Selenium
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| iso | NA | half-life | DM | DE (MeV) | DP
Template:Elementbox isotopes decay2 Template:Elementbox isotopes stable Template:Elementbox isotopes decay2 Template:Elementbox isotopes stable Template:Elementbox isotopes stable Template:Elementbox isotopes stable Template:Elementbox isotopes decay Template:Elementbox isotopes stable Template:Elementbox isotopes decay Template:Elementbox isotopes end Template:Elementbox footer Editor-In-Chief: C. Michael Gibson, M.S., M.D. [2]
 Selenium (Template:PronEng) is a chemical element with atomic number 34, with the chemical symbol Se. It is a nonmetal, chemically related to sulfur and tellurium, and rarely occurring in its elemental state in nature. It is toxic in large amounts, but trace amounts of it are necessary for cellular function in most, if not all, animals, forming the active center of the enzymes glutathione peroxidase and thioredoxin reductase (which indirectly reduce certain oxidized molecules in animals and some plants) and three known deiodinase enzymes (which convert one thyroid hormone to another). Selenium requirements in plants differ by species, with some plants apparently requiring none.[1] Isolated selenium occurs in several different forms, the most stable of which is a dense purplish-gray semimetal (semiconductor) form that is structurally a trigonal polymer chain. It conducts electricity better in the light than in the dark, and is used in photocells (see allotropic section below). Selenium also exists in many nonconductive forms: a black glass-like allotrope, as well as several red crystalline forms built of eight-membered ring molecules, like its lighter chemical cousin sulfur. Selenium is found in economic quantities in sulfide ores such as pyrite, partially replacing the sulfur in the ore matrix. Minerals that are selenide or selenate compounds are also known, but all are rare. OccurrenceSelenium occurs naturally in a number of inorganic forms, including selenide, selenate and selenite. In soils, selenium most often occurs in soluble forms like selenate (analogous to sulfate), which are leached into rivers very easily by runoff. Selenium has a biological role, and is found in organic compounds such as dimethyl selenide, selenomethionine and selenocysteine. In these compounds selenium plays an analogous role to sulfur. Selenium is most commonly produced from selenide in many sulfide ores, such as those of copper, silver, or lead. It is obtained as a byproduct of the processing of these ores, from the anode mud of copper refineries and the mud from the lead chambers of sulfuric acid plants. These muds can be processed by a number of means to obtain free selenium. Natural sources of selenium include certain selenium-rich soils, and selenium that has been bioconcentrated by certain toxic plants such as locoweed. Anthropogenic sources of selenium include coal burning and the mining and smelting of sulfide ores.[2] IsotopesSelenium has six naturally-occurring isotopes, five of which are stable: 74Se, 76Se, 77Se, 78Se, and 80Se. The last three also occur as fission products, along with Se-79 which has a halflife of 295,000 years, and 82Se which has a very long half life (~1020 yr, decaying via double beta decay to 82Kr) and for practical purposes can be considered to be stable. 23 other unstable isotopes have been characterized. History and global demandSelenium (Greek σελήνη selene meaning "Moon") was discovered in 1817 by Jöns Jakob Berzelius who found the element associated with tellurium (named for the Earth). Growth in selenium consumption was historically driven by steady development of new uses, including applications in rubber compounding, steel alloying, and selenium rectifiers. By 1970, selenium in rectifiers had largely been replaced by silicon, but its use as a photoconductor in plain paper copiers had become its leading application. During the 1980s, the photoconductor application declined (although it was still a large end-use) as more and more copiers using organic photoconductors were produced. Presently, the largest use of selenium world-wide is in glass manufacturing, followed by uses in chemicals and pigments. Electronic use, despite a number of continued applications, continues to decline.[3] In 1996, continuing research showed a positive correlation between selenium supplementation and cancer prevention in humans, but widespread direct application of this important finding would not add significantly to demand owing to the small doses required. In the late 1990s, the use of selenium (usually with bismuth) as an additive to plumbing brasses to meet no-lead environmental standards, became important. At present, total world selenium production continues to increase modestly. Selenium and healthAlthough it is toxic in large doses, selenium is an essential micronutrient for animals. In plants, it occurs as a bystander mineral, sometimes in toxic proportions in forage (some plants may accumulate selenium as a defense against being eaten by animals, but other plants such as locoweed require selenium, and their growth indicates the presence of selenium in soil).[4] It is a component of the unusual amino acids selenocysteine and selenomethionine. In humans, selenium is a trace element nutrient which functions as cofactor for reduction of antioxidant enzymes such as glutathione peroxidases and certain forms of thioredoxin reductase found in animals and some plants (this enzyme occurs in all living organisms, but not all forms of it in plants require selenium). Glutathione peroxidase (GSH-Px) catalyzes certain reactions which remove reactive oxygen species such as peroxide:
