magnesium


magnesium
/mag nee"zee euhm, -zheuhm, -shee euhm/, n. Chem.
a light, ductile, silver-white, metallic element that burns with a dazzling white light, used in lightweight alloys, flares, fireworks, in the manufacture of flashbulbs, optical mirrors, and precision instruments, and as a zinc substitute in batteries. Symbol: Mg; at. wt.: 24.312; at. no.: 12; sp. gr.: 1.74 at 20°C.
[1800-10; < NL; see MAGNESIA, -IUM]

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Chemical element, one of the alkaline earth metals, chemical symbol Mg, atomic number 12.

The silvery white metal does not occur free in nature, but compounds such as the sulfate (Epsom salts), oxide (magnesia), and carbonate (magnesite) have long been known. The metal, which burns in air with a bright white light, is used in photographic flash devices, bombs, flares, and pyrotechnics; it is also a component of lightweight alloys for aircraft, spacecraft, cars, machinery, and tools. The compounds, in which it has valence 2, are used as insulators and refractories and in fertilizers, cement, rubber, plastics, foods, and pharmaceuticals (antacids, purgatives, laxatives). Magnesium is an essential element in human nutrition; it is the cofactor in enzymes of carbohydrate metabolism and in chlorophyll.

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Introduction
 chemical element, one of the alkaline-earth metals of main Group 2 (IIa) of the periodic table, the lightest structural metal. Known originally through compounds such as Epsom salts (the sulfate), magnesia (the oxide), and magnesia alba (the carbonate), the silvery white element itself does not occur free. It was first isolated in 1808 by Sir Humphry Davy (Davy, Sir Humphry, Baronet), who evaporated the mercury from a magnesium amalgam made by electrolyzing a mixture of moist magnesia and mercuric oxide.

Occurrence, properties, and uses
      Magnesium is the eighth most abundant element in the Earth's crust (about 2.5 percent) and is, after aluminum and iron, the third most plentiful structural metal. Its cosmic abundance is estimated as 9.1 × 105 atoms (Si = 106 atoms). It occurs as carbonates (magnesite, MgCO3, and dolomite, CaCO3∙MgCO3) and in many common silicates, including talc, olivine, and most kinds of asbestos. It also is found as hydroxide (brucite), chloride (carnallite, KCl∙MgCl2∙6H2O), and sulfate (kieserite). It is also distributed in minerals such as serpentine, chrysolite, and meerschaum. Seawater contains about 0.13 percent magnesium, mostly as the dissolved chloride, which imparts the characteristic bitter taste. Magnesium is about one-sixth as plentiful as potassium in human body cells, where it is required as a catalyst for enzyme reactions in carbohydrate metabolism.

      Magnesium is commercially produced by electrolysis of molten magnesium chloride (MgCl2), processed mainly from seawater and by the direct reduction of its compounds with suitable reducing agents, as from magnesium oxide or calcined dolomite with ferrosilicon (see magnesium processing).

      At one time, magnesium was used predominantly for photographic flash ribbon and powder, incendiary bombs, and pyrotechnic devices, because in finely divided form it burns in air with an intense white light. Because of its low density (only two-thirds that of aluminum) it has found extensive use in the aerospace industry. However, because the pure metal has low structural strength, magnesium is mainly used in the form of alloys—principally with 10 percent or less of aluminum, zinc, and manganese—to improve its hardness, tensile strength, and ability to be cast, welded, and machined. Casting, rolling, extruding, and forging techniques are all employed with the alloys, and further fabrication of the resulting sheet, plate, or extrusion is carried out by normal forming, joining, and machining operations. Magnesium is the easiest structural metal to machine and has often been used when a large number of machining operations are required. Magnesium alloys have a number of applications; they are used for parts of aircraft, spacecraft, machinery, automobiles, portable tools, and household appliances. Because of its ready combustibility, magnesium still finds application in explosive and pyrotechnic devices.

      Its thermal and electrical conductivity and its melting point are very similar to those of aluminum. Whereas aluminum is attacked by alkalies but is resistant to most acids, magnesium is resistant to most alkalies but is readily attacked by most acids (chromic and hydrofluoric acids are important exceptions). At normal temperatures it is stable in air and water owing to the formation of a thin protective skin of oxide (but burns rapidly when heated in air), and it is attacked by steam. Magnesium is a powerful reducing agent and is used to produce other metals from their compounds (e.g., titanium, zirconium, and hafnium). It reacts directly with many elements.

      Magnesium occurs in nature as a mixture of three isotopes: magnesium-24 (78.70 percent), magnesium-26 (11.17 percent), and magnesium-25 (10.13 percent). It is a very strong reducing agent, reacting with most acids or with boiling water to liberate hydrogen, but is resistant to most alkalies. In compounds it always exhibits a +2 oxidation state because of the loss or sharing of its two 3s electrons.

Principal compounds
      Magnesium carbonate, MgCO3, occurs in nature as the mineral magnesite and is an important source of elemental magnesium. It can be produced artificially by the action of carbon dioxide on a variety of magnesium compounds. The odourless white powder has many industrial uses—e.g., as a heat insulator for boilers and pipes and as an additive in food, pharmaceuticals, cosmetics, rubbers, inks, and glass.

      Magnesium hydroxide, Mg(OH)2, is a white powder produced in large quantities from seawater by the addition of milk of lime (calcium hydroxide). It is the primary raw material in the production of magnesium metal. In water it forms a suspension known as milk of magnesia, which has long been used as an antacid and a laxative.

      The action of hydrochloric acid on magnesium hydroxide produces magnesium chloride, MgCl2, a colourless, deliquescent (water-absorbing) substance employed in magnesium metal production, in the manufacture of a cement for heavy-duty flooring, and as an additive in textile manufacture. Roasting either magnesium carbonate or magnesium hydroxide produces the oxygen compound magnesium oxide, commonly called magnesia, MgO, a white solid used in the manufacture of high-temperature refractory bricks, electrical and thermal insulators, cements, fertilizer, rubber, and plastics. It is used medically as a laxative.

      Magnesium sulfate, MgSO4, is a colourless, crystalline substance formed by the reaction of magnesium hydroxide with sulfur dioxide and air. A hydrate form of magnesium sulfate called kieserite, MgSO4∙H2O, occurs as a mineral deposit. Synthetically prepared magnesium sulfate is sold as Epsom salt, MgSO4∙7H2O. In industry magnesium sulfate is used in the manufacture of cements and fertilizers and in tanning and dyeing; in medicine it serves as a purgative.

      Among the organometallic compounds of magnesium are the important Grignard reagents (Grignard reagent), composed of an organic group (e.g., alkyls and aryls), a halogen atom other than fluorine, and magnesium (see Grignard reagent). These are used in the production of many other kinds of organometallic compounds.

      Magnesium also is a constituent of chlorophyll, in which it apparently plays a role similar to that of iron in hemoglobin. The photosynthetic function of plants depends upon the action of chlorophyll pigments, which contain magnesium at the centre of a complex, nitrogen-containing ring system (porphyrin). These magnesium compounds enable the energy of light to be used to convert carbon dioxide and water to carbohydrates and oxygen and thus directly or indirectly provide the key to nearly all living processes.

atomic number
12
atomic weight
24.312
melting point
651° C
boiling point
1,107° C
specific gravity
1.74 (20° C)
oxidation state
+2
electronic config.
1s22s22p63s2

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Universalium. 2010.

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