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Iodine

April 29, 2023

This article is about the chemical element. For other uses, see Iodine (disambiguation).

Iodine is a chemical element with the symbol I and atomic number 53. The heaviest of the stable halogens, it exists as a semi-lustrous, non-metallic solid at standard conditions that melts to form a deep violet liquid at 114 °C (237 °F), and boils to a violet gas at 184 °C (363 °F). The element was discovered by the French chemist Bernard Courtois in 1811 and was named two years later by Joseph Louis Gay-Lussac, after the Ancient Greek Ιώδης 'violet-coloured'.

Iodine, 53I
Iodine
Pronunciation/ˈaɪədaɪn, -dɪn, -diːn/ ​(EYE-ə-dyne, -⁠din, -⁠deen)
Appearancelustrous metallic gray solid, black/violet liquid, violet gas
Standard atomic weight Ar°(I)
  • 126.90447±0.00003
  • 126.90±0.01 (abridged)[1]
Iodine in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
Br
↑
I
↓
At
tellurium ← iodine → xenon
Atomic number (Z)53
Groupgroup 17 (halogens)
Periodperiod 5
Block  p-block
Electron configuration[Kr] 4d10 5s2 5p5
Electrons per shell2, 8, 18, 18, 7
Physical properties
Phase at STPsolid
Melting point(I2) 386.85 K ​(113.7 °C, ​236.66 °F)
Boiling point(I2) 457.4 K ​(184.3 °C, ​363.7 °F)
Density (near r.t.)4.933 g/cm3
Triple point386.65 K, ​12.1 kPa
Critical point819 K, 11.7 MPa
Heat of fusion(I2) 15.52 kJ/mol
Heat of vaporisation(I2) 41.57 kJ/mol
Molar heat capacity(I2) 54.44 J/(mol·K)
Vapour pressure (rhombic)
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 260 282 309 342 381 457
Atomic properties
Oxidation states−1, +1, +2,[2] +3, +4, +5, +6, +7 (a strongly acidic oxide)
ElectronegativityPauling scale: 2.66
Ionisation energies
  • 1st: 1008.4 kJ/mol
  • 2nd: 1845.9 kJ/mol
  • 3rd: 3180 kJ/mol
Atomic radiusempirical: 140 pm
Covalent radius139±3 pm
Van der Waals radius198 pm
Spectral lines of iodine
Other properties
Natural occurrenceprimordial
Crystal structure ​base-centered orthorhombic
Thermal conductivity0.449 W/(m⋅K)
Electrical resistivity1.3×107 Ω⋅m (at 0 °C)
Magnetic orderingdiamagnetic[3]
Molar magnetic susceptibility−88.7×10−6 cm3/mol (298 K)[4]
Bulk modulus7.7 GPa
CAS Number7553-56-2
History
Discovery and first isolationBernard Courtois (1811)
Isotopes of iodine
Main isotopes Decay
abun­dance half-life (t1/2) mode pro­duct
123I synth 13 h β+100% 123Te
124I synth 4.176 d ε 124Te
125I synth 59.40 d ε 125Te
127I 100% stable
129I trace 1.57×107 y β− 129Xe
131I synth 8.02070 d β−100% 131Xe
135I synth 6.57 h β− 135Xe
 Category: Iodine
| references

Iodine occurs in many oxidation states, including iodide (I−), iodate (IO−
3), and the various periodate anions. It is the least abundant of the stable halogens, being the sixty-first most abundant element. As the heaviest essential mineral nutrient, iodine is required for the synthesis of thyroid hormones.[5] Iodine deficiency affects about two billion people and is the leading preventable cause of intellectual disabilities.[6]

The dominant producers of iodine today are Chile and Japan. Due to its high atomic number and ease of attachment to organic compounds, 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 thyroid cancer. Iodine is also used as a catalyst in the industrial production of acetic acid and some polymers.

It is on the World Health Organization's List of Essential Medicines.[7]

History

In 1811, iodine was discovered by French chemist Bernard Courtois,[8][9] who was born to a manufacturer of saltpetre (an essential component of gunpowder). At the time of the Napoleonic Wars, saltpetre was in great demand in France. Saltpetre produced from French nitre beds required sodium carbonate, which could be isolated from seaweed collected on the coasts of Normandy and Brittany. To isolate the sodium carbonate, seaweed was burned and the ash washed with water. The remaining waste was destroyed by adding 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.[10] Courtois suspected that this material was a new element but lacked funding to pursue it further.[11]

Courtois gave samples to his friends, Charles Bernard Desormes (1777–1838) and Nicolas Clément (1779–1841), to continue research. He also gave some of the substance to chemist Joseph Louis Gay-Lussac (1778–1850), and to physicist 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.[12] On 6 December, Gay-Lussac announced that the new substance was either an element or a compound of oxygen.[13][14][15] Gay-Lussac suggested the name "iode", from the Ancient Greek ἰοειδής (ioeidēs, "violet"), because of the colour of iodine vapor.[8][13] Ampère had given some of his sample to English chemist Humphry Davy (1778–1829), who experimented on the substance and noted its similarity to chlorine.[16] Davy sent a letter dated 10 December to the Royal Society of London stating that he had identified a new element.[17] Arguments erupted between Davy and Gay-Lussac over who identified iodine first, but both scientists acknowledged Courtois as the first to isolate the element.[11]

In 1873 the French medical researcher Casimir Joseph Davaine (1812–1882) discovered the antiseptic action of iodine.[18] Antonio Grossich (1849–1926), an Istrian-born surgeon, was among the first to use 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.[19]

In early periodic tables, iodine was often given the symbol J, for Jod, its name in German.[20]

Properties

 
Iodine vapour in a flask.

Iodine is the fourth halogen, being a member of group 17 in the periodic table, below fluorine, chlorine, and bromine; it is the heaviest stable member of its group. (The fifth and sixth halogens, the radioactive astatine and 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]4d105s25p5, with the seven electrons in the fifth and outermost shell being its valence electrons. 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 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 diatomic molecules with chemical formula I2, 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−, is the strongest reducing agent among the stable halogens, being the most easily oxidised back to diatomic I2.[21] (Astatine goes further, being indeed unstable as At− and readily oxidised to At0 or At+.)[22]

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 ml at 20 °C and 1280 ml at 50 °C; potassium iodide may be added to increase solubility via formation of triiodide ions, among other polyiodides.[23] Nonpolar solvents such as hexane and carbon tetrachloride provide a higher solubility.[24] 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.[23] Charge-transfer complexes 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 alcohols and amines, solvents that form charge-transfer adducts.[25]

 
I2•PPh3 charge-transfer complexes in CH2Cl2. From left to right: (1) I2 dissolved in dichloromethane – no CT complex. (2) A few seconds after excess PPh3 was added – CT complex is forming. (3) One minute later after excess PPh3 was added, the CT complex [Ph3PI]+I− has been formed. (4) Immediately after excess I2 was added, which contains [Ph3PI]+[I3]−.[26]

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.[21] Due to this property Iodine is commonly used to demonstrate sublimation directly from solid to gas, which gives rise to a misconception that it does not melt in atmospheric pressure.[27] Because it has the largest atomic radius among the halogens, iodine has the lowest first ionisation energy, lowest electron affinity, lowest electronegativity and lowest reactivity of the halogens.[21]

 
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 °C. Temperatures greater than 750 °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.[21] Gaseous iodine is composed of I2 molecules with an I–I bond length of 266.6 pm. The I–I bond is one of the longest single bonds known. It is even longer (271.5 pm) in solid orthorhombic crystalline iodine, which has the same crystal structure as chlorine and bromine. (The record is held by iodine's neighbour xenon: the Xe–Xe bond length is 308.71 pm.)[28] 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 semiconducting properties.[21] Iodine is a two-dimensional semiconductor with a band gap of 1.3 eV (125 kJ/mol): it is a semiconductor in the plane of its crystalline layers and an insulator in the perpendicular direction.[21]

Isotopes

Main article: Isotopes of iodine

Of the thirty-seven known isotopes of iodine, only one occurs in nature, iodine-127. The others are radioactive and have half-lives too short to be primordial. As such, iodine is both monoisotopic and mononuclidic and its atomic weight is known to great precision, as it is a constant of nature.[21]

The longest-lived of the radioactive isotopes of iodine is iodine-129, which has a half-life of 15.7 million years, decaying via beta decay to stable xenon-129.[29] 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 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 daughter xenon-129.[30][31][32][33][34] Traces of iodine-129 still exist today, as it is also a cosmogenic nuclide, formed from cosmic ray spallation of atmospheric xenon: these traces make up 10−14 to 10−10 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 10−7 of all terrestrial iodine.[35] Excited states of iodine-127 and iodine-129 are often used in Mössbauer spectroscopy.[21]

The other iodine radioisotopes have much shorter half-lives, no longer than days.[29] Some of them have medical applications involving the thyroid gland, where the iodine that enters the body is stored and concentrated. Iodine-123 has a half-life of thirteen hours and decays by electron capture to tellurium-123, emitting gamma radiation; it is used in nuclear medicine imaging, including single photon emission computed tomography (SPECT) and X-ray computed tomography (X-Ray CT) scans.[36] Iodine-125 has a half-life of fifty-nine days, decaying by electron capture to tellurium-125 and emitting low-energy gamma radiation; the second-longest-lived iodine radioisotope, it has uses in biological assays, nuclear medicine imaging and in radiation therapy as brachytherapy to treat a number of conditions, including prostate cancer, uveal melanomas, and brain tumours.[37] Finally, iodine-131, with a half-life of eight days, beta decays to an excited state of stable 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 thyroid cancer in later life. Other risks include the possibility of non-cancerous growths and thyroiditis.[38]

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 potassium iodide tablets, taken daily for optimal prophylaxis.[39] However, iodine-131 may also be used for medicinal purposes in radiation therapy for this very reason, when tissue destruction is desired after iodine uptake by the tissue.[40] Iodine-131 is also used as a radioactive tracer.[41][42][43][44]

