanov B }}</ref> Because it has the largest [[Physics:Atomic radius|atomic radius]] among the halogens, iodine has the lowest first ionisation energy, lowest [[Physics:Electron affinity|electron affinity]], lowest [[Chemistry:Electronegativity|electronegativity]] and lowest reactivity of the halogens.<ref name="Greenwood800" />

[[File:Iodine-unit-cell-3D-balls-B.png|thumb|upright=0.7|right|Structure of solid iodine]]
The interhalogen bond in diiodine is the weakest of all the halogens. As such, 1% of a sample of gaseous iodine at atmospheric pressure is dissociated into iodine atoms at 575&nbsp;°C. Temperatures greater than 750&nbsp;°C are required for fluorine, chlorine, and bromine to dissociate to a similar extent. Most bonds to iodine are weaker than the analogous bonds to the lighter halogens.<ref name="Greenwood800" /> Gaseous iodine is composed of I<sub>2</sub> molecules with an I–I bond length of 266.6&nbsp;pm. The I–I bond is one of the longest single bonds known. It is even longer (271.5&nbsp;pm) in solid [[Chemistry:Orthorhombic|orthorhombic]] crystalline iodine, which has the same crystal structure as chlorine and bromine. (The record is held by iodine's neighbour [[Chemistry:Xenon|xenon]]: the Xe–Xe bond length is 308.71&nbsp;pm.)<ref>{{cite book| title = Advanced Structural Inorganic Chemistry| url = https://archive.org/details/advancedstructur00liwa| url-access = limited| vauthors = Li WK, Zhou GD, Mak TC | publisher = Oxford University Press| date = 2008| isbn = 978-0-19-921694-9| page = [https://archive.org/details/advancedstructur00liwa/page/n696 674]}}</ref> As such, within the iodine molecule, significant electronic interactions occur with the two next-nearest neighbours of each atom, and these interactions give rise, in bulk iodine, to a shiny appearance and [[Physics:Semiconductor|semiconducting]] properties.<ref name="Greenwood800" /> Iodine is a two-dimensional semiconductor with a [[Physics:Band gap|band gap]] of 1.3&nbsp;eV (125&nbsp;kJ/mol): it is a semiconductor in the plane of its crystalline layers and an insulator in the perpendicular direction.<ref name="Greenwood800" />

===Isotopes===
{{main|Physics:Isotopes of iodine}}
Of the thirty-seven known [[Physics:Isotopes of iodine|isotopes of iodine]], only one occurs in nature, iodine-127. The others are radioactive and have half-lives too short to be [[Chemistry:Primordial nuclide|primordial]]. As such, iodine is both [[Monoisotopic element|monoisotopic]] and [[Mononuclidic element|mononuclidic]] and its atomic weight is known to great precision, as it is a constant of nature.<ref name="Greenwood800" />

The longest-lived of the radioactive isotopes of iodine is [[Physics:Iodine-129|iodine-129]], which has a half-life of 15.7&nbsp;million&nbsp;years, decaying via [[Physics:Beta decay|beta decay]] to stable [[Chemistry:Xenon|xenon]]-129.<ref name="NUBASE">{{NUBASE 2003}}</ref> Some iodine-129 was formed along with iodine-127 before the formation of the Solar System, but it has by now completely decayed away, making it an [[Chemistry:Extinct radionuclide|extinct radionuclide]] that is nevertheless still useful in dating the history of the early Solar System or very old groundwaters, due to its mobility in the environment. Its former presence may be determined from an excess of its [[Physics:Decay product|daughter]] xenon-129.<ref>{{cite journal | vauthors = Watson JT, Roe DK, Selenkow HA | title = Iodine-129 as a "nonradioactive" tracer | journal = Radiation Research | volume = 26 | issue = 1 | pages = 159–163 | date = September 1965 | pmid = 4157487 | doi = 10.2307/3571805 | bibcode = 1965RadR...26..159W | jstor = 3571805 }}</ref><ref>{{cite web |url=https://e-reports-ext.llnl.gov/pdf/234761.pdf | vauthors = Santschi PH, Moran JE, Oktay S, Hoehn E, Sharma P |date=1998 |title=129Iodine: A new tracer for surface water/groundwater interaction |publisher=Lawrence Livermore National Laboratory preprint UCRL-JC-132516. |location=Livermore, US|archive-url=https://web.archive.org/web/20161221192732/https://e-reports-ext.llnl.gov/pdf/234761.pdf|archive-date=21 December 2016|url-status=dead}}</ref><ref>{{cite journal | vauthors = Snyder G, Fabryka-Martin J | year = 2007 | title = I-129 and Cl-36 in dilute hydrocarbon waters: Marine-cosmogenic, in situ, and anthropogenic sources | journal = Applied Geochemistry | volume = 22 | issue = 3| pages = 692–714 | doi = 10.1016/j.apgeochem.2006.12.011 | bibcode = 2007ApGC...22..692S }}</ref><ref>{{cite book| vauthors = Clayton DD |year=1983|title=Principles of Stellar Evolution and Nucleosynthesis| url=https://archive.org/details/principlesofstel0000clay|url-access=registration|page=[https://archive.org/details/principlesofstel0000clay/page/75 75]|edition=2nd|publisher=University of Chicago Press|isbn=978-0-226-10953-4}}</ref><ref>{{cite web| vauthors = Bolt BA, Packard RE, Price PB |year=2007|url=http://content.cdlib.org/xtf/view?docId=hb1r29n709&doc.view=content&chunk.id=div00061&toc.depth=1&brand=oac&anchor.id=0|title=John H. Reynolds, Physics: Berkeley|publisher=The University of California, Berkeley|access-date=2007-10-01 }}</ref> Traces of iodine-129 still exist today, as it is also a [[Physics:Cosmogenic nuclide|cosmogenic nuclide]], formed from [[Physics:Cosmic ray spallation|cosmic ray spallation]] of atmospheric xenon: these traces make up 10<sup>−14</sup> to 10<sup>−10</sup> of all terrestrial iodine. It also occurs from open-air nuclear testing, and is not hazardous because of its very long half-life, the longest of all fission products. At the peak of thermonuclear testing in the 1960s and 1970s, iodine-129 still made up only about 10<sup>−7</sup> of all terrestrial iodine.<ref name="SCOPE50">
[http://www.scopenvironment.org/downloadpubs/scope50 SCOPE 50 - Radioecology after Chernobyl] {{webarchive|url=https://web.archive.org/web/20140513065145/http://www.scopenvironment.org/downloadpubs/scope50/ |date=13 May 2014 }}, the [[Organization:Scientific Committee on Problems of the Environment|Scientific Committee on Problems of the Environment]] (SCOPE), 1993. See table 1.9 in Section 1.4.5.2.</ref> Excited states of iodine-127 and iodine-129 are often used in [[Mössbauer spectroscopy]].<ref name="Greenwood800" />

The other iodine radioisotopes have much shorter half-lives, no longer than days.<ref name="NUBASE" /> Some of them have medical applications involving the thyroid gland, where the iodine that enters the body is stored and concentrated. [[Chemistry:Iodine-123|Iodine-123]] has a half-life of thirteen hours and decays by [[Physics:Electron capture|electron capture]] to [[Physics:Tellurium-123|tellurium-123]], emitting gamma radiation; it is used in [[Medicine:Nuclear medicine|nuclear medicine]] imaging, including single photon emission computed tomography (SPECT) and [[Astronomy:X-ray computed tomography|X-ray computed tomography]] (X-Ray CT) scans.<ref>{{cite journal | vauthors = Hupf HB, Eldridge JS, Beaver JE | title = Production of iodine-123 for medical applications | journal = The International Journal of Applied Radiation and Isotopes | volume = 19 | issue = 4 | pages = 345–351 | date = April 1968 | pmid = 5650883 | doi = 10.1016/0020-708X(68)90178-6 }}</ref> [[Physics:Iodine-125|Iodine-125]] has a half-life of fifty-nine days, decaying by electron capture to [[Physics:Tellurium-125|tellurium-125]] and emitting low-energy gamma radiation; the second-longest-lived iodine radioisotope, it has uses in biological assays, [[Medicine:Nuclear medicine|nuclear medicine]] imaging and in [[Physics:Radiation therapy|radiation therapy]] as [[Medicine:Brachytherapy|brachytherapy]] to treat a number of conditions, including [[Medicine:Prostate cancer|prostate cancer]], [[Medicine:Uveal melanoma|uveal melanoma]]s, and brain tumours.<ref>Harper, P.V.; Siemens, W.D.; Lathrop, K.A.; Brizel, H.E.; Harrison, R.W. ''Iodine-125.'' Proc. Japan Conf. Radioisotopes; Vol: 4th Jan 01, 1961</ref> Finally, [[Physics:Iodine-131|iodine-131]], with a half-life of eight days, beta decays to an excited state of stable [[Physics:Xenon-131|xenon-131]] that then converts to the ground state by emitting gamma radiation. It is a common fission product and thus is present in high levels in radioactive fallout. It may then be absorbed through contaminated food, and will also accumulate in the thyroid. As it decays, it may cause damage to the thyroid. The primary risk from exposure to high levels of iodine-131 is the chance occurrence of radiogenic [[Medicine:Thyroid cancer|thyroid cancer]] in later life. Other risks include the possibility of non-cancerous growths and [[Medicine:Thyroiditis|thyroiditis]].<ref name="Rivkees">{{cite journal | vauthors = Rivkees SA, Sklar C, Freemark M | title = Clinical review 99: The management of Graves' disease in children, with special emphasis on radioiodine treatment | journal = The Journal of Clinical Endocrinology and Metabolism | volume = 83 | issue = 11 | pages = 3767–3776 | date = November 1998 | pmid = 9814445 | doi = 10.1210/jcem.83.11.5239 | doi-access = free }}</ref>

