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Magnesium, 12Mg
CSIRO ScienceImage 2893 Crystalised magnesium.jpg
Pronunciation /mæɡˈnziəm/ (mag-NEE-zee-əm)
Appearanceshiny grey solid
Standard atomic weight Ar°(Mg)
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


Atomic number (Z)12
Group group 2 (alkaline earth metals)
Period period 3
Block   s-block
Electron configuration [ Ne ] 3s2
Electrons per shell2, 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/cm3 [3]
when liquid (at m.p.)1.584 g/cm3
Heat of fusion 8.48  kJ/mol
Heat of vaporization 128 kJ/mol
Molar heat capacity 24.869 [4]  J/(mol·K)
Vapor pressure
P (Pa)1101001 k10 k100 k
at T (K)70177386197111321361
Atomic properties
Oxidation states 0, [5] +1, [6] +2 (a strongly basic oxide)
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
Magnesium spectrum visible.png
Spectral lines of magnesium
Other properties
Natural occurrence primordial
Crystal structure hexagonal close-packed (hcp)(hP2)
Lattice constants
Hexagonal close packed.svg
a = 320.91 pm
c = 521.03 pm (at 20 °C) [3]
Thermal expansion 25.91×10−6/K (at 20 °C) [3] [lower-alpha 1]
Thermal conductivity 156 [7]  W/(m⋅K)
Electrical resistivity 43.9 [8]  nΩ⋅m(at 20 °C)
Magnetic ordering paramagnetic
Molar magnetic susceptibility +13.1×10−6 cm3/mol(298 K) [9]
Young's modulus 45 GPa
Shear modulus 17 GPa
Bulk modulus 35.4 [10]  GPa
Speed of sound thin rod4940 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
Namingafter Magnesia, Greece [11]
Discovery Joseph Black (1755 [11] )
First isolation Humphry Davy (1808 [11] )
Isotopes of magnesium
Main isotopes [12] Decay
abun­dance half-life (t1/2) mode pro­duct
24Mg79% stable
Symbol category class.svg  Category: Magnesium
| references

Magnesium is a chemical element; it has symbol  Mg and atomic number  12. It is a shiny gray metal having a low density, low melting point and high chemical reactivity. Like the other alkaline earth metals (group 2 of the periodic table) it occurs naturally only in combination with other elements and it almost always has an oxidation state of +2. It reacts readily with air to form a thin passivation coating of magnesium oxide that inhibits further corrosion of the metal. The free metal burns with a brilliant-white light. The metal is obtained mainly by electrolysis of magnesium salts obtained from brine. It is less dense than aluminium and is used primarily as a component in strong and lightweight alloys that contain aluminium.


In the cosmos, magnesium is produced in large, aging stars by the sequential addition of three helium nuclei to a carbon nucleus. When such stars explode as supernovas, much of the magnesium is expelled into the interstellar medium where it may recycle into new star systems. Magnesium is the eighth most abundant element in the Earth's crust [13] and the fourth most common element in the Earth (after iron, oxygen and silicon), making up 13% of the planet's mass and a large fraction of the planet's mantle. It is the third most abundant element dissolved in seawater, after sodium and chlorine. [14]

This element is the eleventh most abundant element by mass in the human body and is essential to all cells and some 300 enzymes. [15] Magnesium ions interact with polyphosphate compounds such as ATP, DNA, and RNA. Hundreds of enzymes require magnesium ions to function. Magnesium compounds are used medicinally as common laxatives and antacids (such as milk of magnesia), and to stabilize abnormal nerve excitation or blood vessel spasm in such conditions as eclampsia. [15]


Physical properties

Elemental magnesium is a gray-white lightweight metal, two-thirds the density of aluminium. Magnesium has the lowest melting (923 K (650 °C)) and the lowest boiling point (1,363 K (1,090 °C)) of all the alkaline earth metals. [16]

Pure polycrystalline magnesium is brittle and easily fractures along shear bands. It becomes much more malleable when alloyed with small amounts of other metals, such as 1% aluminium. [17] The malleability of polycrystalline magnesium can also be significantly improved by reducing its grain size to ca. 1 micron or less. [18]