Chemistry and compounds

Halogen bond energies (kJ/mol)[23]
X XX HX BX3 AlX3 CX4
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 carbon monoxide, nitric oxide, and sulfur dioxide (to phosgene, nitrosyl chloride, and 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, rhenium metal reacts with chlorine to form rhenium hexachloride, but with bromine it forms only rhenium pentabromide and iodine can achieve only rhenium tetraiodide.[21] By the same token, however, since iodine has the lowest ionisation energy among the halogens and is the most easily oxidised of them, it has a more significant cationic chemistry and its higher oxidation states are rather more stable than those of bromine and chlorine, for example in iodine heptafluoride.[23]

Charge-transfer complexes

The iodine molecule, I2, dissolves in CCl4 and aliphatic hydrocarbons to give bright violet solutions. In these solvents the absorption band maximum occurs in the 520 – 540 nm region and is assigned to a π* to σ* transition. When I2 reacts with Lewis bases in these solvents a blue shift in I2 peak is seen and the new peak (230 – 330 nm) arises that is due to the formation of adducts, which are referred to as charge-transfer complexes.[45]

Hydrogen iodide

The simplest compound of iodine is hydrogen iodide, HI. It is a colourless gas that reacts with oxygen to give water and iodine. Although it is useful in iodination reactions in the laboratory, it does not have large-scale industrial uses, unlike the other hydrogen halides. Commercially, it is usually made by reacting iodine with hydrogen sulfide or hydrazine:[46]

2 I2 + N2H4 H2O⟶ 4 HI + N2

At room temperature, it is a colourless gas, like all of the hydrogen halides except hydrogen fluoride, since hydrogen cannot form strong hydrogen bonds to the large and only mildly electronegative iodine atom. It melts at −51.0 °C and boils at −35.1 °C. It is an endothermic compound that can exothermically dissociate at room temperature, although the process is very slow unless a 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 kJ/mol.[47]

Aqueous hydrogen iodide is known as hydroiodic acid, which is a strong acid. Hydrogen iodide is exceptionally soluble in water: one litre of water will dissolve 425 litres of hydrogen iodide, and the saturated solution has only four water molecules per molecule of hydrogen iodide.[48] Commercial so-called "concentrated" hydroiodic acid usually contains 48–57% HI by mass; the solution forms an azeotrope with boiling point 126.7 °C at 56.7 g HI per 100 g solution. Hence hydroiodic acid cannot be concentrated past this point by evaporation of water.[47]

Unlike 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 dielectric constant is low and it does not dissociate appreciably into H2I+ and HI−
2 ions – the latter, in any case, are much less stable than the bifluoride ions (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 Cs+ and NR+
4 (R = Me, Et, Bun) may still be isolated. Anhydrous hydrogen iodide is a poor solvent, able to dissolve only small molecular compounds such as nitrosyl chloride and phenol, or salts with very low lattice energies such as tetraalkylammonium halides.[47]

Other binary iodine compounds

With the exception of the noble gases, nearly all elements on the periodic table up to einsteinium (EsI3 is known) are known to form binary compounds with iodine. Until 1990, nitrogen triiodide[49] was only known as an ammonia adduct. Ammonia-free NI3 was found to be isolable at –196 °C but spontaneously decomposes at 0 °C.[50] For thermodynamic reasons related to electronegativity of the elements, neutral sulfur and selenium iodides that are stable at room temperature are also nonexistent, although S2I2 and SI2 are stable up to 183 and 9 K, respectively. As of 2022, no neutral binary selenium iodide has been unambiguously identified (at any temperature).[51] Sulfur- and selenium-iodine polyatomic cations (e.g., [S2I42+][AsF6–]2 and [Se2I42+][Sb2F11–]2) have been prepared and characterized crystallographically.[52]

Given the large size of the iodide anion and iodine's weak oxidising power, high oxidation states are difficult to achieve in binary iodides, the maximum known being in the pentaiodides of niobium, tantalum, and protactinium. Iodides can be made by reaction of an element or its oxide, hydroxide, or carbonate with hydroiodic acid, and then dehydrated by mildly high temperatures combined with either low pressure or anhydrous hydrogen iodide gas. These methods work best when the iodide product is stable to hydrolysis; otherwise, the possibilities include high-temperature oxidative iodination of the element with iodine or hydrogen iodide, high-temperature iodination of a metal oxide or other halide by iodine, a volatile metal halide, carbon tetraiodide, or an organic iodide. For example, molybdenum(IV) oxide reacts with aluminium(III) iodide at 230 °C to give molybdenum(II) iodide. An example involving halogen exchange is given below, involving the reaction of tantalum(V) chloride with excess aluminium(III) iodide at 400 °C to give tantalum(V) iodide:[53]

 

Lower iodides may be produced either through thermal decomposition or disproportionation, or by reducing the higher iodide with hydrogen or a metal, for example:[53]

 

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 aluminium iodide is not). Ionic iodides MIn tend to have the lowest melting and boiling points among the halides MXn of the same element, because the electrostatic forces of attraction between the cations and anions are weakest for the large iodide anion. In contrast, covalent iodides tend to instead have the highest melting and boiling points among the halides of the same element, since iodine is the most polarisable of the halogens and, having the most electrons among them, can contribute the most to van der Waals forces. Naturally, exceptions abound in intermediate iodides where one trend gives way to the other. Similarly, solubilities in water of predominantly ionic iodides (e.g. potassium and calcium) are the greatest among ionic halides of that element, while those of covalent iodides (e.g. silver) are the lowest of that element. In particular, silver iodide is very insoluble in water and its formation is often used as a qualitative test for iodine.[53]

Iodine halides

The halogens form many binary, diamagnetic interhalogen compounds with stoichiometries XY, XY3, XY5, and XY7 (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 ICl+
2 and the dark brown or purplish black compounds of I2Cl+. Apart from these, some pseudohalides are also known, such as cyanogen iodide (ICN), iodine thiocyanate (ISCN), and iodine azide (IN3).[54]

 
Iodine monochloride

Iodine monofluoride (IF) is unstable at room temperature and disproportionates very readily and irreversibly to iodine and iodine pentafluoride, and thus cannot be obtained pure. It can be synthesised from the reaction of iodine with fluorine gas in trichlorofluoromethane at −45 °C, with iodine trifluoride in trichlorofluoromethane at −78 °C, or with silver(I) fluoride at 0 °C.[54] Iodine monochloride (ICl) and iodine monobromide (IBr), on the other hand, are moderately stable. The former, a volatile red-brown compound, was discovered independently by Joseph Louis Gay-Lussac and Humphry Davy in 1813–1814 not long after the discoveries of chlorine and iodine, and it mimics the intermediate halogen bromine so well that Justus von Liebig was misled into mistaking bromine (which he had found) for iodine monochloride. Iodine monochloride and iodine monobromide may be prepared simply by reacting iodine with chlorine or bromine at room temperature and purified by fractional crystallisation. Both are quite reactive and attack even platinum and gold, though not boron, carbon, cadmium, lead, zirconium, niobium, molybdenum, and tungsten. Their reaction with organic compounds depends on conditions. Iodine chloride vapour tends to chlorinate phenol and salicyclic acid, since when iodine chloride undergoes homolytic dissociation, chlorine and iodine are produced and the former is more reactive. However, iodine chloride in tetrachloromethane solution results in iodination being the main reaction, since now heterolytic fission of the I–Cl bond occurs and I+ attacks phenol as an electrophile. However, iodine monobromide tends to brominate phenol even in tetrachloromethane solution because it tends to dissociate into its elements in solution, and bromine is more reactive than iodine.[54] When liquid, iodine monochloride and iodine monobromide dissociate into I
2X+
 and IX−
2 anions (X = Cl, Br); thus they are significant conductors of electricity and can be used as ionising solvents.[54]

Iodine trifluoride (IF3) is an unstable yellow solid that decomposes above −28 °C. It is thus little-known. It is difficult to produce because fluorine gas would tend to oxidise iodine all the way to the pentafluoride; reaction at low temperature with xenon difluoride is necessary. Iodine trichloride, which exists in the solid state as the planar dimer I2Cl6, is a bright yellow solid, synthesised by reacting iodine with liquid chlorine at −80 °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 ICl+
2 and ICl−
4 ions.[55]

Iodine pentafluoride (IF5), 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 IF+
4 and IF−
6. The pentagonal bipyramidal iodine heptafluoride (IF7) is an extremely powerful fluorinating agent, behind only chlorine trifluoride, chlorine pentafluoride, and bromine pentafluoride among the interhalogens: it reacts with almost all the elements even at low temperatures, fluorinates Pyrex glass to form iodine(VII) oxyfluoride (IOF5), and sets carbon monoxide on fire.[56]

Iodine oxides and oxoacids

 
Structure of iodine pentoxide

Iodine oxides 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.[25] The stable, white, hygroscopic iodine pentoxide (I2O5) has been known since its formation in 1813 by Gay-Lussac and Davy. It is most easily made by the dehydration of iodic acid (HIO3), of which it is the anhydride. It will quickly oxidise carbon monoxide completely to carbon dioxide at room temperature, and is thus a useful reagent in determining carbon monoxide concentration. It also oxidises nitrogen oxide, ethylene, and hydrogen sulfide. It reacts with sulfur trioxide and peroxydisulfuryl difluoride (S2O6F2) to form salts of the iodyl cation, [IO2]+, and is reduced by concentrated sulfuric acids to iodosyl salts involving [IO]+. It may be fluorinated by fluorine, bromine trifluoride, sulfur tetrafluoride, or chloryl fluoride, resulting iodine pentafluoride, which also reacts with iodine pentoxide, giving iodine(V) oxyfluoride, IOF3. A few other less stable oxides are known, notably I4O9 and I2O4; their structures have not been determined, but reasonable guesses are IIII(IVO3)3 and [IO]+[IO3]− respectively.[57]

Standard reduction potentials for aqueous I species[58]
E°(couple) a(H+) = 1
(acid)		 E°(couple) a(OH−) = 1
(base)
I2/I− +0.535 I2/I− +0.535
HOI/I− +0.987 IO−/I− +0.48
    IO−
3/I−		 +0.26
HOI/I2 +1.439 IO−/I2 +0.42
IO−
3/I2		 +1.195    
IO−
3/HOI		 +1.134 IO−
3/IO−		 +0.15
IO−
4/IO−
3		 +1.653    
H5IO6/IO−
3		 +1.601 H
3IO2−
6/IO−
3		 +0.65

