Iodine

For other uses, see Iodine (disambiguation).

Iodine, ₅₃I

Iodine
Pronunciation	/ˈaɪədaɪn, -dɪn, -diːn/ ​(EYE-ə-dyne, -⁠din, -⁠deen)
Appearance	lustrous metallic gray solid, black/violet liquid, violet gas
Standard atomic weight Ar°(I)
	126.90447±0.00003 126.90±0.01 (abridged)
Iodine in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson Br ↑ I ↓ At tellurium ← iodine → xenon
Atomic number (Z)	53
Group	group 17 (halogens)
Period	period 5
Block	p-block
Electron configuration	[Kr] 4d10 5s2 5p5
Electrons per shell	2, 8, 18, 18, 7
Physical properties
Phase at STP	solid
Melting point	(I2) 386.85 K ​(113.7 °C, ​236.66 °F)
Boiling point	(I2) 457.4 K ​(184.3 °C, ​363.7 °F)
Density (at 20° C)	4.944 g/cm3
Triple point	386.65 K, ​12.1 kPa
Critical point	819 K, 11.7 MPa
Heat of fusion	(I2) 15.52 kJ/mol
Heat of vaporization	(I2) 41.57 kJ/mol
Molar heat capacity	(I2) 54.44 J/(mol·K)
Vapor 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	common: −1, +1, +3, +5, +7 +2, +4,? +6?
Electronegativity	Pauling scale: 2.66
Ionization energies	1st: 1008.4 kJ/mol 2nd: 1845.9 kJ/mol 3rd: 3180 kJ/mol
Atomic radius	empirical: 140 pm
Covalent radius	139±3 pm
Van der Waals radius	198 pm
Spectral lines of iodine
Other properties
Natural occurrence	primordial
Crystal structure	​base-centered orthorhombic (oS8)
Lattice constants	a = 725.79 pm b = 478.28 pm c = 982.38 pm (at 20 °C)
Thermal expansion	74.9×10−6/K (at 20 °C)
Thermal conductivity	0.449 W/(m⋅K)
Electrical resistivity	1.3×107 Ω⋅m (at 0 °C)
Magnetic ordering	diamagnetic
Molar magnetic susceptibility	−88.7×10−6 cm3/mol (298 K)
Bulk modulus	7.7 GPa
CAS Number	7553-56-2
History
Discovery and first isolation	Bernard Courtois (1811)
Isotopes of iodine
Main isotopes Decay abun­dance half-life (t1/2) mode pro­duct 123I synth 13.2232 h β+100% 123Te 124I synth 4.1760 d ε 124Te 125I synth 59.392 d ε 125Te 127I 100% stable 129I trace 1.614×107 y β− 129Xe 131I synth 8.0249 d β−100% 131Xe 135I synth 6.58 h β− 135Xe

Iodine (/ˈaɪədaɪn, ˈaɪədɪn/, or /ˈaɪədiːn/; from Template:Lang-el "violet"), is a chemical element that has the symbol I and atomic number 53. Naturally-occurring iodine is a single isotope with 74 neutrons.

Chemically, iodine is the least reactive of the halogens, and the most electropositive halogen after astatine. However, the element does not occur in the free state in nature. As with all other halogens (members of Group XVII in the Periodic Table), when freed from its compounds iodine forms diatomic molecules (I₂).

Iodine and its compounds are primarily used in medicine, photography, and dyes. Although it is rare in the solar system and Earth's crust, the iodides are very soluble in water, and the element is concentrated in seawater. This mechanism helps to explain how the element came to be required in trace amounts by all animals and some plants, being by far the heaviest element known to be necessary to living organisms.

Characteristics

Iodine under standard conditions is a dark-purple/dark-brown solid. It can be seen apparently sublimating at standard temperatures into a violet-pink gas that has an irritating odor. This halogen forms compounds with many elements, but is less reactive than the other members of its Group VII (halogens) and has some metallic light reflectance.

Elemental iodine dissolves easily in chloroform and carbon tetrachloride. The solubility of elemental iodine in water can be vastly increased by the addition of potassium iodide. The molecular iodine reacts reversibly with the negative ion, creating the triiodide anion, I₃⁻, which dissolves well in water. This is also the formulation of some types of medicinal (antiseptic) iodine, although tincture of iodine classically dissolves the element in alcohol. The deep blue color of starch-iodine complexes is produced only by the free element.

