Iodine is a chemical element with symbol I and atomic number 53. The heaviest of the stable halogens, it exists as a lustrous, purple-black non-metallic solid at standard conditions that sublimes readily to form a violet gas. The elemental form was discovered by the French chemist Bernard Courtois in 1811. It was named two years later by Joseph Louis Gay-Lussac from this property, after the Greek ἰώδης "violet-coloured".

Iodine occurs in many oxidation states, including iodide (I), iodate (IO
), and the various periodate anions. It is the least abundant of the stable halogens, being the sixty-first most abundant element. It is the heaviest essential mineral nutrient. Iodine is essential in the synthesis of thyroid hormones.[4] Iodine deficiency affects about two billion people and is the leading preventable cause of intellectual disabilities.

The dominant producers of iodine today are Chile and Japan. Iodine and its compounds are primarily used in nutrition. 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.

Iodine,  53I
Sample of iodine
Pronunciation/ˈaɪədaɪn, -dɪn, -diːn/ (EYE-ə-dyn, -⁠din, -⁠deen)
Appearancelustrous metallic gray, violet as a gas
Standard atomic weight Ar, std(I)126.90447(3)[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


Atomic number (Z)53
Groupgroup 17 (halogens)
Periodperiod 5
Element category  reactive nonmetal
Electron configuration[Kr] 4d10 5s2 5p5
Electrons per shell
2, 8, 18, 18, 7
Physical properties
Phase at STPsolid
Melting point386.85 K ​(113.7 °C, ​236.66 °F)
Boiling point457.4 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, +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
Color lines in a spectral range
Spectral lines of iodine
Other properties
Natural occurrenceprimordial
Crystal structureorthorhombic
Orthorhombic crystal structure for iodine
Thermal conductivity0.449 W/(m·K)
Electrical resistivity1.3×107 Ω·m (at 0 °C)
Magnetic orderingdiamagnetic[2]
Magnetic susceptibility−88.7·10−6 cm3/mol (298 K)[3]
Bulk modulus7.7 GPa
CAS Number7553-56-2
Discovery and first isolationBernard Courtois (1811)
Main isotopes of iodine
Iso­tope Abun­dance Half-life (t1/2) Decay mode Pro­duct
123I syn 13 h ε, γ 123Te
124I syn 4.176 d ε 124Te
125I syn 59.40 d ε 125Te
127I 100% stable
129I trace 1.57×107 y β 129Xe
131I syn 8.02070 d β, γ 131Xe
135I syn 6.57 h β 135Xe


In 1811, iodine was discovered by French chemist Bernard Courtois,[5][6] 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.[7] Courtois suspected that this material was a new element but lacked funding to pursue it further.[8]

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.[9] On 6 December, Gay-Lussac announced that the new substance was either an element or a compound of oxygen.[10][11][12] It was Gay-Lussac who suggested the name "iode", from the Greek word ἰοειδής[13] (ioeidēs) for violet (because of the colour of iodine vapor).[5][10] 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.[14] Davy sent a letter dated 10 December to the Royal Society of London stating that he had identified a new element.[15] Arguments erupted between Davy and Gay-Lussac over who identified iodine first, but both scientists acknowledged Courtois as the first to isolate the element.[8]

In early periodic tables, iodine is often given the symbol J, for jod, its name in German.[16]


Violet iodine vapour in a round-bottomed 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 scarce and fugitive fifth halogen, the radioactive astatine, is not well-studied due to its expense and inaccessibility in large quantities, but appears to show various unusual properties 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 a strong 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.[17] (Astatine goes further, being indeed unstable as At and readily oxidised to At0 or At+, although the existence of At2 is not settled.)[18]

The halogens darken in colour as the group is descended: fluorine is a very pale yellow gas, chlorine is greenish-yellow, and bromine is a reddish-brown volatile liquid. Iodine conforms to the prevailing trend, being a shiny black crystalline solid that melts at 114 °C and boils at 183 °C to form a violet gas. This trend occurs because the wavelengths of visible light absorbed by the halogens increase down the group (though astatine may not conform to it, depending on how metallic it turns out to be).[17] Specifically, the violet colour of iodine gas results from the electron transition between the highest occupied antibonding πg molecular orbital and the lowest vacant antibonding σu molecular orbital.[19]

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.[19] Nonpolar solvents such as hexane and carbon tetrachloride provide a higher solubility.[20] 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.[19] 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.[21]

Iodine-triphenylphosphine charge-transfer complex in dichloromethane
I2PPh3 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].[22]

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.[17] 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.[17]

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.[17] 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.)[23] 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.[17] 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.[17]


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 monoisotopic and its atomic weight is known to great precision, as it is a constant of nature.[17]

