Magnesium

Magnesium
12Mg
Be ↑ Mg ↓ Ca sodium ← magnesium → aluminium
Magnesium in the periodic table
Appearance
shiny grey solid Spectral lines of Magnesium
General properties
Name, symbol, number	magnesium, Mg, 12
Pronunciation	/mæɡˈniːziəm/ mag-NEE-zee-əm
Element category	alkaline earth metal
Group, period, block	2 (alkaline earth metals), 3, s
Standard atomic weight	24.305(1)
Electron configuration	[Ne] 3s2 2, 8, 2
History
Naming	after Magnesia, Greece
Discovery	Joseph Black (1755)
First isolation	Humphry Davy (1808)
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
Molar heat capacity	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
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	160 pm
Covalent radius	141±7 pm
Van der Waals radius	173 pm
Miscellanea
Crystal structure	hexagonal close-packed
Magnetic ordering	paramagnetic
Electrical resistivity	(20 °C) 43.9 nΩ·m
Thermal conductivity	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 10.00% 25Mg is stable with 13 neutrons 26Mg 11.01% 26Mg is stable with 14 neutrons
v t e

Magnesium is a chemical element with the symbol Mg and atomic number 12. Its common oxidation number is +2. It is an alkaline earth metal and the eighth-most-abundant element in the Earth's crust^[2] and ninth in the known universe as a whole.^[3][4] Magnesium is the fourth-most-common element in the Earth as a whole (behind iron, oxygen and silicon), making up 13% of the planet's mass and a large fraction of the planet's mantle. The relative abundance of magnesium is related to the fact that it easily builds up in supernova stars from a sequential addition of three helium nuclei to carbon (which in turn is made from three helium nuclei).^{[citation needed]} Due to magnesium ion's high solubility in water, it is the third-most-abundant element dissolved in seawater.^[5] Magnesium is produced in stars larger than 3 solar masses by fusing helium and neon in the alpha process at temperatures above 600 megakelvins.^{[citation needed]}

The free element (metal) is not found naturally on Earth, as it is highly reactive (though once produced, it is coated in a thin layer of oxide (see passivation), 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 obtained mainly by electrolysis of magnesium salts obtained from brine. In commerce, 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 aluminium, these alloys are prized for their relative lightness and strength.

In human biology, magnesium is the eleventh-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 ions to function. Magnesium compounds are used medicinally as common laxatives, antacids (e.g., milk of magnesia), and in a number of situations where stabilization of abnormal nerve excitation and blood vessel spasm is required (e.g., to treat eclampsia). Magnesium ions are sour to the taste, and in low concentrations they help to impart a natural tartness to fresh mineral waters.

In vegetation, magnesium is the metallic ion at the center of chlorophyll, and is, thus, a common additive to fertilizers.^[6]

Characteristics

Physical properties

Elemental magnesium is a rather strong, silvery-white, light-weight metal (two-thirds the density of aluminium). It tarnishes slightly when exposed to air, although, unlike the alkali metals, an oxygen-free environment is unnecessary for storage because magnesium is protected by a thin layer of oxide that is fairly impermeable and difficult 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 submerged in water, hydrogen bubbles almost unnoticeably begin to form on the surface of the metal—though, if powdered, it reacts much more rapidly. The reaction occurs faster with higher temperatures (see precautions). Magnesium's ability to react with water can be harnessed to produce energy and run a magnesium-based engine. Magnesium also reacts exothermically with most acids, such as hydrochloric acid (HCl). As with aluminium, zinc, and many other metals, the reaction with HCl produces the chloride of the metal and releases hydrogen gas.

Chemical properties

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 nitrogen (forming magnesium nitride), carbon dioxide (forming magnesium oxide, and carbon) and water (forming magnesium oxide and hydrogen). This property was used in incendiary weapons used in the firebombing of cities in World War II, the only practical civil defense being to smother a burning flare under dry sand to exclude the atmosphere. On burning in air, magnesium produces a brilliant-white light that includes strong ultraviolet. 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 flashbulbs. 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 3,100 °C (3,370 K; 5,610 °F),^[7] although flame height above the burning metal is usually less than 300 mm (12 in).^[8] Magnesium may be used as an ignition source for thermite, a mixture of aluminium and iron oxide powder that is otherwise difficult to ignite. Those properties are due to magnesium's high specific heat, the fourth-highest specific heat among the metals.

Magnesium compounds are typically white crystals. Most are soluble in water, providing the sour-tasting magnesium ion Mg²⁺. Small amounts of dissolved magnesium ion contribute to the tartness and taste of natural waters. Magnesium ion in large amounts is an ionic laxative, and magnesium sulfate (common name: Epsom salt) 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, which has some laxative side-effect.

