Known originally through compounds such as Epsom salts (the sulfate), magnesia or magnesia alba (the oxide), and magnesite (the carbonate), the silvery white element itself does not occur free in nature. It was first isolated in by Sir Humphry Davy, who evaporated the mercury from a magnesium amalgam made by electrolyzing a mixture of moist magnesia and mercuric oxide. The name magnesium comes from Magnesia, a district of Thessaly (Greece) where the mineral magnesia alba was first found.
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Magnesium is the eighth most abundant element in Earth’s crust (about 2.5 percent) and is, after aluminum and iron, the third most plentiful structural metal. Its cosmic abundance is estimated as 9.1 × 105 atoms (on a scale where the abundance of silicon = 106 atoms). It occurs as carbonates—magnesite, MgCO3, and dolomite, CaMg(CO3)2—and in many common silicates, including talc, olivine, and most kinds of asbestos. It also is found as hydroxide (brucite), chloride (carnallite, KMgCl3∙6H2O), and sulfate (kieserite). It is distributed in minerals such as serpentine, chrysolite, and meerschaum. Seawater contains about 0.13 percent magnesium, mostly as the dissolved chloride, which imparts its characteristic bitter taste.
Britannica Quiz Facts You Should Know: The Periodic Table QuizMagnesium is commercially produced by electrolysis of molten magnesium chloride (MgCl2), processed mainly from seawater and by the direct reduction of its compounds with suitable reducing agents—e.g., from the reaction of magnesium oxide or calcined dolomite with ferrosilicon (the Pidgeon process). (See magnesium processing.)
At one time, magnesium was used for photographic flash ribbon and powder, because in finely divided form it burns in air with an intense white light; it still finds application in explosive and pyrotechnic devices. Because of its low density (only two-thirds that of aluminum), it has found extensive use in the aerospace industry. However, because the pure metal has low structural strength, magnesium is mainly used in the form of alloys—principally with 10 percent or less of aluminum, zinc, and manganese—to improve its hardness, tensile strength, and ability to be cast, welded, and machined. Casting, rolling, extruding, and forging techniques are all employed with the alloys, and further fabrication of the resulting sheet, plate, or extrusion is carried out by normal forming, joining, and machining operations. Magnesium is the easiest structural metal to machine and has often been used when a large number of machining operations are required. Magnesium alloys have a number of applications: they are used for parts of aircraft, spacecraft, machinery, automobiles, portable tools, and household appliances.
The thermal and electrical conductivity of magnesium and its melting point are very similar to those of aluminum. Whereas aluminum is attacked by alkalies but is resistant to most acids, magnesium is resistant to most alkalies but is readily attacked by most acids to liberate hydrogen (chromic and hydrofluoric acids are important exceptions). At normal temperatures it is stable in air and water because of the formation of a thin protective skin of oxide, but it is attacked by steam. Magnesium is a powerful reducing agent and is used to produce other metals from their compounds (e.g., titanium, zirconium, and hafnium). It reacts directly with many elements.
Magnesium occurs in nature as a mixture of three isotopes: magnesium-24 (79.0 percent), magnesium-26 (11.0 percent), and magnesium-25 (10.0 percent). Nineteen radioactive isotopes have been prepared; magnesium-28 has the longest half-life, at 20.9 hours, and is a beta emitter. Although magnesium-26 is not radioactive, it is the daughter nuclide of aluminum-26, which has a half-life of 7.2 × 105 years. Elevated levels of magnesium-26 have been found in some meteorites, and the ratio of magnesium-26 to magnesium-24 has been used in determining their age.
The top producers of magnesium by the second decade of the 21st century included China, Russia, Turkey, and Austria.
In compounds, magnesium virtually always exhibits a +2 oxidation state because of the loss or sharing of its two 3s electrons. There are, however, a small number of coordination compounds known with magnesium-magnesium bonds, LMg―MgL, in which the magnesium centres have a formal +1 oxidation state. Magnesium carbonate, MgCO3, occurs in nature as the mineral magnesite and is an important source of elemental magnesium. It can be produced artificially by the action of carbon dioxide on a variety of magnesium compounds. The odourless white powder has many industrial uses—e.g., as a heat insulator for boilers and pipes and as an additive in food, pharmaceuticals, cosmetics, rubbers, inks, and glass. As magnesium carbonate is both hygroscopic and insoluble in water, it was the original additive used to make table salt free-flowing even in high-humidity conditions.
Magnesium hydroxide, Mg(OH)2, is a white powder produced in large quantities from seawater by the addition of milk of lime (calcium hydroxide). It is the primary raw material in the production of magnesium metal and has been used as a fire-retardant additive. In water it forms a suspension known as milk of magnesia, which has long been used as an antacid and a laxative.
