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Magnesium, ₁₂Mg

Magnesium
Pronunciation	/mæɡˈniːziəm/ ​(mag-NEE-zee-əm)
Appearance	shiny grey solid
Standard atomic weight Ar°(Mg)
	24.305±0.002 (abridged)
Magnesium in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson Be ↑ Mg ↓ Ca sodium ← magnesium → aluminium
Atomic number (Z)	12
Group	group 2 (alkaline earth metals)
Period	period 3
Block	s-block
Electron configuration	[Ne] 3s
Electrons per shell	2, 8, 2
Physical properties
Phase at STP	solid
Melting point	923 K ​(650 °C, ​1202 °F)
Boiling point	1363 K ​(1091 °C, ​1994 °F)
Density (at 20° C)	1.737 g/cm
when liquid (at m.p.)	1.584 g/cm
Heat of fusion	8.48 kJ/mol
Heat of vaporization	128 kJ/mol
Molar heat capacity	24.869 J/(mol·K)
Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 701 773 861 971 1132 1361
Atomic properties
Oxidation states	common: +2 0, +1
Electronegativity	Pauling scale: 1.31
Ionization energies	1st: 737.7 kJ/mol 2nd: 1450.7 kJ/mol 3rd: 7732.7 kJ/mol (more)
Atomic radius	empirical: 160 pm
Covalent radius	141±7 pm
Van der Waals radius	173 pm
Other properties
Natural occurrence	primordial
Crystal structure	​hexagonal close-packed (hcp) (hP2)
Lattice constants	a = 320.91 pm c = 521.03 pm (at 20 °C)
Thermal expansion	25.91×10/K (at 20 °C)
Thermal conductivity	156 W/(m⋅K)
Electrical resistivity	43.9 nΩ⋅m (at 20 °C)
Magnetic ordering	paramagnetic
Molar magnetic susceptibility	+13.1×10 cm/mol (298 K)
Young's modulus	45 GPa
Shear modulus	17 GPa
Bulk modulus	35.4 GPa
Speed of sound thin rod	4940 m/s (at r.t.) (annealed)
Poisson ratio	0.290
Mohs hardness	1–2.5
Brinell hardness	44–260 MPa
CAS Number	7439-95-4
History
Naming	after Magnesia, Greece
Discovery	Joseph Black (1755)
First isolation	Humphry Davy (1808)
Isotopes of magnesium v e
Main isotopes Decay abun­dance half-life (t1/2) mode pro­duct Mg 79% stable Mg 10% stable Mg 11% stable
Category: Magnesium view talk edit | references

Magnesium (Template:PronEng) 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. It constitutes about 2% of the Earth's crust by mass, and it is the third most abundant element dissolved in seawater.

Magnesium is the 11th most abundant element by mass in the human body; its ions are essential to all living cells. The free element (metal) is not found in nature. Once produced from magnesium salts, it is now mainly obtained by electrolysis of brine and is used as an alloying agent to make aluminium-magnesium alloys, sometimes called "magnalium" or "magnelium".

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.

Magnesium compounds are typically white crystals. Most are soluble in water, providing the sour-tasting magnesium ion Mg. 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, 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. 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 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).
• 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. 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: Mg + 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. 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 Mg 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 carbonate is added to sea water to form magnesium carbonate precipitate.

MgCl₂ + CaCO₃ → MgCO₃ + CaCl₂

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

MgCO₃ + 2HCl → MgCl₂ + CO₂ + H₂O

From magnesium chloride, electrolysis produces magnesium.

Biological role

Magnesium ions are essential to the basic nucleic acid chemistry of life, and thus are 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 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, and plays a part in over 300 enzymes. Human magnesium deficiency is relatively common, with only 32% of the United States meeting the RDA-DRI, and has been implicated in the development of a number of human illnesses such as asthma, osteoporosis, and ADHD. Adult human bodies have 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. 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. It is largely excreted through the feces.

Spices, nuts, cereals, coffee, cocoa, tea, and vegetables (especially green leafy ones) are rich sources of magnesium. Refining of food can reduce magnesium substantially, however, and fertilizers use less magnesium. This has led to observations of reduced dietary magnesium intake as compared to earlier generations.

There are a number of magnesium supplements available. Magnesium oxide, one of the most common, has been reported as the least bioavailable. Magnesium citrate has been reported as more bioavailable than oxide or amino-acid chelate (glycinate) forms. Magnesium improves the absorption of calcium into

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. With supplements overdose is possible, however, particularly in people with poor renal function, but severe hypermagnesemia can also occur without renal dysfunction. Alcoholism can produce a magnesium deficiency which is easily reversed by oral or parenteral administration, depending on the degree of deficiency.

Isotopes

Magnesium has three stable isotopes: Mg, Mg, Mg. All are present in significant amounts (see table of isotopes above). About 79% of Mg is Mg. The isotope Mg 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.

Mg has found application in isotopic geology, similar to that of aluminium. Mg is a radiogenic daughter product of Al, which has a half-life of 717,000 years. Large enrichments of stable Mg have been observed in the Ca-Al-rich inclusions of some carbonaceous chondrite meteorites. The anomalous abundance of Mg is attributed to the decay of its parent Al in the inclusions. Therefore, the meteorite must have formed in the solar nebula before the Al 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 Mg/Mg against an Al/Mg ratio. In an isochron dating plot, the Al/Mg ratio plotted isAl/Mg. The slope of the isochron has no age significance, but indicates the initial Al/Al 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. One should wear safety glasses while working with magnesium, and if burning it, these should include a heavy U.V. filter, similar to welding eye protection. The bright white light (including ultraviolet) produced by burning magnesium can permanently damage the retinas of the eyes, similar to welding arc burns.