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.[24] 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.[25][26][27][28][29] 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 incredibly 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.[30] Excited states of iodine-127 and iodine-129 are often used in Mössbauer spectroscopy.[17]

The other iodine radioisotopes have much shorter half-lives, no longer than days.[24] 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.[31] 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.[32] 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.[33]

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.[34] 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.[35] Iodine-131 is also used as a radioactive tracer.[36][37][38][39]

Chemistry and compounds

Halogen bond energies (kJ/mol)[19]
F 159 574 645 582 456
Cl 243 428 444 427 327
Br 193 363 368 360 272
I 151 294 272 285 239

Though it is the least reactive of the halogens, iodine is still one of the more reactive elements. 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.[17] 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.[19]

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:[40]

2 I2 + N2H4 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.[41]

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.[42] 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.[41]

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
ions – the latter, in any case, are much less stable than the bifluoride ions (HF
) 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+
(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.[41]

Other binary iodides

Nearly all elements in the periodic table form binary iodides. The exceptions are decidedly in the minority and stem in each case from one of three causes: extreme inertness and reluctance to participate in chemical reactions (the noble gases); extreme nuclear instability hampering chemical investigation before decay and transmutation (many of the heaviest elements beyond bismuth); and having an electronegativity higher than iodine's (oxygen, nitrogen, and the first three halogens), so that the resultant binary compounds are formally not iodides but rather oxides, nitrides, or halides of iodine. (Nonetheless, nitrogen triiodide is named as an iodide as it is analogous to the other nitrogen trihalides.)[43]

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:[43]

Lower iodides may be produced either through thermal decomposition or disproportionation, or by reducing the higher iodide with hydrogen or a metal, for example:[43]

Most of the iodides of the pre-transition metals (groups 1, 2, and 3, along with the lanthanides and actinides in the +2 and +3 oxidation states) are mostly ionic, while nonmetals tend to form covalent molecular iodides, as do metals in high oxidation states from +3 and above. 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.[43]

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+
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).[44]

Iodine monochloride1
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.[44] 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 1813–4 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.[44] When liquid, iodine monochloride and iodine monobromide dissociate into I
and IX
anions (X = Cl, Br); thus they are significant conductors of electricity and can be used as ionising solvents.[44]

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+
and ICl

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+
and IF
. 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.[46]

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.[21] 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.[47]

Standard reduction potentials for aqueous I species[48]
E°(couple) a(H+) = 1
E°(couple) a(OH) = 1
I2/I +0.535 I2/I +0.535
HOI/I +0.987 IO/I +0.48
HOI/I2 +1.439 IO/I2 +0.42
+1.134 IO
+1.601 H

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:[48]

I2 + H2O ⇌ HIO + H+ + I Kac = 2.0 × 10−13 mol2 l−2
I2 + 2 OH ⇌ IO + H2O + I Kalk = 30 mol−1 l

Hypoiodous acid is unstable to disproportionation. The hypoiodite ions thus formed disproportionate immediately to give iodide and iodate:[48]

3 IO ⇌ 2 I + IO
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.[48] 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.[49]

Many periodates are known, including not only the expected tetrahedral IO
, but also square-pyramidal IO3−
, octahedral orthoperiodate IO5−
, [IO3(OH)3]2−, [I2O8(OH2)]4−, and I
. They are usually made by oxidising alkaline sodium iodate electrochemically (with lead(IV) oxide as the anode) or by chlorine gas:[50]

+ 6 OHIO5−
+ 3 H2O + 2 e
+ 6 OH + Cl2IO5−
+ 2 Cl + 3 H2O

They are thermodymically and kinetically powerful oxidising agents, quickly oxidising Mn2+ to MnO
, and cleaving glycols, α-diketones, α-ketols, α-aminoalcohols, and α-diamines.[50] 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)+
cation, isoelectronic to Te(OH)6 and Sb(OH)
, and giving salts with bisulfate and sulfate.[21]

Polyiodine compounds

When iodine dissolves in strong acids, such as fuming sulfuric acid, a bright blue paramagnetic solution including I+
cations is formed. A solid salt of the diiodine cation may be obtained by oxidising iodine with antimony pentafluoride:[21]

2 I2 + 5 SbF5 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+
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+
reversibly dimerises below −60 °C, forming red rectangular diamagnetic I2+
. Other polyiodine cations are not as well-characterised, including bent dark-brown or black I+
and centrosymmetric C2h green or black I+
, known in the AsF
and AlCl
salts among others.[21][51]

The only important polyiodide anion in aqueous solution is linear triiodide, I
. Its formation explains why the solubility of iodine in water may be increased by the addition of potassium iodide solution:[21]

I2 + II
(Keq = ~700 at 20 °C)