The usual means of protection against the negative effects of iodine-131 is by saturating the thyroid gland with stable iodine-127 in the form of [[Biology:Potassium iodide|potassium iodide]] tablets, taken daily for optimal prophylaxis.<ref>{{cite journal | vauthors = Zanzonico PB, Becker DV | title = Effects of time of administration and dietary iodine levels on potassium iodide (KI) blockade of thyroid irradiation by 131I from radioactive fallout | journal = Health Physics | volume = 78 | issue = 6 | pages = 660–667 | date = June 2000 | pmid = 10832925 | doi = 10.1097/00004032-200006000-00008 | s2cid = 30989865 }}</ref> However, iodine-131 may also be used for medicinal purposes in [[Physics:Radiation therapy|radiation therapy]] for this very reason, when tissue destruction is desired after iodine uptake by the tissue.<ref>{{cite news |title=Medical isotopes the likely cause of radiation in Ottawa waste |url=http://www.cbc.ca/news/canada/medical-isotopes-the-likely-cause-of-radiation-in-ottawa-waste-1.852645| date=4 February 2009 |publisher=CBC News|access-date=30 September 2015}}</ref> Iodine-131 is also used as a [[Biology:Radioactive tracer|radioactive tracer]].<ref>{{cite book| vauthors = Moser H, Rauert W |title=Isotopes in the water cycle : past, present and future of a developing science|year=2007|publisher=Springer|location=Dordrecht|isbn=978-1-4020-6671-9 | veditors = Aggarwal PK, Gat JR, Froehlich KF |access-date=6 May 2012|page=11|chapter=Isotopic Tracers for Obtaining Hydrologic Parameters|chapter-url=https://books.google.com/books?id=XKk6V_IeJbIC&pg=PA11}}</ref><ref>{{cite book| vauthors = Rao SM |title=Practical isotope hydrology|year=2006|publisher=New India Publishing Agency|location=New Delhi|isbn=978-81-89422-33-2|chapter-url=https://books.google.com/books?id=E7TVDVVji0EC&q=isotope%20hydrology%20iodine&pg=PA11|access-date=6 May 2012|pages=12–13|chapter=Radioisotopes of hydrological interest}}</ref><ref>{{cite web|title=Investigating leaks in Dams & Reservoirs|url=http://www.iaea.org/technicalcooperation/documents/sheet20dr.pdf|work=IAEA.org|access-date=6 May 2012|archive-url=https://web.archive.org/web/20130730053205/http://www.iaea.org/technicalcooperation/documents/sheet20dr.pdf|archive-date=30 July 2013|url-status=dead}}</ref><ref>{{cite book| vauthors = Araguás LA, Bedmar AP |title=Detection and prevention of leaks from dams|year=2002|publisher=Taylor & Francis|isbn=978-90-5809-355-4|chapter-url=https://books.google.com/books?id=FXB-HMzfBnkC&pg=PA179 |access-date=6 May 2012|pages=179–181|chapter=Artificial radioactive tracers}}</ref>

== Chemistry and compounds ==
{{Main|Chemistry:Iodine compounds}}
{| class="wikitable" style="float:right; margin-top:0; margin-left:1em; text-align:center; font-size:10pt; line-height:11pt; 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 much less reactive than the other halogens. For example, while chlorine gas will halogenate [[Chemistry:Carbon monoxide|carbon monoxide]], [[Chemistry:Nitric oxide|nitric oxide]], and [[Chemistry:Sulfur dioxide|sulfur dioxide]] (to [[Chemistry:Phosgene|phosgene]], [[Chemistry:Nitrosyl chloride|nitrosyl chloride]], and [[Chemistry:Sulfuryl chloride|sulfuryl chloride]] respectively), iodine will not do so. Furthermore, iodination of metals tends to result in lower oxidation states than chlorination or bromination; for example, [[Chemistry:Rhenium|rhenium]] metal reacts with chlorine to form rhenium hexachloride, but with bromine it forms only rhenium pentabromide and iodine can achieve only rhenium tetraiodide.<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 [[Chemistry:Iodine heptafluoride|iodine heptafluoride]].<ref name="Greenwood804" />

===Charge-transfer complexes ===
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. 806-7</ref>

===Hydrogen iodide===
The simplest compound of iodine is [[Chemistry:Hydrogen iodide|hydrogen iodide]], HI. It is a colourless gas that reacts with oxygen to give water and iodine. Although it is useful in 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 [[Chemistry:Hydrogen sulfide|hydrogen sulfide]] or [[Chemistry:Hydrazine|hydrazine]]:<ref name="Greenwood809">Greenwood and Earnshaw, pp. 809–12</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 [[Chemistry:Hydrogen fluoride|hydrogen fluoride]], since hydrogen cannot form strong [[Chemistry:Hydrogen bond|hydrogen bond]]s to the large and only mildly electronegative iodine atom. It melts at −51.0&nbsp;°C and boils at −35.1&nbsp;°C. It is an endothermic compound that can exothermically dissociate at room temperature, although the process is very slow unless a [[Chemistry:Catalyst|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 is likewise the smallest of the hydrogen halides, at 295&nbsp;kJ/mol.<ref name="Greenwood812">Greenwood and Earnshaw, pp. 812–9</ref>

Aqueous hydrogen iodide is known as [[Chemistry:Hydroiodic acid|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>Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. {{ISBN|0-12-352651-5}}.</ref> Commercial so-called "concentrated" hydroiodic acid usually contains 48–57% HI by mass; the solution forms an [[Physics:Azeotrope|azeotrope]] with boiling point 126.7&nbsp;°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 [[Chemistry:Acetic acid|acetic acid]] by the [[Chemistry:Cativa process|Cativa process]].<ref name = "Cativa">{{cite journal|title = The Cativa Process for the Manufacture of Acetic Acid|author =  Jones, J. H.|journal = Platinum Metals Rev.|year = 2000|volume = 44|issue = 3|pages = 94–105|url = http://www.platinummetalsreview.com/pdf/pmr-v44-i3-094-105.pdf}}</ref><ref>{{cite journal| title = High productivity methanol carbonylation catalysis using iridium - The Cativa process for the manufacture of acetic acid |author1=Sunley, G. J. |author2=Watson, D. J. | journal =  Catalysis Today | year = 2000 | volume = 58 | issue = 4 | pages = 293–307 | doi = 10.1016/S0920-5861(00)00263-7}}</ref>

Unlike [[Chemistry:Hydrogen fluoride|hydrogen fluoride]], anhydrous liquid hydrogen iodide is difficult to work with as a solvent, because its boiling point is low, it has a small liquid range, its [[Physics:Permittivity|permittivity]] is low and it does not dissociate appreciably into H<sub>2</sub>I<sup>+</sup> and {{chem|HI|2|-}} ions – the latter, in any case, are much less stable than the [[Chemistry:Bifluoride|bifluoride]] ions ({{chem|HF|2|-}}) due to the very weak hydrogen bonding between hydrogen and iodine, though its salts with very large and weakly polarising cations such as [[Chemistry:Caesium|Cs<sup>+</sup>]] and [[Chemistry:Quaternary ammonium cation|{{chem|NR|4|+}}]] (R = [[Chemistry:Methyl group|Me]], [[Chemistry:Ethyl group|Et]], [[Chemistry:Butyl group|Bu<sup>''n''</sup>]]) may still be isolated. Anhydrous hydrogen iodide is a poor solvent, able to dissolve only small molecular compounds such as [[Chemistry:Nitrosyl chloride|nitrosyl chloride]] and [[Chemistry:Phenol|phenol]], or salts with very low [[Physics:Lattice energy|lattice energies]] such as tetraalkylammonium halides.<ref name="Greenwood812" />

===Other binary iodine compounds===
With the exception of the [[Chemistry:Noble gas|noble gases]], nearly all elements on the periodic table up to einsteinium (EsI<sub>3</sub> is known) are known to form binary compounds with iodine.  Until 1990, [[Chemistry:Nitrogen triiodide|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 International Edition in English |language=en |volume=29 |issue=6 |pages=677–679 |doi=10.1002/anie.199006771 |issn=0570-0833}}</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=2022-11-17 |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}}</ref>   Sulfur- 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 characterized crystallographically.<ref>{{cite journal |last1=Klapoetke |first1=T. |last2=Passmore |first2=J. |date=1989-07-01 |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}}</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 [[Chemistry:Niobium|niobium]], [[Chemistry:Tantalum|tantalum]], and [[Chemistry:Protactinium|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, [[Chemistry:Carbon tetraiodide|carbon tetraiodide]], or an organic iodide. For example, molybdenum(IV) oxide reacts with [[Chemistry:Aluminium(III) iodide|aluminium(III) iodide]] at 230&nbsp;°C to give [[Chemistry:Molybdenum(II) iodide|molybdenum(II) iodide]]. An example involving halogen exchange is given below, involving the reaction of [[Chemistry:Tantalum(V) chloride|tantalum(V) chloride]] with excess aluminium(III) iodide at 400&nbsp;°C to give [[Chemistry:Tantalum(V) iodide|tantalum(V) iodide]]:<ref name="Greenwood821">Greenwood and Earnshaw, pp. 821–4</ref>

: <chem>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>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 iodide is mostly ionic, but [[Chemistry:Aluminium iodide|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. [[Chemistry:Potassium|potassium]] and [[Chemistry:Calcium|calcium]]) are the greatest among ionic halides of that element, while those of covalent iodides (e.g. [[Chemistry:Silver|silver]]) are the lowest of that element. In particular, [[Chemistry:Silver iodide|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, diamagnetic [[Chemistry:Interhalogen|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 [[Chemistry:Pseudohalogen|pseudohalides]] are also known, such as [[Chemistry:Cyanogen iodide|cyanogen iodide]] (ICN), iodine [[Chemistry:Thiocyanate|thiocyanate]] (ISCN), and iodine [[Chemistry:Azide|azide]] (IN<sub>3</sub>).<ref name="Greenwood824">Greenwood and Earnshaw, pp. 824–8</ref>