Water should not be used to extinguish magnesium fires, because it can produce hydrogen which will feed the fire, according to the reaction:

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 should not be used, because magnesium can burn in carbon dioxide (forming magnesium oxide, MgO, and carbon). A Class D dry chemical fire extinguisher should be used, or the fire can be covered with sand or magnesium foundry flux. An easy way to extinguish small metal fires is to place a polyethylene bag filled with dry sand atop the fire. The heat of the fire will melt the bag, releasing the sand onto the fire.

See also

• Magnesium minerals
• Magnesium compounds

References

1. "Standard Atomic Weights: Magnesium". CIAAW. 2011.
2. Prohaska, Thomas; Irrgeher, Johanna; Benefield, Jacqueline; Böhlke, John K.; Chesson, Lesley A.; Coplen, Tyler B.; Ding, Tiping; Dunn, Philip J. H.; Gröning, Manfred; Holden, Norman E.; Meijer, Harro A. J. (2022-05-04). "Standard atomic weights of the elements 2021 (IUPAC Technical Report)". Pure and Applied Chemistry. doi:10.1515/pac-2019-0603. ISSN 1365-3075.
3. ^ Arblaster, John W. (2018). Selected Values of the Crystallographic Properties of Elements. Materials Park, Ohio: ASM International. ISBN 978-1-62708-155-9.
4. Rumble, p. 4.61
5. Mg(0) has been synthesized in a compound containing a Na₂Mg₂ cluster coordinated to a bulky organic ligand; see Rösch, B.; Gentner, T. X.; Eyselein, J.; Langer, J.; Elsen, H.; Li, W.; Harder, S. (2021). "Strongly reducing magnesium(0) complexes". Nature. 592 (7856): 717–721. Bibcode:2021Natur.592..717R. doi:10.1038/s41586-021-03401-w. PMID 33911274. S2CID 233447380
6. Bernath, P. F.; Black, J. H. & Brault, J. W. (1985). "The spectrum of magnesium hydride" (PDF). Astrophysical Journal. 298: 375. Bibcode:1985ApJ...298..375B. doi:10.1086/163620.. See also Low valent magnesium compounds.
7. Rumble, p. 12.137
8. Rumble, p. 12.28
9. Rumble, p. 4.70
10. Gschneider, K. A. (1964). Physical Properties and Interrelationships of Metallic and Semimetallic Elements. Solid State Physics. Vol. 16. p. 308. doi:10.1016/S0081-1947(08)60518-4. ISBN 9780126077162.
11. ^ Rumble, p. 4.19
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14. ^ "Abundance and form of the most abundant elements in Earth's continental crust" (PDF). Retrieved 2008-02-15. {{cite journal}}: Cite journal requires |journal= (help)
15. Alan A. Luo and Bob R. Powell (2001). "Tensile and Compressive Creep of Magnesium-Aluminum-Calcium Based Alloys" (PDF). Materials & Processes Laboratory, General Motors Research & Development Center. Retrieved 2007-08-21. {{cite journal}}: Cite journal requires |journal= (help)
16. Vardi, Nathan (February 22 2007). "Man With Many Enemies". Forbes.com. Retrieved 2006-06-26. {{cite web}}: Check date values in: |date= (help)
17. "Lack Energy? Maybe It's Your Magnesium Level". United States Department of Agriculture. Retrieved 2008-09-18. Last paragraph
18. ^ University of Maryland Medical Center. Magnesium
19. ^ Wester PO (1987). "Magnesium". Am. J. Clin. Nutr. 45 (5 Suppl): 1305–12. PMID 3578120. {{cite journal}}: Unknown parameter |month= ignored (help)
20. Firoz M, Graber M (2001). "Bioavailability of US commercial magnesium preparations". Magnes Res. 14 (4): 257–62. PMID 11794633. {{cite journal}}: Unknown parameter |month= ignored (help)
21. Lindberg JS, Zobitz MM, Poindexter JR, Pak CY (1990). "Magnesium bioavailability from magnesium citrate and magnesium oxide". J Am Coll Nutr. 9 (1): 48–55. PMID 2407766. {{cite journal}}: Unknown parameter |month= ignored (help)CS1 maint: multiple names: authors list (link)
22. Walker AF, Marakis G, Christie S, Byng M (2003). "Mg citrate found more bioavailable than other Mg preparations in a randomised, double-blind study". Magnes Res. 16 (3): 183–91. PMID 14596323. {{cite journal}}: Unknown parameter |month= ignored (help)CS1 maint: multiple names: authors list (link)
23. Kontani M, Hara A, Ohta S, Ikeda T (2005). "Hypermagnesemia induced by massive cathartic ingestion in an elderly woman without pre-existing renal dysfunction". Intern. Med. 44 (5): 448–52. doi:10.2169/internalmedicine.44.448. PMID 15942092.{{cite journal}}: CS1 maint: multiple names: authors list (link)
24. AJ Giannini. Drugs of Abuse--Second Edition. Los Angeles, Physicians Management Information Co., 1997.
25. "Science Safety: Chapter 8". Government of Manitoba. Retrieved 2007-08-21.
26. "Chemistry : Periodic Table : magnesium : chemical reaction data". webelements.com. Retrieved 2006-06-26.
27. "Demo Lab: Reaction Of Magnesium Metal With Carbon Dioxide". Retrieved 2006-06-26.

External links

• 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

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• Chemical elements
• Alkaline earth metals
• Dietary minerals
• Food additives
• Pyrotechnic fuels
• Magnesium
• Desiccants