Occurrence

Magnesium is the eighth-most-abundant element in the Earth's crust by mass and tied in seventh place with iron in terms of molarity.^[2] It is found in large deposits of magnesite, dolomite, and other minerals, and in mineral waters, where magnesium ion is soluble.

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

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 seawater to form magnesium hydroxide precipitate.

MgCl
2 + Ca(OH)
2 → Mg(OH)
2 + CaCl
2

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

Mg(OH)
2 + 2 HCl → MgCl
2 + 2 H
2O

From magnesium chloride, electrolysis produces magnesium.

Forms

Alloy

As of 2013, magnesium alloy consumption is less than a million tons per year, compared with 50 million tons of aluminum alloys. Its use has been historically limited by its tendency to corrode, high-temperature creep, and flammability.^[9]

Research and development eliminated magnesium's tendency toward high-temperature creep by inclusion of scandium and gadolinium. Flammability was greatly reduced by introducing a small amount of calcium into the mix.^[9]

The presence of iron, nickel, copper, and cobalt strongly activates corrosion. This is due to their low solid solubility limits (above a very small percentage, they precipitate out as intermetallic compounds) and because they behave as active cathodic sites that reduce water and cause the loss of magnesium.^[9]

Reducing the quantity of these metals improves corrosion resistance. Sufficient manganese overcomes the corrosive effects of iron. This requires precise control over composition, increasing costs.^[9]

Adding a cathodic poison captures atomic hydrogen within the structure of a metal. This prevents the formation of free hydrogen gas, which is required for corrosive chemical processes. The addition of about one-third of a percent of arsenic reduces its corrosion rate in a salt solution by a factor of nearly ten.^[9][10]

Compounds

Magnesium forms a variety of industrially and biologically important compounds, including magnesium oxide, various salts, and others.

Isotopes

Magnesium has three stable isotopes: ²⁴Mg, ²⁵Mg and 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.

Production

Country	2011 production (tonnes)[11]
China	661,000
U.S. (Capacity. Production figures withheld to avoid disclosing company proprietary data.)	63,500
Russia	37,000
Israel	30,000
Kazakhstan	21,000
Brazil	16,000
Ukraine	2,000
Serbia	1,500

China is the dominant supplier of magnesium, with approximately 80% of the world market share. China is almost completely reliant on the silicothermic Pidgeon process (the reduction of the oxide at high temperatures with silicon) to obtain the metal.^[12]

In the United States, magnesium is obtained principally by electrolysis of fused magnesium chloride from brines, wells, and sea water. At the cathode, the Mg2+
ion is reduced by two electrons to magnesium metal:

Mg2+
+ 2 e− → Mg

At the anode, each pair of Cl−
ions is oxidized to chlorine gas, releasing two electrons to complete the circuit:

2 Cl−
→ Cl
2 (g) + 2 e−

A new process, solid oxide membrane technology, involves the electrolytic reduction of MgO. At the cathode, Mg2+
ion is reduced by two electrons to magnesium metal. The electrolyte is Yttria-stabilized zirconia(YSZ). The anode is a liquid metal. At the YSZ/liquid metal anode O2−
is oxidized. A layer of graphite borders the liquid metal anode, and at this interface carbon and oxygen react to form carbon monoxide. When silver is used as the liquid metal anode, there is no reductant carbon or hydrogen needed, and only oxygen gas is evolved at the anode.^[13] It has been reported that this method provides a 40% reduction in cost per pound over the electrolytic reduction method.^[14] This method is more environmentally sound than others because there is much less carbon dioxide emitted.

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 Renco Group company in Utah born from now-defunct Magcorp.^[15]

History

The name magnesium 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.

In 1618, a farmer at Epsom in England attempted to give his cows water from a well there. The cows refused to drink because of the water's bitter taste, but the farmer noticed that the water seemed to heal scratches and rashes. The substance became known as Epsom salts and its fame spread. It was eventually recognized as hydrated magnesium sulfate, MgSO₄·7 H₂O.

The metal itself was first produced by Sir Humphry Davy in England in 1808. He used electrolysis on a mixture of magnesia and mercuric oxide.^[16] Antoine Bussy prepared it in coherent form in 1831. Davy's first suggestion for a name was magnium,^[16] but the name magnesium is now used.