[[File:Iodine monochloride1.jpg|thumb|right|Iodine monochloride]]
[[Chemistry:Iodine monofluoride|Iodine monofluoride]] (IF) is unstable at room temperature and disproportionates very readily and irreversibly to iodine and [[Chemistry:Iodine pentafluoride|iodine pentafluoride]], and thus cannot be obtained pure. It can be synthesised from the reaction of iodine with fluorine gas in [[Chemistry:Trichlorofluoromethane|trichlorofluoromethane]] at −45&nbsp;°C, with [[Chemistry:Iodine trifluoride|iodine trifluoride]] in trichlorofluoromethane at −78&nbsp;°C, or with [[Chemistry:Silver(I) fluoride|silver(I) fluoride]] at 0&nbsp;°C.<ref name="Greenwood824" /> [[Chemistry:Iodine monochloride|Iodine monochloride]] (ICl) and [[Chemistry:Iodine monobromide|iodine monobromide]] (IBr), on the other hand, are moderately stable. The former, a volatile red-brown compound, was discovered independently by [[Biography:Joseph Louis Gay-Lussac|Joseph Louis Gay-Lussac]] and [[Biography:Humphry Davy|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 [[Chemistry:Fractional crystallization|fractional crystallisation]]. Both are quite reactive and attack even [[Chemistry:Platinum|platinum]] and [[Chemistry:Gold|gold]], though not [[Chemistry:Boron|boron]], [[Chemistry:Carbon|carbon]], [[Chemistry:Cadmium|cadmium]], [[Chemistry:Lead|lead]], [[Chemistry:Zirconium|zirconium]], [[Chemistry:Niobium|niobium]], [[Chemistry:Molybdenum|molybdenum]], and [[Chemistry:Tungsten|tungsten]]. Their reaction with organic compounds depends on conditions. Iodine chloride vapour tends to chlorinate [[Chemistry:Phenol|phenol]] and [[Chemistry:Salicylic acid|salicylic acid]], since when iodine chloride undergoes homolytic fission, chlorine and iodine are produced and the former is more reactive. However, iodine chloride in [[Chemistry:Carbon tetrachloride|carbon tetrachloride]] solution results in iodination being the main reaction, since now 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" />

[[Chemistry:Iodine trifluoride|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 [[Chemistry:Xenon difluoride|xenon difluoride]] is necessary. [[Chemistry:Iodine trichloride|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>

[[Chemistry:Iodine pentafluoride|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 [[Physics:Pentagonal bipyramidal molecular geometry|pentagonal bipyramidal]] [[Chemistry:Iodine heptafluoride|iodine heptafluoride]] (IF<sub>7</sub>) is an extremely powerful fluorinating agent, behind only [[Chemistry:Chlorine trifluoride|chlorine trifluoride]], [[Chemistry:Chlorine pentafluoride|chlorine pentafluoride]], and [[Chemistry:Bromine pentafluoride|bromine pentafluoride]] among the interhalogens: it reacts with almost all the elements even at low temperatures, fluorinates [[Chemistry:Pyrex|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]]
[[Chemistry:Iodine oxide|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, [[Physics:Hygroscopic|hygroscopic]] [[Chemistry:Iodine pentoxide|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 [[Chemistry:Iodic acid|iodic acid]] (HIO<sub>3</sub>), of which it is the anhydride. It will quickly oxidise carbon monoxide completely to [[Chemistry:Carbon dioxide|carbon dioxide]] at room temperature, and is thus a useful reagent in determining carbon monoxide concentration. It also oxidises nitrogen oxide, [[Chemistry:Ethylene|ethylene]], and [[Chemistry:Hydrogen sulfide|hydrogen sulfide]]. It reacts with [[Chemistry:Sulfur trioxide|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 [[Chemistry:Sulfuric acid|sulfuric acid]] to iodosyl salts involving [IO]<sup>+</sup>. It may be fluorinated by fluorine, bromine trifluoride, [[Chemistry:Sulfur tetrafluoride|sulfur tetrafluoride]], or [[Chemistry:Chloryl fluoride|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. 851–3</ref>

{| class="wikitable" style="float:right; margin-top:0; margin-left:1em; text-align:center; font-size:10pt; line-height:11pt; 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
|-
|&nbsp;||&nbsp;||{{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||&nbsp;||&nbsp;
|-
|{{chem|IO|3|-}}/HOI||+1.134||{{chem|IO|3|-}}/IO<sup>−</sup>||+0.15
|-
|{{chem|IO|4|-}}/{{chem|IO|3|-}}||+1.653||&nbsp;||&nbsp;
|-
|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: [[Chemistry:Hypoiodous acid|hypoiodous acid]] (HIO), [[Chemistry:Iodous acid|iodous acid]] (HIO<sub>2</sub>), [[Chemistry:Iodic acid|iodic acid]] (HIO<sub>3</sub>), and [[Chemistry:Periodic acid|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>
:{|
|-
| 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" />
:{|
|-
| 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 [[Chemistry:Alkali metal|alkali metal]] iodides with oxygen at 600&nbsp;°C and high pressure, or by oxidising iodine with [[Chemistry:Chlorate|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 [[Chemistry:Electrolysis|electrolysis]] or fuming [[Chemistry:Nitric acid|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 [[Chemistry:Sodium iodate|sodium iodate]] electrochemically (with [[Chemistry:Lead(IV) oxide|lead(IV) oxide]] as the anode) or by chlorine gas:<ref name="Greenwood872">Greenwood and Earnshaw, pp. 872–5</ref>
:{{chem|IO|3|-}} + 6 OH<sup>−</sup> → {{chem|IO|6|5-}} + 3 H<sub>2</sub>O + 2 e<sup>−</sup>
:{{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 [[Chemistry:Permanganate|{{chem|MnO|4|-}}]], and cleaving glycols, α-[[Chemistry:Diketone|diketone]]s, α-ketols, α-aminoalcohols, and α-[[Chemistry:Diamine|diamine]]s.<ref name="Greenwood872" /> Orthoperiodate especially stabilises high oxidation states among metals because of its very high negative charge of −5. Orthoperiodic acid, H<sub>5</sub>IO<sub>6</sub>, is stable, and dehydrates at 100&nbsp;°C in a vacuum to 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 [[Chemistry:Sulfuric acid|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 [[Physics:Paramagnetic|paramagnetic]] solution including {{chem|I|2|+}} cations is formed. A solid salt of the diiodine cation may be obtained by oxidising iodine with [[Chemistry:Antimony pentafluoride|antimony pentafluoride]]:<ref name="King" />
: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 [[Chemistry:Tantalum|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 [[Chemistry:Fluorosulfuric acid|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 [[Chemistry:Triiodide|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" />
:I<sub>2</sub> + I<sup>−</sup> {{eqm}} {{chem|I|3|-}} (''K''<sub>eq</sub> = ~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 [[Chemistry:Caesium|Cs<sup>+</sup>]] may be isolated.<ref name="King" /><ref name="Greenwood835">Greenwood and Earnshaw, pp. 835–9</ref>

===Organoiodine compounds===
{{main|Chemistry:Organoiodine compound}}
[[File:IBXAcid.png|thumb|right|Structure of the oxidising agent 2-iodoxybenzoic acid]]
Organoiodine compounds have been fundamental in the development of organic synthesis, such as in the [[Chemistry:Hofmann elimination|Hofmann elimination]] of [[Chemistry:Amine|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 }}</ref> the [[Chemistry:Williamson ether synthesis|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 }} ([http://web.lemoyne.edu/~giunta/williamson.html Link to excerpt.])</ref> the [[Chemistry: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 }}</ref> and in [[Chemistry:Grignard reagent|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 = Compt. Rend. | year = 1900 | volume = 130 | pages = 1322–25 | url = http://gallica.bnf.fr/ark:/12148/bpt6k3086n/f1322.table | }}</ref>

The [[Chemistry:Carbon|carbon]]–iodine bond is a common functional group that forms part of core [[Chemistry:Organic chemistry|organic chemistry]]; formally, these compounds may be thought of as organic derivatives of the iodide anion. The simplest organoiodine compounds, alkyl iodides, may be synthesised by the reaction of [[Chemistry:Alcohol|alcohol]]s with [[Chemistry:Phosphorus triiodide|phosphorus triiodide]]; these may then be used in [[Chemistry:Nucleophilic substitution|nucleophilic substitution]] reactions, or for preparing [[Chemistry:Grignard reagent|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 [[Chemistry:Leaving group|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 [[Chemistry:Organic synthesis|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 [[Chemistry:Iodane|iodane]]s contain iodine in a higher oxidation state than −1, such as 2-iodoxybenzoic acid, a common reagent for the oxidation of alcohols to [[Chemistry:Aldehyde|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 | url = http://www.orgsyn.org/orgsyn/pdfs/v77p0141.pdf }}</ref> and [[Chemistry:Iodobenzene dichloride|iodobenzene dichloride]] (PhICl<sub>2</sub>), used for the selective chlorination of [[Chemistry:Alkene|alkene]]s and [[Chemistry:Alkyne|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 iodoform test, where [[Chemistry:Iodoform|iodoform]] (CHI<sub>3</sub>) is produced by the exhaustive iodination of a methyl ketone (or another compound capable of being oxidised to a methyl ketone), as follows:<ref name="March">{{March6th}}</ref>

:[[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}}</ref> For example, [[Chemistry:Iodoacetamide|iodoacetamide]] and [[Chemistry:Iodoacetic acid|iodoacetic acid]] denature proteins by irreversibly alkylating [[Chemistry:Cysteine|cysteine]] residues and preventing the reformation of [[Biology:Disulfide|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 [[Chemistry:Finkelstein reaction|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 alkyl chloride or an alkyl bromide is converted to an alkyl iodide by treatment with a solution of [[Chemistry:Sodium iodide|sodium iodide]] in [[Chemistry:Acetone|acetone]]. Sodium iodide is soluble in acetone and [[Chemistry:Sodium chloride|sodium chloride]] and [[Chemistry:Sodium bromide|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 [[Chemistry: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 = [[Chemistry:Chem. Rev.|Chem. Rev.]] | 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 = J. Am. Chem. Soc. | year = 1964 | volume = 86 | pages = 4645–4650 | doi = 10.1021/ja01075a025 | issue = 21}}</ref>

