|Name, symbol||iodine, I|
|Pronunciation||//, //, or //
EYE-ə-dyn, EYE-ə-dən, or EYE-ə-deen
|Appearance||lustrous metallic gray, violet as a gas|
|Iodine in the periodic table|
|Atomic number (Z)||53|
|Group, block||group 17 (halogens), p-block|
|Element category||diatomic nonmetal|
|Standard atomic weight (±) (Ar)||126.90447(3)|
|Electron configuration||[Kr] 4d10 5s2 5p5|
|2, 8, 18, 18, 7|
|Melting point||386.85 K (113.7 °C, 236.66 °F)|
|Boiling point||457.4 K (184.3 °C, 363.7 °F)|
|Density near r.t.||4.933 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)
|Oxidation states||7, 6, 5, 4, 3, 1, −1 (a strongly acidic oxide)|
|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|
|Thermal conductivity||0.449 W/(m·K)|
|Electrical resistivity||1.3×107 Ω·m (at 0 °C)|
|Bulk modulus||7.7 GPa|
|Discovery and first isolation||Bernard Courtois (1811)|
|Most stable isotopes of iodine|
|Decay modes in parentheses are predicted, but have not yet been observed|
Iodine and its compounds are primarily used in nutrition, and industrially in the production of acetic acid and certain polymers. Iodine's relatively high atomic number, low toxicity, and ease of attachment to organic compounds have made iodine radioisotopes, such as 131I, a part of many X-ray contrast materials in modern medicine. Iodine has only one stable isotope.
Iodine is found on Earth mainly as the highly water-soluble iodide ion I−, which is concentrated in oceans and brine pools. Like the other halogens, free iodine occurs mainly as a diatomic molecule I2, and then only momentarily after being oxidized from iodide by an oxidant like free oxygen. In the universe and on Earth, iodine's high atomic number makes it a relatively rare element. Present in sea water, it is the heaviest essential element used widely by life in biological functions (only tungsten, employed in enzymes by a few species of bacteria, is heavier). Iodine is rare in many soils, has low abundance generally as a crust-element, and is leached by rainwater, leading to many deficiency problems in land animals and inland human populations. Iodine deficiency affects about two billion people and is the leading preventable cause of intellectual disabilities.
Iodine is required by higher animals as a component of thyroid hormones. Because of this function, radioisotopes of iodine are concentrated in the thyroid gland along with nonradioactive iodine. If inhaled, the radioisotope iodine-131, which has a high fission product yield, concentrates in the thyroid, and can be prevented with non-radioactive potassium iodide treatment if taken before exposure.
- 1 Characteristics
- 2 Chemistry
- 3 Production
- 4 History
- 5 Applications
- 6 Biological role
- 7 See also
- 8 References
- 9 External links
Under standard conditions, iodine is a bluish-black solid with a metallic lustre, appearing to sublimate into a noxious violet-pink gas. The colour is due to absorption of visible light by electronic transitions between the highest occupied and lowest unoccupied molecular orbitals. Melting at 113.7 °C (236.7 °F), it forms compounds with many elements but is less reactive than other halogens, and has some metallic light reflectance.
Elemental iodine is slightly soluble in water, with one gram dissolving in 3450 ml at 20 °C and 1280 ml at 50 °C; potassium iodide may be added to increase solubility via formation of triiodide ions. Nonpolar solvents such as hexane and carbon tetrachloride improve solubility. Polar solutions are brown, reflecting the role of these solvents as Lewis bases, while nonpolar solutions are violet, the color of iodine vapor. Charge-transfer complexes form when iodine is dissolved in polar solvents, modifying the energy distribution of iodine's molecular orbitals, hence changing the colour. A metal ion may replace the solvent, in which case the two species exchange electrons, the ion undergoing π backbonding.
Structure and bonding
Iodine normally exists as a diatomic molecule with an I-I bond length of 270 pm, one of the longest single bonds known. The I2 molecules tend to interact via the weak van der Waals forces called the London dispersion forces, and this interaction is responsible for a melting point that is higher than more compact halogens, which are also diatomic. Since the atomic size of iodine is larger, its melting point is higher. The solid crystallizes as orthorhombic crystals. The crystal motif in the Hermann–Mauguin notation is Cmca (No 64), Pearson symbol oS8. The I-I bond is relatively weak, with a bond dissociation energy of 36 kcal/mol, and most iodine bonds are weaker than those of lighter halides. One consequence of this weak bonding is that I2 molecules have a relatively high tendency to dissociate into atomic iodine.