When finely powdered, magnesium reacts with water to produce hydrogen gas:

Mg(s) + 2H2O(g) → Mg(OH)2(aq) + H2(g) + 1203.6 kJ/mol

However, this reaction is much less dramatic than the reactions of the alkali metals with water, because the magnesium hydroxide builds up on the surface of the magnesium metal and inhibits further reaction. [19]

Chemical properties

General chemistry

It tarnishes slightly when exposed to air, although, unlike the heavier alkaline earth 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. [20]

Direct reaction of magnesium with air or oxygen at ambient pressure forms only the "normal" oxide MgO. However, this oxide may be combined with hydrogen peroxide to form magnesium peroxide, MgO2, and at low temperature the peroxide may be further reacted with ozone to form magnesium superoxide Mg(O2)2. [21]

Magnesium reacts with water at room temperature, though it reacts much more slowly than calcium, a similar group 2 metal. [20] When submerged in water, hydrogen bubbles form slowly on the surface of the metal; this reaction happens much more rapidly with powdered magnesium. [20] The reaction also occurs faster with higher temperatures (see § Safety precautions). Magnesium's reversible reaction with water can be harnessed to store energy and run a magnesium-based engine. Magnesium also reacts exothermically with most acids such as hydrochloric acid (HCl), producing magnesium chloride and hydrogen gas, similar to the HCl reaction with aluminium, zinc, and many other metals. [22]


Magnesium is highly flammable, especially when powdered or shaved into thin strips, though it is difficult to ignite in mass or bulk. [20] Flame temperatures of magnesium and magnesium alloys can reach 3,100 °C (5,610 °F), [23] although flame height above the burning metal is usually less than 300 mm (12 in). [24] Once ignited, such fires are difficult to extinguish because they resist several substances commonly used to put out fires; combustion continues in nitrogen (forming magnesium nitride), in carbon dioxide (forming magnesium oxide and carbon), and in water (forming magnesium oxide and hydrogen, which also combusts due to heat in the presence of additional oxygen). This property was used in incendiary weapons during the firebombing of cities in World War II, where the only practical civil defense was to smother a burning flare under dry sand to exclude atmosphere from the combustion.

Magnesium may also be used as an igniter for thermite, a mixture of aluminium and iron oxide powder that ignites only at a very high temperature.

Organic chemistry

Organomagnesium compounds are widespread in organic chemistry. They are commonly found as Grignard reagents, formed by reaction of magnesium with haloalkanes. Examples of Grignard reagents are phenylmagnesium bromide and ethylmagnesium bromide. The Grignard reagents function as a common nucleophile, attacking the electrophilic group such as the carbon atom that is present within the polar bond of a carbonyl group.

A prominent organomagnesium reagent beyond Grignard reagents is magnesium anthracene, with magnesium forming a 1,4-bridge over the central ring. It is used as a source of highly active magnesium. The related butadiene-magnesium adduct serves as a source for the butadiene dianion.

Magnesium in organic chemistry also appears as low valent magnesium compounds, primarily with the magnesium forming diatomic ions in the +1 oxidation state but more recently also with zero oxidation state or a mixture of +1 and zero states. [25] Such compounds find synthetic application as reducing agents and sources of nucleophilic metal atoms.

Source of light

When burning in air, magnesium produces a brilliant white light that includes strong ultraviolet wavelengths. Magnesium powder (flash powder) was used for subject illumination in the early days of photography. [26] [27] Later, magnesium filament was used in electrically ignited single-use photography flashbulbs. Magnesium powder is used in fireworks and marine flares where a brilliant white light is required. It was also used for various theatrical effects, [28] such as lightning, [29] pistol flashes, [30] and supernatural appearances. [31]

Detection in solution

The presence of magnesium ions can be detected by the addition of ammonium chloride, ammonium hydroxide and monosodium phosphate to an aqueous or dilute HCl solution of the salt. The formation of a white precipitate indicates the presence of magnesium ions.

Azo violet dye can also be used, turning deep blue in the presence of an alkaline solution of magnesium salt. The color is due to the adsorption of azo violet by Mg(OH)2.


