BOILING POINT: 685°C
DENSITY: 4,819 kg/m 3
MOST COMMON IONS: SeO 3 2− , SeO 4 2− , Se 2−
Selenium (from the Greek word selēnē —the Moon), discovered by Swedish chemist Jöns Jakob Berzelius in 1817, ranks thirty-fourth among elements in Earth's crust. It has six naturally occurring isotopes , a large number of allotropes (elemental forms), and in compounds has oxidation states −2, +4, and +6. The gray elemental form has the unique photoelectric property of exhibiting lowered electrical resistance when struck by light, and it is used in photovoltaic cells and photocells (e.g., light meters) and in xerography. It conducts electricity in a "unipolar" manner, hence it is commonly used in electrical rectifiers. It is also used to tint glass red and to decolorize green glass.
Selenium substitutes for sulfur in amino acids to form seleno-cysteine, cystine, and methionine. The selenium-containing antioxidant glutathione peroxidase is biologically important, and selenium is a necessary trace nutrient in warm-blooded animals. Grazing animals develop a form of muscular dystrophy and other disorders when grazing in areas in which the selenium has been depleted; with selenium-depleted diets, people develop
Keshan disease, a form of cardiomyopathy. When its intake is too high, selenium disrupts enzyme function, causing poor health in mammals and birth defects and reproductive failure in birds and fish. Good sources of selenium in human diets include wheat, garlic, Brazil nuts, and walnuts.
SEE ALSO Chalcogens .
Barceloux, Donald G. (1999). "Selenium." Journal of Toxicology: Clinical Toxicology 37(2): 145–172.
Frankenberger, William T., and Engberg, Richard A., eds. (1998). Environmental Chemistry of Selenium. New York: Marcel Dekker.