Of the 37 known (characterized) isotopes of iodine, only one, iodine-127, is stable.
The longest-lived radioisotope, iodine-129, has a half-life of 15.7 million years. This is long enough to be a permanent fixture of the environment on human time scales, but far too short for it to survive as a primordial isotope. Instead, iodine-129 is an extinct radionuclide, and its presence in the early Solar System is inferred from the excess of the decay daughter, xenon-129. This nuclide is also a product cosmic rays and a byproduct of artificial nuclear fission, a long-lived bellwether of nuclear contamination in the environment.
The next-longest-lived radioisotope, iodine-125, has a half-life of 59 days. It is used as a convenient gamma-emitting tag for proteins in biological assays, and a few nuclear medicine imaging tests where a longer half-life is required. It is also commonly used in brachytherapy implanted capsules that kill tumors by local short-range gamma radiation (but where the isotope is never released into the body).
Iodine-131 (half-life 8 days), a beta-emitting isotope, is a common nuclear fission product. It is preferably administered to humans only in very high doses that destroy all tissues where it accumulates (usually the thyroid), which in turn prevents these tissues from developing cancer from a lower dose (paradoxically, a high dose of this isotope appears safer for the thyroid than a low dose). Like other radioiodines, I-131 accumulates in the thyroid gland, but unlike the others, in small amounts it is highly carcinogenic there, it seems, owing to the high local cell mutation from beta decay damage. Because of this tendency of 131I to cause high damage to cells that collect it and others near them (0.6 to 2 mm away, the range of the beta rays), it is the only iodine radioisotope used as direct therapy to kill tissues such as cancers that take up artificially iodinated molecules (example, the compound iobenguane, also known as MIBG). For the same reason, only the iodine isotope I-131 is used to treat Grave's disease and those thyroid cancers (sometimes in metastatic form) where the tissue that requires destruction still functions to accumulate the iodide.
Nonradioactive ordinary potassium iodide (iodine-127), in a number of convenient forms (tablets or solution) may be used to saturate the thyroid gland with iodine, and thereby protect it against accidental contamination from iodine-131 generated by nuclear fission accidents, such as the Chernobyl disaster, the Fukushima I nuclear accidents, and nuclear fallout from nuclear weapons.
Iodine is rare in the Solar System and Earth's crust (47–60th in abundance); however, iodide salts are often very soluble in water. Iodine occurs in slightly greater concentrations in seawater than in rocks, 0.05 vs. 0.04 ppm. Minerals containing iodine include caliche, found in Chile. The brown algae Laminaria and Fucus found in temperate zones of the Northern Hemisphere contain 0.028–1.0% iodine, dry weight. Aside from tungsten, iodine is the heaviest essential element in living organisms. About 19,000 tonnes are produced annually from natural sources.
Organoiodine compounds are produced by marine life forms, the most notable being iodomethane (commonly called methyl iodide). About 214 kilotonnes/year of iodomethane is produced by the marine environment, by microbial activity in rice paddies and by the burning of biological material. The volatile iodomethane is broken up in the atmosphere as part of a global iodine cycle.
Iodine adopts a variety of oxidation states, commonly ranging from (formally) I(VII) to I(-I), and including the intermediate states of I(V), I(III) and I(I). Practically, only the −1 oxidation state is of significance, being the form found in iodide salts and organoiodine compounds. Iodine is a Lewis acid. With electron donors such as triphenylphosphine and pyridine it forms a charge-transfer complex. With the iodide anion it forms the triiodide ion. Iodine and the iodide ion form a redox couple. I2 is easily reduced and I− is easily oxidized.
Commonly in Earth's atmosphere, iodides are slowly oxidized by atmospheric oxygen to produce free iodine. Evidence for this conversion is the yellow tint of certain aged samples of iodide salts and some organoiodine compounds. The oxidation of iodide to iodine in air is also responsible for the slow loss of iodide content in iodized salt if exposed to air. Some salts use iodate (IO−
3) to prevent the loss of iodine.