Many other polyiodides may be found when solutions containing iodine and iodide crystallise, such as I
, I
, I2−
, and I2−
, whose salts with large, weakly polarising cations such as Cs+ may be isolated.[21][52]

Organoiodine compounds

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,[53] the Williamson ether synthesis,[54] the Wurtz coupling reaction,[55] and in Grignard reagents.[56]

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.[57] They are also significantly denser than the other organohalogen compounds thanks to the high atomic weight of iodine.[58] 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,[59] and iodobenzene dichloride (PhICl2), used for the selective chlorination of alkenes and alkynes.[60] 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:[61]

Iodoform synthesis
Iodoform synthesis

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.[62] For example, iodoacetamide and iodoacetic acid denature proteins by irreversibly alkylating cysteine residues and preventing the reformation of disulfide linkages.[63]

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.[61] 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.[64] The reaction is driven toward products by mass action due to the precipitation of the insoluble salt.[65][66]

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). Among the eighty-four elements which occur in significant quantities (elements 1–42, 44–60, 62–83, and 90–92), it ranks sixty-first in abundance.[67] 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.[67] 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 weight.[68] 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.[17]

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),[69] 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.[68][70]

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.[67] 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.[71]


Unlike chlorine and bromine, which have one significant main use dwarfing all others, iodine is used in many applications of varying importance. 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. The remaining 5% is for minor uses. 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.[17]

Chemical analysis

Testing seed for starch
Testing a seed for starch with a solution of iodine

Potassium tetraiodomercurate(II), K2HgI4, is also known as Nessler's reagent. It is often used as a sensitive spot test for ammonia. Similarly, Cu2HgI4 is used as a precipitating reagent to test for alkaloids. The iodide and iodate anions are often used for quantitative volumetric analysis, for example in iodometry and the iodine clock reaction (in which iodine also serves as a test for starch, forming a dark blue complex),[17] and aqueous alkaline iodine solution is used in the iodoform test for methyl ketones.[61] The iodine test for starch is still used to detect counterfeit banknotes printed on starch-containing paper.[72]


The spectra 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.


Elemental iodine

Elemental iodine is used as a disinfectant 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.[73]

In the alternative, iodine may be produced from iodophors, which contain iodine complexed with a solubilizing agent (iodide ion may be thought of loosely as the iodophor in triiodide water solutions). Examples of such preparations include:[74]

The antimicrobial action of iodine is quick and works at low concentrations, and thus it is used in operating theatres.[75] 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.[76]

Other formulations

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.[77][78]

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.[79]

Diatrizoic acid
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.[80] 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.


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.[81]

Inorganic iodides find specialised uses. Titanium, zirconium, hafnium, and thorium are purified by the van Arkel 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.[81]

The organoiodine compound erythrosine is an important food coloring agent. Perfluoroalkyl iodides are precursors to important surfactants, such as perfluorooctanesulfonic acid.[81]

Biological role

Thyroid system
The thyroid system of the thyroid hormones T3 and T4
Carte iodurie france µg par jour d'après Mornex 1987 Le Guen 2000
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)[82]

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, are used by a few microorganisms.)[83][84] 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.[85]

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. 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.[86]

Dietary intake

Recommendations by the United States Institute 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.[4][87] The Tolerable Upper Intake Level (UL) for adults is 1,100 μg/day.[88] This upper limit was assessed by analyzing the effect of supplementation on thyroid-stimulating hormone.[86]

The thyroid gland needs no more than 70 μg/day to synthesise the requisite daily amounts of T4 and T3.[4] The higher recommended daily allowance levels of iodine seem necessary for optimal function of a number of body systems, including lactating breast, gastric mucosa, salivary glands, brain cells, choroid plexus, thymus, and arterial walls.[4][89][90][91]

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.[92][93] Iodised salt is fortified with iodine in the form of sodium iodide.[93][94]

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.[88] The general US population has adequate iodine nutrition,[95][96] with women of childbearing age and pregnant women having a possible mild risk of deficiency.[96] In Japan, consumption was considered much higher, ranging between 5,280 μg/day to 13,800 μg/day from dietary seaweed or kombu kelp,[86] 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.[97] The adult UL in Japan was last revised to 3,000 µg/day in 2015.[98]

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.[99]


In areas where there is little iodine in the diet,[100] 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.[101] 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 the wealthier nations, but iodine deficiency remains a serious public health problem today in the developing world.[102] 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.[103]


GHS pictograms The exclamation-mark pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)The health hazard pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)The environment pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)
GHS signal word Danger
H312, H332, H315, H319, H335, H372, H400
P261, P273, P280, P305, P351, P338, P314[104]
NFPA 704