Magnesium is the eighth-most-abundant element in the Earth's crust by mass and tied in seventh place with iron in molarity. [13] 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 more than 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 (about 18 the mass of sodium ions in a given sample), which makes seawater and sea salt attractive commercial sources for Mg. To extract the magnesium, calcium hydroxide is added to seawater to form magnesium hydroxide precipitate.

+ Ca(OH)
+ CaCl

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

+ 2 HCl → MgCl
+ 2 H

From magnesium chloride, electrolysis produces magnesium.



Magnesium is brittle, and fractures along shear bands when its thickness is reduced by only 10% by cold rolling (top). However, after alloying Mg with 1% Al and 0.1% Ca, its thickness could be reduced by 54% using the same process (bottom). Cold rolling of Mg and Mg-1Al-0.1Ca.jpg
Magnesium is brittle, and fractures along shear bands when its thickness is reduced by only 10% by cold rolling (top). However, after alloying Mg with 1% Al and 0.1% Ca, its thickness could be reduced by 54% using the same process (bottom).

As of 2013, magnesium alloys consumption was less than one million tonnes per year, compared with 50 million tonnes of aluminium alloys. Their use has been historically limited by the tendency of Mg alloys to corrode, [32] creep at high temperatures, and combust. [33]


In magnesium alloys, the presence of iron, nickel, copper, or cobalt strongly activates corrosion. In more than trace amounts, these metals precipitate as intermetallic compounds, and the precipitate locales function as active cathodic sites that reduce water, causing the loss of magnesium. [33] Controlling the quantity of these metals improves corrosion resistance. Sufficient manganese overcomes the corrosive effects of iron. This requires precise control over composition, increasing costs. [33] Adding a cathodic poison captures atomic hydrogen within the structure of a metal. This prevents the formation of free hydrogen gas, an essential factor of corrosive chemical processes. The addition of about one in three hundred parts arsenic reduces the corrosion rate of magnesium in a salt solution by a factor of nearly ten. [33] [34]

High-temperature creep and flammability

Magnesium's tendency to creep (gradually deform) at high temperatures is greatly reduced by alloying with zinc and rare-earth elements. [35] Flammability is significantly reduced by a small amount of calcium in the alloy. [33] By using rare-earth elements, it may be possible to manufacture magnesium alloys that are able to not catch fire at higher temperatures compared to magnesium's liquidus and in some cases potentially pushing it close to magnesium's boiling point. [36]


Magnesium forms a variety of compounds important to industry and biology, including magnesium carbonate, magnesium chloride, magnesium citrate, magnesium hydroxide (milk of magnesia), magnesium oxide, magnesium sulfate, and magnesium sulfate heptahydrate (Epsom salts).


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

The nuclide 26
has found application in isotopic geology, similar to that of aluminium. 26
is a radiogenic daughter product of 26
, which has a half-life of 717,000 years. Excessive quantities of stable 26
have been observed in the Ca-Al-rich inclusions of some carbonaceous chondrite meteorites. This anomalous abundance is attributed to the decay of its parent 26
in the inclusions, and researchers conclude that such meteorites were formed in the solar nebula before the 26
had decayed. These are among the oldest objects in the Solar System and contain preserved information about its early history.

It is conventional to plot 26
against an Al/Mg ratio. In an isochron dating plot, the Al/Mg ratio plotted is 27
. The slope of the isochron has no age significance, but indicates the initial 26
ratio in the sample at the time when the systems were separated from a common reservoir.


Magnesium sheets and ingots Mg sheets and ingots.jpg
Magnesium sheets and ingots

World production was approximately 1,100 kt in 2017, with the bulk being produced in China (930 kt) and Russia (60 kt). [37] The United States was in the 20th century the major world supplier of this metal, supplying 45% of world production even as recently as 1995. Since the Chinese mastery of the Pidgeon process the US market share is at 7%, with a single US producer left: US Magnesium, a Renco Group company in Utah born from now-defunct Magcorp. [38]

In September 2021, China took steps to reduce production of magnesium as a result of a government initiative to reduce energy availability for manufacturing industries, leading to a significant price increase. [39]