Iodine is easily reduced. Most common is the interconversion of I− and I2. Molecular iodine can be prepared by oxidizing iodides with chlorine:
- 2 I− + Cl2 → I2 + 2 Cl−
- 2 I− + 4 H+ + MnO2 → I2 + 2 H2O + Mn2+
- 8 I2 + 8 H2S → 16 HI + S8
- 2 I2 + N2H4 → 4 HI + N2
- 2 I
2 + 2 SO
3 + H
4 → 2 I+
2 + SO
2 + 2 HSO−
2 cation is also formed in the oxidation of iodine by SbF
5 or TaF
5. The resulting I+
11 or I+
11 can be isolated as deep blue crystals. The solutions of these salts turn red when cooled below −60 °C, owing to the formation of the I2+
- 2 I+
2 ⇌ I2+
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+
The best-known oxides are the anions, IO−
3 and IO−
4, but several other oxides are known, such as the strong oxidant iodine pentoxide.
In 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) In basic solutions (such as aqueous sodium hydroxide), iodine converts in a two stage reaction to iodide and iodate:
I2 + 2 OH− → I− + IO− + H2O (K = 30) 3 IO− → 2 I− + IO−
(K = 1020)
- I2 + 10 HNO3 → 2 HIO3 + 10 NO2 + 4 H2O
- I2 + 2 ClO−
3 → 2 IO−
3 + Cl2
Other inorganic compounds
Iodine forms compounds with all the elements except for the noble gases. From the perspective of commercial applications, an important compound is hydroiodic acid, used as a co-catalyst in the Cativa process for the production of acetic acid. Titanium and aluminium iodides are used in the production of butadiene, a precursor to rubber tires.
Alkali metal salts are common colourless solids that are highly soluble in water. Potassium iodide is a convenient source of the iodide anion; it is easier to handle than sodium iodide because it is not hygroscopic. Both salts are used mainly in the production of iodized salt. Sodium iodide is especially useful in the Finkelstein reaction because it is soluble in acetone, whereas potassium iodide is less 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)
Despite having the lowest electronegativity of the common halogens, iodine reacts violently with some metals, such as aluminium:
- 3 I2 + 2 Al → 2 AlI3
This reaction produces 314 kJ per mole of aluminium, comparable to thermite's 425 kJ. The reaction initiates spontaneously, and if unconfined, boils the liquid iodine into a gaseous cloud.
Organoiodine compounds can be produced in many ways. For example, methyl iodide can be prepared from methanol, red phosphorus, and iodine. The iodinating reagent is phosphorus triiodide that is formed in situ:
- 3 CH3OH + PI3 → 3 CH3I + H3PO3
The simplest organoiodine compound is iodomethane, approved as a soil fumigant. 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. Aryl and alkyl iodides both form Grignard reagents. Iodine is sometimes used to activate magnesium when preparing Grignard reagents. Alkyl iodides such as iodomethane are good alkylating agents.
Organoiodine compounds have some drawbacks in chemical synthesis:
- Iodine compounds are more expensive than the corresponding bromides and chlorides, in that order.
- Iodides are much stronger alkylating agents, and thus more toxic (e.g., methyl iodide is very toxic (T+)).
- Low-molecular-weight iodides tend to weigh more than other alkylating agents (e.g., methyl iodide versus dimethyl carbonate) from the atomic mass of iodine.
Of the several iodine sources in nature, only two are useful commercially: the caliche, found in Chile, and the iodine-containing brines of gas and oil fields, especially in Japan and the United States. The caliche contains sodium nitrate, which is the main product of the mining activities, contaminated with small amounts of various calcium iodate minerals. The calcium iodates are reduced to the corresponding iodides. The high concentration of iodine in the caliche and extensive mining made Chile the largest producer in 2007.
Most other iodine producers use naturally occurring brine. The Japanese Minami Kanto gas field east of Tokyo and the American Anadarko Basin gas field in northwest Oklahoma are the two largest such sources. The brine is hotter than 60 °C from 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 to evaporate the iodine, which is passed into an absorbing tower where sulfur dioxide reduces 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 + 2 H2O + SO2 → 2 HI + H2SO4
- 2 HI + Cl2 → I2↓ + 2 HCl
Electrolysis is not used to produce of iodine from seawater the iodine-rich brine are sufficient. Kelp was a common source of iodine in the 18th and 19th centuries, but it is no longer economically viable.