Applications

As metal

Magnesium is the third-most-commonly-used structural metal, following iron and aluminium. It has been called the lightest useful metal by The Periodic Table of Videos.^[17]

The main applications of magnesium are, in order: component of aluminium alloys, in die-casting (alloyed with zinc),^[18] to remove sulfur in the production of iron and steel, the production of titanium in the Kroll process.^[19]

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", although the term is often more broadly misapplied to include aluminium wheels.

The high magnesium content used for the crankcase of the late-World War II Wright Duplex Cyclone eighteen-cylinder radial aviation engine was a serious problem for the earliest examples of the Boeing B-29 heavy bomber, as engine fires in flight could ignite the engine crankcases, literally "torching" the wing spar apart in under a minute, causing wing failure.^[20][21]

An earlier Mercedes-Benz race car model, the Mercedes-Benz 300 SLR, had a body made from Elektron, a magnesium alloy; these cars ran (with successes) at Le Mans, the Mille Miglia, and other world-class race events in 1955 (though one was involved in the single worst accident in auto racing history, in terms of human casualties, at the Le Mans race.) Porsche's all-out quest to decrease the weight of their race cars led to the use of magnesium alloy frames in the famous 917/053 that won Le Mans in 1971.

Volkswagen Group 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 alloy 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 these alloys are recent developments in high-temperature low creep magnesium alloys.

Mitsubishi Motors also uses magnesium (branded magnesium alloy) for its paddle shifters. The general strategy for such alloys is to form intermetallic precipitates at the grain boundaries, for example by adding mischmetal or calcium.^[22] New alloy development and lower costs that make magnesium competitive with aluminium will increase the number of automotive applications.^{[citation needed]}

The second application field of magnesium is electronic devices. Because of low weight, and good mechanical and electrical properties, magnesium is widely used for manufacturing of mobile phones, laptop and tablet 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. The term is still used today. The application of magnesium in the commercial aerospace industry was generally restricted to engine related components, due either to perceived hazards with magnesium parts in the event of fire or to corrosion. 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.^[23] The development and testing of new magnesium alloys continues, notably Elektron 21, which has successfully undergone extensive aerospace testing for suitability in engine and internal and airframe components.^[24] The European Community runs three R&D magnesium projects in the Aerospace priority of Six Framework Program.

Niche uses of the metal

Magnesium, being readily available and relatively nontoxic, has a variety of uses:

• Magnesium is flammable, burning at a temperature of approximately 3,100 °C (3,370 K; 5,610 °F),^[7] and the autoignition temperature of magnesium ribbon is approximately 473 °C (746 K; 883 °F).^[25] It produces intense, bright, white light when it burns. Magnesium's high combustion temperature makes it a useful tool for starting emergency fires. Other uses include flash photography, flares, pyrotechnics, and fireworks sparklers. Magnesium is also often used to ignite thermite, or other materials that require a high ignition temperature.
  

  Magnesium firestarter (in left hand), used with a pocket knife and flint to create sparks that ignite the shavings
• In the form of turnings or ribbons, to prepare Grignard reagents, which are useful in organic synthesis.
• 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 sacrificial (galvanic) anode to protect underground tanks, pipelines, buried structures, and water heaters.
• Alloyed with zinc to produce the zinc sheet used in photoengraving plates in the printing industry, dry-cell battery walls, and roofing.^[18]
• 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, sports equipment such as golf clubs, fishing reels, and archery bows and arrows.

In compounds

Magnesium compounds, primarily magnesium oxide (MgO), are used as a refractory material in furnace linings for producing iron, steel, nonferrous metals, glass, and cement. Magnesium oxide and other magnesium compounds are also used in the agricultural, chemical, and construction industries. Magnesium oxide from calcination is used as an electrical insulator in fire-resistant cables.^[26]

Magnesium reacted with an alkyl halide gives a Grignard reagent, which is a very useful tool for preparing alcohols.

Magnesium salts are frequently included in various foods, fertilizers (magnesium is a component of chlorophyll), and culture media.

Magnesium sulfite is used in the manufacture of paper (sulfite process).

Magnesium phosphate is used to fireproof wood used in construction.

Magnesium hexafluorosilicate is used in mothproofing of textiles.

In the form of turnings or ribbons, Mg is useful in purification of solvents, for example the preparation of super-dry ethanol.

Biological

Pharmaceutical preparations of magnesium are used to treat magnesium deficiency and hypomagnesemia, as well as eclampsia.^[27] Usually in lower dosages, magnesium is commonly included in dietary mineral preparations, including many multivitamin preparations.