Pidgeon process

China is almost completely reliant on the silicothermic Pidgeon process (the reduction of the oxide at high temperatures with silicon, often provided by a ferrosilicon alloy in which the iron is but a spectator in the reactions) to obtain the metal. [40] The process can also be carried out with carbon at approx 2300 °C:

+ Si
+ 2CaO
+ Ca
+ C
+ CO

Dow process

In the United States, magnesium is obtained principally with the Dow process, by electrolysis of fused magnesium chloride from brine and sea water. A saline solution containing Mg2+
ions is first treated with lime (calcium oxide) and the precipitated magnesium hydroxide is collected:

(aq) + CaO(s) + H
(l) → Ca2+
(aq) + Mg(OH)

The hydroxide is then converted to magnesium chloride by treating the hydroxide with hydrochloric acid and heating of the product to eliminate water:

(s) + 2HCl(aq) → MgCl
(aq) + 2H

The salt is then electrolyzed in the molten state. At the cathode, the Mg2+
ion is reduced by two electrons to magnesium metal:

+ 2
→ Mg

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

(g) + 2

YSZ process

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. [41] It has been reported that this method provides a 40% reduction in cost per pound over the electrolytic reduction method. [42]


The name magnesium originates from the Greek word for locations related to the tribe of the Magnetes, either a district in Thessaly called Magnesia [43] or Magnesia ad Sipylum, now in Turkey. [44] 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 local well. 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 obtained by evaporating the water became known as Epsom salts and its fame spread. [45] It was eventually recognized as hydrated magnesium sulfate, MgSO
.[ citation needed ]

The metal itself was first isolated by Sir Humphry Davy in England in 1808. He used electrolysis on a mixture of magnesia and mercuric oxide. [46] Antoine Bussy prepared it in coherent form in 1831. Davy's first suggestion for a name was 'magnium', [46] but the name magnesium is now used in most European languages.[ citation needed ]


Magnesium metal

An unusual application of magnesium as an illumination source while wakeskating in 1931 Bundesarchiv Bild 102-12062, Wasserreiter mit Magnesiumfackeln.jpg
An unusual application of magnesium as an illumination source while wakeskating in 1931

Magnesium is the third-most-commonly-used structural metal, following iron and aluminium. [47] The main applications of magnesium are, in order: aluminium alloys, die-casting (alloyed with zinc), [48] removing sulfur in the production of iron and steel, and the production of titanium in the Kroll process. [49]

Magnesium is used in lightweight materials and alloys. For example, when infused with silicon carbide nanoparticles, it has extremely high specific strength. [50]

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, a term which is still used today. In the commercial aerospace industry, magnesium was generally restricted to engine-related components, due to fire and corrosion hazards. Magnesium alloy use in aerospace is increasing in the 21st century, driven by the importance of fuel economy. [51] Recent developments in metallurgy and manufacturing have allowed for the potential for magnesium alloys to act as replacements for aluminium and steel alloys in certain applications. [52] [53]



The Bugatti Type 57 Aerolithe featured a lightweight body made of Elektron, a trademarked magnesium alloy. Bugatti Aerolithe AV.jpg
The Bugatti Type 57 Aérolithe featured a lightweight body made of Elektron, a trademarked magnesium alloy.

Both AJ62A and AE44 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. [63]


Because of low density and good mechanical and electrical properties, magnesium is used for manufacturing of mobile phones, laptop and tablet computers, cameras, and other electronic components. [64] It was used as a premium feature because of its light weight in some 2020 laptops. [65]

Products made of magnesium: firestarter and shavings, sharpener, magnesium ribbon Magnesium-products.jpg
Products made of magnesium: firestarter and shavings, sharpener, magnesium ribbon

Magnesium materials in medicine

Recent research promises a high development potential of magnesium materials as resorbable implant material (e.g. as stent) for the human body [66] [67] . Common magnesium alloy elements are calcium and zinc, [68] but also rare earth and silver are investigated. [69] The biodegradable behavior of magnesium is a decisive advantage when used as an implant material to be used for a limited period of time, as it would dissolve without risk after a certain period of time. This would eliminate the risks and costs of an operation to remove the implant [70] .