The action of hydrochloric acid on magnesium hydroxide produces magnesium chloride, MgCl2, a colourless, deliquescent (water-absorbing) substance employed in magnesium metal production, in the manufacture of a cement for heavy-duty flooring, and as an additive in textile manufacture. It is also used to coagulate soy milk in the production of tofu.
Roasting either magnesium carbonate or magnesium hydroxide produces the oxygen compound magnesium oxide, commonly called magnesia, MgO. It is a white solid used in the manufacture of high-temperature refractory bricks, electrical and thermal insulators, cements, fertilizer, rubber, and plastics. It is also used medically as a laxative and antacid.
Magnesium sulfate, MgSO4, is a colourless crystalline substance formed by the reaction of magnesium hydroxide with sulfur dioxide and air. A hydrate form of magnesium sulfate called kieserite, MgSO4∙H2O, occurs as a mineral deposit. Synthetically prepared magnesium sulfate is sold as Epsom salt, MgSO4∙7H2O. In industry, magnesium sulfate is used in the manufacture of cements and fertilizers and in tanning and dyeing; in medicine it serves as a purgative. Because of its ability to absorb water readily, the anhydrous form is used as a desiccant (drying agent).
Among the organometallic compounds of magnesium are the important Grignard reagents, composed of an organic group (e.g., alkyls and aryls), a halogen atom other than fluorine, and magnesium. These are used in the production of many other kinds of organic and organometallic compounds.
Magnesium is essential to all living cells, as the Mg2+ ion is involved with the critically important biological polyphosphate compounds DNA, RNA, and adenosine triphosphate (ATP). Many enzymes depend on magnesium for their functioning. About one-sixth as plentiful as potassium in human body cells, magnesium is required as a catalyst for enzyme reactions in carbohydrate metabolism. Magnesium also is an essential constituent of the green pigment chlorophyll, found in virtually all plants, algae, and cyanobacteria. The photosynthetic function of plants depends upon the action of chlorophyll pigments, which contain magnesium at the centre of a complex, nitrogen-containing ring system (porphyrin). These magnesium compounds enable light energy to drive the conversion of carbon dioxide and water to carbohydrates and oxygen and thus directly or indirectly provide the key to nearly all living processes.
Timothy P. HanusaMagnesium 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 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]
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 about 1 μm or less.[18]
When finely powdered, magnesium reacts with water to produce hydrogen gas:
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]
The principal property of magnesium metal is its reducing power. One hint is that 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 nitrogen in the solid state if it is powdered and heated to just below the melting point, forming Magnesium nitride Mg3N2.[22]
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.[23] Although it is difficult to ignite in mass or bulk, magnesium metal will ignite.
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.
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, which is used as a source of highly active magnesium. The related butadiene-magnesium adduct serves as a source for the butadiene dianion.
Complexes of dimagnesium(I) have been observed.[24]
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.
As of , consumption of magnesium alloys 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,[25] creep at high temperatures, and combust.[26]
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.[26] 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.[26] 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.[26][27]
Magnesium's tendency to creep (gradually deform) at high temperatures is greatly reduced by alloying with zinc and rare-earth elements.[28] Flammability is significantly reduced by a small amount of calcium in the alloy.[26] 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.[29]
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).[30][31]
As recently as , magnesium hydride was under investigation as a way to store hydrogen.[32][33]
Magnesium has three stable isotopes: 24
Mg, 25
Mg and 26
Mg. All are present in significant amounts in nature (see table of isotopes above). About 79% of Mg is 24
Mg. The isotope 28
Mg is radioactive and in the s to s 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
Mg has found application in isotopic geology, similar to that of aluminium. 26
Mg is a radiogenic daughter product of 26
Al, which has a half-life of 717,000 years. Excessive quantities of stable 26
Mg 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
Al in the inclusions, and researchers conclude that such meteorites were formed in the solar nebula before the 26
Al 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
Mg/24
Mg against an Al/Mg ratio. In an isochron dating plot, the Al/Mg ratio plotted is 27
Al/24
Mg. The slope of the isochron has no age significance, but indicates the initial 26
Al/27
Al ratio in the sample at the time when the systems were separated from a common reservoir.
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.[34]
Although magnesium is found in more than 60 minerals, only dolomite, magnesite, brucite, carnallite, talc, and olivine are of commercial importance.[35]
The Mg2+
cation is the second-most-abundant cation in seawater (about 1⁄8 the mass of sodium ions in a given sample), which makes seawater and sea salt attractive commercial sources for Mg.