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.[106] Iodine supplementation in selenium-deficient populations is, in theory, problematic, partly for this reason.[86] The toxicity derives from its oxidizing properties, through which it denaturates proteins (including enzymes).[107]

Elemental iodine is also a skin irritant, and 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.[108]

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.[109]

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.[110] 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.[111] 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;[112] 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.) or asthma are more likely to suffer reactions to contrast media containing iodine.[112] As with all medications, the patient's allergy history should be questioned and consulted before any containing iodine are administered.[113]


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Astatine is a radioactive chemical element with symbol At and atomic number 85. It is the rarest naturally occurring element in the Earth's crust, occurring only as the decay product of various heavier elements. All of astatine's isotopes are short-lived; the most stable is astatine-210, with a half-life of 8.1 hours. A sample of the pure element has never been assembled, because any macroscopic specimen would be immediately vaporized by the heat of its own radioactivity.

The bulk properties of astatine are not known with any certainty. Many of them have been estimated based on the element's position on the periodic table as a heavier analog of iodine, and a member of the halogens (the group of elements including fluorine, chlorine, bromine, and iodine). Astatine is likely to have a dark or lustrous appearance and may be a semiconductor or possibly a metal; it probably has a higher melting point than that of iodine. Chemically, several anionic species of astatine are known and most of its compounds resemble those of iodine. It also shows some metallic behavior, including being able to form a stable monatomic cation in aqueous solution (unlike the lighter halogens).

The first synthesis of the element was in 1940 by Dale R. Corson, Kenneth Ross MacKenzie, and Emilio G. Segrè at the University of California, Berkeley, who named it from the Greek astatos (ἄστατος), meaning "unstable". Four isotopes of astatine were subsequently found to be naturally occurring, although much less than one gram is present at any given time in the Earth's crust. Neither the most stable isotope astatine-210, nor the medically useful astatine-211, occur naturally; they can only be produced synthetically, usually by bombarding bismuth-209 with alpha particles.


Bromine is a chemical element with symbol Br and atomic number 35. It is the third-lightest halogen, and is a fuming red-brown liquid at room temperature that evaporates readily to form a similarly coloured gas. Its properties are thus intermediate between those of chlorine and iodine. Isolated independently by two chemists, Carl Jacob Löwig (in 1825) and Antoine Jérôme Balard (in 1826), its name was derived from the Ancient Greek βρῶμος ("stench"), referencing its sharp and disagreeable smell.

Elemental bromine is very reactive and thus does not occur free in nature, but in colourless soluble crystalline mineral halide salts, analogous to table salt. While it is rather rare in the Earth's crust, the high solubility of the bromide ion (Br−) has caused its accumulation in the oceans. Commercially the element is easily extracted from brine pools, mostly in the United States, Israel and China. The mass of bromine in the oceans is about one three-hundredth that of chlorine.

At high temperatures, organobromine compounds readily dissociate to yield free bromine atoms, a process that stops free radical chemical chain reactions. This effect makes organobromine compounds useful as fire retardants, and more than half the bromine produced worldwide each year is put to this purpose. The same property causes ultraviolet sunlight to dissociate volatile organobromine compounds in the atmosphere to yield free bromine atoms, causing ozone depletion. As a result, many organobromide compounds—such as the pesticide methyl bromide—are no longer used. Bromine compounds are still used in well drilling fluids, in photographic film, and as an intermediate in the manufacture of organic chemicals.

Large amounts of bromide salts are toxic from the action of soluble bromide ion, causing bromism. However, a clear biological role for bromide ion and hypobromous acid has recently been elucidated, and it now appears that bromine is an essential trace element in humans. The role of biological organobromine compounds in sea life such as algae has been known for much longer. As a pharmaceutical, the simple bromide ion (Br−) has inhibitory effects on the central nervous system, and bromide salts were once a major medical sedative, before replacement by shorter-acting drugs. They retain niche uses as antiepileptics.

Congenital iodine deficiency syndrome

Congenital iodine deficiency syndrome, previously known as cretinism, is a condition of severely stunted physical and mental growth owing to untreated congenital deficiency of thyroid hormone (congenital hypothyroidism) usually owing to maternal hypothyroidism.


A goitre, or goiter, is a swelling in the neck resulting from an enlarged thyroid gland. A goitre can be associated with a thyroid that is not functioning properly.

Worldwide, over 90% of goitre cases are caused by iodine deficiency. The term is from the Latin gutturia, meaning throat. Most goitres are of a benign nature.


The halogens () are a group in the periodic table consisting of five chemically related elements: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). The artificially created element 117 (tennessine, Ts) may also be a halogen. In the modern IUPAC nomenclature, this group is known as group 17. The symbol X is often used generically to refer to any halogen.