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), [23] and the autoignition temperature of magnesium ribbon is approximately 473 °C (746 K; 883 °F). [71] 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, fireworks sparklers, and trick birthday candles. Magnesium is also often used to ignite thermite or other materials that require a high ignition temperature. Magnesium continues to be used as an incendiary element in warfare. [72]
    Magnesium firestarter (in left hand), used with a pocket knife and flint to create sparks that ignite the shavings Magnesium Sparks.jpg
    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.[ citation needed ]
  • As an additive agent in conventional propellants and the production of nodular graphite in cast iron.[ citation needed ]
  • As a reducing agent to separate uranium and other metals from their salts. [73]
  • As a sacrificial (galvanic) anode to protect boats, underground tanks, pipelines, buried structures, and water heaters.[ citation needed ]
  • Alloyed with zinc to produce the zinc sheet used in photoengraving plates in the printing industry, dry-cell battery walls, and roofing. [48]
  • 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.[ citation needed ]
  • Specialty, high-grade car wheels of magnesium alloy are called "mag wheels", although the term is often misapplied to aluminium wheels. Many car and aircraft manufacturers have made engine and body parts from magnesium.[ citation needed ]
  • Magnesium batteries have been commercialized as primary batteries, and are an active topic of research for rechargeable batteries. [74]


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. [75]

Magnesium hydride is under investigation as a way to store hydrogen.[ citation needed ]

Magnesium reacts with haloalkanes to give Grignard reagents, which are used for a wide variety of organic reactions forming carbon–carbon bonds. [76]

Magnesium salts are included in various foods, fertilizers (magnesium is a component of chlorophyll), and microbe culture media.[ citation needed ]

Magnesium sulfite is used in the manufacture of paper (sulfite process).[ citation needed ]

Magnesium phosphate is used to fireproof wood used in construction.[ citation needed ]

Magnesium hexafluorosilicate is used for moth-proofing textiles.[ citation needed ]

Biological roles

Mechanism of action

The important interaction between phosphate and magnesium ions makes magnesium essential to the basic nucleic acid chemistry of all cells of all known living organisms. More than 300 enzymes require magnesium ions for their catalytic action, including all enzymes using or synthesizing ATP and those that use other nucleotides to synthesize DNA and RNA. The ATP molecule is normally found in a chelate with a magnesium ion. [77]



Examples of food sources of magnesium (clockwise from top left): bran muffins, pumpkin seeds, barley, buckwheat flour, low-fat vanilla yogurt, trail mix, halibut steaks, garbanzo beans, lima beans, soybeans, and spinach FoodSourcesOfMagnesium.jpg
Examples of food sources of magnesium (clockwise from top left): bran muffins, pumpkin seeds, barley, buckwheat flour, low-fat vanilla yogurt, trail mix, halibut steaks, garbanzo beans, lima beans, soybeans, and spinach

Spices, nuts, cereals, cocoa and vegetables are good sources of magnesium. [15] Green leafy vegetables such as spinach are also rich in magnesium. [78]

Dietary recommendations

In the UK, the recommended daily values for magnesium are 300 mg for men and 270 mg for women. [79] In the U.S. the Recommended Dietary Allowances (RDAs) are 400 mg for men ages 19–30 and 420 mg for older; for women 310 mg for ages 19–30 and 320 mg for older. [80]


Numerous pharmaceutical preparations of magnesium and dietary supplements are available. In two human trials magnesium oxide, one of the most common forms in magnesium dietary supplements because of its high magnesium content per weight, was less bioavailable than magnesium citrate, chloride, lactate or aspartate. [81] [82]


An adult body has 22–26 grams of magnesium, [15] [83] with 60% in the skeleton, 39% intracellular (20% in skeletal muscle), and 1% extracellular. [15] Serum levels are typically 0.7–1.0 mmol/L or 1.8–2.4 mEq/L. Serum magnesium levels may be normal even when intracellular magnesium is deficient. The mechanisms for maintaining the magnesium level in the serum are varying gastrointestinal absorption and renal excretion. Intracellular magnesium is correlated with intracellular potassium. Increased magnesium lowers calcium [84] and can either prevent hypercalcemia or cause hypocalcemia depending on the initial level. [84] Both low and high protein intake conditions inhibit magnesium absorption, as does the amount of phosphate, phytate, and fat in the gut. Unabsorbed dietary magnesium is excreted in feces; absorbed magnesium is excreted in urine and sweat. [85]

