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12 sodium ← magnesium → aluminium Be ↑ Mg ↓ Ca Periodic table
General
Name, symbol, number	magnesium, Mg, 12
Element category	alkaline earth metal
Group, period, block	2, 3, s
Appearance	silvery white solid at room temp
Standard atomic weight	24.3050(6) g·mol−1
Electron configuration	[Ne] 3s2
Electrons per shell	2, 8, 2 (Image)
Physical properties
Phase	solid
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 (1091 °C, 1994 °F)
Heat of fusion	8.48 kJ·mol−1
Heat of vaporization	128 kJ·mol−1
Specific heat capacity	(25 °C) 24.869 J·mol−1·K−1
Vapor pressure P/Pa 1 10 100 1 k 10 k 100 k at T/K 701 773 861 971 1132 1361
Atomic properties
Crystal structure	hexagonal
Oxidation states	2, 1 [1] (strongly basic oxide)
Electronegativity	1.31 (Pauling scale)
Ionization energies (more)	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
Miscellaneous
Magnetic ordering	paramagnetic
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) 4940 m·s−1
Young's modulus	45 GPa
Shear modulus	17 GPa
Bulk modulus	45 GPa
Poisson ratio	0.290
Mohs hardness	2.5
Brinell hardness	260 MPa
CAS registry number	7439-95-4
Most stable isotopes
Main article: Isotopes of magnesium iso NA half-life DM DE (MeV) DP 24Mg 78.99% 24Mg is stable with 12 neutrons 25Mg 10% 25Mg is stable with 13 neutrons 26Mg 11.01% 26Mg is stable with 14 neutrons
References
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Magnesium (pronounced /mæɡˈniːziəm/) is a chemical element with the symbol Mg, atomic number 12, atomic weight 24.3050 and common oxidation number +2.

Magnesium, an alkaline earth metal, is the ninth most abundant element in the universe by mass.^[1] The commonness of magnesium is related to the fact that it is easily built up in supernova stars from a sequential addition of three helium nuclei to carbon (which in turn is made from a single reaction between three helium nuclei at once). Magnesium constitutes about 2% of the Earth's crust by mass, which makes it the eighth most abundant element in the crust.^[2] Magnesium ion's high solubility in water helps ensure that it is the third most abundant element dissolved in seawater.^[3]

Magnesium is the 11th most abundant element by mass in the human body; its ions are essential to all living cells, where they play a major role in manipulating important biological polyphosphate compounds like ATP, DNA, and RNA. Hundreds of enzymes thus require magnesium ion in order to function. Magnesium is also the metallic ion at the center of chlorophyll, and is thus a common additive to fertilizers.^[4] Magnesium compounds are used medicinally as common laxatives, antacids (i.e., milk of magnesia), and in a number of situations where stabalization of abnormal nerve excitation and blood vessel spasm is required (i.e., to treat eclampsia). Magnesium ions are sour to the taste, and in low concentrations help to impart a natural tartness to fresh mineral waters.

The free element (metal) is not found naturally on Earth, since it is highly reactive (though once produced, is coated in a thin layer of oxide which partly masks this reactivity). The free metal burns with a characteristic brilliant white light, making it a useful ingredient in flares. The metal is now mainly obtained by electrolysis of magnesium salts obtained from brine. Commercially, the chief use for the metal is as an alloying agent to make aluminium-magnesium alloys, sometimes called "magnalium" or "magnelium". Since magnesium is less dense than aluminum, these alloys are prized for their relative lightness and strength.

Contents 1 Notable characteristics 2 Applications 2.1 As the metal 2.2 In magnesium compounds 3 History 4 Sources 4.1 Magnesium from sea water 5 Biological role 6 Isotopes 7 Precautions 8 See also 9 References 10 External links

Notable characteristics

Elemental magnesium is a fairly strong, silvery-white, light-weight metal (two thirds the density of aluminium). It tarnishes slightly 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 unnoticeably 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 also reacts exothermically with most acids, such as hydrochloric acid (HCl). As with aluminum, zinc and many other metals, the reaction with hydrochloric acid produces the chloride of the metal and releases hydrogen gas.