Students who have seen the classroom demonstration in which iodine crystals are gently heated in a test tube to violet vapor may gain the impression that liquid iodine does not exist at atmospheric pressure. This misconception arises because the vapor produced has such a deep colour that the liquid appears not to form. In fact, if iodine crystals are heated carefully to just above their melting point of 113.7 °C, the crystals melt into a liquid which is present under a dense blanket of the vapor.

Occurrence

Iodomethane

Iodine naturally occurs in the environment chiefly as a dissolved iodide in seawater, although it is also found in some minerals and soils. This element also exists in small amounts in the mineral caliche, found in Chile, between the Andes and the sea. A type of seaweed, kelp, tends to be high in iodine as well.

Organoiodine compound are produced by marine life forms, the most notable of it is iodomethane (commonly called methyl iodide). The sum of produced iodomethane by in the marine environment, microbial activitiy in rice paddies and burning of biological material is estimated to be 214 kilotonnes. The volatile iodomethane is broken up by oxidation reactions in the atmosphere and a global iodine cycle is established. Although the element is actually quite rare, kelp and certain plants and other algae have some ability to concentrate iodine, which helps introduce the element into the food chain.

Production

File:Iodine.PNG

Iodine output in 2005

From the several places in which iodine occures in nature only two are used as source for iodine: the caliche, found in Chile and the iodine containing brines of gas and oil fields, especially in Japan and the United States.

The caliche, found in Chile contains sodium nitrate, which is the main product of the mining activities and small amounts of sodium iodate and sodium iodide. During leaching and production of pure sodium nitrate the sodium iodate and iodide is extracted. The high concentration of iodine in the caliche and the extensive mining made Chile the largest producer of iodine in 2007.

Most other producers use natural occurring brine for the production of iodine. The Japanese Minami Kanto gas field east of Tokio and the American Anadarko Basin gas field in northwest Oklahoma are the two largest sources for iodine from brine. The brine has a temperature of over 60°C due to the depth it came from. The brine is first purified and acidified using sulfuric acid, then the iodide present is oxidized to iodine with chlorine. An iodine solution is produced, but it is yet too dilute and has to be concentrated. Air is blown into the solution, causing the iodine to evaporate, then it is passed into an absorbing tower containing acid where sulfur dioxide is added to reduce the iodine. The hydrogen iodide (HI) is reacted with chlorine to precipitate the iodine. After filtering and purification the iodine is packed.

2 HI + Cl₂ → I₂↑ + 2 HCl

I₂ + H₂O + SO₂ → 2 HI + H₂SO₄

2 HI + Cl₂ → I₂↓ + 2 HCl

The production of iodine from via electrolysis of seawater is not used due to the sufficent abundance of iodine rich brine. Another source of iodine was kelp, a kind of brown alga. This source was used in the 18th and 19th centuries but is no longer economically viable.

Commercial samples often contain a large amount of impurities; they may be removed by sublimation. The element may also be prepared in an ultrapure form through the reaction of potassium iodide with copper(II) sulfate, which gives copper(II) iodide initially. That decomposes spontaneously to copper(I) iodide and iodine:

Cu²⁺ + 2 I⁻ → CuI₂

2 CuI₂ → 2 CuI + I₂

There are also a few other methods of isolating this element in the laboratory, for example the method used to isolate other halogens: oxidation of the iodide in hydroiodic acid (often made in situ with an iodide and sulfuric acid) by manganese dioxide (see below in Descriptive chemistry).

Isotopes

Main article: isotopes of iodine

There are 37 isotopes of iodine, but only one, ¹²⁷I, is stable.

In many ways, ¹²⁹I is similar to ³⁶Cl. It is a soluble halogen, fairly non-reactive, exists mainly as a non-sorbing anion, and is produced by cosmogenic, thermonuclear, and in-situ reactions. In hydrologic studies, ¹²⁹I concentrations are usually reported as the ratio of ¹²⁹I to total I (which is virtually all ¹²⁷I). As is the case with ³⁶Cl/Cl, ¹²⁹I/I ratios in nature are quite small, 10⁻¹⁴ to 10⁻¹⁰ (peak thermonuclear ¹²⁹I/I during the 1960s and 1970s reached about 10⁻⁷). ¹²⁹I differs from ³⁶Cl in that its halflife is longer (15.7 vs. 0.301 million years), it is highly biophilic, and occurs in multiple ionic forms (commonly, I⁻ and IO₃⁻) which have different chemical behaviors. This makes it fairly easy for ¹²⁹I to enter the biosphere as it becomes incorporated into vegetation, soil, milk, animal tissue, etc.

