|Name, Symbol, Number|| magnesium, Mg, 12
|Chemical series||alkaline earth metals|
|Group, Period, Block||2, 3, s|
|Appearance|| silvery white |
|Atomic mass||24.3050(6) g/mol|
|Electron configuration||[Ne] 3s2|
|Electrons per shell||2, 8, 2|
|Density (near r.t.)||1.738 g·cm−3|
|Liquid density at m.p.||1.584 g·cm−3|
|Melting point|| 923 K|
(650 °C, 1202 °F)
|Boiling point|| 1363 K|
(1090 °C, 1994 °F)
|Heat of fusion||8.48 kJ·mol−1|
|Heat of vaporization||128 kJ·mol−1|
|Heat capacity||(25 °C) 24.869 J·mol−1·K−1|
|Oxidation states|| 2|
(strongly basic oxide)
|Electronegativity||1.31 (Pauling scale)|
| Ionization energies|
|1st: 737.7 kJ·mol−1|
|2nd: 1450.7 kJ·mol−1|
|3rd: 7732.7 kJ·mol−1|
|Atomic radius||150 pm|
|Atomic radius (calc.)||145 pm|
|Covalent radius||130 pm|
|Van der Waals radius||173 pm|
|Electrical resistivity||(20 °C) 43.9 nΩ·m|
|Thermal conductivity||(300 K) 156 W·m−1·K−1|
|Thermal expansion||(25 °C) 24.8 µm·m−1·K−1|
|Speed of sound (thin rod)|| (r.t.) (annealed)|
|Young's modulus||45 GPa|
|Shear modulus||17 GPa|
|Bulk modulus||45 GPa|
|Brinell hardness||260 MPa|
|CAS registry number||7439-95-4|
Magnesium (IPA: /magˈniːziəm, -ʒəm/) is the chemical element in the periodic table that has the symbol Mg and atomic number 12 and an atomic mass of 24.31. Magnesium is the eighth most abundant element and constitutes about 2% of the Earth's crust by weight, and it is the third most plentiful element dissolved in seawater. Magnesium ion is essential to all living cells. The free element (metal) is not found in nature. Once produced from magnesium salts, this alkaline earth metal is primarily used as an alloying agent to make aluminium-magnesium alloys, sometimes called "magnalium" or "magnelium".
26Mg is a stable isotope that has found application in isotopic geology, similar to that of aluminium. 26Mg is a radiogenic daughter product of 26Al, which has a half-life of 717,000 years. Large enrichments of stable 26Mg have been observed in the Ca-Al-rich inclusions of some carbonaceous chondrite meteorites. The anomalous abundance of 26Mg is attributed to the decay of its parent 26Al in the inclusions. Therefore, the meteorite must have formed in the solar nebula before the 26Al had decayed. Hence, these fragments are among the oldest objects in the solar system and have preserved information about it early history.
It is conventional to plot 26Mg/24Mg against an Al/Mg ratio. In an isochrone plot, the Al/Mg ratio plotted is27Al/24Mg. The slope of the isochron has no age significance, but indicates the initial 26Al/27Al ratio in the sample at the time when the systems were separated from a common reservoir.
Notable characteristics of element and compoundsEdit
Elemental magnesium is a fairly strong, silvery-white, light-weight metal (two thirds the density of aluminium). It slightly tarnishes when exposed to air, although unlike the alkaline metals, storage in an oxygen free environment is unnecessary because magnesium is protected by a thin layer of oxide which is fairly impermeable and hard to remove. Like its lower periodic table group neighbor calcium, magnesium reacts with water at room temperature, though it reacts much more slowly than calcium. When it is submerged in water hydrogen bubbles will almost unnoticably begin to form on the surface of the metal, though if powdered it will react much more rapidly. The reaction will occur faster with higher temperatures (see precautions). Magnesium is a highly flammable metal, but while it is easy to ignite when powdered or shaved into thin strips, it is difficult to ignite in mass or bulk. Once ignited it is difficult to extinguish, being able to burn in both nitrogen (forming magnesium nitride), and carbon dioxide (forming magnesium oxide and carbon).