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). On burning in air, magnesium produces a brilliant white light. Thus magnesium powder (flash powder) was used as a source of illumination in the early days of photography. Later, magnesium ribbon was used in electrically ignited flash bulbs. Magnesium powder is used in the manufacture of fireworks and marine flares where a brilliant white light is required. Flame temperatures of magnesium and magnesium alloys can reach 1,371 °C (2,500 °F), although flame height above the burning metal is usually less than 300 mm (12 in).^[5]

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 laxative, 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 commonly used as an antacid.

Applications

As the metal

Magnesium is the third most commonly used structural metal, following steel and aluminium.

Magnesium compounds, primarily magnesium oxide (MgO), 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 with 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 magnesium body panels. Porsche's all-out quest to decrease the weight of their racecars led to the use of magnesium frames in the famous 917/053 which won Le Mans in 1971, and still holds the absolute distance record. The 917/30 Can-Am car also featured a magnesium spaceframe, helping it to make the most of its prodigious 1100-1500hp. 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. There is renewed interest in magnesium engine blocks, as featured in the 2006 BMW 325i and 330i models. The 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.^[6] New alloy development and lower costs, which are becoming competitive to aluminium, will further the number of automotive applications.

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, laptop computers, cameras, 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. 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 and the need to reduce weight. The development and testing of new magnesium alloys continues, notably Elektron 21 which has successfully undergone extensive aerospace testing for suitability in engine, internal and airframe components. The European Community runs three R&D magnesium projects in the Aerospace priority of Six Framework Program.

• Incendiary use: Magnesium is flammable, burning at a temperature of approximately 2500 K (2200 °C, 4000 °F), and the autoignition temperature of magnesium is approximately 744 K (473 °C, 883 °F) in air. The extremely high temperature at which magnesium burns makes it a handy tool for starting emergency fires during outdoor recreation. Other related uses include flashlight photography, flares, pyrotechnics, fireworks sparklers, and incendiary bombs.

Magnesium is also used:

• To remove sulfur from iron and steel.
• To refine titanium in the Kroll process.
• To photoengrave plates in the printing industry.
• To combine in alloys, where this metal is essential for airplane and missile construction.
• In the form of turnings or ribbons, to prepare Grignard reagents, which are useful in organic synthesis.
• As an alloying agent, improving the mechanical, fabrication and welding characteristics of aluminium.
• As an additive agent in conventional propellants and the production of nodular graphite in cast iron.
• As a reducing agent for the production of uranium and other metals from their salts.
• As a desiccant, since it easily reacts with water.
• As a sacrificial (galvanic) anode to protect underground tanks, pipelines, buried structures, and water heaters.

In magnesium compounds

• The magnesium ion is necessary for all life (see magnesium in biology), so magnesium salts are an additive for foods, fertilizers (Mg is a component of chlorophyll), and culture media.
• Magnesium hydroxide is used in milk of magnesia, its chloride, oxide, gluconate, malate, orotate and citrate used as oral magnesium supplements, and its sulfate (Epsom salts) for various purposes in medicine, and elsewhere (see the article for more). Oral magnesium supplements have been claimed to be therapeutic for some individuals who suffer from Restless Leg Syndrome (RLS).^{[citation needed]}
• Magnesium borate, magnesium salicylate and magnesium sulfate are used as antiseptics.
• Magnesium bromide is used as a mild sedative (this action is due to the bromide, not the magnesium).
• 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).
• Magnesium sulfite is used in the manufacture of paper (sulfite process).
• Magnesium phosphate is used to fireproof wood for construction.
• Magnesium hexafluorosilicate is used in mothproofing of textiles.

History

The name originates from the Greek word for a district in Thessaly called Magnesia. It is related to magnetite and manganese, which also originated from this area, and required differentiation as separate substances. See manganese for this history.

Magnesium is the seventh most abundant element in the Earth's crust by mass and eighth by molarity.^[2] It is found in large deposits of magnesite, dolomite, and other minerals, and in mineral waters, where magnesium ion is soluble. In 1618 a farmer at Epsom in England attempted to give his cows water from a well. They refused to drink because of the water's bitter taste. However the farmer noticed that the water seemed to heal scratches and rashes. The fame of Epsom salts spread. Eventually they were recognized to be hydrated magnesium sulfate, MgSO₄.