Excesses of stable ¹²⁹Xe in meteorites have been shown to result from decay of "primordial" iodine-129 produced newly by the supernovas which created the dust and gas from which the solar system formed. ¹²⁹I was the first extinct radionuclide to be identified as present in the early solar system. Its decay is the basis of the I-Xe Iodine-xenon radiometric dating scheme, which covers the first 85 million years of solar system evolution.

Effects of various radioiodine isotopes in biology are discussed below.

History

Iodine was discovered by Bernard Courtois in 1811. He was born to a manufacturer of saltpeter (a vital part of gunpowder). At the time of the Napoleonic Wars, France was at war and saltpeter was in great demand. Saltpeter produced from French niter beds required sodium carbonate, which could be isolated from seaweed washed up on the coasts of Normandy and Brittany. To isolate the sodium carbonate, seaweed was burned and the ash then washed with water. The remaining waste was destroyed by adding sulfuric acid. One day Courtois added too much sulfuric acid and a cloud of purple vapor rose. Courtois noted that the vapor crystallized on cold surfaces making dark crystals. Courtois suspected that this was a new element but lacked the money to pursue his observations.

However he gave samples to his friends, Charles Bernard Desormes (1777–1862) and Nicolas Clément (1779–1841), to continue research. He also gave some of the substance to Joseph Louis Gay-Lussac (1778–1850), a well-known chemist at that time, and to physicist André-Marie Ampère (1775–1836). On 29 November 1813, Dersormes and Clément made public Courtois’ discovery. They described the substance to a meeting of the Imperial Institute of France. On December 6, Gay-Lussac announced that the new substance was either an element or a compound of oxygen. Ampère had given some of his sample to Humphry Davy (1778–1829). Davy did some experiments on the substance and noted its similarity to chlorine. Davy sent a letter dated December 10 to the Royal Society of London stating that he had identified a new element. A large argument erupted between Davy and Gay-Lussac over who identified iodine first but both scientists acknowledged Courtois as the first to isolate the chemical element.

Applications

Disinfectant

Elemental iodine is used as a disinfectant in various forms. The iodine exists as the element, or as the water soluble triiodide anion generated in situ by adding iodide to poorly-soluble iodine. Alternatively, iodine may come from iodophors, which contain iodine complexed with a solubilizing agent. Examples of such preparations include:

• Tincture of iodine (iodine in ethanol, or iodine and sodium iodide in a mixture of ethanol and water)
• Lugol's iodine (iodine and iodide in water)
• Povidone iodine (an iodophor)

Staining

See also: staining

Testing a seed for starch with a solution of iodine

Iodine is a common general stain used in thin-layer chromatography. It is also used in the Gram stain as a mordant, after the sample is treated with crystal violet.

In particular, iodine forms an intense blue complex with starch. Several applications rely on this property:

• Iodometry. The concentration of an oxidant can be determined by adding it to an excess of iodide, to give elemental iodine/triiodide. Starch is used as an indicator close to the end-point to increase the visual contrast (dark blue/colorless instead of yellow of dilute triiodide/colorless).
• Iodine clock reaction is an extension of the techniques in iodometry.
• Iodine may be used to test a sample substance for the presence of starch.
• Iodine solutions are used in counterfeit banknote detection pens; the premise being that counterfeit banknotes made using commercially available paper contain starch.
• Starch-iodide paper are used to test for the presence of oxidants such as peroxides. The oxidants convert iodide to iodine, which shows up as blue. A solution of starch and iodide can perform the same function.

Radiocontrast agent

Diatrizoic acid, a radiocontrast

Iodine, as a heavy element, is quite radio-opaque. Organic compounds of a certain type (typically iodine-substituted benzene derivatives) are thus used in medicine as X-ray radiocontrast agents for intravenous injection. This is often in conjunction with advanced X-ray techniques such as angiography and CT scanning

Radioiodine

Some radioactive iodine isotopes can be used to treat thyroid cancer. The body accumulates iodine in the thyroid, thus radioactive iodine can selectively damage growing thyroid cancer cells while the radioactive dose to the rest of the body remains small.