Commercial iodine often contains impurities, which can be removed by sublimation. The element may also be prepared in an ultra-pure form through the reaction of potassium iodide with copper(II) sulfate, which gives copper(II) iodide initially, which then decomposes spontaneously to copper(I) iodide and iodine:
- Cu2+ + 2 I− → CuI2
- 2 CuI2 → 2 CuI + I2
Other methods for isolating iodine in the laboratory include oxidation of the iodide in hydrogen iodide (often made in situ with an iodide and sulfuric acid) with manganese dioxide, as is done with other halides.
In 1811, iodine was discovered by French chemist Bernard Courtois, who was born to a manufacturer of saltpeter (an essential component of gunpowder). At that time of the Napoleonic Wars, saltpeter was in great demand in France. Saltpeter produced from French nitre beds required sodium carbonate, which could be isolated from seaweed collected on the coasts of Normandy and Brittany. To isolate the sodium carbonate, seaweed was burned and the ash washed with water. The remaining waste was destroyed by adding sulfuric acid. Courtois once added excessive sulfuric acid and a cloud of purple vapour rose. He noted that the vapour crystallized on cold surfaces, making dark crystals. Courtois suspected that this was a new element but lacked funding to pursue it further.
Courtois gave samples to his friends, Charles Bernard Desormes (1777–1838) and Nicolas Clément (1779–1841), to continue research. He also gave some of the substance to chemist Joseph Louis Gay-Lussac (1778–1850), and to physicist André-Marie Ampère (1775–1836). On 29 November 1813, Desormes and Clément made Courtois' discovery public. They described the substance to a meeting of the Imperial Institute of France. On 6 December, Gay-Lussac announced that the new substance was either an element or a compound of oxygen. It was Gay-Lussac who suggested the name "iode", from the Greek word ἰοειδής (ioeidēs) for violet (because of the colour of iodine vapor). Ampère had given some of his sample to English chemist Humphry Davy (1778–1829). Davy did some experiments on the substance and noted its similarity to chlorine. Davy sent a letter dated 10 December to the Royal Society of London stating that he had identified a new element. Arguments erupted between Davy and Gay-Lussac over who identified iodine first, but both scientists acknowledged Courtois as the first to isolate the element.
The production of ethylenediamine dihydroiodide, provided as a nutritional supplement for livestock, consumes a large portion of available iodine. Another significant use is a catalyst for the production of acetic acid by the Monsanto and Cativa processes. In these technologies, which support the world's demand for acetic acid, hydroiodic acid converts the methanol feedstock into methyl iodide, which undergoes carbonylation. Hydrolysis of the resulting acetyl iodide regenerates hydroiodic acid and gives acetic acid.
Elemental iodine is used as a disinfectant either as the element, or as the water-soluble triiodide anion I3− generated in situ by adding iodide to poorly water-soluble elemental iodine (the reverse chemical reaction makes some free elemental iodine available for antisepsis). In the alternative, iodine may be produced 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:
- 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 alone, forming mostly triiodide. Unlike tincture of iodine, Lugol's has a minimized amount of the free iodine (I2) component.
- Povidone iodine (an iodophor).
The antimicrobial action of iodine is quick and works at low concentrations. Its specific mode of action is unknown. It penetrates into microorganisms and attacks particular amino acids (such as cysteine and methionine), nucleotides, and fatty acids, ultimately resulting in cell death. It also has an antiviral action, but nonlipid viruses and parvoviruses are less sensitive than lipid enveloped viruses. Iodine probably attacks surface proteins of enveloped viruses, and it may also destabilize membrane fatty acids by reacting with unsaturated carbon bonds.
Iodine is useful in analytical chemistry because it reacts with alkenes, starch, and oxidizing and reducing agents, producing highly colored species that are ready analytical indicators. Iodine is a common general stain used in thin-layer chromatography. Iodine forms an intense blue complex with the glucose polymers starch and glycogen. Several analytical methods rely on this property:
- Iodometry. The concentration of an oxidant can be determined by adding it to an excess of iodide; the test sample oxidizes the elemental iodine/triiodide. A starch indicator is then used close to the end-point, to increase the visual contrast (dark blue becomes colorless, instead of the yellow of dilute triiodide becoming colorless).