Magnesium, when it burns in air, produces a brilliant white light. This was used in the early days of photography when magnesium powder was used as a source of illumination (flash powder). Later, magnesium ribbon was used in electrically ignited flash bulbs. Magnesium powder is still used in the manufacture of fireworks and marine flares where a brilliant white light is required.
Magnesium, when glowing white, has many chemical properties that it does not possess at lower temperatures. It also becomes more toxic, although this is of little practical import, because the high temperature alone generally prevents human contact.
Magnesium compounds are typically white crystals. Most are soluble in water, providing the sour-tasting magnesium ion Mg2+. Small amounts of dissolved magnesium ion contributes to the tartness and taste of natural waters. Magnesium ion in large amounts is an ionic laxitive, and magnesium sulfate (Epsom salts) is sometimes used for this purpose. So-called "milk of magnesia" is a water suspension of one of the few insoluble magnesium compounds, magnesium hydroxide; the undissolved particles give rise to its appearance and name. Milk of magnesia is a mild base, and is commonly used as an antacid.
The name originates from the Greek word for a district in Thessaly called Magnesia. Joseph Black in Scotland recognized magnesium as being an element in 1755, Sir Humphry Davy electrolytically isolated pure magnesium metal in 1808 from a mix of magnesia and HgO, and A. A. B. Bussy prepared it in coherent form in 1831. Magnesium is the eighth most abundant element in the earth's crust. It is an alkaline earth metal and therefore does not occur uncombined with other elements. It is found in large deposits of magnesite, dolomite, and other minerals.
The United States has traditionally been the major world supplier of this metal, supplying 45% of world production even as recently as 1995. Today, the US market share is at 7%, with a single domestic producer left, US Magnesium, a company born from now-defunct Magcorp. As of 2005 China has taken over as the dominant supplier, pegged at 60% world market share, which increased from 4% in 1995. Unlike the above described electrolytic process, China is almost completely reliant on a different method of obtaining the metal from its ores, the silicothermic Pidgeon process (the reduction of the oxide at high temperatures with silicon).
Compounds in living organismsEdit
Magnesium ion is essential to the basic nucleic acid chemistry of life, and thus is essential to all cells of all known living organisms. Plants have an additional use for magnesium in that chlorophylls are magnesium-centered porphyrins. Many enzymes require the presence of magnesium ions for their catalytic action, especially enzymes utilizing ATP, or those which use other nucleotides to synthesize DNA and RNA.
Magnesium deficiency in humans was first described in the medical literature in 1934. The adult human daily nutritional requirement, which is affected by various factors including sex, weight and size, is 300-400 mg/day. Inadequate magnesium intake frequently causes muscle spasms, and has been associated with cardiovascular disease, diabetes, high blood pressure, anxiety disorders and osteoporosis. Acute deficiency (see hypomagnesemia) is rare, and is more common as a drug side effect (such as chronic alcohol or diuretic use) than from low food intake per se. The incidence of chronic deficiency resulting in less than optimal health, is debated.
The DRI upper tolerated limit for supplemental magnesium is 350 mg/day (calculated as mg of Mg elemental in the salt). (Supplements based on Amino Acid Chelates, Glycinate, Lysinate etc. are much better tolerated by the digestive system and do not have the side effects of the older compunds used.) The most common symptom of excess oral magnesium intake is diarrhea. Since the kidneys of adult humans excrete excess magnesium efficiently, oral magnesium poisoning in adults with normal renal function, is very rare. Infants, which have less ability to excrete excess magnesium even when healthy, should not be given magnesium supplements, except under a physician's care.
Magnesium salts (usually in the form of magnesium sulfate or chloride when given parenterally) are used therapeutically for a number of medical conditions, especially the hypertension of eclampsia. See epsom salts for a list of conditions which have been treated with supplemental magnesium ion. Magnesium is aborbed with reasonable efficiently (30% to 40%) by the body from any soluble magnesium salt, such as the chloride or citrate. Magnesium is similarly absorbed from epsom salts, although the sulfate in these salts adds to their laxative effect at higher doses. Magnesium absorption from the insoluble oxide and hydroxide salts (milk of magnesia) is erratic and of poorer efficiency, since it depends on the neutralization and solution of the salt by the acid of the stomach, which may not be (and usually is not), complete.
