|Name, symbol||Iodine, 53I|
|Appearance||lustrous metallic gray, violet as a gas|
|Pronounciation||/ˈaɪ.ədaɪn/, /ˈaɪ.ədɨn/, or /ˈaɪ.ədiːn/
eye-ə-dyn, eye-ə-dən, or eye-ə-deen
|Standard atomic weight (±) (Ar)||126.90447(3)|
|Element category||Diatomic nonmetal|
|Group, block||group 17 halogens, p-block|
|Electron configuration||(Kr) 4d10 5s2 5p5|
|Per shell||2, 8, 18, 18, 7|
|Melting point||0K, 114°C, 236.7 °F|
|Boiling point||0K, 114°C, 236.7 °F|
|Density near r.t.||0 g/cm3|
|Triple point||0K, 0 kPa|
|Critical point||0K, 0 MPa|
|Heat of fusion||0kJ/mol|
|Heat of vaporization||0kJ/mol|
|Molar heat capacity||0/(mol-K)
Iodine and its compounds are primarily used in nutrition, and industrially in the production of acetic acid and certain polymers. Iodine's relatively high atomic number, low toxicity, and ease of attachment to organic compounds have made it a part of many X-ray contrast materials in modern medicine. Iodine has only one stable isotope. A number of iodine radioisotopes, such as 131I, are also used in medical applications.
Iodine is found on Earth mainly as the highly water-soluble iodide ion
I−, which concentrates it in oceans and brine pools. Like the other halogens, free iodine occurs mainly as a diatomic molecule I2, and then only momentarily after being oxidized from iodide by an oxidant like free oxygen. In the universe and on Earth, iodine's high atomic number makes it a relatively rare element. However, its presence in ocean water has given it a role in biology. It is the heaviest essential element utilized widely by life in biological functions (only tungsten, employed in enzymes by a few species of bacteria, is heavier). Iodine's rarity in many soils, due to initial low abundance as a crust-element, and also leaching of soluble iodide by rainwater, has led to many deficiency problems in land animals and inland human populations. Iodine deficiency affects about two billion people and is the leading preventable cause of intellectual disabilities.
Under standard conditions, iodine is a bluish-black solid with a metallic lustre, appearing to sublimate into a noxious violet-pink gas, the colour due to absorption of visible light by electronic transitions between the highest occupied and lowest unoccupied molecular orbitals. Melting at 113.7 °C (236.7 °F), it forms compounds with many elements but is less reactive than the other members of its group, the halogens, and has some metallic light reflectance.
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. Nonpolar solvents such as hexane and carbon tetrachloride provide a higher solubility. Polar solutions are brown, reflecting the role of these solvents as Lewis bases, while nonpolar solutions are violet, the color of iodine vapor. Charge-transfer complexes form when iodine is dissolved in polar solvents, modifying the energy distribution of iodine's molecular orbitals, hence changing the colour. A metal ion may replace the solvent, in which case the two species exchange electrons, the ion undergoing π backbonding.
Structure and bonding
Iodine normally exists as a diatomic molecule with an I-I bond length of 270 pm, one of the longest single bonds known. The I2 molecules tend to interact via the weak van der Waals forces called the London dispersion forces, and this interaction is responsible for the higher melting point compared to more compact halogens, which are also diatomic. Since the atomic size of iodine is larger, its melting point is higher. The solid crystallizes as orthorhombic crystals. The crystal motif in the Hermann–Mauguin notation is Cmca (No 64), Pearson symbol oS8. The I-I bond is relatively weak, with a bond dissociation energy of 36 kcal/mol, and most bonds to iodine are weaker than for the lighter halides. One consequence of this weak bonding is the relatively high tendency of I2 molecules to dissociate into atomic iodine.
Of the 37 known (characterized) isotopes of iodine, only one, 127I, is stable.
The longest-lived radioisotope, 129I, has a half-life of 15.7 million years. This is long enough to make it a permanent fixture of the environment on human time scales, but far too short for it to exist as a primordial isotope today. Instead, iodine-129 is an extinct radionuclide, and its presence in the early Solar System is inferred from the observation of an excess of its daughter xenon-129. This nuclide is also newly made by cosmic rays and as a byproduct of artificial nuclear fission, which it is used to monitor as a very long-lived environmental contaminant.
The next-longest-lived radioisotope, iodine-125, has a half-life of 59 days. It is used as a convenient gamma-emitting tag for proteins in biological assays, and a few nuclear medicine imaging tests where a longer half-life is required. It is also commonly used in brachytherapy implanted capsules, which kill tumors by local short-range gamma radiation (but where the isotope is never released into the body).
Iodine-131 (half-life 8 days) is a beta-emitting isotope, which is a common nuclear fission product. It is preferably administered to humans only in very high doses that destroy all tissues that accumulate it (usually the thyroid), which in turn prevents these tissues from developing cancer from a lower dose (paradoxically, a high dose of this isotope appears safer for the thyroid than a low dose). Like other radioiodines, I-131 accumulates in the thyroid gland, but unlike the others, in small amounts it is highly carcinogenic there, it seems, owing to the high local cell mutation due to damage from beta decay. Because of this tendency of 131I to cause high damage to cells that accumulate it and other cells near them (0.6 to 2 mm away, the range of the beta rays), it is the only iodine radioisotope used as direct therapy, to kill tissues such as cancers that take up artificially iodinated molecules (example, the compound iobenguane, also known as MIBG). For the same reason, only the iodine isotope I-131 is used to treat Grave's disease and those types of thyroid cancers (sometimes in metastatic form) where the tissue that requires destruction, still functions to naturally accumulate iodide.
Nonradioactive ordinary potassium iodide (iodine-127), in a number of convenient forms (tablets or solution) may be used to saturate the thyroid gland's ability to take up further iodine, and thus protect against accidental contamination from iodine-131 generated by nuclear fission accidents, such as the Chernobyl disaster and more recently the Fukushima I nuclear accidents, as well as from contamination from this isotope in nuclear fallout from nuclear weapons.
Iodine is rare in the Solar System and Earth's crust (47–60th in abundance); however, iodide salts are often very soluble in water. Iodine occurs in slightly greater concentrations in seawater than in rocks, 0.05 vs. 0.04 ppm. Minerals containing iodine include caliche, found in Chile. The brown algae Laminaria and Fucus found in temperate zones of the Northern Hemisphere contain 0.028–0.454 dry weight percent of iodine. Aside from tungsten, iodine is the heaviest element to be essential in living organisms. About 19,000 tonnes are produced annually from natural sources.
Organoiodine compounds are produced by marine life forms, the most notable being iodomethane (commonly called methyl iodide). About 214 kilotonnes/year of iodomethane is produced by the marine environment, by microbial activity in rice paddies and by the burning of biological material. The volatile iodomethane is broken up in the atmosphere as part of a global iodine cycle.
- Online Etymology Dictionary, s.v. iodine. Retrieved 7 February 2012.
- Raising the World's I.Q., the Secret's in the Salt McNeil, Donald G. Jr. 16 December 2006, The New York Times, accessed 4 December 2008
- Lua error in Module:Citation/CS1 at line 925: attempt to concatenate local 'str' (a table value).
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- InorgChem - Housecroft
- Wells, A. F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.
- Bell, N.; Hsu, L.; Jacob, D.J.; Schultz, M. G.; Blake, D. R.; Butler, J. H.; King, D. B.; Lobert, J. M. et al. (2002). "Methyl iodide: Atmospheric budget and use as a tracer of marine convection in global models". Journal of GeophysicalResearch 107: 4340. Bibcode 2002JGRD..107.4340B. doi:10.1029/2001JD001151.
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