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==Occurrence==
==Occurrence==
[[Image:Iodomethane-3D-vdW.png|thumb|left|150px|Iodomethane]]
[[Image:Iodomethane-3D-vdW.png|thumb|left|150px|[[Iodomethane]]]]
Iodine naturally occurs in the environment chiefly as a dissolved [[iodide]] in [[seawater]], although it is also found in some minerals and soils.<ref name="Dissanayake">{{cite journal|last = Dissanayake|first = C. B.|coauthors = Chandrajith, Rohana; Tobschall, H. J.|title = The iodine cycle in the tropical environment — implications on iodine deficiency disorders|journal = International Journal of Environmental Studies|volume = 56 |page= 357| doi = 10.1080/00207239908711210|year = 1999}}</ref> This element also exists in small amounts in the mineral [[caliche (mineral)|caliche]], found in [[Chile]], between the [[Andes]] and the sea. A type of seaweed, [[kelp]], tends to be high in iodine as well.
Iodine naturally occurs in the environment chiefly as a dissolved [[iodide]] in [[seawater]], although it is also found in some minerals and soils.<ref name="Dissanayake">{{cite journal|last = Dissanayake|first = C. B.|coauthors = Chandrajith, Rohana; Tobschall, H. J.|title = The iodine cycle in the tropical environment — implications on iodine deficiency disorders|journal = International Journal of Environmental Studies|volume = 56 |page= 357| doi = 10.1080/00207239908711210|year = 1999}}</ref> This element also exists in small amounts in the mineral [[caliche (mineral)|caliche]], found in [[Chile]], between the [[Andes]] and the sea. A type of seaweed, [[kelp]], tends to be high in iodine as well.


Revision as of 01:40, 19 February 2010

For other uses, see Iodine (disambiguation).
Iodine, 53I
Iodine
Pronunciation/ˈaɪədaɪn, -dɪn, -diːn/ ​(EYE-ə-dyne, -⁠din, -⁠deen)
Appearancelustrous metallic gray solid, black/violet liquid, violet gas
Standard atomic weight Ar°(I)
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
Groupgroup 17 (halogens)
Periodperiod 5
Block  p-block
Electron configuration[Kr] 4d10 5s2 5p5
Electrons per shell2, 8, 18, 18, 7
Physical properties
Phase at STPsolid
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[3]
Triple point386.65 K, ​12.1 kPa
Critical point819 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 statescommon: −1, +1, +3, +5, +7
+2,[4] +4,? +6?
ElectronegativityPauling scale: 2.66
Ionization energies
  • 1st: 1008.4 kJ/mol
  • 2nd: 1845.9 kJ/mol
  • 3rd: 3180 kJ/mol
Atomic radiusempirical: 140 pm
Covalent radius139±3 pm
Van der Waals radius198 pm
Color lines in a spectral range
Spectral lines of iodine
Other properties
Natural occurrenceprimordial
Crystal structure ​base-centered orthorhombic (oS8)
Lattice constants
Base-centered orthorhombic crystal structure for iodine
a = 725.79 pm
b = 478.28 pm
c = 982.38 pm (at 20 °C)[3]
Thermal expansion74.9×10−6/K (at 20 °C)[a]
Thermal conductivity0.449 W/(m⋅K)
Electrical resistivity1.3×107 Ω⋅m (at 0 °C)
Magnetic orderingdiamagnetic[5]
Molar magnetic susceptibility−88.7×10−6 cm3/mol (298 K)[6]
Bulk modulus7.7 GPa
CAS Number7553-56-2
History
Discovery and first isolationBernard 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
 Category: Iodine
| references

Iodine (Template:Pron-en EYE-o-dyne, /ˈaɪ.ɵdɨn/ EYE-o-dən, or in chemistry /ˈaɪ.ɵdiːn/ EYE-o-deen; from [ιώδης iodes] Error: {{Lang-xx}}: text has italic markup (help) meaning violet (or purple), is a chemical element that has the symbol I and the atomic number 53.

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 17 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. Iodine is rare in the solar system and Earth's crust; however, the iodides are very soluble in water and the element concentrates in seawater, where it occurs in far higher amounts than in rocks. 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[7][8]).

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.

Round bottom flasked filled with violet iodine vapour
In the gas phase iodine shows its violet color.

Elemental iodine dissolves easily in most organic solvents such as hexane or chloroform due to its lack of polarity, but is only slightly soluble in water. However, the solubility of elemental iodine in water can be increased by the addition of potassium iodide. The molecular iodine reacts reversibly with the negative ion, generating the triiodide anion I3− in equilibrium, which is soluble in water. This is also the formulation of some types of medicinal (antiseptic) iodine, although tincture of iodine classically dissolves the element in aqueous ethanol.