==Occurrence and production==
Iodine is the least abundant of the stable halogens, comprising only 0.46&nbsp;[[Parts per million|parts per million]] of Earth's crustal rocks (compare: fluorine 544&nbsp;ppm, chlorine 126&nbsp;ppm, bromine 2.5&nbsp;ppm).<ref name="Greenwood795">Greenwood and Earnshaw, pp. 795–796.</ref> Among the 83 [[Chemistry:Chemical element#Occurrence and origin on Earth|elements which occur in significant quantities]] (elements 1–42, 44–60, 62–83, 90 and 92), it ranks 61st in abundance. Iodide minerals are rare, and most deposits that are concentrated enough for economical extraction are [[Chemistry:Iodate|iodate]] minerals instead. Examples include lautarite, Ca(IO<sub>3</sub>)<sub>2</sub>, and dietzeite, 7Ca(IO<sub>3</sub>)<sub>2</sub>·8CaCrO<sub>4</sub>.<ref name="Greenwood795" /> These are the minerals that occur as trace impurities in the [[Earth:Caliche|caliche]], found in {{wipe|Chile}}, whose main product is [[Chemistry:Sodium nitrate|sodium nitrate]]. In total, they can contain at least 0.02% and at most 1% iodine by mass.<ref name="Elzea">{{cite book |title = Industrial Minerals & Rocks: Commodities, Markets, and Uses |publisher = SME |date = 2006 |isbn = 978-0-87335-233-8 |url = https://books.google.com/books?id=zNicdkuulE4C |pages = 541–552 | veditors = Kogel JE, Trivedi NC, Barker JM, Krukowski ST }}</ref> [[Chemistry:Sodium iodate|Sodium iodate]] is extracted from the caliche and reduced to iodide by [[Chemistry:Sodium bisulfite|sodium bisulfite]]. This solution is then reacted with freshly extracted iodate, resulting in comproportionation to iodine, which may be filtered off.<ref name="Greenwood800" />

The caliche was the main source of iodine in the 19th century and continues to be important today, replacing [[Biology:Kelp|kelp]] (which is no longer an economically viable source),<ref>{{cite journal |url = https://books.google.com/books?id=wW8KAAAAIAAJ&pg=PA185 | vauthors = Stanford EC |journal = Journal of the Society of Arts |title = On the Economic Applications of Seaweed |date = 1862 |pages = 185–189}}</ref> but in the late 20th century [[Earth:Brine|brine]]s emerged as a comparable source. The Japanese Minami Kanto gas field east of [[Place:Tokyo|Tokyo]] and the American Anadarko Basin gas field in northwest Oklahoma are the two largest such sources. The brine is hotter than 60&nbsp;°C from the depth of the source. The [[Earth:Brine|brine]] is first [[Chemistry:List of purification methods in chemistry|purified]] and acidified using [[Chemistry:Sulfuric acid|sulfuric acid]], then the iodide present is oxidised to iodine with [[Chemistry:Chlorine|chlorine]]. An iodine solution is produced, but is dilute and must be concentrated. Air is blown into the solution to evaporate the iodine, which is passed into an absorbing tower, where [[Chemistry:Sulfur dioxide|sulfur dioxide]] reduces the iodine. The [[Chemistry:Hydrogen iodide|hydrogen iodide]] (HI) is reacted with chlorine to precipitate the iodine. After filtering and purification the iodine is packed.<ref name="Elzea" /><ref>{{cite journal |journal = Geochemical Journal |volume = 40 |page = 475 |date = 2006 |title = Chemical and isotopic compositions of brines from dissolved-in-water type natural gas fields in Chiba, Japan | vauthors = Maekawa T, Igari SI, Kaneko N |doi = 10.2343/geochemj.40.475 |issue = 5 |bibcode = 2006GeocJ..40..475M|doi-access = free }}</ref>

: 2 HI + Cl<sub>2</sub> → I<sub>2</sub>↑ + 2 HCl
: I<sub>2</sub> + 2 H<sub>2</sub>O + SO<sub>2</sub> → 2 HI + H<sub>2</sub>SO<sub>4</sub>
: 2 HI + Cl<sub>2</sub> → I<sub>2</sub>↓ + 2 HCl

These sources ensure that Chile and Japan are the largest producers of iodine today.<ref name="Greenwood795" /> Alternatively, the brine may be treated with [[Chemistry:Silver nitrate|silver nitrate]] to precipitate out iodine as [[Chemistry:Silver iodide|silver iodide]], which is then decomposed by reaction with iron to form metallic silver and a solution of [[Chemistry:Iron(II) iodide|iron(II) iodide]]. The iodine may then be liberated by displacement with [[Chemistry:Chlorine|chlorine]].<ref name="Greenwood799">Greenwood and Earnshaw, p. 799.</ref>

==Applications==
About half of all produced iodine goes into various organoiodine compounds, another 15% remains as the pure element, another 15% is used to form [[Biology:Potassium iodide|potassium iodide]], and another 15% for other inorganic iodine compounds.<ref name="Greenwood800" /> Among the major uses of iodine compounds are [[Chemistry:Catalyst|catalyst]]s, animal feed supplements, stabilisers, dyes, colourants and pigments, pharmaceutical, sanitation (from [[Chemistry:Tincture of iodine|tincture of iodine]]), and photography; minor uses include smog inhibition, cloud seeding, and various uses in analytical chemistry.<ref name="Greenwood800" />

===Chemical analysis===
[[File:Testing seed for starch.jpg|thumb|Testing a seed for starch with a solution of iodine]]
The iodide and iodate anions are often used for quantitative volumetric analysis, for example in [[Chemistry:Iodometry|iodometry]]. Iodine and starch form a blue complex, and this reaction is often used to test for either starch or iodine and as an [[Physics:Redox indicator|indicator]] in iodometry. The iodine test for starch is still used to detect counterfeit banknotes printed on starch-containing paper.<ref name="Emsley">{{cite book | vauthors = Emsley J | title = Nature's Building Blocks | edition = Hardcover, First | publisher = [[Organization:Oxford University Press|Oxford University Press]] | date = 2001 | pages = [https://archive.org/details/naturesbuildingb0000emsl/page/244 244–250] | isbn = 978-0-19-850340-8 | url = https://archive.org/details/naturesbuildingb0000emsl/page/244 }}</ref>

The [[Chemistry:Iodine value|iodine value]] is the mass of iodine in grams that is consumed by 100 grams of a [[Chemistry:Chemical substance|chemical substance]] typically fats or oils. Iodine numbers are often used to determine the amount of unsaturation in [[Chemistry:Fatty acid|fatty acid]]s. This unsaturation is in the form of [[Chemistry:Double bond|double bond]]s, which react with iodine compounds.

[[Chemistry:Potassium tetraiodomercurate(II)|Potassium tetraiodomercurate(II)]], K<sub>2</sub>HgI<sub>4</sub>, is also known as Nessler's reagent. It is often used as a sensitive spot test for [[Chemistry:Ammonia|ammonia]]. Similarly, [[Chemistry:Mayer's reagent|Mayer's reagent]] (potassium tetraiodomercurate(II) solution) is used as a precipitating reagent to test for [[Chemistry:Alkaloid|alkaloid]]s.<ref>{{cite journal | last1=Szász | first1=György | last2=Buda | first2=László | title=Contribution to the reaction of alkaloids with potassium tetraiodomercurate | journal=Fresenius' Zeitschrift für Analytische Chemie| publisher=Springer Science and Business Media LLC | volume=253 | issue=5 | year=1971 | issn=0016-1152 | doi=10.1007/bf00426350 | pages=361–363| s2cid=91439011 }}</ref> Aqueous alkaline iodine solution is used in the iodoform test for methyl ketones.<ref name="March" />

=== Spectroscopy ===
The spectrum of the iodine molecule, I<sub>2</sub>, consists of (not exclusively) tens of thousands of sharp spectral lines in the wavelength range 500–700&nbsp;nm. It is therefore a commonly used wavelength reference (secondary standard). By measuring with a [[Physics:Saturated absorption spectroscopy|spectroscopic Doppler-free technique]] while focusing on one of these lines, the [[Physics:Hyperfine structure|hyperfine]] structure of the iodine molecule reveals itself. A line is now resolved such that either 15 components (from even rotational quantum numbers, ''J''<sub>even</sub>), or 21 components (from odd rotational quantum numbers, ''J''<sub>odd</sub>) are measurable.<ref>{{cite journal| vauthors = Sansonetti CJ |title=Precise measurements of hyperfine components in the spectrum of molecular iodine | journal = Journal of the Optical Society of America B | date = August 1997 | volume = 14 | issue = 8 | pages = 1913–1920 |language=en |osti=464573 |doi=10.2172/464573|url=https://digital.library.unt.edu/ark:/67531/metadc684462/}}</ref>

Caesium iodide and thallium-doped sodium iodide are used in crystal [[Physics:Scintillator|scintillator]]s for the detection of gamma rays. The efficiency is high and energy dispersive spectroscopy is possible, but the resolution is rather poor.

===Spacecraft propulsion===
Propulsion systems employing iodine as the [[Physics:Propellant|propellant]] can be built more compactly, with less mass (and cost), and operate more efficiently than the [[Physics:Gridded ion thruster|gridded ion thruster]]s that were utilized to propel previous spacecraft, such as Japan's [[Astronomy:Hayabusa|Hayabusa]] probes, ESA's GOCE satellite, or NASA's [[Astronomy:Double Asteroid Redirection Test|DART]] mission, all of which used xenon as the [[Physics:Working mass|reaction mass]]. Iodine's [[Physics:Standard atomic weight|atomic weight]] is only 3.3% less than that of xenon, while its first two [[Physics:Ionization energy|ionisation energies]] average 12% less; together, these make iodine ions a promising substitute.<ref name="ThrustMe1">{{cite journal |vauthors=Rafalskyi D, Martínez JM, Habl L, Zorzoli Rossi E, Proynov P, Boré A, Baret T, Poyet A, Lafleur T, Dudin S, Aanesland A |date=November 2021 |title=In-orbit demonstration of an iodine electric propulsion system |journal=Nature |volume=599 |issue=7885 |pages=411–415 |bibcode=2021Natur.599..411R |doi=10.1038/s41586-021-04015-y |pmc=8599014 |pmid=34789903 |quote=''Both atomic and molecular iodine ions are accelerated by high-voltage grids to generate thrust, and a highly collimated beam can be produced with substantial iodine dissociation.''}}</ref><ref name="ThrustMe2">{{cite web |url=https://www.cnet.com/news/in-a-space-first-scientists-test-ion-thrusters-powered-by-iodine/ |title=In a space first, scientists test ion thrusters powered by iodine  | vauthors = Ravisetti M |date=18 November 2021 |website=CNET |publisher=Red Ventures |access-date=2021-11-29}}</ref> However, iodine introduces chemical reactivity issues not present in xenon plasmas.<ref>Rogers, James Daniel. ''Hollow Cathode Materials in an Iodine Plasma Enviroment''. Diss. The University of Alabama, 2023.</ref>