The metal itself was first produced in England by Sir Humphry Davy in 1808 using electrolysis of a mixture of magnesia and HgO. A. A. B. Bussy prepared it in coherent form in 1831. Davy's first suggestion for a name was magnium, but the name magnesium is now used.

Sources

Although magnesium is found in over 60 minerals, only dolomite, magnesite, brucite, carnallite, talc, and olivine are of commercial importance.

In the United States this metal is principally obtained by electrolysis of fused magnesium chloride from brines, wells, and sea water:

cathode: Mg²⁺ + 2 e^- → Mg

anode: 2 Cl^- → Cl_{2 (gas)} + 2 e^-

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.^[7] 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).

Magnesium from sea water

The Mg²⁺ cation is the second most abundant cation in seawater (occurring at about 12% of the mass of sodium there), which makes seawater and sea-salt an attractive commercial source of Mg. To extract the magnesium, calcium hydroxide is added to sea water to form magnesium hydroxide precipitate.

MgCl₂ + Ca(OH)₂ → Mg(OH)₂ + CaCl₂

Magnesium hydroxide is insoluble in water so it can be filtered out, and reacted with hydrochloric acid to obtain concentrated magnesium chloride.

Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O

From magnesium chloride, electrolysis produces magnesium.

Biological role

Due to the important interaction between phosphate and magnesium ions, magnesium ions are essential to the basic nucleic acid chemistry of life, and thus are essential to all cells of all known living organisms. Over 300 enzymes require the presence of magnesium ions for their catalytic action, including all enzymes utilizing or synthesizing ATP, or those which use other nucleotides to synthesize DNA and RNA. ATP exists in cells normally as a chelate of ATP and a magnesium ion.

Plants have an additional use for magnesium in that chlorophylls are magnesium-centered porphyrins.Magnesium deficiency in plants causes late-season yellowing between leaf veins, especially in older leaves, and can be corrected by applying epsom salts (which is rapidly leached), or else crushed dolomitic limestone to the soil.

Magnesium is a vital component of a healthy human diet. Human magnesium deficiency (including conditions which show few overt symptoms) is relatively common, with only 32% of the United States meeting the RDA-DRI,^[8] and has been implicated in the development of a number of human illnesses such as asthma, osteoporosis, and ADHD.^[9]

Adult human bodies contain about 24 grams of magnesium, with 60% in the skeleton, 39% intracellular (20% in skeletal muscle), and 1% extracellular. Serum levels are typically 0.7 – 1.0 mmol/L. Serum magnesium levels may appear normal even in cases of underlying intracellular deficiency, although no known mechanism maintains a homeostatic level in the blood other than renal excretion of high blood levels. Intracellular magnesium is correlated with intracellular potassium. Magnesium is absorbed in the gastrointestinal tract, with more absorbed when status is lower. In humans, magnesium appears to facilitate calcium absorption. Low and high protein intake inhibit magnesium absorption, and other factors such as phosphate, phytate, and fat affect absorption. Absorbed dietary magnesium is largely excreted through the urine, although most magnesium "administered orally" is excreted through the feces.^[10] Magnesium status may be assessed roughly through serum and erythrocyte Mg concentrations and urinary and fecal excretion, but intravenous magnesium loading tests are likely the most accurate and practical in most people.^[11] In these tests, magnesium is injected intravenously; a retention of 20% or more indicates deficiency.^[12] Other nutrient deficiencies are identified through biomarkers, but none are established for magnesium.^[13]

Spices, nuts, cereals, coffee, cocoa, tea, and vegetables (especially green leafy ones) are rich sources of magnesium. Observations of reduced dietary magnesium intake in modern Western countries as compared to earlier generations may be related to food refining and modern fertilizers which contain no magnesium.^[10]

There are a number of magnesium dietary supplements available. Magnesium oxide, one of the most common because it has a high magnesium content per weight, has been reported to be the least bioavailable.^[14][15] Magnesium citrate has been reported as more bioavailable than oxide or amino-acid chelate (glycinate) forms.^[16]