- Iodine 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 on the premise that counterfeit banknotes are printed on commercially available paper containing starch.
- Starch-iodide paper is used to test for the presence of oxidants such as peroxides. The oxidants convert iodide to iodine, which becomes blue. A solution of starch and iodide can perform the same function.
- During colposcopy, Lugol's iodine is applied to the vagina and cervix. Normal vaginal tissue stains brown from high glycogen content (a color-reaction similar to that with starch), while abnormal tissue (possibly cancerous) does not stain and appears paler than the surrounding tissue. Suspicious tissue can then be biopsied. This is called a Schiller's Test.
Potassium iodide has been used as an expectorant, although this use is increasingly uncommon. In medicine, a saturated solution of potassium iodide called SSKI is used to treat acute thyrotoxicosis. It is also used to block uptake of iodine-131 in the thyroid gland (see isotopes section above), when this isotope is used as part of radiopharmaceuticals (such as iobenguane) that are not targeted to the thyroid or thyroid-type tissues.
Iodine-131 (usually in the chemical form of iodide) is a component of nuclear fallout, and is particularly dangerous owing to the thyroid gland's propensity to concentrate ingested iodine and retain it for periods longer than this isotope's radiological half-life of eight days. For this reason, people who at risk of exposure to environmental radioactive iodine (iodine-131) in fallout may be instructed to take non-radioactive potassium iodide tablets. The typical adult dose is one 130 mg tablet per 24 hours, supplying 100 mg (100,000 micrograms) of ionic iodine. (The typical daily dose of iodine for normal health is of order 100 micrograms; see "Dietary Intake" below.) Ingestion of this large dose of non-radioactive iodine minimizes the uptake of radioactive iodine by the thyroid gland. See the article on potassium iodide for more on this topic.
As an element with high electron density and atomic number, iodine absorbs X-rays weaker than 33.3 keV due to the photoelectric effect of the innermost electrons. Organic compounds of a certain type (typically iodine-substituted benzene derivatives) are used with intravenous injection as X-ray radiocontrast agents. This is often in conjunction with advanced X-ray techniques such as angiography and CT scanning. At present, all water-soluble radiocontrast agents rely on iodine. It is on the World Health Organization's List of Essential Medicines, a list of the most important medications needed in a basic health system.
Iodine and cancer risk
- Breast cancer. The mammary gland actively concentrates iodine into milk for the benefit of the developing infant, and may develop a goiter-like hyperplasia, sometimes manifesting as fibrocystic breast disease, when iodine level is low. Studies 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, with elemental iodine (I2) having been found to be more effective in reducing ductal hyperplasias and perilobular fibrosis in iodine-deficient rats than iodide (I−). On the observation that Japanese women who consume iodine-rich seaweed have a relatively low rate of breast cancer, iodine is suggested as a protection against breast cancer. Iodine is known to induce apoptosis in breast cancer cells. Laboratory evidence has demonstrated an effect of iodine on breast cancer that is in part independent of thyroid function, with iodine inhibiting cancer through modulation of the estrogen pathway. Gene array profiling of the 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.
- Iodine and stomach cancer. Some researchers have found an epidemiologic correlation between iodine deficiency, iodine-deficient goitre, and gastric cancer; a decrease in the death incidence from stomach cancer after iodine-prophylaxis. In the proposed mechanism, the iodide ion functions in gastric mucosa as an antioxidant reducing species that detoxifies poisonous reactive oxygen species, such as hydrogen peroxide.