Use of iodine should allow more widespread application of ion-thrust technology, particularly with smaller-scale space vehicles.<ref name="ThrustMe2" /> According to the [[Organization:European Space Agency|European Space Agency]], "This small but potentially disruptive innovation could help to clear the skies of [[Astronomy:Space debris|space junk]], by enabling tiny satellites to self-destruct cheaply and easily at the end of their missions, by steering themselves into the atmosphere where they would burn up."<ref name="ThrustMe3">{{cite web |url=https://www.esa.int/ESA_Multimedia/Images/2021/01/Iodine_thruster_used_to_change_the_orbit_of_a_small_satellite_for_the_first_time_ever#.YaUuCq-kYyQ.link |title=Iodine thruster used to change the orbit of a small satellite for the first time ever |author=<!--Not stated--> |date=22 January 2021 |website=www.esa.int |publisher=The European Space Agency |access-date=2021-11-29}}</ref>

In early 2021, the French group [[Company:ThrustMe|ThrustMe]] performed an in-orbit demonstration of an electric-powered [[Engineering:Ion thruster|ion thruster]] for spacecraft, where iodine was used in lieu of [[Chemistry:Xenon|xenon]] as the source of [[Physics:Plasma|plasma]], in order to generate [[Thrust|thrust]] by accelerating [[Physics:Ion|ion]]s with an [[Physics:Electric field#Electrostatic fields|electrostatic]] field.<ref name="ThrustMe1" />

===Medicine===
{{Main|Medicine:Iodine (medical use)}}
====Elemental iodine====
Elemental iodine is used as an [[Chemistry:Antiseptic|antiseptic]] either as the element, or as the water-soluble [[Chemistry:Triiodide|triiodide]] anion I<sub>3</sub><sup>−</sup> generated ''in situ'' by adding [[Chemistry:Iodide|iodide]] to poorly water-soluble elemental iodine (the reverse chemical reaction makes some free elemental iodine available for antisepsis). Elemental iodine may also be used to treat [[Medicine:Iodine deficiency|iodine deficiency]].<ref name="WHO2008">{{cite book | title = WHO Model Formulary 2008 | year = 2009 | isbn = 9789241547659 | vauthors = ((World Health Organization)) | veditors = Stuart MC, Kouimtzi M, Hill SR | hdl = 10665/44053 |  publisher = World Health Organization | hdl-access=free | page=499 }}</ref>

In the alternative, iodine may be produced from [[Chemistry:Iodophor|iodophor]]s, which contain iodine complexed with a solubilizing agent (the iodide ion may be thought of loosely as the iodophor in triiodide water solutions). Examples of such preparations include:<ref>{{cite book | vauthors = Block SS |title=Disinfection, sterilization, and preservation |publisher=Lippincott Williams & Wilkins |location=Hagerstwon, MD |date=2001 |page=159 |isbn=978-0-683-30740-5}}</ref>

* [[Chemistry:Tincture of iodine|Tincture of iodine]]: iodine in ethanol, or iodine and [[Chemistry:Sodium iodide|sodium iodide]] in a mixture of ethanol and water.
* [[Chemistry:Lugol's iodine|Lugol's iodine]]: iodine and iodide in water alone, forming mostly triiodide. Unlike tincture of iodine, Lugol's iodine has a minimised amount of the free iodine (I<sub>2</sub>) component.
* Povidone iodine (an [[Chemistry:Iodophor|iodophor]]).
* Iodine-V: iodine (I<sub>2</sub>)  and [[Chemistry:Fulvic acid|fulvic acid]] form a [[Chemistry:Clathrate compound|clathrate compound]] (iodine molecules are "caged" by fulvic acid in this host-guest complex). A water-soluble, solid, stable, crystalline complex. Unlike other iodophors, Iodine-V only contains iodine in molecular (I<sub>2</sub>) form.<ref>{{cite journal | vauthors = Köntös Z | title = Efficacy of "Essential Iodine Drops" against Severe Acute Respiratory Syndrome-Coronavirus 2 (SARS-CoV-2) | journal = PLOS ONE | volume = 16 | issue = 7 | pages = e0254341 | date = 2021-07-09 | pmid = 34242340 | pmc = 8270147 | doi = 10.1371/journal.pone.0254341 | bibcode = 2021PLoSO..1654341K | doi-access = free }}</ref>

The antimicrobial action of iodine is quick and works at low concentrations, and thus it is used in operating theatres.<ref>{{cite journal | vauthors = Patwardhan N, Kelkar U | title = Disinfection, sterilization and operation theater guidelines for dermatosurgical practitioners in India | journal = Indian Journal of Dermatology, Venereology and Leprology | volume = 77 | issue = 1 | pages = 83–93 | date = 2011 | pmid = 21220895 | doi = 10.4103/0378-6323.74965 | doi-access = free }}</ref> Its specific mode of action is unknown. It penetrates into microorganisms and attacks particular amino acids (such as [[Chemistry:Cysteine|cysteine]] and [[Chemistry:Methionine|methionine]]), nucleotides, and [[Chemistry:Fatty acid|fatty acid]]s, ultimately resulting in [[Biology:Cell death|cell death]]. It also has an antiviral action, but nonlipid viruses and [[Biology:Parvovirus|parvovirus]]es are less sensitive than lipid enveloped viruses. Iodine probably attacks surface proteins of [[Biology:Viral envelope|enveloped]] viruses, and it may also destabilise membrane fatty acids by reacting with [[Chemistry:Saturated and unsaturated compounds|unsaturated]] [[Chemistry:Carbon–carbon bond|carbon bonds]].<ref name="pmid9880479">{{cite journal | vauthors = McDonnell G, Russell AD | title = Antiseptics and disinfectants: activity, action, and resistance | journal = Clinical Microbiology Reviews | volume = 12 | issue = 1 | pages = 147–179 | date = January 1999 | pmid = 9880479 | pmc = 88911 | doi = 10.1128/CMR.12.1.147 }}</ref>

====Other formulations====
Before the advent of organic chelating agents, salts of iodide were given orally in the treatment of [[Medicine:Lead poisoning|lead]] or [[Medicine:Mercury poisoning|mercury poisoning]], such as heavily popularized by Louis Melsens and many nineteenth and early twentieth century doctors.<ref>[https://books.google.com/books?id=Cj9FAAAAcAAJ&pg=PP1 "Sur l'emploi de l'iodure de potassium pour combattre les affections saturnines et mercurielles"], in ''Annales de chimie et de physique'', t. 26, 3<sup>e</sup> série, 1849.</ref><ref name=KI_en>[https://www.ncbi.nlm.nih.gov/pmc/articles/PMC5192934/ "On the Employment of Iodide of Potassium as a Remedy for the Affections Caused by Lead and Mercury"], in ''Br Foreign Med Chir Rev.'' 1853 Jan; 11(21): 201–224.</ref>

In medicine, a saturated solution of [[Biology:Potassium iodide|potassium iodide]] is used to treat acute thyrotoxicosis. It is also used to block uptake of [[Physics:Iodine-131|iodine-131]] in the thyroid gland (see isotopes section above), when this isotope is used as part of radiopharmaceuticals (such as [[Chemistry:Iobenguane|iobenguane]]) that are not targeted to the thyroid or thyroid-type tissues.<ref>{{cite web|url=http://hazard.com/msds/mf/baker/baker/files/p5906.htm |title=Solubility of KI in water |publisher=Hazard.com |date=1998-04-21 |access-date=2013-01-21}}</ref><ref>{{cite web|url=http://www.eanm.org/scientific_info/guidelines/gl_radio_ther_benzyl.pdf|archive-url=https://web.archive.org/web/20090617073253/http://www.eanm.org/scientific_info/guidelines/gl_radio_ther_benzyl.pdf | title=EANM procedure guidelines for 131I-meta-iodobenzylguanidine (131I-mIBG) therapy|url-status=dead|archive-date=17 June 2009|date=17 June 2009}}</ref>

Iodine-131 (usually as iodide) is a component of [[Physics:Nuclear fallout|nuclear fallout]], and is particularly dangerous owing to the thyroid gland's propensity to concentrate ingested iodine and retain it for periods longer than this isotope's radiological half-life of eight days. For this reason, people at risk of exposure to environmental radioactive iodine (iodine-131) in fallout may be instructed to take non-radioactive potassium iodide tablets. The typical adult dose is one 130&nbsp;mg tablet per 24 hours, supplying 100&nbsp;mg (100,000 micrograms) of ionic iodine. (The typical daily dose of iodine for normal health is of order 100 micrograms; see "Dietary Intake" below.) Ingestion of this large dose of non-radioactive iodine minimises the uptake of radioactive iodine by the thyroid gland.<ref>[http://www.emergency.cdc.gov/radiation/ki.asp "CDC Radiation Emergencies"], U.S. Centers for Disease Control, 11 October 2006, accessed 14 November 2010.</ref>

[[File:Diatrizoic acid.svg|thumb|right|Diatrizoic acid, an iodine-containing radiocontrast agent]]

As an element with high [[Physics:Electron density|electron density]] and atomic number, iodine absorbs X-rays weaker than 33.3&nbsp;keV due to the [[Physics:Photoelectric effect|photoelectric effect]] of the innermost electrons.<ref>{{cite book | vauthors = Lancaster JL | chapter-url = http://ric.uthscsa.edu/personalpages/lancaster/DI-II_Chapters/DI_chap4.pdf | chapter = Chapter 4: Physical Determinants of Contrast | archive-url =https://web.archive.org/web/20151010172937/http://ric.uthscsa.edu/personalpages/lancaster/DI-II_Chapters/DI_chap4.pdf | archive-date=10 October 2015 | title = Physics of Medical X-Ray Imaging | publisher = The University of Texas Health Science Center }}</ref> Organoiodine compounds are used with intravenous injection as X-ray radiocontrast agents. This application is often in conjunction with advanced X-ray techniques such as [[Medicine:Angiography|angiography]] and [[CT scan]]ning. At present, all water-soluble radiocontrast agents rely on [[Chemistry:Iodinated contrast|iodine-containing compounds]].