More important are the four oxoacids: hypoiodous acid (HIO), iodous acid (HIO2), iodic acid (HIO3), and periodic acid (HIO4 or H5IO6). When iodine dissolves in aqueous solution, the following reactions occur:[58]

I2 + H2O ⇌ HIO + H+ + I− Kac = 2.0 × 10−13 mol2 l−2
I2 + 2 OH− ⇌ IO− + H2O + I− Kalk = 30 mol2 l−2

Hypoiodous acid is unstable to disproportionation. The hypoiodite ions thus formed disproportionate immediately to give iodide and iodate:[58]

3 IO− ⇌ 2 I− + IO−
3		 K = 1020

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.[58] Iodates are by far the most important of these compounds, which can be made by oxidising alkali metal iodides with oxygen at 600 °C and high pressure, or by oxidising iodine with chlorates. Unlike chlorates, which disproportionate very slowly to form chloride and perchlorate, iodates are stable to disproportionation in both acidic and alkaline solutions. From these, salts of most metals can be obtained. Iodic acid is most easily made by oxidation of an aqueous iodine suspension by electrolysis or fuming nitric acid. Iodate has the weakest oxidising power of the halates, but reacts the quickest.[59]

Many periodates are known, including not only the expected tetrahedral IO−
4, but also square-pyramidal IO3−
5, octahedral orthoperiodate IO5−
6, [IO3(OH)3]2−, [I2O8(OH2)]4−, and I
2O4−
9. They are usually made by oxidising alkaline sodium iodate electrochemically (with lead(IV) oxide as the anode) or by chlorine gas:[60]

IO−
3 + 6 OH− → IO5−
6 + 3 H2O + 2 e−
IO−
3 + 6 OH− + Cl2 → IO5−
6 + 2 Cl− + 3 H2O

They are thermodymically and kinetically powerful oxidising agents, quickly oxidising Mn2+ to MnO−
4, and cleaving glycols, α-diketones, α-ketols, α-aminoalcohols, and α-diamines.[60] Orthoperiodate especially stabilises high oxidation states among metals because of its very high negative charge of −5. Orthoperiodic acid, H5IO6, is stable, and dehydrates at 100 °C in a vacuum to metaperiodic acid, HIO4. Attempting to go further does not result in the nonexistent iodine heptoxide (I2O7), but rather iodine pentoxide and oxygen. Periodic acid may be protonated by sulfuric acid to give the I(OH)+
6 cation, isoelectronic to Te(OH)6 and Sb(OH)−
6, and giving salts with bisulfate and sulfate.[25]

Polyiodine compounds

When iodine dissolves in strong acids, such as fuming sulfuric acid, a bright blue paramagnetic solution including I+
2 cations is formed. A solid salt of the diiodine cation may be obtained by oxidising iodine with antimony pentafluoride:[25]

2 I2 + 5 SbF5 SO2⟶20 °C 2 I2Sb2F11 + SbF3

The salt I2Sb2F11 is dark blue, and the blue tantalum analogue I2Ta2F11 is also known. Whereas the I–I bond length in I2 is 267 pm, that in I+
2 is only 256 pm as the missing electron in the latter has been removed from an antibonding orbital, making the bond stronger and hence shorter. In fluorosulfuric acid solution, deep-blue I+
2 reversibly dimerises below −60 °C, forming red rectangular diamagnetic I2+
4. Other polyiodine cations are not as well-characterised, including bent dark-brown or black I+
3 and centrosymmetric C2h green or black I+
5, known in the AsF−
6 and AlCl−
4 salts among others.[25][61]

The only important polyiodide anion in aqueous solution is linear triiodide, I−
3. Its formation explains why the solubility of iodine in water may be increased by the addition of potassium iodide solution:[25]

I2 + I− ⇌ I−
3 (Keq = ~700 at 20 °C)

Many other polyiodides may be found when solutions containing iodine and iodide crystallise, such as I−
5, I−
9, I2−
4, and I2−
8, whose salts with large, weakly polarising cations such as Cs+ may be isolated.[25][62]

Organoiodine compounds

Main article: Organoiodine compound
 
Structure of the oxidising agent 2-iodoxybenzoic acid

Organoiodine compounds have been fundamental in the development of organic synthesis, such as in the Hofmann elimination of amines,[63] the Williamson ether synthesis,[64] the Wurtz coupling reaction,[65] and in Grignard reagents.[66]

The carbon–iodine bond is a common functional group that forms part of core organic chemistry; formally, these compounds may be thought of as organic derivatives of the iodide anion. The simplest organoiodine compounds, alkyl iodides, may be synthesised by the reaction of alcohols with phosphorus triiodide; these may then be used in nucleophilic substitution reactions, or for preparing Grignard reagents. The C–I bond is the weakest of all the carbon–halogen bonds due to the minuscule difference in electronegativity between carbon (2.55) and iodine (2.66). As such, iodide is the best leaving group among the halogens, to such an extent that many organoiodine compounds turn yellow when stored over time due to decomposition into elemental iodine; as such, they are commonly used in organic synthesis, because of the easy formation and cleavage of the C–I bond.[67] They are also significantly denser than the other organohalogen compounds thanks to the high atomic weight of iodine.[68] A few organic oxidising agents like the iodanes contain iodine in a higher oxidation state than −1, such as 2-iodoxybenzoic acid, a common reagent for the oxidation of alcohols to aldehydes,[69] and iodobenzene dichloride (PhICl2), used for the selective chlorination of alkenes and alkynes.[70] One of the more well-known uses of organoiodine compounds is the so-called iodoform test, where iodoform (CHI3) is produced by the exhaustive iodination of a methyl ketone (or another compound capable of being oxidised to a methyl ketone), as follows:[71]

 

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.[72] For example, iodoacetamide and iodoacetic acid denature proteins by irreversibly alkylating cysteine residues and preventing the reformation of disulfide linkages.[73]

Halogen exchange to produce iodoalkanes by the Finkelstein reaction is slightly complicated by the fact that iodide is a better leaving group than chloride or bromide. The difference is nevertheless small enough that the reaction can be driven to completion by exploiting the differential solubility of halide salts, or by using a large excess of the halide salt.[71] In the classic Finkelstein reaction, an alkyl chloride or an alkyl bromide is converted to an alkyl iodide by treatment with a solution of sodium iodide in acetone. Sodium iodide is soluble in acetone and sodium chloride and sodium bromide are not.[74] The reaction is driven toward products by mass action due to the precipitation of the insoluble salt.[75][76]

Occurrence and production

Iodine is the least abundant of the stable halogens, comprising only 0.46 parts per million of Earth's crustal rocks (compare: fluorine 544 ppm, chlorine 126 ppm, bromine 2.5 ppm).[77] Among the 84 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 iodate minerals instead. Examples include lautarite, Ca(IO3)2, and dietzeite, 7Ca(IO3)2·8CaCrO4.[77] These are the minerals that occur as trace impurities in the caliche, found in Chile, whose main product is sodium nitrate. In total, they can contain at least 0.02% and at most 1% iodine by mass.[78] Sodium iodate is extracted from the caliche and reduced to iodide by sodium bisulfite. This solution is then reacted with freshly extracted iodate, resulting in comproportionation to iodine, which may be filtered off.[21]

The caliche was the main source of iodine in the 19th century and continues to be important today, replacing kelp (which is no longer an economically viable source),[79] but in the late 20th century brines emerged as a comparable source. The Japanese Minami Kanto gas field east of Tokyo and the American Anadarko Basin gas field in northwest Oklahoma are the two largest such sources. The brine is hotter than 60 °C from the depth of the source. The brine is first purified and acidified using sulfuric acid, then the iodide present is oxidised to iodine with 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 sulfur dioxide reduces the iodine. The hydrogen iodide (HI) is reacted with chlorine to precipitate the iodine. After filtering and purification the iodine is packed.[78][80]

2 HI + Cl2 → I2↑ + 2 HCl
I2 + 2 H2O + SO2 → 2 HI + H2SO4
2 HI + Cl2 → I2↓ + 2 HCl

These sources ensure that Chile and Japan are the largest producers of iodine today.[77] Alternatively, the brine may be treated with silver nitrate to precipitate out iodine as silver iodide, which is then decomposed by reaction with iron to form metallic silver and a solution of iron(II) iodide. The iodine may then be liberated by displacement with chlorine.[81]

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 potassium iodide, and another 15% for other inorganic iodine compounds.[21] Among the major uses of iodine compounds are catalysts, animal feed supplements, stabilisers, dyes, colourants and pigments, pharmaceutical, sanitation (from tincture of iodine), and photography; minor uses include smog inhibition, cloud seeding, and various uses in analytical chemistry.[21]

Chemical analysis

 
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 iodometry. Iodine and starch form a blue complex, and this reaction is often used to test for either starch or iodine and as an indicator in iodometry. The iodine test for starch is still used to detect counterfeit banknotes printed on starch-containing paper.[82]

The iodine value is the mass of iodine in grams that is consumed by 100 grams of a chemical substance typically fats or oils. Iodine numbers are often used to determine the amount of unsaturation in fatty acids. This unsaturation is in the form of double bonds, which react with iodine compounds.