The name "halogen" means "salt-producing". When halogens react with metals they produce a wide range of salts, including calcium fluoride, sodium chloride (common table salt), silver bromide and potassium iodide.

The group of halogens is the only periodic table group that contains elements in three of the main states of matter at standard temperature and pressure. All of the halogens form acids when bonded to hydrogen. Most halogens are typically produced from minerals or salts. The middle halogens, that is chlorine, bromine and iodine, are often used as disinfectants. Organobromides are the most important class of flame retardants. Elemental halogens are dangerous and can be lethally toxic.


Hyperthyroidism is the condition that occurs due to excessive production of thyroid hormone by the thyroid gland. Thyrotoxicosis is the condition that occurs due to excessive thyroid hormone of any cause and therefore includes hyperthyroidism. Some, however, use the terms interchangeably. Signs and symptoms vary between people and may include irritability, muscle weakness, sleeping problems, a fast heartbeat, heat intolerance, diarrhea, enlargement of the thyroid, hand tremor, and weight loss. Symptoms are typically less in the old and during pregnancy. An uncommon complication is thyroid storm in which an event such as an infection results in worsening symptoms such as confusion and a high temperature and often results in death. The opposite is hypothyroidism, when the thyroid gland does not make enough thyroid hormone.Graves' disease is the cause of about 50% to 80% of the cases of hyperthyroidism in the United States. Other causes include multinodular goiter, toxic adenoma, inflammation of the thyroid, eating too much iodine, and too much synthetic thyroid hormone. A less common cause is a pituitary adenoma. The diagnosis may be suspected based on signs and symptoms and then confirmed with blood tests. Typically blood tests show a low thyroid stimulating hormone (TSH) and raised T3 or T4. Radioiodine uptake by the thyroid, thyroid scan, and TSI antibodies may help determine the cause.Treatment depends partly on the cause and severity of disease. There are three main treatment options: radioiodine therapy, medications, and thyroid surgery. Radioiodine therapy involves taking iodine-131 by mouth which is then concentrated in and destroys the thyroid over weeks to months. The resulting hypothyroidism is treated with synthetic thyroid hormone. Medications such as beta blockers may control the symptoms, and anti-thyroid medications such as methimazole may temporarily help people while other treatments are having effect. Surgery to remove the thyroid is another option. This may be used in those with very large thyroids or when cancer is a concern. In the United States hyperthyroidism affects about 1.2% of the population. It occurs between two and ten times more often in women. Onset is commonly between 20 and 50 years of age. Overall the disease is more common in those over the age of 60 years.


Hypothyroidism, also called underactive thyroid or low thyroid, is a disorder of the endocrine system in which the thyroid gland does not produce enough thyroid hormone. It can cause a number of symptoms, such as poor ability to tolerate cold, a feeling of tiredness, constipation, depression, and weight gain. Occasionally there may be swelling of the front part of the neck due to goiter. Untreated hypothyroidism during pregnancy can lead to delays in growth and intellectual development in the baby or congenital iodine deficiency syndrome.Worldwide, too little iodine in the diet is the most common cause of hypothyroidism. In countries with enough iodine in the diet, the most common cause of hypothyroidism is the autoimmune condition Hashimoto's thyroiditis. Less common causes include: previous treatment with radioactive iodine, injury to the hypothalamus or the anterior pituitary gland, certain medications, a lack of a functioning thyroid at birth, or previous thyroid surgery. The diagnosis of hypothyroidism, when suspected, can be confirmed with blood tests measuring thyroid-stimulating hormone (TSH) and thyroxine levels.Salt iodization has prevented hypothyroidism in many populations. Hypothyroidism can be treated with levothyroxine. The dose is adjusted according to symptoms and normalization of the thyroxine and TSH levels. Thyroid medication is safe in pregnancy. While a certain amount of dietary iodine is important, excessive amounts can worsen certain types of hypothyroidism.Worldwide about one billion people are estimated to be iodine deficient; however, it is unknown how often this results in hypothyroidism. In the United States, hypothyroidism occurs in 0.3–0.4% of people. Subclinical hypothyroidism, a milder form of hypothyroidism characterized by normal thyroxine levels and an elevated TSH level, is thought to occur in 4.3–8.5% of people in the United States. Hypothyroidism is more common in women than men. People over the age of 60 are more commonly affected. Dogs are also known to develop hypothyroidism and in rare cases cats and horses. The word "hypothyroidism" is from Greek hypo- meaning "reduced", thyreos for "shield", and eidos for "form."