Sorted by type of magnesium salt, biological applications of magnesium include:

• Magnesium sulfate, as the heptahydrate called Epsom salts, is used as bath salts, as a laxative, and as a highly soluble fertilizer.^[28]
• Magnesium hydroxide, suspended in water, is used in milk of magnesia antacids and laxatives.
• Magnesium chloride, oxide, gluconate, malate, orotate, glycinate and citrate are all used as oral magnesium supplements. Some have claimed that oral magnesium supplements are therapeutic for Restless Leg Syndrome (RLS) in some individuals.^[29]
• 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).
• Magnesium stearate is a slightly flammable white powder with lubricating properties. In pharmaceutical technology, it is used in the manufacturing of numerous kinds 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 carbonate powder is used by athletes such as gymnasts, weightlifters, and climbers to eliminate moisture and improving the grip on a gymnastic apparatus, lifting bar, and climbing rocks.

Biological roles

Because of 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 that 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 that show few overt symptoms) is relatively rare^[30] although only 32% of people in the United States meet the RDA-DRI;^[31] low levels of magnesium in the body have been associated with the development of a number of human illnesses such as asthma, diabetes, and osteoporosis.^[32] Taken in the proper amount, magnesium plays a role in preventing both stroke and heart attack. The symptoms of people with fibromyalgia, migraines, and premenstrual syndrome are less severe, and magnesium can shorten the length of the migraine symptoms.^[33][34]

Adult human bodies contain about 24 grams of magnesium, with 60% in the skeleton, 39% intracellular (20% in skeletal muscle), and 1% extracellular.^[35] Serum levels are typically 0.7–1.0 mmol/L or 1.8–2.4 mEq/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. Magnesium competes with calcium in the human body,^[36] in this way it actually keeps calcium in check. However, this can cause a calcium deficiency if calcium levels are already low.^[36] Low and high protein intake inhibit magnesium absorption, and other factors such as phosphate, phytate, and fat affect absorption. Excess dietary magnesium is excreted in feces, urine, and sweat.^[37] 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.^[38] In these tests, magnesium is injected intravenously; a retention of 20% or more indicates deficiency.^[39] Other nutrient deficiencies are identified through biomarkers, but none are established for magnesium.^[40]

The UK recommended daily values for magnesium is 300 mg for men and 270 mg for women.^[41] Spices, nuts, cereals, coffee, cocoa, tea, and vegetables are rich sources of magnesium.^[42] Green leafy vegetables such as spinach are also rich in magnesium as they contain chlorophyll. Observations of reduced dietary magnesium intake in modern Western countries compared to earlier generations may be related to food refining and modern fertilizers that contain no magnesium.^[37]

Numerous pharmaceutical preparations of magnesium, as well as magnesium dietary supplements are available. Magnesium oxide, one of the most common forms in magnesium dietary supplements because it has high magnesium content per weight, has been reported the least bioavailable.^[43][44] Magnesium citrate has been reported as more bioavailable than oxide or amino-acid chelate (glycinate) forms.^[45]

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.^[32] With supplements, overdose is possible, in particular 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.^[46] Alcoholism can produce a magnesium deficiency, which is reversible by oral or parenteral administration, depending on the degree of deficiency.^[47]

Detection in biological fluids

Magnesium concentrations in plasma or serum may be measured to monitor for efficacy and safety in those receiving the drug therapeutically, to confirm the diagnosis in potential poisoning victims or to assist in the forensic investigation in a case of fatal overdosage. The newborn children of mothers having received parenteral magnesium sulfate during labor may exhibit toxicity at serum magnesium levels that were considered appropriate for the mothers.^[48]

Disease

Results from a meta-analysis of randomized clinical trials demonstrated that a magnesium supplement can lower high blood pressure in a dose-dependent manner.^[49] Low serum magnesium levels are associated with metabolic syndrome, diabetes mellitus type 2 and hypertension.^[50] Low serum magnesium levels have been associated with a higher risk of developing metabolic syndrome.^[51] Magnesium therapy is recommended by the ACC/AHA/ESC 2006 Guidelines for Management of Patients With Ventricular Arrhythmias and the Prevention of Sudden Cardiac Death for patients with ventricular arrhythmia associated with torsade de pointes who present with long QT syndrome as well as for the treatment of patients with digoxin intoxication-induced arrhythmias.^[52] Magnesium is also the drug of choice in the management of pre-eclampsia and eclampsia.^[53]