Food sources Edit
Green vegetables such as spinach provide magnesium because the center of the chlorophyll molecule contains magnesium. Nuts (especially almonds), seeds, and some whole grains are also good sources of magnesium.
Although magnesium is present in many foods, it usually occurs in dilute form. As with most nutrients, daily needs for magnesium are unlikely to be met from a single serving of any single food. Eating a wide variety of foods, including five servings of fruits and vegetables daily and plenty of whole grains, helps to ensure an adequate intake of magnesium.
Because magnesium readily dissolves in the water used to refine foods, and also in the water-rich parts of certain foods which are removed during refining, the magnesium content of many refined foods is low. Whole-wheat bread, for example, has twice as much magnesium as white bread because the magnesium-rich germ and bran are removed when white flour is processed. The table of food sources of magnesium suggests many dietary sources of magnesium.
Water can provide magnesium, but the amount varies according to the water supply. "Hard" water contains more magnesium than "soft" water. Dietary surveys do not estimate magnesium intake from water, which may lead to underestimating total magnesium intake and its variability.
Too much magnesium may make it difficult for the body to absorb calcium. Not enough magnesium can lead to hypomagnesemia as described above, with irregular heartbeats, high blood pressure (a sign in humans but not some experimental animals such as rodents), insomnia and muscle spasms. However, as noted, symptoms of low magnesium from pure dietary deficiency are thought to be rarely encountered.
Following are some foods and the amount of magnesium in them:
- spinach (1/2 cup) = 80 milligrams (mg)
- peanut butter (2 tablespoons) = 50 mg
- black-eyed peas (1/2 cup) = 45 mg
- milk: low fat (1 cup) = 40 mg
The U.S. RDA/RDV is 400 mg of magnesium.
As the metalEditMagnesium is the third most commonly used structural metal, following steel and aluminium.
Magnesium compounds, primarily magnesium oxide, are used mainly as refractory material in furnace linings for producing iron, steel, nonferrous metals, glass and cement. Magnesium oxide and other compounds also are used in agricultural, chemical and construction industries. As a metal, this element's principal use is as an alloying additive to aluminium with these aluminium-magnesium alloys being used mainly for beverage cans.
Magnesium, in its purest form, can be compared to aluminium, and is strong and light, so it is used in several high volume part manufacturing applications, including automotive and truck components. Specialty, high-grade car wheels of magnesium alloy are called "mag wheels". In 1957, a Corvette SS, designed for racing, was constructed, with completely magnesium body panels. Volkswagen has used magnesium in its engine components for many years. For a long time, Porsche used magnesium alloy for its engine blocks due to the weight advantage. However, there is renewed interest in magnesium engine blocks, as featured in the 2006 BMW 325i and 330i models. The award-winning BMW engine uses an aluminium alloy insert for the cylinder walls and cooling jackets surrounded by a high temperature magnesium alloy AJ62A. The application of magnesium AE44 alloy in the 2006 Corvette Z06 engine cradle has advanced the technology of designing robust automotive parts in magnesium. Both of these alloys are recent developments in high temperature low creep magnesium alloys. The general strategy for such alloys is to form intermetallic precipitates at the grain boundaries, for example by adding mischmetal or calcium. New alloy development and lower costs, which are becoming competitive to aluminium, will further the number of automotive applications.
In December 2005, for the first time on record, the automotive grade magnesium alloy price per cm³ dropped below the A380 aluminum alloy price per cm³. 
The second application field of magnesium is electronic devices. Due to low weight, good mechanical and electrical properties, magnesium is widely used for manufacturing of mobile phones, laptops, cameras, etc. housings and other electronic components.
Historically, magnesium was one of the main aerospace construction metals and was used for German military aircraft as early as World War I and extensively for German aircraft in World War II. The Germans coined the name 'Elektron' for magnesium alloy which is still used today. However, due to perceived hazards with magnesium parts in the event of fire, the application of magnesium in the commercial aerospace industry was generally restricted to engine related components. Currently the use of magnesium alloys in aerospace is increasing, mostly driven by the increasing importance of fuel economy in aerospace and the need to reduce weight. The development and testing of new magnesium alloys notably Elektron 21 which has successfully undergone extensive aerospace testing for suitability in both engine, internal and airframe components. European Community runs three R&D magnesium projects in Aerospace priority of Six Framework Program.