Iodine-131 (131I) is an important radioisotope of iodine discovered by Glenn Seaborg and John Livingood in 1938 at the University of California, Berkeley. It has a radioactive decay half-life of about eight days. It is associated with nuclear energy, medical diagnostic and treatment procedures, and natural gas production. It also plays a major role as a radioactive isotope present in nuclear fission products, and was a significant contributor to the health hazards from open-air atomic bomb testing in the 1950s, and from the Chernobyl disaster, as well as being a large fraction of the contamination hazard in the first weeks in the Fukushima nuclear crisis. This is because I-131 is a major fission product of uranium and plutonium, comprising nearly 3% of the total products of fission (by weight). See fission product yield for a comparison with other radioactive fission products. I-131 is also a major fission product of uranium-233, produced from thorium.

Due to its mode of beta decay, iodine-131 is notable for causing mutation and death in cells that it penetrates, and other cells up to several millimeters away. For this reason, high doses of the isotope are sometimes less dangerous than low doses, since they tend to kill thyroid tissues that would otherwise become cancerous as a result of the radiation. For example, children treated with moderate dose of I-131 for thyroid adenomas had a detectable increase in thyroid cancer, but children treated with a much higher dose did not. Likewise, most studies of very-high-dose I-131 for treatment of Graves disease have failed to find any increase in thyroid cancer, even though there is linear increase in thyroid cancer risk with I-131 absorption at moderate doses. Thus, iodine-131 is increasingly less employed in small doses in medical use (especially in children), but increasingly is used only in large and maximal treatment doses, as a way of killing targeted tissues. This is known as "therapeutic use".

Iodine-131 can be "seen" by nuclear medicine imaging techniques (i.e., gamma cameras) whenever it is given for therapeutic use, since about 10% of its energy and radiation dose is via gamma radiation. However, since the other 90% of radiation (beta radiation) causes tissue damage without contributing to any ability to see or "image" the isotope, other less-damaging radioisotopes of iodine such as iodine-123 (see isotopes of iodine) are preferred in situations when only nuclear imaging is required. The isotope I-131 is still occasionally used for purely diagnostic (i.e., imaging) work, due to its low expense compared to other iodine radioisotopes. Very small medical imaging doses of I-131 have not shown any increase in thyroid cancer. The low-cost availability of I-131, in turn, is due to the relative ease of creating I-131 by neutron bombardment of natural tellurium in a nuclear reactor, then separating I-131 out by various simple methods (i.e., heating to drive off the volatile iodine). By contrast, other iodine radioisotopes are usually created by far more expensive techniques, starting with reactor radiation of expensive capsules of pressurized xenon gas.

Iodine-131 is also one of the most commonly used gamma-emitting radioactive industrial tracer. Radioactive tracer isotopes are injected with hydraulic fracturing fluid to determine the injection profile and location of fractures created by hydraulic fracturing.Much smaller incidental doses of iodine-131 than those used in medical therapeutic procedures, are supposed by some studies to be the major cause of increased thyroid cancers after accidental nuclear contamination. These studies suppose that cancers happen from residual tissue radiation damage caused by the I-131, and should appear mostly years after exposure, long after the I-131 has decayed. Other studies can't find a correlation.

Iodine deficiency

Iodine deficiency is a lack of the trace element iodine, an essential nutrient in the diet. It may result in a goiter, sometimes as an endemic goiter as well as cretinism due to untreated congenital hypothyroidism, which results in developmental delays and other health problems. Iodine deficiency is an important public health issue as it is a preventable cause of intellectual disability.

Iodine heptafluoride

Iodine heptafluoride, also known as iodine(VII) fluoride or iodine fluoride, is an interhalogen compound with the chemical formula IF7. It has an unusual pentagonal bipyramidal structure, as predicted by VSEPR theory. The molecule can undergo a pseudorotational rearrangement called the Bartell mechanism, which is like the Berry mechanism but for a heptacoordinated system.

It forms colourless crystals, which melt at 4.5 °C: the liquid range is extremely narrow, with the boiling point at 4.77 °C. The dense vapor has a mouldy, acrid odour. The molecule has D5h symmetry.

Iodine monofluoride

Iodine monofluoride is an interhalogen compound of iodine and fluorine with formula IF. It is a chocolate-brown solid that decomposes at 0 C, disproportionating to elemental iodine and iodine pentafluoride:

5 IF → 2 I2 + IF5However, its molecular properties can still be precisely determined by spectroscopy: the iodine-fluorine distance is 190.9 pm and the I−F bond dissociation energy is around 277 kJ mol−1. At 298 K, its standard enthalpy change of formation is ΔHf° = −95.4 kJ mol−1, and its Gibbs free energy is ΔGf° = −117.6 kJ mol−1.