Potassium tetraiodomercurate(II), K2HgI4, is also known as Nessler's reagent. It is often used as a sensitive spot test for ammonia. Similarly, Mayer's reagent (potassium tetraiodomercurate(II) solution) is used as a precipitating reagent to test for alkaloids.[83] Aqueous alkaline iodine solution is used in the iodoform test for methyl ketones.[71]

Spectroscopy

The spectrum of the iodine molecule, I2, consists of (not exclusively) tens of thousands of sharp spectral lines in the wavelength range 500–700 nm. It is therefore a commonly used wavelength reference (secondary standard). By measuring with a spectroscopic Doppler-free technique while focusing on one of these lines, the hyperfine structure of the iodine molecule reveals itself. A line is now resolved such that either 15 components (from even rotational quantum numbers, Jeven), or 21 components (from odd rotational quantum numbers, Jodd) are measurable.[84]

Caesium iodide and thallium-doped sodium iodide are used in crystal scintillators 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 propellant can be built more compactly, with less mass (and cost), and operate more efficiently than the gridded ion thrusters that were utilized to propel previous spacecraft, such as Japan's Hayabusa probes, ESA's GOCE satellite, or NASA's DART mission, all of which used xenon as the reaction mass. Iodine's atomic weight is only 3.3% less than that of xenon, while its first two ionisation energies average 12% less; together, these make iodine ions a promising substitute.[85][86]

Use of iodine should allow more widespread application of ion-thrust technology, particularly with smaller-scale space vehicles.[86] According to the European Space Agency, "This small but potentially disruptive innovation could help to clear the skies of 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."[87]

In early 2021, the French group ThrustMe performed an in-orbit demonstration of an electric-powered ion thruster for spacecraft, where iodine was used in lieu of xenon as the source of plasma, in order to generate thrust by accelerating ions with an electrostatic field.[85]

Medicine

Main article: Iodine (medical use)

Elemental iodine

Elemental iodine is used as an antiseptic either as the element, or as the water-soluble triiodide anion I3− generated in situ by adding 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 iodine deficiency.[88]

In the alternative, iodine may be produced from iodophors, 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:[89]

The antimicrobial action of iodine is quick and works at low concentrations, and thus it is used in operating theatres.[91] Its specific mode of action is unknown. It penetrates into microorganisms and attacks particular amino acids (such as cysteine and methionine), nucleotides, and fatty acids, ultimately resulting in cell death. It also has an antiviral action, but nonlipid viruses and parvoviruses are less sensitive than lipid enveloped viruses. Iodine probably attacks surface proteins of enveloped viruses, and it may also destabilise membrane fatty acids by reacting with unsaturated carbon bonds.[92]

Other formulations

Before the advent of organic chelating agents, salts of iodide were given orally in the treatment of lead or mercury poisoning, such as heavily popularized by Louis Melsens and many nineteenth and early twentieth century doctors.[93][94]

In medicine, a saturated solution of potassium iodide is used to treat acute thyrotoxicosis. It is also used to block uptake of iodine-131 in the thyroid gland (see isotopes section above), when this isotope is used as part of radiopharmaceuticals (such as iobenguane) that are not targeted to the thyroid or thyroid-type tissues.[95][96]

Iodine-131 (usually as iodide) is a component of 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 mg tablet per 24 hours, supplying 100 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.[97]

 
Diatrizoic acid, an iodine-containing radiocontrast agent

As an element with high electron density and atomic number, iodine absorbs X-rays weaker than 33.3 keV due to the photoelectric effect of the innermost electrons.[98] 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 angiography and CT scanning. At present, all water-soluble radiocontrast agents rely on iodine-containing compounds.

Others

Main article: Iodine value

The production of ethylenediamine dihydroiodide, provided as a nutritional supplement for livestock, consumes a large portion of available iodine. Another significant use is a catalyst for the production of acetic acid by the Monsanto and Cativa processes. In these technologies, which support the world's demand for acetic acid, hydroiodic acid converts the methanol feedstock into methyl iodide, which undergoes carbonylation. Hydrolysis of the resulting acetyl iodide regenerates hydroiodic acid and gives acetic acid.[99]

Inorganic iodides find specialised uses. Titanium, zirconium, hafnium, and thorium are purified by the 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.[99]

The organoiodine compound erythrosine is an important food coloring agent. Perfluoroalkyl iodides are precursors to important surfactants, such as perfluorooctanesulfonic acid.[99]

The iodine clock reaction (in which iodine also serves as a test for starch, forming a dark blue complex),[21] 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 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 agricultural treatment due to its relatively natural chemistry.[100]

125I is used as the radiolabel in investigating which ligands go to which plant pattern recognition receptors (PRRs).[101]

Biological role

Main article: Iodine in biology
 
The thyroid system of the thyroid hormones T3 and T4
 
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))[102]

Iodine is an essential element for life and, at atomic number Z = 53, is the heaviest element commonly needed by living organisms. (Lanthanum and the other lanthanides, as well as tungsten with Z = 74 and uranium with Z = 92, are used by a few microorganisms.[103][104][105]) It is required for the synthesis of the growth-regulating thyroid hormones thyroxine and triiodothyronine (T4 and T3 respectively, named after their number of iodine atoms). A deficiency of iodine leads to decreased production of T3 and T4 and a concomitant enlargement of the 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 (T4), which has a longer half-life than T3. In humans, the ratio of T4 to T3 released into the blood is between 14:1 and 20:1. T4 is converted to the active T3 (three to four times more potent than T4) within cells by deiodinases (5'-iodinase). These are further processed by decarboxylation and deiodination to produce iodothyronamine (T1a) and thyronamine (T0a'). All three isoforms of the deiodinases are selenium-containing enzymes; thus dietary selenium is essential for T3 production.[106]

Iodine accounts for 65% of the molecular weight of T4 and 59% of T3. Fifteen to 20 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.[107][108] In the cells of those tissues, iodide enters directly by 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.[109]

Dietary intake

The daily levels of intake recommended by the United States National Academy of Medicine are between 110 and 130 µg for infants up to 12 months, 90 µg for children up to eight years, 130 µg for children up to 13 years, 150 µg for adults, 220 µg for pregnant women and 290 µg for lactation.[5][110] The Tolerable Upper Intake Level (UL) for adults is 1,100 μg/day.[111] This upper limit was assessed by analyzing the effect of supplementation on thyroid-stimulating hormone.[109]

The thyroid gland needs no more than 70 μg/day to synthesise the requisite daily amounts of T4 and T3.[5] The higher recommended daily allowance levels of iodine seem necessary for optimal function of a number of body systems, including lactation, gastric mucosa, salivary glands, brain cells, choroid plexus, thymus, and arterial walls.[5][112][113][114]

Natural sources of dietary iodine include seafood, such as fish, seaweeds (such as kelp) and shellfish, dairy products and eggs so long as the animals received enough iodine, and plants grown on iodine-rich soil.[115][116] Iodised salt is fortified with iodine in the form of sodium iodide.[116][117]

As of 2000, the median intake of iodine from food in the United States was 240 to 300 μg/day for men and 190 to 210 μg/day for women.[111] The general US population has adequate iodine nutrition,[118][119] with women of childbearing age and pregnant women having a possible mild risk of deficiency.[119] In Japan, consumption was considered much higher, ranging between 5,280 μg/day to 13,800 μg/day from dietary seaweed or kombu kelp,[109] 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 μg/day.[120] The adult UL in Japan was last revised to 3,000 µg/day in 2015.[121]

After iodine fortification programs such as iodisation of salt have been implemented, some cases of iodine-induced hyperthyroidism have been observed (so-called 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.[122]

Deficiency

Main article: Iodine deficiency

In areas where there is little iodine in the diet,[123] typically remote inland areas and semi-arid equatorial climates where no marine foods are eaten, iodine deficiency gives rise to hypothyroidism, symptoms of which are extreme fatigue, goitre, mental slowing, depression, weight gain, and low basal body temperatures.[124] Iodine deficiency is the leading cause of preventable 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.[125] 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.[126]

Precautions

Toxicity

Iodine
Hazards
GHS labelling:
  
Danger
H312, H315, H319, H332, H335, H372, H400
P261, P273, P280, P305, P314, P338, P351[127]
NFPA 704 (fire diamond)
3
0
0

Elemental iodine (I2) is toxic if taken orally undiluted. The lethal dose for an adult human is 30 mg/kg, which is about 2.1–2.4 grams for a human weighing 70 to 80 kg (even if experiments on rats demonstrated that these animals could survive after eating a 14000 mg/kg dose). Excess iodine can be more cytotoxic in the presence of selenium deficiency.[129] Iodine supplementation in selenium-deficient populations is, in theory, problematic, partly for this reason.[109] The toxicity derives from its oxidizing properties, through which it denaturates proteins (including enzymes).[130]

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 tincture of iodine and Lugol's solution, are capable of causing tissue damage if used in prolonged cleaning or antisepsis; similarly, liquid Povidone-iodine (Betadine) trapped against the skin resulted in chemical burns in some reported cases.[131]

Occupational exposure

People can be exposed to iodine in the workplace by inhalation, ingestion, skin contact, and eye contact. The Occupational Safety and Health Administration (OSHA) has set the legal limit (Permissible exposure limit) for iodine exposure in the workplace at 0.1 ppm (1 mg/m3) during an 8-hour workday. The National Institute for Occupational Safety and Health (NIOSH) has set a Recommended exposure limit (REL) of 0.1 ppm (1 mg/m3) during an 8-hour workday. At levels of 2 ppm, iodine is immediately dangerous to life and health.[132]

Allergic reactions

Some people develop a hypersensitivity to products and foods containing iodine. Applications of tincture of iodine or Betadine can cause rashes, sometimes severe.[133] Parenteral use of iodine-based contrast agents (see above) can cause reactions ranging from a mild rash to fatal 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.[134] 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;[135] 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.[136][135] As with all medications, the patient's allergy history should be questioned and consulted before any containing iodine are administered.[137]

US DEA List I status

Phosphorus can reduce elemental iodine to hydroiodic acid, which is a reagent effective for reducing ephedrine or pseudoephedrine to methamphetamine.[138] For this reason, iodine was designated by the United States Drug Enforcement Administration as a List I precursor chemical under 21 CFR 1310.02.[139]