Solutions of elemental iodine have the unique property of exhibiting dramatically different colors depending on the polarity of the solvent. When dissolved in nonpolar solvents like hexane, the solution appears deep violet; in moderately polar dichloromethane the solution is dark crimson, and in strongly polar solvents like acetone or ethanol, it appears dark orange or brown. This is due to ligand field interactions of solvent molecules with the d-orbitals of iodine, which is the only halogen with a sufficiently occupied electronic configuration to allow such interactions. This same property allows the formation of hypervalent iodine compounds, which have expanded bonding orbitals beyond the generally allowed octet rule.

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.[9]

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.[10] 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 activity in rice paddies and by the burning of biological material is estimated to be 214 kilotonnes.[11] The volatile iodomethane is broken up by oxidation reactions in the atmosphere and a global iodine cycle is established.[10][11] 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.

Structure

Structure of solid iodine

Iodine crystallizes in the orthorhombic 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.

Production

From the several places in which iodine occurs 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.[12] The high concentration of iodine in the caliche and the extensive mining made Chile the largest producer of iodine in 2007.

File:Iodine.PNG
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.[12][13]

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 it 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 other methods of isolating this element in the laboratory, for example the method used to isolate other halogens: oxidation of the iodide in hydrogen iodide (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 known (characterized) 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 half-life 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.[14][15] 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.[16][17][18] 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.[19] Davy sent a letter dated December 10 to the Royal Society of London stating that he had identified a new element.[20] 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 and water treatment

Elemental iodine is used as a disinfectant in various forms. The iodine exists as the element, or as the water soluble triiodide anion I3- generated in situ by adding iodide to poorly-soluble iodine (the reverse chemical reaction makes some free elemental iodine available for antisepsis). Alternatively, iodine may come from iodophors, which contain iodine complexed with a solubilizing agent (iodide ion may be thought of loosely as the iodophor in triiodide water solutions). Examples of such preparations include:[21]

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 the glucose polymers starch and glycogen. Many applications rely on this property:

  • Iodometry. The concentration of an oxidant can be determined by adding it to an excess of iodide with a little free iodine, to destroy elemental iodine/triiodide as a result of oxidation by the oxidant. A starch indicator is then used as the indicator close to the end-point, in order to increase the visual contrast (dark blue becomes colorless, instead of the yellow of dilute triiodide becoming colorless).
  • An Iodine test may be used to test a sample substance for the presence of starch.
  • The Iodine clock reaction is an extension of the techniques in iodometry.
  • 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.[22]
  • During colposcopy, Lugol's iodine is applied to the vagina and cervix. Normal vaginal tissue stains brown due to its high glycogen content (a color-reaction similar to that with starch), while abnormal tissue suspicious for cancer does not stain, and thus appears pale compared to the surrounding tissue. Biopsy of suspicious tissue can then be performed. This is called a Schiller's Test.

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.

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.

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.[23]

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)[24]

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).[25] Aqueous and ethanol solutions are brown. Solutions in chloroform, carbon tetrachloride, and carbon disulfide are violet.

Molecular iodine can be prepared by oxidizing iodides with chlorine:

2 I− + Cl2 → I2 + 2 Cl−

or with manganese dioxide in acid solution:[24]

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

Iodine is reduced to hydroiodic acid by hydrogen sulfide:[26]

I2 + H2S → 2 HI + S↓

or by hydrazine:

2 I2 + N2H4 → 4 HI + N2

Iodine is oxidized to iodate by nitric acid:[27]

I2 + 10 HNO3 → 2 HIO3 + 10 NO2 + 4 H2O

or by chlorates:[27]

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):[24]

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.

When dissolved in fuming sulfuric acid (65% oleum), iodine forms an intense blue solution. This has been shown to be due to the formation of the I+
2 cation, the result of iodine being oxidised by SO
3:[28]

2 I
2 + 2 SO
3 + H
2SO
4 → 2 I+
2 + SO
2 + 2 HSO−
4

The I+
2 cation is also formed in the oxidation of iodine by SbF
5 or TaF
5. The resulting I+
2Sb
2F−
11 or I+
2Ta
2F−
11 can be isolated as deep blue crystals. The solutions of these salts turn red when cooled below −60 °C, due to the formation of the I2+
4 cation:[28]

2 I+
2 ⇌ I2+
4

Under slightly more alkaline conditions, I2+
4 disproportionates into I+
3 and an iodine(III) compound. Excess iodine can then react with I+
3 to form I+
5 (green) and I3+
15 (black).[28]

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.[29] 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+)[30]
  • 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.

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.[31][32]

Biological role

Main article: Iodine in biology

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.[citation needed] 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%.[citation needed] T4 acts largely as a precursor to T3, which is (with minor exceptions) the biologically active hormone.