It can be generated, albeit only fleetingly, by the reaction of the elements at −45 °C in CCl3F:

I2 + F2 → 2 IFIt can also be generated by the reaction of iodine with iodine trifluoride at −78 °C in CCl3F:

I2 + IF3 → 3 IFThe reaction of iodine with silver(I) fluoride at 0 °C also yields iodine monofluoride:

I2 + AgF → IF + AgI

Iodised salt

Iodised salt (also spelled iodized salt) is table salt mixed with a minute amount of various salts of the element iodine. The ingestion of iodine prevents iodine deficiency. Worldwide, iodine deficiency affects about two billion people and is the leading preventable cause of intellectual and developmental disabilities. Deficiency also causes thyroid gland problems, including "endemic goitre". In many countries, iodine deficiency is a major public health problem that can be cheaply addressed by purposely adding small amounts of iodine to the sodium chloride salt.

Iodine is a micronutrient and dietary mineral that is naturally present in the food supply in some regions, especially near sea coasts, but is generally quite rare in the Earth's crust, since iodine is a so-called heavy element, and abundance of chemical elements generally declines with greater atomic mass. Where natural levels of iodine in the soil are low and the iodine is not taken up by vegetables, iodine added to salt provides the small but essential amount of iodine needed by humans.

An opened package of table salt with iodide may rapidly lose its iodine content through the process of oxidation and iodine sublimation.

Iofetamine (123I)

Iofetamine (iodine-123, 123I), brand names Perfusamine, SPECTamine), or N-isopropyl-(123I)-p-iodoamphetamine (IMP), is a lipid-soluble amine and radiopharmaceutical drug used in cerebral blood perfusion imaging with single photon emission computed tomography (SPECT). Labeled with the radioactive isotope iodine-123, it is approved for use in the United States as a diagnostic aid in determining the localization of and in the evaluation of non-lacunar stroke and complex partial seizures, as well as in the early diagnosis of Alzheimer's disease.An analogue of amphetamine, iofetamine has shown to inhibit the reuptake of serotonin and norepinephrine as well as induce the release of these neurotransmitters and of dopamine with similar potencies to other amphetamines like d-amphetamine and p-chloroamphetamine. In addition, on account of its high lipophilicity, iofetamine rapidly penetrates the blood-brain-barrier. Accordingly, though not known to have been reported in the medical literature, iofetamine likely possesses psychostimulant and possibly entactogenic effects. However, based on structure-activity relationships, it may also be highly neurotoxic to serotonergic and dopaminergic neurons similarly to most other para-halogenated amphetamines.

Isotopes of iodine

There are 37 known isotopes of iodine (53I) from 108I to 144I; all undergo radioactive decay except 127I, which is stable. Iodine is thus a monoisotopic element.

Its longest-lived radioactive isotope, 129I, has a half-life of 15.7 million years, which is far too short for it to exist as a primordial nuclide. Cosmogenic sources of 129I produce very tiny quantities of it that are too small to affect atomic weight measurements; iodine is thus also a mononuclidic element—one that is found in nature only as a single nuclide. Most 129I derived radioactivity on Earth is man-made, an unwanted long-lived byproduct of early nuclear tests and nuclear fission accidents.

All other iodine radioisotopes have half-lives less than 60 days, and four of these are used as tracers and therapeutic agents in medicine. These are 123I, 124I, 125I, and 131I. All industrial production of radioactive iodine isotopes involves these four useful radionuclides.

The isotope 135I has a half-life less than seven hours, which is too short to be used in biology. Unavoidable in situ production of this isotope is important in nuclear reactor control, as it decays to 135Xe, the most powerful known neutron absorber, and the nuclide responsible for the so-called iodine pit phenomenon.

In addition to commercial production, 131I (half-life 8 days) is one of the common radioactive fission-products of nuclear fission, and is thus produced inadvertently in very large amounts inside nuclear reactors. Due to its volatility, short half-life, and high abundance in fission products, 131I, (along with the short-lived iodine isotope 132I from the longer-lived 132Te with a half-life of 3 days) is responsible for the largest part of radioactive contamination during the first week after accidental environmental contamination from the radioactive waste from a nuclear power plant.

Potassium iodide

Potassium iodide is a chemical compound, medication, and dietary supplement. As a medication it is used to treat hyperthyroidism, in radiation emergencies, and to protect the thyroid gland when certain types of radiopharmaceuticals are used. In the developing world it is also used to treat skin sporotrichosis and phycomycosis. As a supplement it is used in those who have low intake of iodine in the diet. It is given by mouth.Common side effects include vomiting, diarrhea, abdominal pain, rash, and swelling of the salivary glands. Other side effects include allergic reactions, headache, goitre, and depression. While use during pregnancy may harm the baby, its use is still recommended in radiation emergencies. Potassium iodide has the chemical formula KI. Commercially it is made by mixing potassium hydroxide with iodine.Potassium iodide has been used medically since at least 1820. It is on the World Health Organization's List of Essential Medicines, the most effective and safe medicines needed in a health system. Potassium iodide is available as a generic medication and over the counter. In the United States a course of treatment is less than 25 USD. Potassium iodide is also used for the iodization of salt.