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Bibliography

April 29, 2023
iodine, this, article, about, chemical, element, other, uses, disambiguation, chemical, element, with, symbol, atomic, number, heaviest, stable, halogens, exists, semi, lustrous, metallic, solid, standard, conditions, that, melts, form, deep, violet, liquid, b. This article is about the chemical element For other uses see Iodine disambiguation Iodine is a chemical element with the symbol I and atomic number 53 The heaviest of the stable halogens it exists as a semi lustrous non metallic solid at standard conditions that melts to form a deep violet liquid at 114 C 237 F and boils to a violet gas at 184 C 363 F The element was discovered by the French chemist Bernard Courtois in 1811 and was named two years later by Joseph Louis Gay Lussac after the Ancient Greek Iwdhs violet coloured Iodine 53IIodinePronunciation ˈ aɪ e d aɪ n d ɪ n d iː n wbr EYE e dyne din deen Appearancelustrous metallic gray solid black violet liquid violet gasStandard atomic weight Ar I 126 90447 0 00003126 90 0 01 abridged 1 Iodine in the periodic tableHydrogen HeliumLithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine NeonSodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine ArgonPotassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine KryptonRubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine XenonCaesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury element Thallium Lead Bismuth Polonium Astatine RadonFrancium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson Br I Attellurium iodine xenonAtomic number Z 53Groupgroup 17 halogens Periodperiod 5Block p blockElectron configuration Kr 4d10 5s2 5p5Electrons per shell2 8 18 18 7Physical propertiesPhase at STPsolidMelting point I2 386 85 K 113 7 C 236 66 F Boiling point I2 457 4 K 184 3 C 363 7 F Density near r t 4 933 g cm3Triple point386 65 K 12 1 kPaCritical point819 K 11 7 MPaHeat of fusion I2 15 52 kJ molHeat of vaporisation I2 41 57 kJ molMolar heat capacity I2 54 44 J mol K Vapour pressure rhombic P Pa 1 10 100 1 k 10 k 100 kat T K 260 282 309 342 381 457Atomic propertiesOxidation states 1 1 2 2 3 4 5 6 7 a strongly acidic oxide ElectronegativityPauling scale 2 66Ionisation energies1st 1008 4 kJ mol2nd 1845 9 kJ mol3rd 3180 kJ molAtomic radiusempirical 140 pmCovalent radius139 3 pmVan der Waals radius198 pmSpectral lines of iodineOther propertiesNatural occurrenceprimordialCrystal structure base centered orthorhombicThermal conductivity0 449 W m K Electrical resistivity1 3 107 W m at 0 C Magnetic orderingdiamagnetic 3 Molar magnetic susceptibility 88 7 10 6 cm3 mol 298 K 4 Bulk modulus7 7 GPaCAS Number7553 56 2HistoryDiscovery and first isolationBernard Courtois 1811 Isotopes of iodineveMain isotopes Decayabun dance half life t1 2 mode pro duct123I synth 13 h b 100 123Te124I synth 4 176 d e 124Te125I synth 59 40 d e 125Te127I 100 stable129I trace 1 57 107 y b 129Xe131I synth 8 02070 d b 100 131Xe135I synth 6 57 h b 135Xe Category Iodineviewtalkedit referencesIodine occurs in many oxidation states including iodide I iodate IO 3 and the various periodate anions It is the least abundant of the stable halogens being the sixty first most abundant element As the heaviest essential mineral nutrient iodine is required for the synthesis of thyroid hormones 5 Iodine deficiency affects about two billion people and is the leading preventable cause of intellectual disabilities 6 The dominant producers of iodine today are Chile and Japan Due to its high atomic number and ease of attachment to organic compounds 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 thyroid cancer Iodine is also used as a catalyst in the industrial production of acetic acid and some polymers It is on the World Health Organization s List of Essential Medicines 7 Contents 1 History 2 Properties 2 1 Isotopes 3 Chemistry and compounds 3 1 Charge transfer complexes 3 2 Hydrogen iodide 3 3 Other binary iodine compounds 3 4 Iodine halides 3 5 Iodine oxides and oxoacids 3 6 Polyiodine compounds 3 7 Organoiodine compounds 4 Occurrence and production 5 Applications 5 1 Chemical analysis 5 2 Spectroscopy 5 3 Spacecraft propulsion 5 4 Medicine 5 4 1 Elemental iodine 5 4 2 Other formulations 5 5 Others 6 Biological role 6 1 Dietary intake 6 2 Deficiency 7 Precautions 7 1 Toxicity 7 1 1 Occupational exposure 7 1 2 Allergic reactions 7 2 US DEA List I status 8 References 9 BibliographyHistory EditIn 1811 iodine was discovered by French chemist Bernard Courtois 8 9 who was born to a manufacturer of saltpetre an essential component of gunpowder At the time of the Napoleonic Wars saltpetre was in great demand in France Saltpetre produced from French nitre beds required sodium carbonate which could be isolated from seaweed collected on the coasts of Normandy and Brittany To isolate the sodium carbonate seaweed was burned and the ash washed with water The remaining waste was destroyed by adding 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 10 Courtois suspected that this material was a new element but lacked funding to pursue it further 11 Courtois gave samples to his friends Charles Bernard Desormes 1777 1838 and Nicolas Clement 1779 1841 to continue research He also gave some of the substance to chemist Joseph Louis Gay Lussac 1778 1850 and to physicist Andre Marie Ampere 1775 1836 On 29 November 1813 Desormes and Clement made Courtois discovery public They described the substance to a meeting of the Imperial Institute of France 12 On 6 December Gay Lussac announced that the new substance was either an element or a compound of oxygen 13 14 15 Gay Lussac suggested the name iode from the Ancient Greek ἰoeidhs ioeides violet because of the colour of iodine vapor 8 13 Ampere had given some of his sample to English chemist Humphry Davy 1778 1829 who experimented on the substance and noted its similarity to chlorine 16 Davy sent a letter dated 10 December to the Royal Society of London stating that he had identified a new element 17 Arguments erupted between Davy and Gay Lussac over who identified iodine first but both scientists acknowledged Courtois as the first to isolate the element 11 In 1873 the French medical researcher Casimir Joseph Davaine 1812 1882 discovered the antiseptic action of iodine 18 Antonio Grossich 1849 1926 an Istrian born surgeon was among the first to use 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 19 In early periodic tables iodine was often given the symbol J for Jod its name in German 20 Properties Edit Iodine vapour in a flask Iodine is the fourth halogen being a member of group 17 in the periodic table below fluorine chlorine and bromine it is the heaviest stable member of its group The fifth and sixth halogens the radioactive astatine and 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 4d105s25p5 with the seven electrons in the fifth and outermost shell being its valence electrons 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 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 diatomic molecules with chemical formula I2 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 is the strongest reducing agent among the stable halogens being the most easily oxidised back to diatomic I2 21 Astatine goes further being indeed unstable as At and readily oxidised to At0 or At 22 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 ml at 20 C and 1280 ml at 50 C potassium iodide may be added to increase solubility via formation of triiodide ions among other polyiodides 23 Nonpolar solvents such as hexane and carbon tetrachloride provide a higher solubility 24 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 23 Charge transfer complexes 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 alcohols and amines solvents that form charge transfer adducts 25 I2 PPh3 charge transfer complexes in CH2Cl2 From left to right 1 I2 dissolved in dichloromethane no CT complex 2 A few seconds after excess PPh3 was added CT complex is forming 3 One minute later after excess PPh3 was added the CT complex Ph3PI I has been formed 4 Immediately after excess I2 was added which contains Ph3PI I3 26 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 21 Due to this property Iodine is commonly used to demonstrate sublimation directly from solid to gas which gives rise to a misconception that it does not melt in atmospheric pressure 27 Because it has the largest atomic radius among the halogens iodine has the lowest first ionisation energy lowest electron affinity lowest electronegativity and lowest reactivity of the halogens 21 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 C Temperatures greater than 750 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 21 Gaseous iodine is composed of I2 molecules with an I I bond length of 266 6 pm The I I bond is one of the longest single bonds known It is even longer 271 5 pm in solid orthorhombic crystalline iodine which has the same crystal structure as chlorine and bromine The record is held by iodine s neighbour xenon the Xe Xe bond length is 308 71 pm 28 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 semiconducting properties 21 Iodine is a two dimensional semiconductor with a band gap of 1 3 eV 125 kJ mol it is a semiconductor in the plane of its crystalline layers and an insulator in the perpendicular direction 21 Isotopes Edit Main article Isotopes of iodine Of the thirty seven known isotopes of iodine only one occurs in nature iodine 127 The others are radioactive and have half lives too short to be primordial As such iodine is both monoisotopic and mononuclidic and its atomic weight is known to great precision as it is a constant of nature 21 The longest lived of the radioactive isotopes of iodine is iodine 129 which has a half life of 15 7 million years decaying via beta decay to stable xenon 129 29 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 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 daughter xenon 129 30 31 32 33 34 Traces of iodine 129 still exist today as it is also a cosmogenic nuclide formed from cosmic ray spallation of atmospheric xenon these traces make up 10 14 to 10 10 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 7 of all terrestrial iodine 35 Excited states of iodine 127 and iodine 129 are often used in Mossbauer spectroscopy 21 The other iodine radioisotopes have much shorter half lives no longer than days 29 Some of them have medical applications involving the thyroid gland where the iodine that enters the body is stored and concentrated Iodine 123 has a half life of thirteen hours and decays by electron capture to tellurium 123 emitting gamma radiation it is used in nuclear medicine imaging including single photon emission computed tomography SPECT and X ray computed tomography X Ray CT scans 36 Iodine 125 has a half life of fifty nine days decaying by electron capture to tellurium 125 and emitting low energy gamma radiation the second longest lived iodine radioisotope it has uses in biological assays nuclear medicine imaging and in radiation therapy as brachytherapy to treat a number of conditions including prostate cancer uveal melanomas and brain tumours 37 Finally iodine 131 with a half life of eight days beta decays to an excited state of stable 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 thyroid cancer in later life Other risks include the possibility of non cancerous growths and thyroiditis 38 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 potassium iodide tablets taken daily for optimal prophylaxis 39 However iodine 131 may also be used for medicinal purposes in radiation therapy for this very reason when tissue destruction is desired after iodine uptake by the tissue 40 Iodine 