Extrathyroidal iodine

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 directly 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.[33] It has been shown to act as an antioxidant in these tissues.[33]

Iodine may have a relationship with selenium, and iodine supplementation in selenium-deficient populations may pose risks for thyroid function.[33]

The US Food and Nutrition Board and Institute of Medicine recommended daily allowance of iodine ranges from 150 micrograms /day for adult humans to 290 micrograms /day for lactating mothers. However, the thyroid gland needs no more than 70 micrograms /day to synthesize the requisite daily amounts of T4 and T3. These higher recommended daily allowance levels of iodine seem necessary for optimal function of a number of body systems, including lactating breast, gastric mucosa, salivary glands, oral mucosa, thymus, epidermis, choroid plexus, etc.[34][35][36]

Iodine and the development of cancer

*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.[37][38][39] 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. Whether iodine/iodide will be useful as an adjuvant therapy in the pharmacologic manipulation of the estrogen pathway in women with breast cancer has not been determined clinically.[37]

*Iodine and stomach cancer. Some researchers have found an epidemiologic correlation between iodine deficiency, iodine-deficient goitre and gastric cancer;[40] [41][42] a decrease of the incidence of death rate from stomach cancer after implementation of the effective iodine-prophylaxis was reported too.[43] The proposed mechanism of action is that iodide ion can function in gastric mucosa as an antioxidant reducing species that can detoxify poisonous reactive oxygen species, such as hydrogen peroxide.

Iodine and immunity

Iodine has important actions in the immune system. The high iodide-concentration of thymus suggests an anatomical rationale for this role of iodine in immune system.[44][45][46][47][48][49]

Iodine in salivary glands and oral health

The trophic, antioxidant and apoptosis-inductor actions and the presumed anti-tumour activity of iodides might also be important for prevention of oral and salivary glands diseases.[50][51][52][53][54][55]

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).[56] The tolerable upper limit was assessed by analyzing the effect of supplementation on thyroid-stimulating hormone.[33]

Natural sources of iodine include sea life, such as kelp and certain seafood, as well as plants grown on iodine-rich soil.[57][58] Iodized salt is fortified with iodine.[58]

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.[56] In Japan, consumption is much higher due to the frequent consumption of seaweed or kombu kelp.[33]

After iodine fortification programs (e.g. iodized salt) have been implemented, some cases of iodine-induced hyperthyroidism have been observed (so called Jod-Basedow disease). 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.[59]

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.[60]

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.[61] Iodine deficiency is also a problem in certain areas of Europe. In Germany it has been estimated to cause a billion dollars in health care costs per year.[33]

Radioiodine in biology

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 as I131 can be used for thyroid ablation, a procedure in which radioactive iodine is administered intravenously or orally 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 performed 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.

Potassium iodide has been distributed to populations exposed to nuclear fission accidents such as the Chernobyl disaster. The iodide solution SSKI, a saturated solution of potassium (K) iodide in water, has been used to block absorption of the radioiodine (it has no effect on other radioisotopes from fission). Tablets containing potassium iodide are now also manufactured and stocked in central disaster sites by the governments for this purpose. In theory, many harmful late-cancer effects of nuclear fallout might be prevented in this way, since an excess of thyroid cancers, presumably due to radioiodine uptake, is the only proven radioisotope contamination effect after a fission accident, or from contamination by fallout from an atomic bomb (prompt radiation from the bomb also cases other cancers, such as leukemias, directly). Taking large amounts of thyroid saturates iodide receptors prevents uptake of most radioactive iodine-131 that may be present from fission product exposure (although it does not protect from other radioisotopes, nor from any other form of direct radiation). The protective effect of KI lasts approximately 24 hours, so must be dosed daily until a risk of significant exposure to radioiodines from fission products no longer exists.[62][63] Iodine-131 (the most common radioiodine contaminant in fallout) also decays relatively rapidly with a half-life of 8 days, so that 99.95% of the original radioiodine is gone after three months.

Iodine 125

Iodine-125 is also commonly used by radiation oncologists in low dose rate brachytherapy in the treatment of cancer at sites other than the thyroid, especially in prostate cancer. The radioiodine is encapsulated in titanium seeds and implanted in the area of tumor involvement. In contrast to the blood-borne dissemination of radioiodine used in the thyroid, the radioiodine acts only locally in the area where it is implanted.

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. Artificial nuclear processes, in particular nuclear fuel reprocessing and atmospheric nuclear weapons tests, have now swamped the natural signal for this isotope. Nevertheless, it now serves as a groundwater tracer as indicator of nuclear waste dispersion into the natural environment. In a similar fashion, 129I was used in rainwater studies to track fission products following the Chernobyl disaster.

Radioiodine and the kidney

In the 1970s imaging techniques were developed to employ radioiodine in diagnostics for renal hypertension; however methods using other chemical compounds, such as DMSA, are more popular in clinics nowadays.

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 is a lethal dose for an adult human.

Iodine vapor is very irritating to the eye, to mucous membranes, and in 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.

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]

Iodine sensitivity

Some people develop a sensitivity to iodine. Application of tincture of iodine can cause a rash. Some cases of reaction to Povidone-iodine (Betadine) have been documented to be a chemical burn.[64] Eating iodine-containing foods can cause hives. Medical use of iodine (i.e. as a contrast agent, see above) can cause anaphylactic shock in highly iodine sensitive patients. Some cases of sensitivity to iodine can be formally classified as iodine allergies. Iodine sensitivity is rare but has a considerable effect given the extremely widespread use of iodine-based contrast media[65].

See also

References

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