===Others===
{{main|Chemistry:Iodine value}}
The production of [[Chemistry:Ethylenediamine dihydroiodide|ethylenediamine dihydroiodide]], provided as a [[Biology:Nutrition|nutritional supplement]] for livestock, consumes a large portion of available iodine. Another significant use is a catalyst for the production of [[Chemistry:Acetic acid|acetic acid]] by the [[Chemistry:Monsanto process|Monsanto]] and [[Chemistry:Cativa process|Cativa process]]es. In these technologies, which support the world's demand for acetic acid, [[Chemistry:Hydroiodic acid|hydroiodic acid]] converts the [[Chemistry:Methanol|methanol]] feedstock into methyl iodide, which undergoes [[Chemistry:Carbonylation|carbonylation]]. Hydrolysis of the resulting acetyl iodide regenerates hydroiodic acid and gives acetic acid.<ref name="Ullmann">{{cite book | vauthors = Lyday PA, Kaiho T | chapter = Iodine and Iodine Compounds | title = Ullmann's Encyclopedia of Industrial Chemistry | date = 2015 | publisher = Wiley-VCH | location = Weinheim | doi = 10.1002/14356007.a14_381.pub2 | volume = A14 | pages = 382–390 | isbn = 9783527306732 }}</ref>

Inorganic iodides find specialised uses. Titanium, zirconium, hafnium, and [[Chemistry:Thorium|thorium]] are purified by the [[Physics:Van Arkel–de Boer process|van Arkel–de Boer process]], which involves the reversible formation of the tetraiodides of these elements. Silver iodide is a major ingredient to traditional photographic film. Thousands of kilograms of silver iodide are used annually for cloud seeding to induce rain.<ref name = Ullmann/>

The organoiodine compound [[Chemistry:Erythrosine|erythrosine]] is an important food coloring agent. Perfluoroalkyl iodides are precursors to important surfactants, such as [[Chemistry:Perfluorooctanesulfonic acid|perfluorooctanesulfonic acid]].<ref name = Ullmann/>

The [[Chemistry:Iodine clock reaction|iodine clock reaction]] (in which iodine also serves as a test for starch, forming a dark blue complex),<ref name="Greenwood800" /> is a popular educational demonstration experiment and example of a seemingly oscillating reaction (it is only the concentration of an intermediate product that oscillates).

Although iodine has widespread roles in many species, agents containing it can exert a ''differential'' effect upon different species in an agricultural system. The growth of all strains of ''[[Biology:Fusarium verticillioides|Fusarium verticillioides]]'' is significantly inhibited by an iodine-containing fungistatic (AJ1629-34EC) at concentrations that do not harm the crop. This might be a less toxic anti-fungal [[Biology:Pesticide|agricultural treatment]] due to its relatively natural chemistry.<ref name="Yates-et-al-2004" />

{{sup|125}}I is used as the radiolabel in investigating which [[Biology:Ligand (biochemistry)|ligand]]s go to which plant pattern recognition receptors (PRRs).<ref name="Boutrot-Zipfel-2017">{{cite journal | vauthors = Boutrot F, Zipfel C | title = Function, Discovery, and Exploitation of Plant Pattern Recognition Receptors for Broad-Spectrum Disease Resistance | journal = Annual Review of Phytopathology | volume = 55 | issue = 1 | pages = 257–286 | date = August 2017 | pmid = 28617654 | doi = 10.1146/annurev-phyto-080614-120106 | publisher = Annual Reviews | doi-access = free }}</ref>

==Biological role==
{{main|Biology:Iodine in biology}}
[[File:Thyroid system.svg|thumb|upright=1.2|The [[Biology:Thyroid|thyroid]] system of the thyroid hormones T<sub>3</sub> and T<sub>4</sub>]]
[[File:Carte iodurie france µg par jour d'après Mornex 1987 Le Guen 2000.jpg|thumb|upright=1.2|Comparison of the iodine content in urine in France (in microgramme/day), for some regions and departments (average levels of urine iodine, measured in micrograms per liter at the end of the twentieth century (1980 to 2000))<ref>Mornex, 1987 and Le Guen et al., 2000, cited by {{cite journal | vauthors = Le Guen B, Hemidy PY, Gonin M, Bailloeuil C, Van Boxsom D, Renier S, Garcier Y | year = 2001 | title = Arguments et retour d'expérience sur la distribution d'iode stable autour des centrales nucléaires françaises | url = https://www.researchgate.net/publication/245276139 | journal = Radioprotection | volume = 36 | issue = 4| pages = 417–430 | doi = 10.1051/radiopro:2001101 | doi-access = free }}</ref>]]
Iodine is an essential element for life and, at atomic number ''Z'' = 53, is the heaviest element commonly needed by living organisms. ([[Chemistry:Lanthanum|Lanthanum]] and the other [[Chemistry:Lanthanide|lanthanide]]s, as well as [[Chemistry:Tungsten|tungsten]] with ''Z'' = 74 and [[Chemistry:Uranium|uranium]] with ''Z'' = 92, are used by a few microorganisms.<ref>{{cite journal | vauthors = Pol A, Barends TR, Dietl A, Khadem AF, Eygensteyn J, Jetten MS, Op den Camp HJ | title = Rare earth metals are essential for methanotrophic life in volcanic mudpots | journal = Environmental Microbiology | volume = 16 | issue = 1 | pages = 255–264 | date = January 2014 | pmid = 24034209 | doi = 10.1111/1462-2920.12249 | bibcode = 2014EnvMi..16..255P | url = https://repository.ubn.ru.nl//bitstream/handle/2066/128108/128108.pdf }}</ref><ref>{{cite journal| title = Identification of molybdopterin as the organic component of the tungsten cofactor in four enzymes from hyperthermophilic Archaea | author = Johnson JL, Rajagopalan KV, Mukund S, Adams MW. | journal = [[Biology:Journal of Biological Chemistry|Journal of Biological Chemistry]] |date = 5 March 1993 |volume =268 |issue=7 | pages = 4848–52| doi = 10.1016/S0021-9258(18)53474-8 |pmid= 8444863 | doi-access = free }}</ref><ref>{{cite journal | vauthors = Koribanics NM, Tuorto SJ, Lopez-Chiaffarelli N, McGuinness LR, Häggblom MM, Williams KH, Long PE, Kerkhof LJ | title = Spatial distribution of an uranium-respiring betaproteobacterium at the Rifle, CO field research site | journal = PLOS ONE | volume = 10 | issue = 4 | pages = e0123378 | year = 2015 | pmid = 25874721 | pmc = 4395306 | doi = 10.1371/journal.pone.0123378 | doi-access = free | bibcode = 2015PLoSO..1023378K }}</ref>) It is required for the synthesis of the growth-regulating thyroid hormones thyroxine and [[Chemistry:Triiodothyronine|triiodothyronine]] (T<sub>4</sub> and T<sub>3</sub> respectively, named after their number of iodine atoms). A deficiency of iodine leads to decreased production of T<sub>3</sub> and T<sub>4</sub> and a concomitant enlargement of the [[Biology:Thyroid|thyroid tissue]] in an attempt to obtain more iodine, causing the disease known as simple goitre. The major form of thyroid hormone in the blood is thyroxine (T<sub>4</sub>), which has a longer half-life than T<sub>3</sub>. In humans, the ratio of T<sub>4</sub> to T<sub>3</sub> released into the blood is between 14:1 and 20:1. T<sub>4</sub> is converted to the active T<sub>3</sub> (three to four times more potent than T<sub>4</sub>) within [[Biology:Cell|cells]] by [[Biology:Deiodinase|deiodinase]]s (5'-iodinase). These are further processed by [[Chemistry:Decarboxylation|decarboxylation]] and deiodination to produce iodothyronamine (T<sub>1</sub>a) and [[Chemistry:Thyronamine|thyronamine]] (T<sub>0</sub>a'). All three isoforms of the deiodinases are [[Chemistry:Selenium|selenium]]-containing enzymes; thus dietary selenium is essential for T<sub>3</sub> production.<ref>{{cite web|url=http://emedicine.medscape.com/article/819692-overview#showall| vauthors = Irizarry L |title=Thyroid Hormone Toxicity|website=Medscape|publisher=WedMD LLC|date=23 April 2014|access-date=2 May 2014}}</ref>

Iodine accounts for 65% of the molecular weight of T<sub>4</sub> and 59% of T<sub>3</sub>. Fifteen to 20&nbsp;mg of iodine is concentrated in thyroid tissue and hormones, but 70% of all iodine in the body is found in other tissues, including mammary glands, eyes, gastric mucosa, fetal thymus, cerebro-spinal fluid and choroid plexus, arterial walls, the cervix, and salivary glands. During pregnancy, the placenta is able to store and accumulate iodine.<ref>{{cite journal |last1=Burns |first1=R |last2=O'Herlihy |first2=C |last3=Smyth |first3=PP |title=The placenta as a compensatory iodine storage organ. |journal=Thyroid |date=May 2011 |volume=21 |issue=5 |pages=541–6 |doi=10.1089/thy.2010.0203 |pmid=21417918}}</ref><ref>{{cite journal |last1=Neven |first1=KY |last2=Marien |first2=CBD |last3=Janssen |first3=BG |last4=Roels |first4=HA |last5=Waegeneers |first5=N |last6=Nawrot |first6=TS |last7=Ruttens |first7=A |title=Variability of iodine concentrations in the human placenta. |journal=Scientific Reports |date=13 January 2020 |volume=10 |issue=1 |pages=161 |doi=10.1038/s41598-019-56775-3 |pmid=31932629|pmc=6957482 |bibcode=2020NatSR..10..161N }}</ref> In the cells of those tissues, iodide enters directly by [[Biology:Sodium-iodide symporter|sodium-iodide symporter]] (NIS). The action of iodine in mammary tissue is related to fetal and neonatal development, but in the other tissues, it is (at least) partially unknown.<ref name="Patrick2008" /> 

