Inadequate intake or metabolism of iodine. It directly affects thyroid secretions, which influence heart action, nerve response, growth rate, and metabolism. Simple goitre, the most frequent result, is most common in areas without access to salt water and is rare along seacoasts. Severe, prolonged deficiency can cause hypothyroidism. Eating seafood regularly or using iodized table salt will prevent iodine deficiency. Some countries have made dietary iodine additives mandatory.
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Nonmetallic chemical element, chemical symbol I, atomic number 53. The heaviest nonradioactive halogen, it is a nearly black crystalline solid (diatomic molecule I2) that sublimes (see sublimation) to a deep violet, irritating vapour. It is never found in nature uncombined. Its sources (mostly in brines and seaweeds) and compounds are usually iodides; iodates (small amounts in saltpeter) and periodates also occur. Dietary iodine is essential for thyroid gland function; in areas of the world where food contains insufficient iodine, an iodine compound such as potassium iodide (KI) is added to table salt (sodium chloride) to prevent iodine deficiency. Elemental iodine is used in medicine, in synthesizing some organic chemicals, in manufacturing dyes, in analytical chemistry (see analysis) to measure fat saturation (see hydrogenation) and to detect starch, and in photography. The radioactive isotope iodine-131 (see radioactivity), with an eight-day half-life, is very useful in medicine (see nuclear medicine) and other applications.
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Chemically, iodine is the least reactive of the halogens, and the most electropositive halogen after astatine. However, the element does not occur in the free state in nature. As with all other halogens (members of Group VII in the Periodic Table), when freed from its compounds iodine forms diatomic molecules (I2).
Iodine and its compounds are primarily used in medicine, photography, and dyes. Although it is rare in the solar system and Earth's crust, the iodides are very soluble in water, and the element is concentrated in seawater. This mechanism helps to explain how the element came to be required in trace amounts by all animals and some plants, being by far the heaviest element known to be necessary to living organisms.
Elemental iodine dissolves easily in chloroform and carbon tetrachloride. The solubility of elementary iodine 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.
However he gave samples to his friends, Charles Bernard Desormes (1777–1862) and Nicolas Clément (1779–1841), to continue research. He also gave some of the substance to Joseph Louis Gay-Lussac (1778–1850), a well-known chemist at that time, and to physicist André-Marie Ampère (1775–1836). On 29 November 1813, Dersormes and Clément made public Courtois’ discovery. They described the substance to a meeting of the Imperial Institute of France. On December 6, Gay-Lussac announced that the new substance was either an element or a compound of oxygen. Ampère had given some of his sample to Humphry Davy (1778–1829). Davy did some experiments on the substance and noted its similarity to chlorine. Davy sent a letter dated December 10 to the Royal Society of London stating that he had identified a new element. A large argument erupted between Davy and Gay-Lussac over who identified iodine first but both scientists acknowledged Courtois as the first to isolate the chemical element.
Extraction from seawater involves electrolysis. The brine is first purified and acidified using sulphuric acid and is then reacted with chlorine. An iodine solution is produced but it is yet too dilute and has to be concentrated. To do this air is blown into the solution which causes the iodine to evaporate, then it is passed into an absorbing tower containing acid where sulfur dioxide is added to reduce the iodine. The solution is then added to chlorine again to concentrate the solution more, and the final solution is at a level of about 99%.
Another source is from kelp,a kind of brown alga. This source was used in the 18th and 19th centuries but is no longer economically viable.
Elemental iodine can be prepared by oxidizing iodides with chlorine:
or with manganese dioxide in acid solution:
or by hydrazine:
or by chlorates:
|I2 + 2OH− → I− + IO− + H2O||(K = 30)|
|3IO− → 2I− + IO3−||(K = 1020)|
Thyroid hormones play a basic role in biology, acting on gene transcription to regulate the basal metabolic rate. The total deficiency of thyroid hormones can reduce basal metabolic rate up to 50%, while in excessive production of thyroid hormones the basal metabolic rate can be increased by 100%. T4 acts largely as a precursor to T3, which is (with minor exceptions) the biologically active hormone.
Iodine accounts for 65% of the molecular weight of T4 and 59% of the T3. 15-20 mg of iodine is concentrated in thyroid tissue and hormones, but 70% of the body's iodine is distributed in other tissues, including mammary glands, eyes, gastric mucosa, the cervix, and salivary glands. Its role in mammary tissue is related to fetal and neonatal development, but its role in the other tissues is unknown. It has been shown to act as an antioxidant in these tissues.
Iodine may have a relationship with selenium, and iodine supplementation in selenium-deficient populations may pose risks for thyroid function.
Natural sources of iodine include sea life, such as kelp and certain seafood, as well as plants grown on iodine-rich soil. Iodized salt is fortified with iodine.
As of 2000, the median intake of iodine from food in the United States was 240 to 300 μg/day for men and 190 to 210 μg/day for women. In Japan, consumption is much higher due to the frequent consumption of seaweed or kombu kelp. Estimates range from 5,280 to 13,800 μg/day. Although some Chinese data associates excess iodine with autoimmune thyroiditis and hypothyroidism, these effects have not been observed in Japanese populations, and a protective effect on breast cancer has been hypothesized.
Donald W. Miller believes that the daily FDA intake recommendation may be 100 times too low.
Iodine deficiency is the leading cause of preventable mental retardation, a result which occurs primarily when babies or small children are rendered hypothyroidic by a lack of the element. The addition of iodine to table salt has largely eliminated this problem in the wealthier nations, but as of March 2006, iodine deficiency remained a serious public health problem in the developing world. Iodine deficiency is also a problem in certain areas of Europe. In Germany it has been estimated to cause a billion dollars in healthcare costs per year.
The artificial radioisotope 131I (a beta emitter), has a half-life of 8.0207 days. Also known as radioiodine, 131I has been used in treating cancer and other pathologies of the thyroid glands. 123I is the radioisotope most often used in nuclear imaging of the kidney and thyroid as well as thyroid uptake scans (used for the evaluation of Graves' Disease). The most common compounds of iodine are the iodides of sodium and potassium (KI) and the iodates (KIO3).
Potassium iodide (KI tablets, or "SSKI" = "Saturated Solution of KI" liquid drops) can be given to people in a nuclear disaster area when fission has taken place, to block the uptake of iodine-131 by the thyroid. The protective effect of KI lasts approximately 24 hours, so it should be dosed daily until a risk of significant exposure to radioiodines no longer exists. The exposure can be reduced by evacuation, sheltering, and by control of the food supply. Iodine-131 also decays rapidly, with a half-life of 8 days, so that 99.95% of the original radioiodine is gone after three months.
Iodine-129 129I (half-life 15.7 million years) is a product of cosmic ray spallation on various isotopes of xenon in the atmosphere, in cosmic ray muon interaction with tellurium-130, and also and 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.
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 above.