Organic synthesis

With phosphorus, iodine is able to replace hydroxyl groups on alcohols with iodide. For example, the synthesis of methyl iodide from methanol, red phosphorus, and iodine. The iodinating reagent is phosphorus triiodide that is formed in situ:

3 CH₃OH + PI₃ → 3 CH₃I + H₃PO₃

The iodoform test uses an alkaline solution of iodine to react with methyl ketones to give the labile triiodomethide leaving group, forming iodoform which precipitates.

Iodine is often used to activate magnesium when preparing Grignard reagents.

Clandestine use

In the United States, the Drug Enforcement Agency (DEA) regards iodine and compounds containing iodine (ionic iodides, iodoform, ethyl iodide, and so on) as reagents useful for the clandestine manufacture of methamphetamine. Persons who attempt to purchase significant quantities of such chemicals without establishing a legitimate use are likely to find themselves the target of a DEA investigation. Persons selling such compounds without doing due diligence to establish that the materials are not being diverted to clandestine use may be subject to stiff penalties, such as expensive fines or even imprisonment.

Chemistry

Elemental iodine is poorly soluble in water, with one gram dissolving in 3450 ml at 20 °C and 1280 ml at 50 °C. By contrast with chlorine, the formation of the hypohalite ion (IO^–) in neutral aqueous solutions of iodine is negligible.

I₂+ H₂O H⁺ + I⁻ + HIO   (K = 2.0×10⁻¹³)

Solubility in water is greatly improved if the solution contains dissolved iodides such as hydroiodic acid, potassium iodide, or sodium iodide; this extra solubility results from the high solubility of the I₃^- ion. Dissolved bromides also improve water solubility of iodine. Iodine is soluble in a number of organic solvents, including ethanol (20.5 g/100 ml at 15 °C, 21.43 g/100 ml at 25 °C), diethyl ether (20.6 g/100 ml at 17 °C, 25.20 g/100 ml at 25 °C), chloroform, acetic acid, glycerol, benzene (14.09 g/100 ml at 25 °C), carbon tetrachloride (2.603 g/100 ml at 35 °C), and carbon disulfide (16.47 g/100 ml at 25 °C). Aqueous and ethanol solutions are brown. Solutions in chloroform, carbon tetrachloride, and carbon disulfide are violet.

Elemental iodine can be prepared by oxidizing iodides with chlorine:

2I⁻ + Cl₂ → I₂ + 2Cl⁻

or with manganese dioxide in acid solution:

2I⁻ + 4H⁺ + MnO₂ → I₂ + 2H₂O + Mn²⁺

Iodine is reduced to hydroiodic acid by hydrogen sulfide:

I₂ + H₂S → 2HI + S↓

or by hydrazine:

2I₂ + N₂H₄ → 4HI + N₂

Iodine is oxidized to iodate by nitric acid:

I₂ + 10HNO₃ → 2HIO₃ + 10NO₂ + 4H₂O

or by chlorates:

I₂ + 2ClO₃⁻ → 2IO₃⁻ + Cl₂

Iodine is converted in a two stage reaction to iodide and iodate in solutions of alkali hydroxides (such as sodium hydroxide):

I2 + 2OH− → I− + IO− + H2O		(K = 30)
3IO− → 2I− + IO3−		(K = 1020)

Biological role

Iodine is an essential trace element, the heaviest element known to be needed by living organisms. Its main role in animal biology is as constituents of the thyroid hormones, thyroxine (T4) and triiodothyronine (T3). These are made from addition condensation products of the amino acid tyrosine, and are stored prior to release in an iodine-containing protein called thyroglobulin. T4 and T3 contain four and three atoms of iodine per molecule, respectively. The thyroid gland actively absorbs iodide from the blood to make and release these hormones into the blood, actions which are regulated by a second hormone TSH from the pituitary. Thyroid hormones are phylogenetically very old molecules which are synthesized by most multicellular organisms, and which even have some effect on unicellular organisms.