131 is also used as a radioactive tracer 41 42 43 44 Chemistry and compounds EditHalogen bond energies kJ mol 23 X XX HX BX3 AlX3 CX4F 159 574 645 582 456Cl 243 428 444 427 327Br 193 363 368 360 272I 151 294 272 285 239Iodine is quite reactive but it is much less reactive than the other halogens For example while chlorine gas will halogenate carbon monoxide nitric oxide and sulfur dioxide to phosgene nitrosyl chloride and 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 rhenium metal reacts with chlorine to form rhenium hexachloride but with bromine it forms only rhenium pentabromide and iodine can achieve only rhenium tetraiodide 21 By the same token however since iodine has the lowest ionisation energy among the halogens and is the most easily oxidised of them it has a more significant cationic chemistry and its higher oxidation states are rather more stable than those of bromine and chlorine for example in iodine heptafluoride 23 Charge transfer complexes Edit The iodine molecule I2 dissolves in CCl4 and aliphatic hydrocarbons to give bright violet solutions In these solvents the absorption band maximum occurs in the 520 540 nm region and is assigned to a p to s transition When I2 reacts with Lewis bases in these solvents a blue shift in I2 peak is seen and the new peak 230 330 nm arises that is due to the formation of adducts which are referred to as charge transfer complexes 45 Hydrogen iodide Edit The simplest compound of iodine is hydrogen iodide HI It is a colourless gas that reacts with oxygen to give water and iodine Although it is useful in iodination reactions in the laboratory it does not have large scale industrial uses unlike the other hydrogen halides Commercially it is usually made by reacting iodine with hydrogen sulfide or hydrazine 46 2 I2 N2H4 H2O 4 HI N2At room temperature it is a colourless gas like all of the hydrogen halides except hydrogen fluoride since hydrogen cannot form strong hydrogen bonds to the large and only mildly electronegative iodine atom It melts at 51 0 C and boils at 35 1 C It is an endothermic compound that can exothermically dissociate at room temperature although the process is very slow unless a 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 kJ mol 47 Aqueous hydrogen iodide is known as hydroiodic acid which is a strong acid Hydrogen iodide is exceptionally soluble in water one litre of water will dissolve 425 litres of hydrogen iodide and the saturated solution has only four water molecules per molecule of hydrogen iodide 48 Commercial so called concentrated hydroiodic acid usually contains 48 57 HI by mass the solution forms an azeotrope with boiling point 126 7 C at 56 7 g HI per 100 g solution Hence hydroiodic acid cannot be concentrated past this point by evaporation of water 47 Unlike 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 dielectric constant is low and it does not dissociate appreciably into H2I and HI 2 ions the latter in any case are much less stable than the bifluoride ions 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 Cs and NR 4 R Me Et Bun may still be isolated Anhydrous hydrogen iodide is a poor solvent able to dissolve only small molecular compounds such as nitrosyl chloride and phenol or salts with very low lattice energies such as tetraalkylammonium halides 47 Other binary iodine compounds Edit With the exception of the noble gases nearly all elements on the periodic table up to einsteinium EsI3 is known are known to form binary compounds with iodine Until 1990 nitrogen triiodide 49 was only known as an ammonia adduct Ammonia free NI3 was found to be isolable at 196 C but spontaneously decomposes at 0 C 50 For thermodynamic reasons related to electronegativity of the elements neutral sulfur and selenium iodides that are stable at room temperature are also nonexistent although S2I2 and SI2 are stable up to 183 and 9 K respectively As of 2022 no neutral binary selenium iodide has been unambiguously identified at any temperature 51 Sulfur and selenium iodine polyatomic cations e g S2I42 AsF6 2 and Se2I42 Sb2F11 2 have been prepared and characterized crystallographically 52 Given the large size of the iodide anion and iodine s weak oxidising power high oxidation states are difficult to achieve in binary iodides the maximum known being in the pentaiodides of niobium tantalum and protactinium Iodides can be made by reaction of an element or its oxide hydroxide or carbonate with hydroiodic acid and then dehydrated by mildly high temperatures combined with either low pressure or anhydrous hydrogen iodide gas These methods work best when the iodide product is stable to hydrolysis otherwise the possibilities include high temperature oxidative iodination of the element with iodine or hydrogen iodide high temperature iodination of a metal oxide or other halide by iodine a volatile metal halide carbon tetraiodide or an organic iodide For example molybdenum IV oxide reacts with aluminium III iodide at 230 C to give molybdenum II iodide An example involving halogen exchange is given below involving the reaction of tantalum V chloride with excess aluminium III iodide at 400 C to give tantalum V iodide 53 3 TaCl 5 5 AlI 3 excess 3 TaI 5 5 AlCl 3 displaystyle ce 3TaCl5 underset excess 5AlI3 gt 3TaI5 5AlCl3 Lower iodides may be produced either through thermal decomposition or disproportionation or by reducing the higher iodide with hydrogen or a metal for example 53 TaI 5 Ta 630 C 575 C thermal gradient Ta 6 I 14 displaystyle ce TaI5 Ta gt text thermal gradient ce 630 circ C gt 575 circ C Ta6I14 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 aluminium iodide is not Ionic iodides MIn tend to have the lowest melting and boiling points among the halides MXn of the same element because the electrostatic forces of attraction between the cations and anions are weakest for the large iodide anion In contrast covalent iodides tend to instead have the highest melting and boiling points among the halides of the same element since iodine is the most polarisable of the halogens and having the most electrons among them can contribute the most to van der Waals forces Naturally exceptions abound in intermediate iodides where one trend gives way to the other Similarly solubilities in water of predominantly ionic iodides e g potassium and calcium are the greatest among ionic halides of that element while those of covalent iodides e g silver are the lowest of that element In particular silver iodide is very insoluble in water and its formation is often used as a qualitative test for iodine 53 Iodine halides Edit The halogens form many binary diamagnetic interhalogen compounds with stoichiometries XY XY3 XY5 and XY7 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 ICl 2 and the dark brown or purplish black compounds of I2Cl Apart from these some pseudohalides are also known such as cyanogen iodide ICN iodine thiocyanate ISCN and iodine azide IN3 54 Iodine monochloride Iodine monofluoride IF is unstable at room temperature and disproportionates very readily and irreversibly to iodine and iodine pentafluoride and thus cannot be obtained pure It can be synthesised from the reaction of iodine with fluorine gas in trichlorofluoromethane at 45 C with iodine trifluoride in trichlorofluoromethane at 78 C or with silver I fluoride at 0 C 54 Iodine monochloride ICl and iodine monobromide IBr on the other hand are moderately stable The former a volatile red brown compound was discovered independently by Joseph Louis Gay Lussac and Humphry Davy in 1813 1814 not long after the discoveries of chlorine and iodine and it mimics the intermediate halogen bromine so well that Justus von Liebig was misled into mistaking bromine which he had found for iodine monochloride Iodine monochloride and iodine monobromide may be prepared simply by reacting iodine with chlorine or bromine at room temperature and purified by fractional crystallisation Both are quite reactive and attack even platinum and gold though not boron carbon cadmium lead zirconium niobium molybdenum and tungsten Their reaction with organic compounds depends on conditions Iodine chloride vapour tends to chlorinate phenol and salicyclic acid since when iodine chloride undergoes homolytic dissociation chlorine and iodine are produced and the former is more reactive However iodine chloride in tetrachloromethane solution results in iodination being the main reaction since now heterolytic fission of the I Cl bond occurs and I attacks phenol as an electrophile However iodine monobromide tends to brominate phenol even in tetrachloromethane solution because it tends to dissociate into its elements in solution and bromine is more reactive than iodine 54 When liquid iodine monochloride and iodine monobromide dissociate into I2 X and IX 2 anions X Cl Br thus they are significant conductors of electricity and can be used as ionising solvents 54 Iodine trifluoride IF3 is an unstable yellow solid that decomposes above 28 C It is thus little known It is difficult to produce because fluorine gas would tend to oxidise iodine all the way to the pentafluoride reaction at low temperature with xenon difluoride is necessary Iodine trichloride which exists in the solid state as the planar dimer I2Cl6 is a bright yellow solid synthesised by reacting iodine with liquid chlorine at 80 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 ICl 2 and ICl 4 ions 55 Iodine pentafluoride IF5 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 IF 4 and IF 6 The pentagonal bipyramidal iodine heptafluoride IF7 is an extremely powerful fluorinating agent behind only chlorine trifluoride chlorine pentafluoride and bromine pentafluoride among the interhalogens it reacts with almost all the elements even at low temperatures fluorinates Pyrex glass to form iodine VII oxyfluoride IOF5 and sets carbon monoxide on fire 56 Iodine oxides and oxoacids Edit Structure of iodine pentoxide Iodine oxides 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 25 The stable white hygroscopic iodine pentoxide I2O5 has been known since its formation in 1813 by Gay Lussac and Davy It is most easily made by the dehydration of iodic acid HIO3 of which it is the anhydride It will quickly oxidise carbon monoxide completely to carbon dioxide at room temperature and is thus a useful reagent in determining carbon monoxide concentration It also oxidises nitrogen oxide ethylene and hydrogen sulfide It reacts with sulfur trioxide and peroxydisulfuryl difluoride S2O6F2 to form salts of the iodyl cation IO2 and is reduced by concentrated sulfuric acids to iodosyl salts involving IO It may be fluorinated by fluorine bromine trifluoride sulfur tetrafluoride or chloryl fluoride resulting iodine pentafluoride which also reacts with iodine pentoxide giving iodine V oxyfluoride IOF3 A few other less stable oxides are known notably I4O9 and I2O4 their structures have not been determined but reasonable guesses are IIII IVO3 3 and IO IO3 respectively 57 Standard reduction potentials for aqueous I species 58 E couple a H 1 acid E couple a OH 1 base I2 I 0 535 I2 I 0 535HOI I 0 987 IO I 0 48 IO 3 I 0 26HOI I2 1 439 IO I2 0 42IO 3 I2 1 195 IO 3 HOI 1 134 IO 3 IO 0 15IO 4 IO 3 1 653 H5IO6 IO 3 1 601 H3 IO2 6 IO 3 0 65More important are the four oxoacids hypoiodous acid HIO iodous acid HIO2 iodic acid HIO3 and periodic acid HIO4 or H5IO6 When iodine dissolves in aqueous solution the following reactions occur 58 I2 H2O HIO H I Kac 2 0 10 13 mol2 l 2I2 2 OH IO H2O I Kalk 30 mol2 l 2Hypoiodous acid is unstable to disproportionation The hypoiodite ions thus formed disproportionate immediately to give iodide and iodate 58 3 IO 2 I IO 3 K 1020Iodous acid and iodite are even less stable and exist only as a