===Dietary recommendations and intake===
The daily levels of intake recommended by the United States National Academy of Medicine are between 110 and 130 [[Microgram|µg]] for infants up to 12 months, 90&nbsp;µg for children up to eight years, 130&nbsp;µg for children up to 13 years, 150&nbsp;µg for adults, 220&nbsp;µg for pregnant women and 290&nbsp;µg for lactation.<ref name="lpi" /><ref>{{cite web|url=http://iom.edu/en/Global/News%20Announcements/~/media/Files/Activity%20Files/Nutrition/DRIs/DRISummaryListing2.ashx |archive-url=https://web.archive.org/web/20091030004039/http://iom.edu/en/Global/News%20Announcements/~/media/Files/Activity%20Files/Nutrition/DRIs/DRISummaryListing2.ashx |url-status=dead |archive-date=30 October 2009 |title=Dietary Reference Intakes (DRIs): Recommended Intakes for Individuals, Vitamins |publisher=Institute of Medicine |date=2004 |access-date=9 June 2010 }}</ref> The Tolerable Upper Intake Level (UL) for adults is 1,100&nbsp;μg/day.<ref name="InstituteofMedicine">{{cite book| author = United States National Research Council| date = 2000|title = Dietary Reference Intakes for Vitamin A, Vitamin K, Arsenic, Boron, Chromium, Copper, Iodine, Iron, Manganese, Molybdenum, Nickel, Silicon, Vanadium, and Zinc|pages = 258–259| publisher = National Academies Press| url = http://books.nap.edu/openbook.php?record_id=10026&page=258| doi = 10.17226/10026| pmid = 25057538| isbn = 978-0-309-07279-3}}</ref> This upper limit was assessed by analyzing the effect of supplementation on [[Biology:Thyroid-stimulating hormone|thyroid-stimulating hormone]].<ref name="Patrick2008" />

The European Food Safety Authority (EFSA) refers to the collective set of information as Dietary Reference Values, with Population Reference Intake (PRI) instead of RDA, and Average Requirement instead of EAR; AI and UL are defined the same as in the United States. For women and men ages 18 and older, the PRI for iodine is set at 150&nbsp;μg/day; the PRI during pregnancy or lactation is 200&nbsp;μg/day. For children aged 1–17 years, the PRI increases with age from 90 to 130&nbsp;μg/day. These PRIs are comparable to the U.S. RDAs with the exception of that for lactation.<ref>{{cite web | title = Overview on Dietary Reference Values for the EU population as derived by the EFSA Panel on Dietetic Products, Nutrition and Allergies| year = 2017| url = https://www.efsa.europa.eu/sites/default/files/assets/DRV_Summary_tables_jan_17.pdf}}</ref>

The thyroid gland needs no more than 70&nbsp;μg/day to synthesise the requisite daily amounts of T4 and T3.<ref name="lpi" /> The higher recommended daily allowance levels of iodine seem necessary for optimal function of a number of body systems, including [[Biology:Lactation|lactation]], [[Biology:Gastric mucosa|gastric mucosa]], salivary glands, brain cells, choroid plexus, [[Biology:Thymus|thymus]], and [[Biology:Artery|arterial walls]].<ref name="lpi" /><ref>{{cite journal | vauthors = Venturi S, Venturi M | title = Iodine, thymus, and immunity | journal = Nutrition | volume = 25 | issue = 9 | pages = 977–979 | date = September 2009 | pmid = 19647627 | doi = 10.1016/j.nut.2009.06.002 }}</ref><ref>{{cite journal | vauthors = Ullberg S, Ewaldsson B | title = Distribution of radio-iodine studied by whole-body autoradiography | journal = Acta Radiologica | volume = 2 | pages = 24–32 | date = February 1964 | pmid = 14153759 | doi = 10.3109/02841866409134127 }}</ref><ref name="Venturi, Sebastiano 2014 185–205">{{cite journal| vauthors = Venturi S |title=Iodine, PUFAs and Iodolipids in Health and Disease: An Evolutionary Perspective|journal=Human Evolution|volume= 29 |issue= 1–3|pages=185–205|year=2014|issn=0393-9375}}</ref>

Natural sources of dietary iodine include seafood, such as fish, seaweeds (such as [[Biology:Kelp|kelp]]) and shellfish, dairy products and eggs so long as the animals received enough iodine, and plants grown on iodine-rich soil.<ref>{{cite web| publisher =Iodine Global Network|url =http://ign.org/p142002146.html|archive-url=https://web.archive.org/web/20150813130042/http://ign.org/p142002146.html|archive-date=13 August 2015|title=Where do we get iodine from?|url-status=live}}</ref><ref name = medlineplus002421/> [[Chemistry:Iodised salt|Iodised salt]] is fortified with iodine in the form of [[Chemistry:Sodium iodide|sodium iodide]] or [[Chemistry:Potassium iodate|potassium iodate]].<ref name="medlineplus002421">{{cite encyclopedia| url = https://www.nlm.nih.gov/medlineplus/ency/article/002421.htm|title = Iodine in diet|encyclopedia=MedlinePlus Medical Encyclopedia}}</ref><ref name="American Thyroid Association">{{cite web|title=American Thyroid Association|url=http://www.thyroid.org/iodine-deficiency/|work=thyroid.org|publisher=American Thyroid Association|access-date=4 April 2014}}</ref><ref>{{cite web | url = https://www.waitrose.com/ecom/products/cerebos-iodised-table-salt/559124-79136-79137 | title = Cerebos iodised table salt | year = 2023 | website = Waitrose | access-date = 2023-05-30 | archive-url = https://web.archive.org/web/20230328192627/https://www.waitrose.com/ecom/products/cerebos-iodised-table-salt/559124-79136-79137 | archive-date = 2023-03-28}}</ref>

As of 2000, the median intake of iodine from food in the United States was 240 to 300&nbsp;μg/day for men and 190 to 210&nbsp;μg/day for women.<ref name="InstituteofMedicine" /> The general US population has adequate iodine nutrition,<ref>{{cite journal | vauthors = Caldwell KL, Makhmudov A, Ely E, Jones RL, Wang RY | title = Iodine status of the U.S. population, National Health and Nutrition Examination Survey, 2005–2006 and 2007–2008 | journal = Thyroid | volume = 21 | issue = 4 | pages = 419–427 | date = April 2011 | pmid = 21323596 | doi = 10.1089/thy.2010.0077 | url = https://zenodo.org/record/1235283 }}</ref><ref name="Lueng">{{cite journal | vauthors = Leung AM, Braverman LE, Pearce EN | title = History of U.S. iodine fortification and supplementation | journal = Nutrients | volume = 4 | issue = 11 | pages = 1740–1746 | date = November 2012 | pmid = 23201844 | pmc = 3509517 | doi = 10.3390/nu4111740 | doi-access = free }}</ref> with women of childbearing age and pregnant women having a possible mild risk of deficiency.<ref name="Lueng" /> In Japan, consumption was considered much higher, ranging between 5,280&nbsp;μg/day to 13,800&nbsp;μg/day from dietary seaweed or [[Biology:Kombu|kombu]] kelp,<ref name="Patrick2008">{{cite journal | vauthors = Patrick L | title = Iodine: deficiency and therapeutic considerations | journal = Alternative Medicine Review | volume = 13 | issue = 2 | pages = 116–127 | date = June 2008 | pmid = 18590348 | url = http://www.thorne.com/altmedrev/.fulltext/13/2/116.pdf | url-status = dead | archive-url = https://web.archive.org/web/20130531112100/http://www.thorne.com/altmedrev/.fulltext/13/2/116.pdf | archive-date = 31 May 2013 }}</ref> often in the form of kombu umami extracts for soup stock and potato chips. However, new studies suggest that Japan's consumption is closer to 1,000–3,000&nbsp;μg/day.<ref>{{cite journal | vauthors = Zava TT, Zava DT | title = Assessment of Japanese iodine intake based on seaweed consumption in Japan: A literature-based analysis | journal = Thyroid Research | volume = 4 | pages = 14 | date = October 2011 | pmid = 21975053 | pmc = 3204293 | doi = 10.1186/1756-6614-4-14 | doi-access = free }}</ref> The adult UL in Japan was last revised to 3,000&nbsp;µg/day in 2015.<ref>{{cite web |title=Overview of Dietary Reference Intakes for Japanese (2015) |publisher=Minister of Health, Labour and Welfare, Japan |url = http://www.mhlw.go.jp/file/06-Seisakujouhou-10900000-Kenkoukyoku/Overview.pdf |access-date=14 March 2022}}</ref>

After iodine fortification programs such as iodisation of salt have been implemented, some cases of iodine-induced [[Medicine:Hyperthyroidism|hyperthyroidism]] have been observed (so-called [[Medicine:Jod-Basedow phenomenon|Jod-Basedow phenomenon]]). The condition seems to occur mainly in people over forty, and the risk appears higher when iodine deficiency is severe and the initial rise in iodine intake is high.<ref>{{cite journal | vauthors = Wu T, Liu GJ, Li P, Clar C | title = Iodised salt for preventing iodine deficiency disorders | journal = The Cochrane Database of Systematic Reviews | volume = 2010 | issue = 3 | pages = CD003204 | date = 2002 | pmid = 12137681 | pmc = 9006116 | doi = 10.1002/14651858.CD003204 | veditors = Wu T }}</ref>