Thyroid hormones play a basic role in biology, acting on gene transcription to regulate the basal metabolic rate. The total deficiency of thyroid hormones can reduce basal metabolic rate up to 50%, while in excessive production of thyroid hormones the basal metabolic rate can be increased by 100%. T4 acts largely as a precursor to T3, which is (with minor exceptions) the biologically active hormone.

Iodine accounts for 65% of the molecular weight of T4 and 59% of the T3. 15-20 mg of iodine is concentrated in thyroid tissue and hormones, but 70% of the body's iodine is distributed in other tissues, including mammary glands, eyes, gastric mucosa, the cervix, and salivary glands. Its role in mammary tissue is related to fetal and neonatal development, but its role in the other tissues is unknown. It has been shown to act as an antioxidant in these tissues.

Iodine may have a relationship with selenium, and iodine supplementation in selenium-deficient populations may pose risks for thyroid function.

Human dietary intake

The United States Recommended Daily Allowance (RDA) is 150 micrograms per day (μg/day) for both men and women, with a Tolerable Upper Intake Level (UL) for adults is 1,100 μg/day (1.1 mg/day). The tolerable upper limit was assessed by analyzing the effect of supplementation on thyroid-stimulating hormone.

Natural sources of iodine include sea life, such as kelp and certain seafood, as well as plants grown on iodine-rich soil. Iodized salt is fortified with iodine.

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. In Japan, consumption is much higher due to the frequent consumption of seaweed or kombu kelp. Estimates range from 5,280 to 13,800 μg/day. Although some Chinese data associate excess iodine with autoimmune thyroiditis and hypothyroidism, these effects have not been observed in Japanese populations, and a protective effect on breast cancer has been hypothesized.

Donald W. Miller believes that the daily FDA intake recommendation may be 100 times too low.

Deficiency

Main article: Iodine deficiency

In areas where there is little iodine in the diet, 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.

Iodine deficiency is the leading cause of preventable mental retardation, a result which 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 as of March 2006, iodine deficiency remained a serious public health problem in the developing world. Iodine deficiency is also a problem in certain areas of Europe. In Germany it has been estimated to cause a billion dollars in healthcare costs per year.

Radioiodine and biology

Radioiodine and the thyroid

Human exposure to radioactive iodine will cause thyroid uptake, as with all iodine, leading to elevated chances of thyroid cancer. Isotopes with shorter half-lives such as I¹³¹ present a greater risk than those with longer half-lives since they generate more radiation per unit of time. Taking large amounts of regular iodine will saturate the thyroid and prevent uptake. Iodine pills are sometimes distributed to persons living close to nuclear establishments, for use in case of accidents that could lead to releases of radioactive iodine.

• Iodine-123 and iodine-125 are used in medicine as tracers for imaging and evaluating the function of the thyroid.
• Noncombined (elemental) iodine is mildly toxic to all living things.

The artificial radioisotope ¹³¹I (a beta emitter), has a half-life of 8.0207 days. Also known as radioiodine, ¹³¹I has been used in treating cancer and other pathologies of the thyroid glands. ¹²³I is the radioisotope most often used in nuclear imaging of the kidney and thyroid as well as thyroid uptake scans (used for the evaluation of Graves' Disease). The most common compounds of iodine are the iodides of sodium (NaI) and potassium (KI) and the iodates (KIO₃).

Potassium iodide (KI tablets, or "SSKI" = "Saturated Solution of KI" liquid drops) can be given to people in a nuclear disaster area when fission has taken place, to block the uptake of iodine-131 by the thyroid. The protective effect of KI lasts approximately 24 hours, so it should be dosed daily until a risk of significant exposure to radioiodines no longer exists. The exposure can be reduced by evacuation, sheltering, and by control of the food supply. Iodine-131 also decays rapidly, with a half-life of 8 days, so that 99.95% of the original radioiodine is gone after three months.

Iodine-129 (¹²⁹I; half-life 15.7 million years) is a product of cosmic ray spallation on various isotopes of xenon in the atmosphere, in cosmic ray muon interaction with tellurium-130, and also uranium and plutonium fission, both in subsurface rocks and nuclear reactors. Nuclear processes, in particular nuclear fuel reprocessing and atmospheric nuclear weapons tests have now swamped the natural signal for this isotope. ¹²⁹I was used in rainwater studies following the Chernobyl accident. It also has been used as a groundwater tracer and as an indicator of nuclear waste dispersion into the natural environment.