In addition to its therapeutic role, magnesium improves calcification. Patients with chronic kidney disease have a high prevalence of vascular calcification, and cardiovascular disease is the leading cause of death in this population. Several in vitro and animal studies point toward a protective role of magnesium through multiple molecular mechanisms. Magnesium is a natural calcium antagonist, and both human and animal studies have shown that low circulating magnesium levels are associated with vascular calcification.^[54] Results from an observational study conducted in the general Japanese population demonstrated that lower serum magnesium levels were significantly and independently associated with a greater average intima-media thickness and the risk of at least two carotid plaques.^[55] Magnesium supplementation might be useful in reducing the progression of atherosclerosis in chronic dialysis patients.^[56] Low serum magnesium may be an independent risk factor for death in patients with chronic kidney disease,^[57] and patients with mildly elevated serum magnesium levels could have a survival advantage over those with lower magnesium levels.^[58]

Magnesium overdose

Since the kidneys are responsible for the excretion of magnesium, anyone with a heart or kidney disorder should not take any extra magnesium except under their doctor's supervision. It is very rare to overdose on magnesium from food,^[36] however, people that ingest large amounts of milk of magnesia (as a laxative or antacid), epsom salts (as a laxative or tonic), or magnesium supplements may overdose, especially if they suffer from kidney problems. Too much magnesium can cause several serious health problems, including nausea, vomiting, severely lowered blood pressure, confusion, slowed heart rate, respiratory paralysis, deficiencies of other minerals, coma, cardiac arrhythmia, cardiac arrest, and eventually death.^[36] The most common side effects of magnesium toxicity are stomach upset and diarrhea.^[36]

Safety precautions for magnesium metal

Magnesium metal and its alloys are explosive hazards; they are highly flammable in their pure form when molten or in powder or ribbon form. Burning or molten magnesium metal reacts violently with water. When working with powdered magnesium, safety glasses with welding eye protection are employed, because the bright-white light produced by burning magnesium contains ultraviolet light that can permanently damage the retinas of the eyes.^[59]

Magnesium is capable of reducing water to highly flammable hydrogen gas:^[60]

Mg (s) + 2 H
2O (l) → Mg(OH)
2 (s) + H
2 (g)

As a result, water cannot extinguish magnesium fires. The hydrogen gas produced only intensifies the fire. Dry sand is an effective smothering agent, but only on relatively level and flat surfaces.

Magnesium also reacts with carbon dioxide to form magnesium oxide and carbon:

2 Mg (s) + CO
2 → 2 MgO (s) + C (s)

hence, carbon dioxide fire extinguishers cannot be used for extinguishing magnesium fires, either.^[61]

Burning magnesium is usually quenched by using a Class D dry chemical fire extinguisher, or by covering the fire with sand or magnesium foundry flux to remove its air source.

See also

References

External links

• Magnesium at The Periodic Table of Videos (University of Nottingham)
• Chemistry in its element podcast (MP3) from the Royal Society of Chemistry's Chemistry World: Magnesium
• Magnesium – a versatile and often overlooked element: new perspectives with a focus on chronic kidney disease

v t e Magnesium compounds MgB2 MgBr2 MgCO3 MgC6H6O7 MgC14H10O4 MgCl2 Mg(ClO4)2 MgF2 MgH2 Mg(HCO3)2 MgI2 Mg(NO3)2 MgO MgO2 Mg(OH)2 MgPo MgS MgSO3 MgSO4 MgU2O7 Mg2Al3 Mg2Si Mg2Si3O8 Mg3N2 Mg3(PO4)2 Mg2(CrO4)2

v t e Periodic table
	1	2															3	4	5	6	7	8	9	10	11	12	13	14	15	16	17	18
1	H																															He
2	Li	Be																									B	C	N	O	F	Ne
3	Na	Mg																									Al	Si	P	S	Cl	Ar
4	K	Ca															Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn	Ga	Ge	As	Se	Br	Kr
5	Rb	Sr															Y	Zr	Nb	Mo	Tc	Ru	Rh	Pd	Ag	Cd	In	Sn	Sb	Te	I	Xe
6	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
7	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	Cn	Uut	Fl	Uup	Lv	Uus	Uuo
Alkali metal Alkaline earth metal Lan­thanide Actinide Transition metal Poor metal Metalloid Polyatomic nonmetal Diatomic nonmetal Noble gas Unknown chemical properties
Large version

v t e Chemical elements named after places Named after terrestrial places Americium Berkelium Californium Copper Darmstadtium Dubnium Erbium Europium Francium Gallium Germanium Hafnium Hassium Holmium Livermorium Lutetium Magnesium Polonium Rhenium Ruthenium Scandium Strontium Terbium Thulium Ytterbium Yttrium Named after astronomical objects Cerium Helium Neptunium Palladium Plutonium Selenium Tellurium Uranium See also Chemical element name etymologies Chemical elements naming controversies Scientists whose names are used in chemical element names