Detection in serum and plasma

Magnesium status may be assessed by measuring serum and erythrocyte magnesium concentrations coupled with urinary and fecal magnesium content, but intravenous magnesium loading tests are more accurate and practical. [86] A retention of 20% or more of the injected amount indicates deficiency. [87] As of 2004, no biomarker has been established for magnesium. [88]

Magnesium concentrations in plasma or serum may be monitored 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 overdose. The newborn children of mothers who received parenteral magnesium sulfate during labor may exhibit toxicity with normal serum magnesium levels. [89]


Low plasma magnesium (hypomagnesemia) is common: it is found in 2.5–15% of the general population. [90] From 2005 to 2006, 48 percent of the United States population consumed less magnesium than recommended in the Dietary Reference Intake. [91] Other causes are increased renal or gastrointestinal loss, an increased intracellular shift, and proton-pump inhibitor antacid therapy. Most are asymptomatic, but symptoms referable to neuromuscular, cardiovascular, and metabolic dysfunction may occur. [90] Alcoholism is often associated with magnesium deficiency. Chronically low serum magnesium levels are associated with metabolic syndrome, diabetes mellitus type 2, fasciculation, and hypertension. [92]


Sorted by type of magnesium salt, other therapeutic applications include:


Overdose from dietary sources alone is unlikely because excess magnesium in the blood is promptly filtered by the kidneys, [90] and overdose is more likely in the presence of impaired renal function. In spite of this, megadose therapy has caused death in a young child, [99] and severe hypermagnesemia in a woman [100] and a young girl [101] who had healthy kidneys. The most common symptoms of overdose are nausea, vomiting, and diarrhea; other symptoms include hypotension, confusion, slowed heart and respiratory rates, deficiencies of other minerals, coma, cardiac arrhythmia, and death from cardiac arrest. [84]

Function in plants

Plants require magnesium to synthesize chlorophyll, essential for photosynthesis. [102] Magnesium in the center of the porphyrin ring in chlorophyll functions in a manner similar to the iron in the center of the porphyrin ring in heme. Magnesium deficiency in plants causes late-season yellowing between leaf veins, [103] especially in older leaves, and can be corrected by either applying epsom salts (which is rapidly leached), or crushed dolomitic limestone, to the soil.

Safety precautions

GHS labelling:
H228, H251, H261
P210, P231, P235, P410, P422 [104]
NFPA 704 (fire diamond)
Magnesium block heated with blowtorch to self-combustion, emitting intense white light

Magnesium metal and its alloys can be explosive hazards; they are highly flammable in their pure form when molten or in powder or ribbon form. Burning or molten magnesium reacts violently with water. When working with powdered magnesium, safety glasses with eye protection and UV filters (such as welders use) are employed because burning magnesium produces ultraviolet light that can permanently damage the retina of a human eye. [106]

Magnesium is capable of reducing water and releasing highly flammable hydrogen gas: [107]

Mg(s) + 2H
(l) → Mg(OH)
(s) + H

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

Magnesium reacts with carbon dioxide exothermically to form magnesium oxide and carbon: [108]

2Mg(s) + CO
(g) → 2MgO(s) + C(s)

Hence, carbon dioxide fuels rather than extinguishes magnesium fires.

Burning magnesium can be 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. [109]

See also


  1. The thermal expansion is anisotropic: the parameters (at 20 °C) for each crystal axis are αa = 25.31×10−6/K, αc = 27.03×10−6/K, and αaverage = αV/3 = 25.91×10−6/K. [3]

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A carbonate is a salt of carbonic acid, H2CO3, characterized by the presence of the carbonate ion, a polyatomic ion with the formula CO2−3. The word "carbonate" may also refer to a carbonate ester, an organic compound containing the carbonate groupO=C(−O−)2.