fleeting intermediate in the oxidation of iodide to iodate if at all 58 Iodates are by far the most important of these compounds which can be made by oxidising alkali metal iodides with oxygen at 600 C and high pressure or by oxidising iodine with chlorates Unlike chlorates which disproportionate very slowly to form chloride and perchlorate iodates are stable to disproportionation in both acidic and alkaline solutions From these salts of most metals can be obtained Iodic acid is most easily made by oxidation of an aqueous iodine suspension by electrolysis or fuming nitric acid Iodate has the weakest oxidising power of the halates but reacts the quickest 59 Many periodates are known including not only the expected tetrahedral IO 4 but also square pyramidal IO3 5 octahedral orthoperiodate IO5 6 IO3 OH 3 2 I2O8 OH2 4 and I2 O4 9 They are usually made by oxidising alkaline sodium iodate electrochemically with lead IV oxide as the anode or by chlorine gas 60 IO 3 6 OH IO5 6 3 H2O 2 e IO 3 6 OH Cl2 IO5 6 2 Cl 3 H2OThey are thermodymically and kinetically powerful oxidising agents quickly oxidising Mn2 to MnO 4 and cleaving glycols a diketones a ketols a aminoalcohols and a diamines 60 Orthoperiodate especially stabilises high oxidation states among metals because of its very high negative charge of 5 Orthoperiodic acid H5IO6 is stable and dehydrates at 100 C in a vacuum to metaperiodic acid HIO4 Attempting to go further does not result in the nonexistent iodine heptoxide I2O7 but rather iodine pentoxide and oxygen Periodic acid may be protonated by sulfuric acid to give the I OH 6 cation isoelectronic to Te OH 6 and Sb OH 6 and giving salts with bisulfate and sulfate 25 Polyiodine compounds Edit When iodine dissolves in strong acids such as fuming sulfuric acid a bright blue paramagnetic solution including I 2 cations is formed A solid salt of the diiodine cation may be obtained by oxidising iodine with antimony pentafluoride 25 2 I2 5 SbF5 SO2 20 C 2 I2Sb2F11 SbF3The salt I2Sb2F11 is dark blue and the blue tantalum analogue I2Ta2F11 is also known Whereas the I I bond length in I2 is 267 pm that in I 2 is only 256 pm as the missing electron in the latter has been removed from an antibonding orbital making the bond stronger and hence shorter In fluorosulfuric acid solution deep blue I 2 reversibly dimerises below 60 C forming red rectangular diamagnetic I2 4 Other polyiodine cations are not as well characterised including bent dark brown or black I 3 and centrosymmetric C2h green or black I 5 known in the AsF 6 and AlCl 4 salts among others 25 61 The only important polyiodide anion in aqueous solution is linear triiodide I 3 Its formation explains why the solubility of iodine in water may be increased by the addition of potassium iodide solution 25 I2 I I 3 Keq 700 at 20 C Many other polyiodides may be found when solutions containing iodine and iodide crystallise such as I 5 I 9 I2 4 and I2 8 whose salts with large weakly polarising cations such as Cs may be isolated 25 62 Organoiodine compounds Edit Main article Organoiodine compound Structure of the oxidising agent 2 iodoxybenzoic acid Organoiodine compounds have been fundamental in the development of organic synthesis such as in the Hofmann elimination of amines 63 the Williamson ether synthesis 64 the Wurtz coupling reaction 65 and in Grignard reagents 66 The carbon iodine bond is a common functional group that forms part of core organic chemistry formally these compounds may be thought of as organic derivatives of the iodide anion The simplest organoiodine compounds alkyl iodides may be synthesised by the reaction of alcohols with phosphorus triiodide these may then be used in nucleophilic substitution reactions or for preparing Grignard reagents The C I bond is the weakest of all the carbon halogen bonds due to the minuscule difference in electronegativity between carbon 2 55 and iodine 2 66 As such iodide is the best leaving group among the halogens to such an extent that many organoiodine compounds turn yellow when stored over time due to decomposition into elemental iodine as such they are commonly used in organic synthesis because of the easy formation and cleavage of the C I bond 67 They are also significantly denser than the other organohalogen compounds thanks to the high atomic weight of iodine 68 A few organic oxidising agents like the iodanes contain iodine in a higher oxidation state than 1 such as 2 iodoxybenzoic acid a common reagent for the oxidation of alcohols to aldehydes 69 and iodobenzene dichloride PhICl2 used for the selective chlorination of alkenes and alkynes 70 One of the more well known uses of organoiodine compounds is the so called iodoform test where iodoform CHI3 is produced by the exhaustive iodination of a methyl ketone or another compound capable of being oxidised to a methyl ketone as follows 71 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 72 For example iodoacetamide and iodoacetic acid denature proteins by irreversibly alkylating cysteine residues and preventing the reformation of disulfide linkages 73 Halogen exchange to produce iodoalkanes by the Finkelstein reaction is slightly complicated by the fact that iodide is a better leaving group than chloride or bromide The difference is nevertheless small enough that the reaction can be driven to completion by exploiting the differential solubility of halide salts or by using a large excess of the halide salt 71 In the classic Finkelstein reaction an alkyl chloride or an alkyl bromide is converted to an alkyl iodide by treatment with a solution of sodium iodide in acetone Sodium iodide is soluble in acetone and sodium chloride and sodium bromide are not 74 The reaction is driven toward products by mass action due to the precipitation of the insoluble salt 75 76 Occurrence and production EditIodine is the least abundant of the stable halogens comprising only 0 46 parts per million of Earth s crustal rocks compare fluorine 544 ppm chlorine 126 ppm bromine 2 5 ppm 77 Among the 84 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 iodate minerals instead Examples include lautarite Ca IO3 2 and dietzeite 7Ca IO3 2 8CaCrO4 77 These are the minerals that occur as trace impurities in the caliche found in Chile whose main product is sodium nitrate In total they can contain at least 0 02 and at most 1 iodine by mass 78 Sodium iodate is extracted from the caliche and reduced to iodide by sodium bisulfite This solution is then reacted with freshly extracted iodate resulting in comproportionation to iodine which may be filtered off 21 The caliche was the main source of iodine in the 19th century and continues to be important today replacing kelp which is no longer an economically viable source 79 but in the late 20th century brines emerged as a comparable source The Japanese Minami Kanto gas field east of Tokyo and the American Anadarko Basin gas field in northwest Oklahoma are the two largest such sources The brine is hotter than 60 C from the depth of the source The brine is first purified and acidified using sulfuric acid then the iodide present is oxidised to iodine with 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 sulfur dioxide reduces the iodine The hydrogen iodide HI is reacted with chlorine to precipitate the iodine After filtering and purification the iodine is packed 78 80 2 HI Cl2 I2 2 HCl I2 2 H2O SO2 2 HI H2SO4 2 HI Cl2 I2 2 HClThese sources ensure that Chile and Japan are the largest producers of iodine today 77 Alternatively the brine may be treated with silver nitrate to precipitate out iodine as silver iodide which is then decomposed by reaction with iron to form metallic silver and a solution of iron II iodide The iodine may then be liberated by displacement with chlorine 81 Applications EditAbout half of all produced iodine goes into various organoiodine compounds another 15 remains as the pure element another 15 is used to form potassium iodide and another 15 for other inorganic iodine compounds 21 Among the major uses of iodine compounds are catalysts animal feed supplements stabilisers dyes colourants and pigments pharmaceutical sanitation from tincture of iodine and photography minor uses include smog inhibition cloud seeding and various uses in analytical chemistry 21 Chemical analysis Edit 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 iodometry Iodine and starch form a blue complex and this reaction is often used to test for either starch or iodine and as an indicator in iodometry The iodine test for starch is still used to detect counterfeit banknotes printed on starch containing paper 82 The iodine value is the mass of iodine in grams that is consumed by 100 grams of a chemical substance typically fats or oils Iodine numbers are often used to determine the amount of unsaturation in fatty acids This unsaturation is in the form of double bonds which react with iodine compounds Potassium tetraiodomercurate II K2HgI4 is also known as Nessler s reagent It is often used as a sensitive spot test for ammonia Similarly Mayer s reagent potassium tetraiodomercurate II solution is used as a precipitating reagent to test for alkaloids 83 Aqueous alkaline iodine solution is used in the iodoform test for methyl ketones 71 Spectroscopy Edit The spectrum of the iodine molecule I2 consists of not exclusively tens of thousands of sharp spectral lines in the wavelength range 500 700 nm It is therefore a commonly used wavelength reference secondary standard By measuring with a spectroscopic Doppler free technique while focusing on one of these lines the hyperfine structure of the iodine molecule reveals itself A line is now resolved such that either 15 components from even rotational quantum numbers Jeven or 21 components from odd rotational quantum numbers Jodd are measurable 84 Caesium iodide and thallium doped sodium iodide are used in crystal scintillators for the detection of gamma rays The efficiency is high and energy dispersive spectroscopy is possible but the resolution is rather poor Spacecraft propulsion Edit Propulsion systems employing iodine as the propellant can be built more compactly with less mass and cost and operate more efficiently than the gridded ion thrusters that were utilized to propel previous spacecraft such as Japan s Hayabusa probes ESA s GOCE satellite or NASA s DART mission all of which used xenon as the reaction mass Iodine s atomic weight is only 3 3 less than that of xenon while its first two ionisation energies average 12 less together these make iodine ions a promising substitute 85 86 Use of iodine should allow more widespread application of ion thrust technology particularly with smaller scale space vehicles 86 According to the European Space Agency This small but potentially disruptive innovation could help to clear the skies of 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 87 In early 2021 the French group ThrustMe performed an in orbit demonstration of an electric powered ion thruster for spacecraft where iodine was used in lieu of xenon as the source of plasma in order to generate thrust by accelerating ions with an electrostatic field 85 Medicine Edit Main article Iodine medical use Elemental iodine Edit Elemental iodine is used as an antiseptic either as the element or as the water soluble triiodide anion I3 generated in situ by adding 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 iodine deficiency 88 In the alternative iodine may be produced from iodophors 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 89 Tincture of iodine iodine in ethanol or iodine and sodium iodide in a mixture of ethanol and water 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 I2 component Povidone iodine an iodophor Iodine V iodine I2 and