Historical medical applications
In the early 1900s, The Encyclopædia Britannica described iodine being "of definite value" for treatment of multiple conditions including "metallic poisonings, as by lead and mercury, asthma, aneurism, arteriosclerosis, angina pectoris, gout, goitre, syphilis, haemophilia, Bright's disease (nephritis) and bronchitis" with "usual doses" of iodide salts ranging from "five to thirty grains or more" (324 mg to 1,944 mg), though this is hundred of times higher than what is considered generally safe by today's tolerable UL. For treatment of syphilis, it states "in its tertiary stages and also earlier this disease yields in the most rapid and unmistakable fashion to iodides; so much so that the administration of these salts is at present the best means of determining whether, for instance, a cranial tumour be syphilitic or not". (Modern treatment for syphilis involves the use of antibiotics to kill syphilis bacteria - see Syphilis.) For the treatment of chronic lead poisoning, it states "the essential part of the medicinal treatment of this condition is the administration of iodides, which are able to decompose the insoluble albuminates of lead which have become locked up in the tissues, rapidly causing their degeneration, and to cause the excretion of the poisonous metal by means of the intestine and the kidneys" (modern treatment for lead poisoning involves the use of a variety of substances - see Lead poisoning).
Inorganic iodides find specialized uses. Hafnium, zirconium, titanium are purified by the van Arkel Process, which involves the reversible formation of the tetraiodides of these elements. Silver iodide is a major ingredient to traditional photographic film. Thousands of kilograms of silver iodide are consumed annually for cloud seeding to induce rain.
In the United States, the Drug Enforcement Administration (DEA) regards iodine and compounds containing iodine (ionic iodides, iodoform, ethyl iodide, and so on) as reagents useful for the clandestine manufacture of methamphetamine.
Iodine is an essential trace element for life, the heaviest element commonly needed by living organisms. Only tungsten, a component of a few bacterial enzymes, has a higher atomic number and atomic weight.
In animal biology, iodine is primarily a constituent of the thyroid hormones thyroxine (T4) and triiodothyronine (T3). These are addition condensation products of the amino acid tyrosine, and are stored prior to release in aprotein 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 that are regulated by a second hormone TSH from the pituitary. Thyroid hormones are phylogenetically very old molecules that are synthesized by most multicellular organisms, and that 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. Total deficiency of thyroid hormones can reduce basal metabolic rate by as much as 50%. Excessive production of thyroid hormones can increase the basal metabolic rate by 100%. T4 acts largely as a precursor to T3, which is (with minor exceptions) the biologically active hormone. In amphibian metamorphosis iodine and thyroid hormones exert a well-studied experimental model of apoptosis on the cells of gills, tail, and fins of tadpoles.
Iodine has a nutritional relationship with selenium. A family of selenium-dependent enzymes called deiodinases converts T4 to T3 (the active hormone) by removing an iodine atom from the outer tyrosine ring. These enzymes also convert T4 to reverse T3 (rT3) by removing an inner ring iodine atom, and convert T3 to 3,3'-diiodothyronine (T2) by the same action. Both of the latter are inactivated hormones that are ready for disposal and have, in essence, no biological effects. A family of non-selenium-dependent enzymes then further deiodinates the products of these reactions.
Iodine accounts for 65% of the molecular weight of T4 and 59% of the T3. Fifteen to 20 mg of iodine is concentrated in thyroid tissue and hormones, but 70% of all iodine in the body is found in other tissues, including mammary glands, eyes, gastric mucosa, fetal thymus, cerebro-spinal fluid and choroid plexus, arterial walls, the cervix, and salivary glands. In the cells of those tissues, iodide enters directly by sodium-iodide symporter (NIS). The action of iodine in mammary tissue is related to fetal and neonatal development, but in the other tissues, it is (at least) partially unknown.
The daily Dietary Reference Intake recommended by the United States Institute of Medicine is between 110 and 130 µg for infants up to 12 months, 90 µg for children up to eight years, 130 µg for children up to 13 years, 150 µg for adults, 220 µg for pregnant women and 290 µg for lactating mothers. The 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. and does not take into account a tolerable upper limit for absorption into other tissues of the body, such as breast tissue in women or prostate tissue in men.
The thyroid gland needs no more than 70 μg/day to synthesize the requisite daily amounts of T4 and T3. The 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, brain cells, choroid plexus, oral mucosa, and arterial walls.