Selenium also plays a role in the functioning of the thyroid gland by participating as a cofactor for the three known thyroid hormone deiodinases.[5] Dietary selenium comes from nuts, cereals, meat, fish, and eggs. Brazil nuts are the richest ordinary dietary source (though this is soil-dependent, since the Brazil nut does not require high levels of the element for its own needs). High levels are found in meats found in kidney, crab and lobster, in that order.[6][7] In the USA, the recommended dietary allowance for adults is 55 micrograms per day. ToxicityAlthough selenium is an essential trace element, it is toxic if taken in excess. Exceeding the Tolerable Upper Intake Level of 400 micrograms per day can lead to selenosis.[8] Symptoms of selenosis include a garlic odour on the breath, gastrointestinal disorders, hair loss, sloughing of nails, fatigue, irritability and neurological damage. Extreme cases of selenosis can result in cirrhosis of the liver, pulmonary edema and death.[9] Elemental selenium and most metallic selenides have relatively low toxicities because of their low bioavailability. By contrast, selenate and selenite are very toxic, and have modes of action similar to that of arsenic. Hydrogen selenide is an extremely toxic, corrosive gas.[10] Selenium also occurs in organic compounds such as dimethyl selenide, selenomethionine and selenocysteine, all of which have high bioavailability and are toxic in large doses. Selenium poisoning of water systems may result whenever new agricultural runoff courses through normally-dry undeveloped lands. This process leaches natural soluble selenium compounds (such as selenates) into the water, which may then be concentrated in new "wetlands" as it evaporates. High selenium levels produced in this fashion have been found to have caused certain congenital disorders in wetland birds.[11] DeficiencySelenium deficiency is relatively rare in healthy well-nourished individuals. It can occur in patients with severely compromised intestinal function, or those undergoing total parenteral nutrition. Alternatively, people dependent on food grown from selenium-deficient soil are also at risk. In the USA, the Dietary Reference Intake for adults is 55 micrograms per day. In the UK it is 75 µg/day for adult males and 60 µg/day for adult females Selenium deficiency can lead to Keshan disease, which is potentially fatal. Selenium deficiency also contributes (along with iodine deficiency) to Kashin-Beck disease.[12] The primary symptom of Keshan disease is myocardial necrosis, leading to weakening of the heart. Kashin-Beck disease results in atrophy, degeneration and necrosis of cartilage tissue.[13] Keshan disease also makes the body more susceptible to illness caused by other nutritional, biochemical, or infectious diseases. These diseases are most common in certain parts of China where the soil is extremely deficient in selenium. Studies in Jiangsu Province of China have indicated a reduction in the prevalence of these diseases by taking selenium supplements. Selenium is also necessary for the conversion of the thyroid hormone thyroxine (T4) into its more active counterpart, triiodothyronine, and as such a deficiency can cause symptoms of hypothyroidism, including extreme fatigue, mental slowing, goitre, cretinism and recurrent miscarriage.[9] Controversial health effects
Production and allotropic formsSelenium is a common byproduct of copper refining, or the production of sulfuric acid.[36][37][38] Isolation of selenium is often complicated by the presence of other compounds and elements. Commonly, production begins by oxidation with sodium carbonate to produce sodium selenite. The sodium selenite is then acidified with sulfuric acid producing selenous acid. The selenous acid is finally bubbled with sulfur dioxide producing elemental red amorphous selenium. Selenium produced in chemical reactions invariably appears as the amorphous red form-- an insoluble brick red powder. When this form is rapidly melted, it forms the black, vitreous form which is usually sold industrially as beads. The most thermodynamically stable and dense form of selenium is the electrically conductive gray (trigonal) form, which is composed of long helical chains of selenium atoms. The conductivity of this form is notably light sensitive. Selenium also exists in three different deep red crystalline monoclinic forms, which are composed of Se8 molecules, similar to many allotropes of sulfur.[39] Non-biological applications
Biological applications
Compounds
Selenium also occurs in the III oxidation state, but only in the Se412+ cation; Se(III) compounds are not otherwise known.[42]
References
External links
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