Radioiodine and the kidney

In the 1970s imaging techniques were developed in California to utilize radioiodine in diagnostics for renal hypertension.

Iodine and breast cancer

It is known that a diet lacking in iodine is connected with adverse health effects collectively referred as iodine deficiency diseases or disorders. Studies also indicate that iodine deficiency, either dietary or pharmacologic, can lead to breast atypia and increased incidence of malignancy in animal models, while iodine treatment can reverse dysplasia. Laboratory evidences demonstrate that the effect of iodine on breast cancer is in part independent of thyroid function and that iodine inhibit cancer promotion through modulation of the estrogen pathway. Gene array profiling of estrogen responsive breast cancer cell line shows that the combination of iodine and iodide alters gene expression and inhibits the estrogen response through up-regulating proteins involved in estrogen metabolism. This suggests that iodine/iodide may be useful as an important adjuvant therapy in the pharmacologic manipulation of the estrogen pathway in women with breast cancer.

Precautions

Direct contact with skin can cause lesions, so it should be handled with care. Iodine vapor is very irritating to the eye and to mucous membranes. Concentration of iodine in the air should not exceed 1 mg/m³ (eight-hour time-weighted average). When mixed with ammonia, it can form nitrogen triiodide which is extremely sensitive and can explode unexpectedly.

Toxicity

Excess iodine has symptoms similar to those of iodine deficiency. Commonly encountered symptoms are abnormal growth of the thyroid gland and disorders in functioning and growth of the organism as a whole. Elemental iodine, I₂, is a deadly poison if taken in larger amounts; if 2–3 grams of it are consumed, it is fatal to humans. Iodides are similar in toxicity to bromides.

Compounds

See also: Category:Iodine compounds

Iodine forms many compounds. Potassium iodide is the most commercially significant iodine compound. It is a convenient source of the iodide anion; it is easier to handle than sodium iodide because it is not hydroscopic. Sodium iodide is especially useful in the Finkelstein reaction, because it is soluble in acetone, while potassium iodide is poorly so. In this reaction, an alkyl chloride is converted to an alkyl iodide. This relies on the insolubility of sodium chloride in acetone to drive the reaction:

R-Cl (acetone) + NaI (acetone) → R-I (acetone) + NaCl (s)

Iodic acid (HIO₃) and its salts are strong oxidizers. Periodic acid (HIO₄) cleaves vicinal diols along the C-C bond to give aldehyde fragments. 2-Iodoxybenzoic acid and Dess-Martin periodinane are hypervalent iodine oxidants used to specifically oxidize alcohols to ketones or aldehydes. Iodine pentoxide is a strong oxidant as well.

Interhalogen compounds are well known; examples include iodine monochloride and trichloride; iodine pentafluoride and heptafluoride.

Organic compounds

See also: Organoiodine compound

Several organoiodine compounds exist, the simplest is iodomethane, approved as an soil fumigant. Alkyl iodides such as iodomethane are good alkylating agents; aryl and alkyl iodides also form Grignard reagents. Some drawbacks are:

• iodine compounds tend to be more expensive than the corresponding bromides and chlorides, in that order
• iodides tend to be much stronger alkylating agents, and so are more toxic (e.g. methyl iodide is very toxic (T+)
• low molecular weight iodides tend to have a much higher equivalent weight, compared with other alkylating agents (e.g. methyl iodide versus dimethyl carbonate), due to the atomic mass of iodine.

Biologigal active substances like the thyroid hormones are naturally occurring organoiodine compounds.

See also

• Iodide as an antioxidant
• Chemical Oxygen Iodine Laser
• Nutrition facts label

External links

• "Micronutrient Research for Optimum Health", Linus Pauling Institute, OSU Oregon State University
• ATSDR - CSEM: Radiation Exposure from Iodine 131 U.S. Department of Health and Human Services (public domain)
• ChemicalElements.com - Iodine
• who.int, WHO Global Database on Iodine Deficiency
• Network for Sustained Elimination of Iodine Deficiency
• Oxidizing Agents > Iodine
• Los Alamos National Laboratory: Iodine
• WebElements.com – Iodine