Natural sources of dietary iodine include seafood, such as fish, seaweeds (e.g. kelp) and shellfish, dairy products and eggs so long as the animals received enough iodine, and 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. The general US population has adequate iodine nutrition, with women of childbearing age and pregnant women having a possible mild risk of deficiency. In Japan, consumption was considered much higher, ranging between 5,280 μg/day to 13,800 μg/day from dietary seaweed or kombu kelp, often in the form of Kombu Umami extracts for soup stock and potato chips. However, new studies suggest that Japan's consumption is closer to 1,000–3,000 μg/day. The tolerable upper intake limit of iodine in Japan is 3,000 µg/day for an adult. Some reports indicate that a large intake of Kombu affected the health in Japan.
After iodine fortification programs (e.g., iodized salt) have been implemented, some cases of iodine-induced hyperthyroidism have been observed (so-called Jod-Basedow phenomenon). The condition seems to occur mainly in people over forty, and the risk appears higher when iodine deficiency is severe and the initial rise in iodine intake is high.
Information processing, fine motor skills, and visual problem solving are improved by iodine repletion in moderately iodine-deficient children.
In an estimated two-thirds of households on Earth, table salt is iodized. However, this still leaves an estimated two billion people iodine-deficient. Iodine is required for the essential thyroxin hormones produced by and concentrated in the thyroid gland.
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 intellectual disability, a result that occurs primarily when babies or small children are rendered hypothyroidic by a lack of the element. The addition of iodine to table salt has largely eliminated this problem in 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.
Investigation continues into other possible health problems caused by iodine deficiency, including:
- Breast cancer. The mammary glands concentrates iodine into milk for the benefit of the developing infant, and may develop a goiter-like hyperplasia, sometimes manifesting as fibrocystic breast disease, when iodine levels are low.
- Stomach cancer. Some researchers have found an epidemiologic correlation between iodine deficiency, iodine-deficient goitre, and gastric cancer. Other studies have reported that iodine supplements were coincident with a decrease in the death rate from stomach cancer.
- Autism. Although no definitive cause of autism in children has been found, research shows that iodine-deficient women are more likely to have autistic children.
Elemental iodine (I2) is toxic if taken orally. The lethal dose for an adult human is 30 mg/kg, which is about 2.1–2.4 grams for a human weighing 70 to 80 kg (even if experiments on rats demonstrated that these animals could survive after eating a 14000 mg/kg dose). Excess iodine can be more cytotoxic in the presence of selenium deficiency. Iodine supplementation in selenium-deficient populations is, in theory, problematic, partly for this reason. The toxicity derives from its oxidizing properties, through which it denaturates proteins (including enzymes).
Elemental iodine is also a skin irritant, and direct contact with skin can cause damage and solid iodine crystals should be handled with care. Solutions with high elemental iodine concentration, such as tincture of iodine and Lugol's solution, are capable of causing tissue damage if used in prolonged cleaning or antisepsis; similarly, liquid Povidone-iodine (Betadine) trapped against the skin resulted in chemical burns in some reported cases.
People can be exposed to iodine in the workplace by inhalation, ingestion, skin contact, and eye contact. The Occupational Safety and Health Administration (OSHA) has set the legal limit (Permissible exposure limit) for iodine exposure in the workplace at 0.1 ppm (1 mg/m3) during an 8-hour workday. The National Institute for Occupational Safety and Health (NIOSH) has set a Recommended exposure limit (REL) of 0.1 ppm (1 mg/m3) during an 8-hour workday. At levels of 2 ppm, iodine is immediately dangerous to life and health.
Some people develop a hypersensitivity to products and foods containing iodine. Applications of tincture of iodine or Betadine can cause rashes, sometimes severe. Parenteral use of iodine-based contrast agents (see above) can cause reactions ranging from a mild rash to fatal anaphylaxis. Such reactions have led to the misconception (widely held, even among physicians) that some people are allergic to iodine itself; even allergies to iodine-rich seafood have been so construed. In fact, there has never been a confirmed report of a true iodine allergy, and an allergy to elemental iodine or simple iodide salts is theoretically impossible. Hypersensitivity reactions to products and foods containing iodine are apparently related to their other molecular components; thus, a person who has demonstrated an allergy to one food or product containing iodine may not be have an allergic reaction to another. Patients with various food allergies (shellfish, egg, milk, etc.) or asthma are more likely to suffer reactions to contrast media containing iodine. As with all medications, the patient's allergy history should be questioned and consulted before any containing iodine are administered.
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