===Deficiency===
{{Main|Medicine:Iodine deficiency}}
In areas where there is little iodine in the diet,<ref name="Dissanayake">{{cite journal| vauthors = Dissanayake CB, Chandrajith R, Tobschall HJ |title = The iodine cycle in the tropical environment – implications on iodine deficiency disorders|journal = International Journal of Environmental Studies|volume = 56 |page= 357| doi = 10.1080/00207239908711210|date = 1999|issue = 3| bibcode=1999IJEnS..56..357D }}</ref> typically remote inland areas and semi-arid equatorial climates where no marine foods are eaten, [[Medicine:Iodine deficiency|iodine deficiency]] gives rise to [[Medicine:Hypothyroidism|hypothyroidism]], symptoms of which are extreme fatigue, [[Medicine:Goitre|goitre]], mental slowing, depression, weight gain, and low basal body temperatures.<ref>{{cite book|chapter = Endemic Goiter|title = Endocrinology & metabolism| vauthors = Felig P, Frohman LA |publisher = McGraw-Hill Professional|date = 2001|isbn = 978-0-07-022001-0|chapter-url = https://books.google.com/books?id=AZUUGrp6yUgC&pg=RA1-PA351}}</ref> Iodine deficiency is the leading cause of preventable [[Medicine:Intellectual disability|intellectual disability]], a result that occurs primarily when babies or small children are rendered hypothyroidic by a lack of the element. The addition of iodine to table salt has largely eliminated this problem in wealthier nations, but iodine deficiency remains a serious public health problem in the developing world today.<ref>{{cite web|url =https://www.who.int/nutrition/topics/idd/en/|archive-url =https://web.archive.org/web/20060930020824/http://www.who.int/nutrition/topics/idd/en/|url-status =dead|archive-date =30 September 2006|title = Micronutrient deficiency: iodine deficiency disorders|publisher = WHO}}</ref> Iodine deficiency is also a problem in certain areas of Europe. Information processing, fine motor skills, and visual problem solving are improved by iodine repletion in moderately iodine-deficient children.<ref>{{cite journal | vauthors = Zimmermann MB, Connolly K, Bozo M, Bridson J, Rohner F, Grimci L | title = Iodine supplementation improves cognition in iodine-deficient schoolchildren in Albania: a randomized, controlled, double-blind study | journal = The American Journal of Clinical Nutrition | volume = 83 | issue = 1 | pages = 108–114 | date = January 2006 | pmid = 16400058 | doi = 10.1093/ajcn/83.1.108 | doi-access = free }}</ref>

==Precautions==
===Toxicity===
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| GHSPictograms = {{GHS07}}{{GHS09}}
| GHSSignalWord = Danger
| HPhrases = {{H-phrases|312|332|315|319|335|372|400}}
| PPhrases = {{P-phrases|261|273|280|305|351|338|314}}<ref>{{cite web|url=https://www.sigmaaldrich.com/catalog/product/sigald/207772|title=Iodine 207772|website=I2}}</ref>
| NFPA-H = 3
| NFPA-F = 0
| NFPA-R = 0
| NFPA-S = 
| NFPA_ref = <ref>[http://periodictable.com/Elements/053/data.html Technical data for Iodine]. periodictable.com</ref> 
 }}
}}
Elemental iodine (I<sub>2</sub>) is [[Medicine:Toxicity|toxic]] if taken orally undiluted. The lethal dose for an adult human is 30&nbsp;mg/kg, which is about 2.1–2.4&nbsp;grams for a human weighing 70 to 80&nbsp;kg (even if experiments on rats demonstrated that these animals could survive after eating a 14000&nbsp;mg/kg dose). Excess iodine can be more [[Biology:Cytotoxicity|cytotoxic]] in the presence of [[Medicine:Selenium deficiency|selenium deficiency]].<ref>{{cite journal | vauthors = Smyth PP | title = Role of iodine in antioxidant defence in thyroid and breast disease | journal = BioFactors | volume = 19 | issue = 3–4 | pages = 121–130 | year = 2003 | pmid = 14757962 | doi = 10.1002/biof.5520190304 | s2cid = 7803619 }}</ref> Iodine supplementation in selenium-deficient populations is, in theory, problematic, partly for this reason.<ref name="Patrick2008" /> The toxicity derives from its oxidizing properties, through which it denaturates proteins (including enzymes).<ref>{{cite web |url=http://butane.chem.uiuc.edu/cyerkes/chem104A_S07/Lecture_Notes_104/lect29c.html |title=Lecture 29: Protein Structure and Denaturation | vauthors = Yerkes C |date=2007 |website=chem.uiuc.edu |publisher=University of Illinois |access-date=23 October 2016}}</ref>

Elemental iodine is also a skin irritant. Direct contact with skin can cause damage, and solid iodine crystals should be handled with care. Solutions with high elemental iodine concentration, such as [[Chemistry:Tincture of iodine|tincture of iodine]] and Lugol's solution, are capable of causing [[Chemistry:Dermatoxin|tissue damage]] if used in prolonged cleaning or antisepsis; similarly, liquid [[Chemistry:Povidone-iodine|Povidone-iodine]] (Betadine) trapped against the skin resulted in chemical burns in some reported cases.<ref name="Lowe">{{cite journal | vauthors = Lowe DO, Knowles SR, Weber EA, Railton CJ, Shear NH | title = Povidone-iodine-induced burn: case report and review of the literature | journal = Pharmacotherapy | volume = 26 | issue = 11 | pages = 1641–1645 | date = November 2006 | pmid = 17064209 | doi = 10.1592/phco.26.11.1641 | s2cid = 25708713 }}</ref>

==== Occupational exposure ====
People can be exposed to iodine in the workplace by inhalation, ingestion, skin contact, and eye contact. The [[Organization:Occupational Safety and Health Administration|Occupational Safety and Health Administration]] (OSHA) has set the legal limit ([[Earth:Permissible exposure limit|Permissible exposure limit]]) for iodine exposure in the workplace at 0.1 ppm (1&nbsp;mg/m<sup>3</sup>) during an 8-hour workday. The [[Organization:National Institute for Occupational Safety and Health|National Institute for Occupational Safety and Health]] (NIOSH) has set a [[Earth:Recommended exposure limit|Recommended exposure limit]] (REL) of 0.1&nbsp;ppm (1&nbsp;mg/m<sup>3</sup>) during an 8-hour workday. At levels of 2 ppm, iodine is immediately dangerous to life and health.<ref>{{cite web|title = CDC - NIOSH Pocket Guide to Chemical Hazards - Iodine|url = https://www.cdc.gov/niosh/npg/npgd0342.html|website = cdc.gov|access-date = 2015-11-06}}</ref>

====Allergic reactions====
Some people develop a [[Medicine:Hypersensitivity|hypersensitivity]] to products and foods containing iodine. Applications of tincture of iodine or Betadine can cause rashes, sometimes severe.<ref name="dermnet">DermNet New Zealand Trust, [http://www.dermnetnz.org/treatments/iodine.html Iodine]</ref> Parenteral use of iodine-based contrast agents (see above) can cause reactions ranging from a mild rash to fatal [[Medicine:Anaphylaxis|anaphylaxis]]. Such reactions have led to the misconception (widely held, even among physicians) that some people are allergic to iodine itself; even allergies to iodine-rich seafood have been so construed.<ref>{{cite journal | vauthors = Boehm I | title = Seafood allergy and radiocontrast media: are physicians propagating a myth? | journal = The American Journal of Medicine | volume = 121 | issue = 8 | pages = e19 | date = August 2008 | pmid = 18691465 | doi = 10.1016/j.amjmed.2008.03.035 | doi-access = free }}</ref> In fact, there has never been a confirmed report of a true iodine allergy, and an allergy to elemental iodine or simple iodide salts is theoretically impossible. Hypersensitivity reactions to products and foods containing iodine are apparently related to their other molecular components;<ref name="ucsf">UCSF Department of Radiology & Biomedical Imaging, [http://www.radiology.ucsf.edu/patient-care/patient-safety/contrast/iodine-allergy Iodine Allergy and Contrast Administration]</ref> thus, a person who has demonstrated an allergy to one food or product containing iodine may not have an allergic reaction to another. Patients with various food allergies (shellfish, egg, milk, etc.) do not have an increased risk for a contrast medium hypersensitivity.<ref name="pmid31153557">{{cite journal | vauthors = Lombardo P, Nairz K, Boehm I | title = Patients' safety and the "iodine allergy" - How should we manage patients with iodine allergy before they receive an iodinated contrast medium? | journal = European Journal of Radiology | volume = 116 | issue = 7 | pages = 150–151 | date = July 2019 | pmid = 31153557 | doi = 10.1016/j.ejrad.2019.05.002 | s2cid = 164898934 }}</ref><ref name="ucsf" /> As with all medications, the patient's allergy history should be questioned and consulted before any containing iodine are administered.<ref>{{cite journal | vauthors = Katelaris C |date=2009 |title='Iodine Allergy' label is misleading |journal=Australian Prescriber |volume=32 |pages=125–128 |issue=5 |doi=10.18773/austprescr.2009.061 |doi-access=free }}</ref>
{{clear|right}}

===US DEA List I status===
[[Chemistry:Phosphorus|Phosphorus]] can reduce elemental iodine to [[Chemistry:Hydroiodic acid|hydroiodic acid]], which is a reagent effective for reducing [[Chemistry:Ephedrine|ephedrine]] or [[Chemistry:Pseudoephedrine|pseudoephedrine]] to [[Chemistry:Methamphetamine|methamphetamine]].<ref>{{cite journal| vauthors = Skinner HF |date = 1990|title = Methamphetamine synthesis via hydriodic acid/red phosphorus reduction of ephedrine|journal = Forensic Science International|volume = 48|issue = 2|pages = 123–134|doi = 10.1016/0379-0738(90)90104-7}}</ref> For this reason, iodine was designated by the United States [[Organization:Drug Enforcement Administration|Drug Enforcement Administration]] as a [[Chemistry:DEA list of chemicals#List I chemicals|List I precursor chemical]] under 21 CFR 1310.02.<ref>{{cite web | url=https://www.deadiversion.usdoj.gov/21cfr/cfr/1310/1310_02.htm | title=PART 1310 - Section 1310.02 Substances covered | access-date=5 December 2019 | archive-date=17 October 2017 | archive-url=https://web.archive.org/web/20171017090223/https://www.deadiversion.usdoj.gov/21cfr/cfr/1310/1310_02.htm | url-status=dead }}</ref>

== References ==
{{reflist|refs=

<ref name="Yates-et-al-2004">{{cite journal | vauthors = Yates IE, Arnold JW, Bacon CW, Hinton DM | title=In vitro assessments of diverse plant pathogenic fungi treated with a novel growth control agent | journal=Crop Protection | publisher=[[Company:Elsevier|Elsevier]] BV | volume=23 | issue=12 | year=2004 | issn=0261-2194 | doi=10.1016/j.cropro.2004.03.019 | pages=1169–1176}}</ref>

}}

==Bibliography==

{{Subject bar
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|book1=Iodine
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[[Category:Iodine| ]]
[[Category:Chemical elements]]
[[Category:Halogens]]
[[Category:Reactive nonmetals]]
[[Category:Diatomic nonmetals]]
[[Category:Dietary minerals]]
[[Category:Oxidizing agents]]
[[Category:Gases with color]]

{{Sourceattribution|Iodine}}