Iodine

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For other uses, see Iodine (disambiguation).
53 tellurium ← iodine → xenon
Br
↑
I
↓
At
General
Name, Symbol, Number iodine, I, 53
Element category halogens
Group, Period, Block 17, 5, p
Appearance violet-dark gray, lustrous
Standard atomic weight 126.90447(3)  g·mol−1
Electron configuration [Kr] 4d10 5s2 5p5
Electrons per shell 2, 8, 18, 18, 7
Physical properties
Phase solid
Density (near r.t.) 4.933  g·cm−3
Melting point 386.85 K
(113.7 °C, 236.66 °F)
Boiling point 457.4 K
(184.3 °C, 363.7 °F)
Triple point 386.65 K, 12.1×103 Pa
Critical point 819 K, 11.7 MPa
Heat of fusion (I2) 15.52  kJ·mol−1
Heat of vaporization (I2) 41.57  kJ·mol−1
Specific heat capacity (25 °C) (I2) 54.44  J·mol−1·K−1
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
Crystal structure orthorhombic
Oxidation states 7, 5, 3, 1, -1
(strongly acidic oxide)
Electronegativity 2.66 (Pauling scale)
Ionization energies 1st: 1008.4 kJ/mol
2nd: 1845.9 kJ/mol
3rd: 3180 kJ/mol
Atomic radius 140  pm
Covalent radius 139±3  pm
Van der Waals radius 198 pm
Miscellaneous
Magnetic ordering diamagnetic[1]
Electrical resistivity (0 °C) 1.3×107 Ω·m
Thermal conductivity (300 K) 0.449  W·m−1·K−1
Bulk modulus 7.7  GPa
CAS registry number 7553-56-2
Most-stable isotopes
Main article: Isotopes of iodine
iso NA half-life DM DE (MeV) DP
123I syn 13 h ε, γ 0.16 123Te
127I 100% 127I is stable with 74 neutrons
129I syn 15.7×106 y β− 0.194 129Xe
131I syn 8.02070 d β−, γ 0.971 131Xe
References

Iodine (pronounced /ˈaɪ.ədaɪn/, /ˈaɪ.ədɨn/, or in chemistry /ˈaɪ.ədiːn/; from Greek: ιώδης iodes "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 second least reactive of the halogens, and the second most electropositive halogen; trailing behind astatine in both of these categories. 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 the heaviest element commonly used by living organisms (only tungsten, used in enzymes by a few bacteria, is heavier).

Contents

[edit] Characteristics

Iodine under standard conditions is a shiny grey 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.

When iodine is encapsulated into carbon nanotubes it forms atomic chains, whose structure depends on the nanotube diameter.[2]

[edit] 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.[3] 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 compounds are produced by marine life forms, the most notable being iodomethane (commonly called methyl iodide). The total iodomethane that is produced by the marine environment, by microbial activitiy in rice paddies and by the burning of biological material is estimated to be 214 kilotonnes.[4] The volatile iodomethane is broken up by oxidation reactions in the atmosphere and a global iodine cycle is established.[3][4] 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.

[edit] Structure

Structure of solid iodine

Iodine crystallizes in the orthorombic space group Cmca No 64, Pearson symbol oS8, the same as black phosphorus. In the solid state, I2 molecules are still represented by a short I-I bond of 270 pm.

[edit] Production

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.[5] The high concentration of iodine in the caliche and the extensive mining made Chile the largest producer of iodine in 2007.

Iodine output in 2005

Most other producers use natural occurring brine for the production of iodine. The Japanese Minami Kanto gas field east of Tokyo 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 of the source. 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 is dilute and must 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.[5][6]

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

The production of iodine from seawater via electrolysis is not used due to the sufficient abundance of iodine-rich brine. Another source of iodine was kelp, used in the 18th and 19th centuries, but 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 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).

[edit] 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.

[edit] History

Iodine was discovered by Bernard Courtois in 1811.[7][8] 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’s 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.[9][10][11] 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.[12] Davy sent a letter dated December 10 to the Royal Society of London stating that he had identified a new element.[13] 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.

[edit] Applications

[edit] 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:[14]

[edit] 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 indicator 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 test 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.[15]

[edit] 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

[edit] 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.

[edit] Iodine 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 hygroscopic. 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.