Other uses include:
- Removal of sulfur from iron and steel.
- Photoengraved plates in the printing industry.
- Combined in alloys, this metal is essential for airplane and missile construction.
- When used as an alloying agent, this metal improves the mechanical, fabrication and welding characteristics of aluminium.
- Additive agent for conventional propellants and used in producing nodular graphite in cast iron.
- Reducing agent for the production of pure uranium and other metals from their salts.
- Magnesium turnings or ribbon are used to prepare Grignard reagents, which are useful in organic synthesis
- Easily reacting with water, it can serve as a desiccant
- Magnesium is also flammable, burning at a temperature of approximately 2500 K (2200 °C, 4000 °F).
- The autoignition temperature of magnesium is approximately 744 K (473 °C, 883 °F).
- The extremely high temperature at which magnesium burns makes it a handy tool for starting emergency fires during outdoor recreation.
- Other uses include flashlight photography, flares, pyrotechnics, sparklers, and incendiary bombs.
In magnesium compoundsEdit
- Magnesium's hydroxide is used in milk of magnesia, its chloride and citrate used as oral magnesium supplements, and its sulfate (Epsom salts) for various purposes in medicine, and elsewhere.
- Dead-burned magnesite is used for refractory purposes such as brick and liners in furnaces and converters.
- Magnesium carbonate (MgCO3) powder is also used by athletes, such as gymnasts and weightlifters, to improve the grip on objects – the apparatus or lifting bar.
- Magnesium stearate is a slightly flammable white powder with lubricative properties. In pharmaceutical technology it is used in the manufacturing of tablets, to prevent the tablets from sticking to the equipment during the tablet compression process (i.e., when the tablet's substance is pressed into tablet form).
- The magnesium ion is necessary for all life (see magnesium in biological systems), so magnesium salts are an additive for foods, fertilizers (Mg is a component of chlorophyll), and culture media.
Magnesium metal and alloys are highly flammable in their pure form when molten, as a powder, or in ribbon form. Burning or molten magnesium metal reacts violently with water. Magnesium powder is an explosive hazard. One should wear safety glasses while working with magnesium. The bright white light (including ultraviolet) produced by burning magnesium can damage the eyes. Water should not be used to extinguish magnesium fires, because it can actually feed the fire, according to the reaction:
- Mg (s) + 2 H2O (g) → Mg(OH)2 (aq) + H2 (g)
- or in words:
- Magnesium (solid) + steam → Magnesium hydroxide (aqueous) + Hydrogen (gas)
Carbon dioxide fire extinguishers should not be used either, because magnesium can burn in carbon dioxide (forming magnesium oxide, MgO, and carbon). A Class D dry chemical fire extinguisher should be used if available, or else the fire should be covered with sand or magnesium foundry flux. An easy way to put out small metal fires is to place a polyethene bag filled with dry sand on top of the fire. The heat of the fire will melt the bag and the sand will flow out onto the fire.
See also Edit
- El Mutún in Bolivia, where 70% of the world's magnesium and iron are located
- Magnesium minerals
- Magnesium compounds
- ↑ Vardi, Nathan (07.22.02). Man With Many Enemies. Forbes.com. Retrieved on 2006-06-26.
- ↑ Chemistry : Periodic Table : magnesium : chemical reaction data. webelements.com. Retrieved on 2006-06-26.
- ↑ Demo Lab: Reaction Of Magnesium Metal With Carbon Dioxide. Retrieved on 2006-06-26.
- WebElements.com – Magnesium
- Magnesium Deficiency
- The Magnesium Website
- Dietary Reference Intake
- Computational Chemistry Wiki
|This page uses content from Wikipedia. The original article was at Magnesium. The list of authors can be seen in the page history. As with Chemistry, the text of Wikipedia is available under the GNU Free Documentation License.|