<span class="mw-page-title-main">Hydroxide</span> Chemical compound

Hydroxide is a diatomic anion with chemical formula OH. It consists of an oxygen and hydrogen atom held together by a single covalent bond, and carries a negative electric charge. It is an important but usually minor constituent of water. It functions as a base, a ligand, a nucleophile, and a catalyst. The hydroxide ion forms salts, some of which dissociate in aqueous solution, liberating solvated hydroxide ions. Sodium hydroxide is a multi-million-ton per annum commodity chemical. The corresponding electrically neutral compound HO is the hydroxyl radical. The corresponding covalently bound group –OH of atoms is the hydroxy group. Both the hydroxide ion and hydroxy group are nucleophiles and can act as catalysts in organic chemistry.

<span class="mw-page-title-main">Rust</span> Type of iron oxide

Rust is an iron oxide, a usually reddish-brown oxide formed by the reaction of iron and oxygen in the catalytic presence of water or air moisture. Rust consists of hydrous iron(III) oxides (Fe2O3·nH2O) and iron(III) oxide-hydroxide (FeO(OH), Fe(OH)3), and is typically associated with the corrosion of refined iron.

<span class="mw-page-title-main">Alkaline earth metal</span> Group of chemical elements

The alkaline earth metals are six chemical elements in group 2 of the periodic table. They are beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). The elements have very similar properties: they are all shiny, silvery-white, somewhat reactive metals at standard temperature and pressure.

<span class="mw-page-title-main">Electrolysis</span> Technique in chemistry and manufacturing

In chemistry and manufacturing, electrolysis is a technique that uses direct electric current (DC) to drive an otherwise non-spontaneous chemical reaction. Electrolysis is commercially important as a stage in the separation of elements from naturally occurring sources such as ores using an electrolytic cell. The voltage that is needed for electrolysis to occur is called the decomposition potential. The word "lysis" means to separate or break, so in terms, electrolysis would mean "breakdown via electricity".

<span class="mw-page-title-main">Base (chemistry)</span> Type of chemical substance

In chemistry, there are three definitions in common use of the word "base": Arrhenius bases, Brønsted bases, and Lewis bases. All definitions agree that bases are substances that react with acids, as originally proposed by G.-F. Rouelle in the mid-18th century.

<span class="mw-page-title-main">Corrosion</span> Gradual destruction of materials by chemical reaction with its environment

Corrosion is a natural process that converts a refined metal into a more chemically stable oxide. It is the gradual deterioration of materials by chemical or electrochemical reaction with their environment. Corrosion engineering is the field dedicated to controlling and preventing corrosion.

In physical chemistry and engineering, passivation is coating a material so that it becomes "passive", that is, less readily affected or corroded by the environment. Passivation involves creation of an outer layer of shield material that is applied as a microcoating, created by chemical reaction with the base material, or allowed to build by spontaneous oxidation in the air. As a technique, passivation is the use of a light coat of a protective material, such as metal oxide, to create a shield against corrosion. Passivation of silicon is used during fabrication of microelectronic devices. Undesired passivation of electrodes, called "fouling", increases the circuit resistance so it interferes with some electrochemical applications such as electrocoagulation for wastewater treatment, amperometric chemical sensing, and electrochemical synthesis.

<span class="mw-page-title-main">Galvanic anode</span> Main component of cathodic protection

A galvanic anode, or sacrificial anode, is the main component of a galvanic cathodic protection system used to protect buried or submerged metal structures from corrosion.

A period 3 element is one of the chemical elements in the third row of the periodic table of the chemical elements. The periodic table is laid out in rows to illustrate recurring (periodic) trends in the chemical behavior of the elements as their atomic number increases: a new row is begun when chemical behavior begins to repeat, meaning that elements with similar behavior fall into the same vertical columns. The third period contains eight elements: sodium, magnesium, aluminium, silicon, phosphorus, sulfur, chlorine and argon. The first two, sodium and magnesium, are members of the s-block of the periodic table, while the others are members of the p-block. All of the period 3 elements occur in nature and have at least one stable isotope.