HI He
LiI BeI2 BI3 CI4 NI3 I2O4, I2O5, I4O9 IF, IF3, IF5, IF7 Ne
NaI MgI2 AlI3 SiI4 PI3, P2I4 S ICl, ICl3 Ar
KI CaI2 Sc TiI4 VI3 Cr MnI2 Fe CoI2 NiI2 CuI ZnI2 Ga2I6 GeI2, GeI4 AsI3 Se IBr Kr
RbI SrI2 Y ZrI4 Nb Mo Tc Ru Rh Pd AgI CdI2 InI3 SnI4, SnI2 SbI3 TeI4 I Xe
CsI BaI2 Hf Ta W Re Os Ir Pt AuI Hg2I2, HgI2 TlI PbI2 Bi Po At Rn
Fr Ra Rf Db Sg Bh Hs Mt Ds Rg Uub Uut Uuq Uup Uuh Uus Uuo
↓
La Ce Pr Nd Pm SmI2 Eu Gd TbI3 Dy Ho Er Tm Yb Lu
Ac ThI4 Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

[edit] Organic compounds

See also: Organoiodine compound

Many organoiodine compounds exist, the simplest is iodomethane, approved as a soil fumigant. Iodinated organics are used as synthetic reagents, and also radiocontrast agents.

Biologically active substances like the thyroid hormones are naturally occurring organoiodine compounds.[16]

[edit] 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 is in equilibrium with H+ + I− + HIO   (K = 2.0×10−13)[17]

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).[18] 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:

2 I− + Cl2 → I2 + 2Cl−

or with manganese dioxide in acid solution:[17]

2 I− + 4 H+ + MnO2 → I2 + 2 H2O + Mn2+

Iodine is reduced to hydroiodic acid by hydrogen sulfide:[19]

I2 + H2S → 2 HI + S↓

or by hydrazine:

2 I2 + N2H4 → 4 HI + N2

Iodine is oxidized to iodate by nitric acid:[20]

I2 + 10 HNO3 → 2 HIO3 + 10 NO2 + 4H2O

or by chlorates:[20]

I2 + 2 ClO3− → 2 IO3− + Cl2

Iodine is converted in a two stage reaction to iodide and iodate in solutions of alkali hydroxides (such as sodium hydroxide):[17]

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

Despite having the lowest electronegativity of the common halogens, iodine reacts violently with some metals, such as aluminum:

3 I2 + 2 Al → 2 AlI3

This reaction produces 314 kJ per mole of aluminum, comparable to thermite's 425 kJ. Yet the reaction initiates spontaneously, and if unconfined, causes a cloud of gaseous iodine due to the high heat.

[edit] 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.[21] The iodinating reagent is phosphorus triiodide that is formed in situ:

3 CH3OH + PI3 → 3 CH3I + H3PO3

Phosphorous acid is formed as a side-product.

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 sometimes used to activate magnesium when preparing Grignard reagents; aryl and alkyl iodides both form Grignard reagents. Alkyl iodides such as iodomethane are good alkylating agents. Some drawbacks to use of iodo-organics in chemical synthesis 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+)[22]
  • 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.

[edit] Clandestine synthetic chemical 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.[23][24]

[edit] Biological role

Iodine is an essential trace element for life, the heaviest element commonly needed by living organisms, and the second-heaviest known to be used by any form of life (only tungsten, a component of a few bacterial enzymes, has a higher atomic number and atomic weight). Iodine's 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. In the cells of these tissues iodide enters direcly by sodium-iodide symporter (NIS). Its role in mammary tissue is related to fetal and neonatal development, but its role in the other tissues is unknown.[25] It has been shown to act as an antioxidant in these tissues.[25]

Iodine may have a relationship with selenium, and iodine supplementation in selenium-deficient populations may pose risks for thyroid function.[25]

[edit] 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.[26][27][28] 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.[26]

[edit] Iodine and immunity

Iodine has an important action on the immune system. The high iodide-concentration of thymus gives the anatomical rationale for this role of iodine in immune system.[29][30][31][32][33]

[edit] 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).[34] The tolerable upper limit was assessed by analyzing the effect of supplementation on thyroid-stimulating hormone.[25]

Natural sources of iodine include sea life, such as kelp and certain seafood, as well as plants grown on iodine-rich soil.[35][36] Iodized salt is fortified with iodine.[36]

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.[34] In Japan, consumption is much higher due to the frequent consumption of seaweed or kombu kelp.[25] 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.[25]

After iodine fortification programs (e.g. iodized salt) have been implemented, iodine-induced hyperthyroidism has been observed. The condition mainly seems to occur in people over forty, and the risk appears higher when iodine deficiency is severe and the initial rise in iodine intake is high.[37]

[edit] 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.[38]

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.[39] 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.[25]

[edit] Radioiodine in biology

[edit] Radioiodine and the thyroid

The most common compounds of iodine are the iodides of sodium (NaI) and potassium (KI) and the iodates (KIO3), as elemental iodine is mildly toxic to all living things. Normal iodine is an essential precursor for the manufacture of thyroid hormone.

Due to preferential uptake of iodine by the thyroid, isotopes with short half lives such was I131 can be used for thyroid ablation, a procedure in which radioactive iodine is administed intravenously following a diagnostic scan. This procedure is generally performed on patients with thyroid cancer or hyperfunctioning thyroid tissue. After uptake, the iodine undergoes degeneration via beta decay, destroying its associated thyroid tissue. Normally thyroidectomy is preformed prior to ablation to avoid side effects of epilation and radiation toxicity. The purpose of radioablation is to destroy remnant tissue that was unable to be removed with surgery.