Povidone-iodine (PVP-I), also known as iodopovidone, is an antiseptic used for skin disinfection before and after surgery. It may be used both to disinfect the skin of the patient and the hands of the healthcare providers. It may also be used for minor wounds. It may be applied to the skin as a liquid or a powder.Side effects include skin irritation. If used on large wounds kidney problems, high blood sodium, and metabolic acidosis may occur. It is not recommended in people who are less than 32 weeks pregnant or are taking lithium. Frequent use is not recommended in people with thyroid problems. Povidone-iodine is a chemical complex of povidone, hydrogen iodide, and elemental iodine. It contains from 9% to 12% available iodine. It works by releasing iodine which results in the death of a range of microorganisms.Povidone-iodine came into commercial use in 1955. It is on the World Health Organization's List of Essential Medicines, the most effective and safe medicines needed in a health system. Povidone-iodine is available over the counter. The wholesale cost in the developing world is about US$3.30 to US$11.40 per liter of 10% solution. This amount in the United Kingdom cost the NHS about £10.86. It is sold under a number of brand names including Betadine.


The thyroid gland, or simply the thyroid, is an endocrine gland in the neck, consisting of two lobes connected by an isthmus. It is found at the front of the neck, below the Adam's apple. The thyroid gland secretes three hormones, namely the thyroid hormones (thyroxine/T4 and triiodothyronine/T3) and calcitonin. The thyroid hormones primarily influence the metabolic rate and protein synthesis, but they also have many other effects including on development. The thyroid hormones are created from iodine and tyrosine. Calcitonin plays a role in calcium homeostasis.Hormonal output from the thyroid is regulated by thyroid-stimulating hormone (TSH) secreted from the anterior pituitary gland, which itself is regulated by thyrotropin-releasing hormone (TRH) produced by the hypothalamus.The thyroid may be affected by several diseases. Hyperthyroidism occurs when the gland produces excessive amounts of thyroid hormones, the most common cause being Graves' disease, an autoimmune disorder. In contrast, hypothyroidism is a state of insufficient thyroid hormone production. Worldwide, the most common cause is iodine deficiency. Thyroid hormones are important for development, and hypothyroidism secondary to iodine deficiency remains the leading cause of preventable intellectual disability. In iodine-sufficient regions, the most common cause of hypothyroidism is Hashimoto's thyroiditis, also an autoimmune disorder. In addition, the thyroid gland may also develop several types of nodules and cancer.

Thyroid hormones

Thyroid hormones are two hormones produced and released by the thyroid gland, namely triiodothyronine (T3) and thyroxine (T4). They are tyrosine-based hormones that are primarily responsible for regulation of metabolism. T3 and T4 are partially composed of iodine. A deficiency of iodine leads to decreased production of T3 and T4, enlarges the thyroid tissue and will cause 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 approximately 14:1. 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. Edward Calvin Kendall was responsible for the isolation of thyroxine in 1915.


Tositumomab is a murine IgG2a lambda monoclonal antibody directed against the CD20 antigen, produced in mammalian cells. Combined with radioisotope iodine 131 it was called Iodine I 131 Tositumomab. Unlabelled tositumomab together with iodine-labelled tositumomab was called Bexxar and a personalized regimen using Bexxar was approved for the treatment of relapsed or chemotherapy/rituxan-refractory Non-Hodgkin lymphoma in 2003.The treatment regimen was personalized for each person receiving the drug in order to maximize the radiation delivered to the tumor and to minimize the amount of radiation to which the person would be exposed. A first dose of labelled antibody was given once, and whole-body radiation was measured with a gamma camera over seven days. Analysis of that imaging data allowed an optimal dose of labelled antibody to be calculated, which was then administered once a day, for up to seven days. Each time the labelled antibody was administered, it was always preceded by unlabelled antibody because early clinical trials had shown that total body residence times of radioactivity were longer in people who first received unlabelled antibody, so that a lower dose of labelled antibody was needed to deliver the required total dose of radiation; additionally labelled antibody targeted tumors better in people pre-treated with unlabelled antibody.This drug combination was developed by Corixa which was purchased by GlaxoSmithKline in 2005. It was sold for about $25,000 for one round of treatment.Sale of Bexxar was discontinued and marketing approval was withdrawn in February 2014 due to the decline in usage (fewer than 75 patients in 2012) even though it had a 70% response rate; the lack of demand was because oncologists could not sell it directly to patients but had to refer patients to third parties, and because of the emergence of other drugs that were as good and could be administered by the oncologists.

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