Iodine

Isotopes

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.^[20]

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.^[28] 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.^{[29][30][31][32][33]} 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.^[34] Excited states of iodine-127 and iodine-129 are often used in Mössbauer spectroscopy.^[20]

The other iodine radioisotopes have much shorter half-lives, no longer than days.^[28] 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.^[35] 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.^[36] 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.^[37]

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.^[38] 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.^[39] Iodine-131 is also used as a radioactive tracer.^{[40][41][42][43]}

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.^[44]

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:^[45]

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.^[46]

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.^[47] 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.^[46]

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.^[46]

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.)^[48]

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:^[48]

${\displaystyle {\ce {3TaCl5 + {\underset {(excess)}{5AlI3}}-> 3TaI5 + 5AlCl3}}}$

Lower iodides may be produced either through thermal decomposition or disproportionation, or by reducing the higher iodide with hydrogen or a metal, for example:^[48]

${\displaystyle {\ce {TaI5{}+Ta->[{\text{thermal gradient}}][{\ce {630^{\circ }C\ ->\ 575^{\circ }C}}]Ta6I14}}}$

Most of the iodides of 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 MI_n tend to have the lowest melting and boiling points among the halides MX_n 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.^[48]

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).^[49]

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.^[49] 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.^[49] 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.^[49]

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.^[50]

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.^[51]

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.^[24] 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 I^III(I^VO₃)₃ and [IO]⁺[IO₃]⁻ respectively.^[52]

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:^[53]

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:^[53]

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.^[53] 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.^[54]

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:^[55]

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.^[55] 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.^[24]

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:^[24]

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 C_2h green or black I⁺
₅, known in the AsF⁻
₆ and AlCl⁻
₄ salts among others.^[24][56]

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:^[24]

I₂ + I⁻ ⇌ I⁻
₃ (K_eq = ~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.^[24][57]

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,^[58] the Williamson ether synthesis,^[59] the Wurtz coupling reaction,^[60] and in Grignard reagents.^[61]

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.^[62] They are also significantly denser than the other organohalogen compounds thanks to the high atomic weight of iodine.^[63] 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,^[64] and iodobenzene dichloride (PhICl₂), used for the selective chlorination of alkenes and alkynes.^[65] 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:^[66]

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.^[67] For example, iodoacetamide and iodoacetic acid denature proteins by irreversibly alkylating cysteine residues and preventing the reformation of disulfide linkages.^[68]

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.^[66] 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.^[69] The reaction is driven toward products by mass action due to the precipitation of the insoluble salt.^[70][71]

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.^[77]

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. In biology, linoleic acid (C18:2 n-6), omega-6 and alpha-linolenic (C18:3 n-3) omega-3, arachidonic acid (AA) – omega-6 (C20: 4n-6), and docosahexaenoic acid (DHA) – omega-3 (C22:6n-3) synthesized with iodine iodolipids developed among cell membranes during the evolution of life, important in the mechanism of apoptosis, carcinogenesis and degenerative diseases.^[78][79]

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. Aqueous alkaline iodine solution is used in the iodoform test for methyl ketones.^[66]

Spectroscopy

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, J_even), or 21 components (from odd rotational quantum numbers, J_odd) are measurable.^[80]

Cesium 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

In early 2021, 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.^[81]

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, the ESA's GOCE satellite, or NASA's DART mission, all of which used xenon as the reaction mass. Yet iodine's atomic weight is only 3.3% less than that of xenon, while its first two ionization energies average 12% less; together, these make iodine ions a promising substitute.^[81][82]

Use of iodine should allow more widespread application of ion-thrust technology, particularly with smaller-scale space vehicles.^[82] 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."^[83]

Medicine

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.^[88][89]

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.^[90]

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.^[91] 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.

Others

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.^[92]

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.^[92]

The organoiodine compound erythrosine is an important food coloring agent. Perfluoroalkyl iodides are precursors to important surfactants, such as perfluorooctanesulfonic acid.^[92]

The iodine clock reaction (in which iodine also serves as a test for starch, forming a dark blue complex),^[20] 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.^[93]

¹²⁵I is used as the radiolabel in investigating which ligands go to which plant pattern recognition receptors (PRRs).^[94]

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.^[4][100] The Tolerable Upper Intake Level (UL) for adults is 1,100 μg/day.^[101] This upper limit was assessed by analyzing the effect of supplementation on thyroid-stimulating hormone.^[99]

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 lactation, gastric mucosa, salivary glands, brain cells, choroid plexus, thymus, and arterial walls.^{[4][102][103][79]}

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.^[104][105] Iodised salt is fortified with iodine in the form of sodium iodide.^[105][106]

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.^[101] The general US population has adequate iodine nutrition,^[107][108] with women of childbearing age and pregnant women having a possible mild risk of deficiency.^[108] In Japan, consumption was considered much higher, ranging between 5,280 μg/day to 13,800 μg/day from dietary seaweed or kombu kelp,^[99] 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.^[109] The adult UL in Japan was last revised to 3,000 µg/day in 2015.^[110]

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.^[111]

Deficiency

In areas where there is little iodine in the diet,^[112] 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.^[113] 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.^[114] 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.^[115]

Toxicity

Iodine

Hazards
GHS labelling:
Pictograms
Signal word	Danger
Hazard statements	H312, H315, H319, H332, H335, H372, H400
Precautionary statements	P261, P273, P280, P305, P314, P338, P351[116]
NFPA 704 (fire diamond)	[117] 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.^[118] Iodine supplementation in selenium-deficient populations is, in theory, problematic, partly for this reason.^[99] The toxicity derives from its oxidizing properties, through which it denaturates proteins (including enzymes).^[119]

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.^[120]

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.^[127] For this reason, iodine was designated by the United States Drug Enforcement Administration as a List I precursor chemical under 21 CFR 1310.02.^[128]