<span class="mw-page-title-main">Calcium hydroxide</span> Inorganic compound of formula Ca(OH)2

Calcium hydroxide (traditionally called slaked lime) is an inorganic compound with the chemical formula Ca(OH)2. It is a colorless crystal or white powder and is produced when quicklime (calcium oxide) is mixed with water. It has many names including hydrated lime, caustic lime, builders' lime, slaked lime, cal, and pickling lime. Calcium hydroxide is used in many applications, including food preparation, where it has been identified as E number E526. Limewater, also called milk of lime, is the common name for a saturated solution of calcium hydroxide.

Neodymium(III) chloride or neodymium trichloride is a chemical compound of neodymium and chlorine with the formula NdCl3. This anhydrous compound is a mauve-colored solid that rapidly absorbs water on exposure to air to form a purple-colored hexahydrate, NdCl3·6H2O. Neodymium(III) chloride is produced from minerals monazite and bastnäsite using a complex multistage extraction process. The chloride has several important applications as an intermediate chemical for production of neodymium metal and neodymium-based lasers and optical fibers. Other applications include a catalyst in organic synthesis and in decomposition of waste water contamination, corrosion protection of aluminium and its alloys, and fluorescent labeling of organic molecules (DNA).

<span class="mw-page-title-main">Ferrosilicon</span>

Ferrosilicon is an alloy of iron and silicon with a typical silicon content by weight of 15–90%. It contains a high proportion of iron silicides.

Water-reactive substances are those that spontaneously undergo a chemical reaction with water, as they are highly reducing in nature. Notable examples include alkali metals, lithium through caesium, and alkaline earth metals, magnesium through barium.

<span class="mw-page-title-main">Boiler water</span>

Boiler water is liquid water within a boiler, or in associated piping, pumps and other equipment, that is intended for evaporation into steam. The term may also be applied to raw water intended for use in boilers, treated boiler feedwater, steam condensate being returned to a boiler, or boiler blowdown being removed from a boiler.

A metal–air electrochemical cell is an electrochemical cell that uses an anode made from pure metal and an external cathode of ambient air, typically with an aqueous or aprotic electrolyte.

<span class="mw-page-title-main">Concrete degradation</span> Damage to concrete affecting its mechanical strength and its durability

Concrete degradation may have many different causes. Concrete is mostly damaged by the corrosion of reinforcement bars due to the carbonatation of hardened cement paste or chloride attack under wet conditions. Chemical damages are caused by the formation of expansive products produced by various chemical reactions, by aggressive chemical species present in groundwater and seawater, or by microorganisms. Other damaging processes can also involve calcium leaching by water infiltration and different physical phenomena initiating cracks formation and propagation. All these detrimental processes and damaging agents adversely affects the concrete mechanical strength and its durability.

<span class="mw-page-title-main">Aluminium compounds</span>

Aluminium (British and IUPAC spellings) or aluminum (North American spelling) combines characteristics of pre- and post-transition metals. Since it has few available electrons for metallic bonding, like its heavier group 13 congeners, it has the characteristic physical properties of a post-transition metal, with longer-than-expected interatomic distances. Furthermore, as Al3+ is a small and highly charged cation, it is strongly polarizing and aluminium compounds tend towards covalency; this behaviour is similar to that of beryllium (Be2+), an example of a diagonal relationship. However, unlike all other post-transition metals, the underlying core under aluminium's valence shell is that of the preceding noble gas, whereas for gallium and indium it is that of the preceding noble gas plus a filled d-subshell, and for thallium and nihonium it is that of the preceding noble gas plus filled d- and f-subshells. Hence, aluminium does not suffer the effects of incomplete shielding of valence electrons by inner electrons from the nucleus that its heavier congeners do. Aluminium's electropositive behavior, high affinity for oxygen, and highly negative standard electrode potential are all more similar to those of scandium, yttrium, lanthanum, and actinium, which have ds2 configurations of three valence electrons outside a noble gas core: aluminium is the most electropositive metal in its group. Aluminium also bears minor similarities to the metalloid boron in the same group; AlX3 compounds are valence isoelectronic to BX3 compounds (they have the same valence electronic structure), and both behave as Lewis acids and readily form adducts. Additionally, one of the main motifs of boron chemistry is regular icosahedral structures, and aluminium forms an important part of many icosahedral quasicrystal alloys, including the Al–Zn–Mg class.


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