Excess magnesium in the blood is freely filtered at the kidneys, and for this reason it is difficult to overdose on magnesium from dietary sources alone.^[9] With supplements, overdose is possible, however, particularly in people with poor renal function; occasionally, with use of high cathartic doses of magnesium salts, severe hypermagnesemia has been reported to occur even without renal dysfunction.^[17] Alcoholism can produce a magnesium deficiency which is easily reversed by oral or parenteral administration, depending on the degree of deficiency.^[18]

Isotopes

Magnesium has three stable isotopes: ²⁴Mg, ²⁵Mg, ²⁶Mg. All are present in significant amounts (see table of isotopes above). About 79% of Mg is ²⁴Mg. The isotope ²⁸Mg is radioactive and in the 1950s to 1970s was made commercially by several nuclear power plants for use in scientific experiments. This isotope has a relatively short half-life (21 hours) and so its use was limited by shipping times.

²⁶Mg 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 its early history.

It is conventional to plot ²⁶Mg/²⁴Mg against an Al/Mg ratio. In an isochron dating 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.

Precautions

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. Wear safety glasses when working with powdered magnesium. Use safety glasses that include a heavy U.V. filter, similar to welding eye protection, when burning magnesium. The bright white light (including ultraviolet) produced by burning magnesium can permanently damage the retinas of the eyes, similar to welding arc burns.^[19]

Do not use water to extinguish magnesium fires because it can produce hydrogen which will feed the fire. The reaction is:^[20]

Mg _(s) + 2 H₂O _(g) → Mg(OH)_{2 (s)} + H_{2 (g)}

or in words:

Magnesium _(solid) + steam → Magnesium hydroxide _(solid) + Hydrogen _(gas)

Carbon dioxide fire extinguishers also are not recommended, because magnesium can burn in carbon dioxide (forming magnesium oxide, and carbon).^[21]

Burning magnesium can be extinguished using a Class D dry chemical fire extinguisher, or the fire can be covered with sand or magnesium foundry flux. Small magnesium metal fires may be conveniently extinguished by placing a polyethylene bag filled with dry sand atop the fire. The heat of the fire melts the bag, releasing the sand onto the fire.

See also

• Magnesium minerals
• Magnesium compounds

References

External links

Wikimedia Commons has media related to: Magnesium

Look up magnesium in Wiktionary, the free dictionary.

• WebElements.com – Magnesium
• Online Resource for industry professionals - Magnesium.com
• The Magnesium Website – Includes full text papers and textbook chapters by leading magnesium authorities Mildred Seelig, Jean Durlach, Burton M. Altura and Bella T. Altura. Links to over 300 articles discussing magnesium and magnesium deficiency.
• Magnesium in Health - Mg12.info

v • d • e   Magnesium compounds MgB2 · MgBr2 · MgCO3 · MgC6H6O7 · MgC14H10O4 · MgCl2 · Mg(ClO4)2 · MgF2 · MgH2 · MgI2 · Mg(NO3)2 · MgO · MgO2 · Mg(OH)2 · MgS · MgSO3 · MgSO4 · MgU2O7 · Mg2Si · Mg3(PO4)2

v • d • e Periodic table
H																															He
Li	Be																									B	C	N	O	F	Ne
Na	Mg																									Al	Si	P	S	Cl	Ar
K	Ca															Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn	Ga	Ge	As	Se	Br	Kr
Rb	Sr															Y	Zr	Nb	Mo	Tc	Ru	Rh	Pd	Ag	Cd	In	Sn	Sb	Te	I	Xe
Cs	Ba	La	Ce	Pr	Nd	Pm	Sm	Eu	Gd	Tb	Dy	Ho	Er	Tm	Yb	Lu	Hf	Ta	W	Re	Os	Ir	Pt	Au	Hg	Tl	Pb	Bi	Po	At	Rn
Fr	Ra	Ac	Th	Pa	U	Np	Pu	Am	Cm	Bk	Cf	Es	Fm	Md	No	Lr	Rf	Db	Sg	Bh	Hs	Mt	Ds	Rg	Uub	Uut	Uuq	Uup	Uuh	Uus	Uuo
Alkali metals Alkaline earth metals Lanthanoids Actinoids Transition metals Other metals Metalloids Other nonmetals Halogens Noble gases