fulvic acid form a 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 I2 form 90 The antimicrobial action of iodine is quick and works at low concentrations and thus it is used in operating theatres 91 Its specific mode of action is unknown It penetrates into microorganisms and attacks particular amino acids such as cysteine and methionine nucleotides and fatty acids ultimately resulting in cell death It also has an antiviral action but nonlipid viruses and parvoviruses are less sensitive than lipid enveloped viruses Iodine probably attacks surface proteins of enveloped viruses and it may also destabilise membrane fatty acids by reacting with unsaturated carbon bonds 92 Other formulations Edit Before the advent of organic chelating agents salts of iodide were given orally in the treatment of lead or mercury poisoning such as heavily popularized by Louis Melsens and many nineteenth and early twentieth century doctors 93 94 In medicine a saturated solution of potassium iodide is used to treat acute thyrotoxicosis It is also used to block uptake of iodine 131 in the thyroid gland see isotopes section above when this isotope is used as part of radiopharmaceuticals such as iobenguane that are not targeted to the thyroid or thyroid type tissues 95 96 Iodine 131 usually as iodide is a component of 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 mg tablet per 24 hours supplying 100 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 97 Diatrizoic acid an iodine containing radiocontrast agent As an element with high electron density and atomic number iodine absorbs X rays weaker than 33 3 keV due to the photoelectric effect of the innermost electrons 98 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 angiography and CT scanning At present all water soluble radiocontrast agents rely on iodine containing compounds Others Edit Main article Iodine value The production of ethylenediamine dihydroiodide provided as a nutritional supplement for livestock consumes a large portion of available iodine Another significant use is a catalyst for the production of acetic acid by the Monsanto and Cativa processes In these technologies which support the world s demand for acetic acid hydroiodic acid converts the methanol feedstock into methyl iodide which undergoes carbonylation Hydrolysis of the resulting acetyl iodide regenerates hydroiodic acid and gives acetic acid 99 Inorganic iodides find specialised uses Titanium zirconium hafnium and thorium are purified by the 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 99 The organoiodine compound erythrosine is an important food coloring agent Perfluoroalkyl iodides are precursors to important surfactants such as perfluorooctanesulfonic acid 99 The iodine clock reaction in which iodine also serves as a test for starch forming a dark blue complex 21 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 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 agricultural treatment due to its relatively natural chemistry 100 125I is used as the radiolabel in investigating which ligands go to which plant pattern recognition receptors PRRs 101 Biological role EditMain article Iodine in biology The thyroid system of the thyroid hormones T3 and T4 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 102 Iodine is an essential element for life and at atomic number Z 53 is the heaviest element commonly needed by living organisms Lanthanum and the other lanthanides as well as tungsten with Z 74 and uranium with Z 92 are used by a few microorganisms 103 104 105 It is required for the synthesis of the growth regulating thyroid hormones thyroxine and triiodothyronine T4 and T3 respectively named after their number of iodine atoms A deficiency of iodine leads to decreased production of T3 and T4 and a concomitant enlargement of the 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 T4 which has a longer half life than T3 In humans the ratio of T4 to T3 released into the blood is between 14 1 and 20 1 T4 is converted to the active T3 three to four times more potent than T4 within cells by deiodinases 5 iodinase These are further processed by decarboxylation and deiodination to produce iodothyronamine T1a and thyronamine T0a All three isoforms of the deiodinases are selenium containing enzymes thus dietary selenium is essential for T3 production 106 Iodine accounts for 65 of the molecular weight of T4 and 59 of T3 Fifteen to 20 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 107 108 In the cells of those tissues iodide enters directly by 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 109 Dietary intake Edit The daily levels of intake recommended by the United States National Academy of Medicine are between 110 and 130 µg for infants up to 12 months 90 µg for children up to eight years 130 µg for children up to 13 years 150 µg for adults 220 µg for pregnant women and 290 µg for lactation 5 110 The Tolerable Upper Intake Level UL for adults is 1 100 mg day 111 This upper limit was assessed by analyzing the effect of supplementation on thyroid stimulating hormone 109 The thyroid gland needs no more than 70 mg day to synthesise the requisite daily amounts of T4 and T3 5 The higher recommended daily allowance levels of iodine seem necessary for optimal function of a number of body systems including lactation gastric mucosa salivary glands brain cells choroid plexus thymus and arterial walls 5 112 113 114 Natural sources of dietary iodine include seafood such as fish seaweeds such as kelp and shellfish dairy products and eggs so long as the animals received enough iodine and plants grown on iodine rich soil 115 116 Iodised salt is fortified with iodine in the form of sodium iodide 116 117 As of 2000 the median intake of iodine from food in the United States was 240 to 300 mg day for men and 190 to 210 mg day for women 111 The general US population has adequate iodine nutrition 118 119 with women of childbearing age and pregnant women having a possible mild risk of deficiency 119 In Japan consumption was considered much higher ranging between 5 280 mg day to 13 800 mg day from dietary seaweed or kombu kelp 109 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 mg day 120 The adult UL in Japan was last revised to 3 000 µg day in 2015 121 After iodine fortification programs such as iodisation of salt have been implemented some cases of iodine induced hyperthyroidism have been observed so called 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 122 Deficiency Edit Main article Iodine deficiency In areas where there is little iodine in the diet 123 typically remote inland areas and semi arid equatorial climates where no marine foods are eaten iodine deficiency gives rise to hypothyroidism symptoms of which are extreme fatigue goitre mental slowing depression weight gain and low basal body temperatures 124 Iodine deficiency is the leading cause of preventable 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 125 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 126 Precautions EditToxicity Edit Iodine HazardsGHS labelling Pictograms Signal word DangerHazard statements H312 H315 H319 H332 H335 H372 H400Precautionary statements P261 P273 P280 P305 P314 P338 P351 127 NFPA 704 fire diamond 128 300 Elemental iodine I2 is toxic if taken orally undiluted The lethal dose for an adult human is 30 mg kg which is about 2 1 2 4 grams for a human weighing 70 to 80 kg even if experiments on rats demonstrated that these animals could survive after eating a 14000 mg kg dose Excess iodine can be more cytotoxic in the presence of selenium deficiency 129 Iodine supplementation in selenium deficient populations is in theory problematic partly for this reason 109 The toxicity derives from its oxidizing properties through which it denaturates proteins including enzymes 130 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 tincture of iodine and Lugol s solution are capable of causing tissue damage if used in prolonged cleaning or antisepsis similarly liquid Povidone iodine Betadine trapped against the skin resulted in chemical burns in some reported cases 131 Occupational exposure Edit People can be exposed to iodine in the workplace by inhalation ingestion skin contact and eye contact The Occupational Safety and Health Administration OSHA has set the legal limit Permissible exposure limit for iodine exposure in the workplace at 0 1 ppm 1 mg m3 during an 8 hour workday The National Institute for Occupational Safety and Health NIOSH has set a Recommended exposure limit REL of 0 1 ppm 1 mg m3 during an 8 hour workday At levels of 2 ppm iodine is immediately dangerous to life and health 132 Allergic reactions Edit Some people develop a hypersensitivity to products and foods containing iodine Applications of tincture of iodine or Betadine can cause rashes sometimes severe 133 Parenteral use of iodine based contrast agents see above can cause reactions ranging from a mild rash to fatal 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 134 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 135 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 136 135 As with all medications the patient s allergy history should be questioned and consulted before any containing iodine are administered 137 US DEA List I status Edit Phosphorus can reduce elemental iodine to hydroiodic acid which is a reagent effective for reducing ephedrine or pseudoephedrine to methamphetamine 138 For this reason iodine was designated by the United States Drug Enforcement Administration as a List I precursor chemical under 21 CFR 1310 02 139 References Edit Standard Atomic Weights Iodine CIAAW 1985 I II is known to exist in monoxide IO see Nikitin I V 31 August 2008 Halogen monoxides Russian Chemical Reviews 77 8 739 749 Bibcode 2008RuCRv 77 739N doi 10 1070 RC2008v077n08ABEH003788 S2CID 250898175 Magnetic susceptibility of the elements and inorganic compounds in Handbook of Chemistry and Physics 81st edition CRC press Weast Robert 1984 CRC Handbook of Chemistry and Physics Boca Raton Florida Chemical Rubber Company Publishing pp E110 ISBN 0 8493 0464 4 a b c d Iodine Micronutrient Information Center Linus Pauling Institute Oregon State University Corvallis 2015 Retrieved 20 November 2017 McNeil Jr DG 16 December 2006 In Raising the World s I Q the Secret s in the Salt The New York Times Archived from the original on 12 July 2010 Retrieved 21 July 2009 World Health Organization 2021 World Health Organization model list of essential medicines 22nd list 2021 Geneva World Health Organization hdl 10665 345533 WHO MHP HPS EML 2021 02 a b Courtois B 1813 Decouverte d une substance nouvelle dans le Vareck Discovery of a new substance in seaweed Annales de chimie in French 88 304 310 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 Swain PA 2005 Bernard Courtois 1777 1838 famed for discovering iodine 1811 and his life in Paris from 1798 PDF Bulletin for the History of Chemistry 30 2 103 Archived from the original PDF on 14 July 2010 Retrieved 2 April 2009 Greenwood and Earnshaw p 794 a b 53 Iodine Elements vanderkrogt net Retrieved 23 October 2016 Desormes and Clement made their announcement at the Institut imperial 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 Staff 2 December 1813 Institut Imperial de France Le Moniteur Universel in French 336 1344 Chattaway FD 23 April 1909 The discovery of iodine Chemical News and Journal of Industrial Science 99 2578 193 195 a b Gay Lussac J 1813 Sur un nouvel acide forme avec la substance decourverte par M Courtois On a new acid formed by the substance discovered by Mr Courtois Annales de Chimie in French 88 311 318 Gay Lussac J 1813 Sur la 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