Lower energy isotopes such as iodine-123, and less commonly iodine-125, are used as tracers to evaluate the anatomic and physiologic function of the thyroid. Abnormal results may be caused by disorders such as Graves' Disease or Hashimoto's thyroiditis.

Normal iodine tablets are sometimes distributed to population near medical establishments or use in case of accidental exposure in the form of potassium iodide tablets (KI or SSKI, saturated solution of potassium iodide drops). Taking large amounts of thyroid saturates its receptors and would prevent uptake of any radioactive iodine-131 that may be present from radiation exposure. 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.[40][41] 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.

[edit] Iodine 129

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.

[edit] Radioiodine and the kidney

In the 1970s imaging techniques were developed in California to utilize radioiodine in diagnostics for renal hypertension, although this is clinically no commonly performed today and has been placaded by other chemical compounds such as DMSA.

[edit] Precautions and toxicity of elemental iodine

Elemental iodine is an oxidizing irritant and direct contact with skin can cause lesions, so iodine crystals should be handled with care. Solutions with high elemental iodine concentration such as tincture of iodine are capable of causing tissue damage if use for cleaning and antisepsis is prolonged.

Elemental iodine (I2) is poisonous if taken orally in larger amounts; 2–3 grams of it are a lethal dose for an adult human.

Iodine vapor is very irritating to the eye, to mucous membranes, andin the respiratory tract. Concentration of iodine in the air should not exceed 1 mg/m³ (eight-hour time-weighted average).

When mixed with ammonia and water, elemental iodine forms nitrogen triiodide which is extremely shock sensitive and can explode unexpectedly.

[edit] Toxicity of iodide ion

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. Iodides are similar in toxicity to bromides.[citation needed]

[edit] See also

[edit] References

  1. ^ Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics 81th edition, CRC press.
  2. ^ Guan, L; Suenaga, K; Shi, Z; Gu, Z; Iijima, S (Jun 2007). "Polymorphic structures of iodine and their phase transition in confined nanospace.". Nano letters 7 (6): 1532–5. doi:10.1021/nl070313t. ISSN 1530-6984. PMID 17477579. 
  3. ^ a b Dissanayake, C. B.; Chandrajith, Rohana; Tobschall, H. J. (1999). "The iodine cycle in the tropical environment — implications on iodine deficiency disorders". International Journal of Environmental Studies 56 (3): 357–372. doi:10.1080/00207239908711210. 
  4. ^ a b N. Bell, L. Hsu, D. J. Jacob, M. G. Schultz, D. R. Blake, J. H. Butler, D. B. King, J. M. Lobert, and E. Maier-Reimer (2002). "Methyl iodide: Atmospheric budget and use as a tracer of marine convection in global models". Journal of GeophysicalResearch 107 (D17): 4340. doi:10.1029/2001JD001151. 
  5. ^ a b Jessica Elzea Kogel, Nikhil C. Trivedi, James M. Barker, Stanley T. Krukowski (2006). Industrial Minerals & Rocks: Commodities, Markets, and Uses. SME. pp. 541–552. ISBN 9780873352338. http://www.google.com/books?id=zNicdkuulE4C. 
  6. ^ Tatsuo Maekawa, Shun-Ichiro Igari and Nobuyuki Kaneko (2006). "Chemical and isotopic compositions of brines from dissolved-in-water type natural gas fields in Chiba, Japan". Geochemical Journal 40: 475–484. doi:10.2343/geochemj.40.475. 
  7. ^ Bernard Courtois (1813). "Découverte d'une substance nouvelle dans le Vareck". Annales de chimie 88: 304–10.  In French, seaweed that had been washed onto the shore was called "varec", "varech", or "vareck", whence the English word "wrack". Later, "varec" also referred to the ashes of such seaweed: the ashes were used as a source of iodine and salts of sodium and potassium.
  8. ^ Patricia A. Swain (2005). "Bernard Courtois (1777-1838) famed for discovering iodine (1811), and his life in Paris from 1798". Bulletin for the History of Chemistry 30 (2): 103–11. http://www.scs.uiuc.edu/~mainzv/HIST/awards/OPA%20Papers/2007-Swain.pdf. 
  9. ^ J. Gay-Lussac (1813). "Sur un nouvel acide formé avec la substance décourverte par M. Courtois". Annales de chimie 88: 311–8. 
  10. ^ J. Gay-Lussac (1813). "Sur la combination de l'iode avec d'oxigène". Annales de chimie 88: 319–21. 
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Diatomic chemical elements

Hydrogen H2 | Nitrogen N2 | Oxygen O2 | Fluorine F2 | Chlorine Cl2 | Bromine Br2 | Iodine I2 | Astatine At2 |
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