World production was approximately 1,100 kt in , with the bulk being produced in China (930 kt) and Russia (60 kt).[36] The United States was in the 20th century the major world supplier of this metal, supplying 45% of world production even as recently as . Since the Chinese mastery of the Pidgeon process the US market share is at 7%, with a single US producer left as of : US Magnesium, a Renco Group company located on the shores of the Great Salt Lake.[37]
In September , 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.[38]
The Pidgeon process and the Bolzano process are similar. In both, magnesium oxide is the precursor to magnesium metal. The magnesium oxide is produced as a solid solution with calcium oxide by calcining the mineral dolomite, which is a solid solution of calcium and magnesium carbonates:
Reduction occurs at high temperatures with silicon. A ferrosilicon alloy is used rather than pure silicon as it is more economical. The iron component has no bearing on the reaction, having the simplified equation:[citation needed]
The calcium oxide combines with silicon as the oxygen scavenger, yielding the very stable calcium silicate. The Mg/Ca ratio of the precursors can be adjusted by the addition of MgO or CaO.[39]
The Pidgeon and the Bolzano process differ in the details of the heating and the configuration of the reactor. Both generate gaseous Mg that is condensed and collected. The Pidgeon process dominates the worldwide production.[40][41] The Pidgeon method is less technologically complex and because of distillation/vapour deposition conditions, a high purity product is easily achievable.[40] China is almost completely reliant on the silicothermic Pidgeon process.
Besides the Pidgeon process, the second most used process for magnesium production is electrolysis. This is a two step process. The first step is to prepare feedstock containing magnesium chloride and the second step is to dissociate the compound in electrolytic cells as magnesium metal and chlorine gas.[41]
To extract the magnesium, calcium hydroxide is added to the seawater to precipitate magnesium hydroxide.[42]
Magnesium hydroxide (brucite) is poorly soluble in water and can be collected by filtration. It reacts with hydrochloric acid to magnesium chloride.[41]
From magnesium chloride, electrolysis produces magnesium.[43]
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The basic reaction is as follows:
The temperatures at which this reaction is operated is between 680 and 750 °C.[41]
The magnesium chloride can be obtained using the Dow process, a process that mixes sea water and dolomite in a flocculator or by dehydration of magnesium chloride brines. The electrolytic cells are partially submerged in a molten salt electrolyte to which the produced magnesium chloride is added in concentrations between 6–18%.[41] This process does have its share of disadvantages including production of harmful chlorine gas and the overall reaction being very energy intensive, creating environmental risks.[44] The Pidgeon process is more advantageous regarding its simplicity, shorter construction period, low power consumption and overall good magnesium quality compared to the electrolysis method.[20]
In the United States, magnesium was once obtained principally with the Dow process in Corpus Christi TX, 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:
The hydroxide is then converted to magnesium chloride by treatment with hydrochloric acid and heating of the product to eliminate water:
The salt is then electrolyzed in the molten state. At the cathode, the Mg2+
ion is reduced by two electrons to magnesium metal:
At the anode, each pair of Cl−
ions is oxidized to chlorine gas, releasing two electrons to complete the circuit:
The carbothermic route to magnesium has been recognized as a low energy, yet high productivity path to magnesium extraction. The chemistry is as follows:
C + MgO → CO + Mg
A disadvantage of this method is that slow cooling the vapour can cause the reaction to quickly revert. To prevent this from happening, the magnesium can be dissolved directly in a suitable metal solvent before reversion starts happening. Rapid quenching of the vapour can also be performed to prevent reversion.[45]
A newer 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.[46] It was reported in that this method provides a 40% reduction in cost per pound over the electrolytic reduction method.[47]
Rieke et al. developed a "general approach for preparing highly reactive metal powders by reducing metal salts in ethereal or hydrocarbon solvents using alkali metals as reducing agents" now known as the Rieke process.[48] Rieke finalized the identification of Rieke metals in ,[49] one of which was Rieke-magnesium, first produced in .[50]
The name magnesium originates from the Greek word for locations related to the tribe of the Magnetes, either a district in Thessaly called Magnesia[51] or Magnesia ad Sipylum, now in Turkey.[52] 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 , 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.[53] It was eventually recognized as hydrated magnesium sulfate, MgSO
4·7 H
2O.[54]
The metal itself was first isolated by Sir Humphry Davy in England in . He used electrolysis on a mixture of magnesia and mercuric oxide.[55] Antoine Bussy prepared it in coherent form in . Davy's first suggestion for a name was 'magnium',[55] but the name magnesium is now used in most European languages.[56]
Further discoveries about magnesium were made by the father of physical chemistry in Imperial Russia, Nikolai Beketov (–), who established that magnesium and zinc displaced other metals from their salts under high temperatures.
Magnesium is the third-most-commonly-used structural metal, following iron and aluminium.[57] The main applications of magnesium are, in order: aluminium alloys, die-casting (alloyed with zinc),[58] removing sulfur in the production of iron and steel, and the production of titanium in the Kroll process.[39]
Magnesium is used in lightweight materials and alloys. For example, when infused with silicon carbide nanoparticles, it has extremely high specific strength.[59]
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.[60] Magnesium alloys can act as replacements for aluminium and steel alloys in structural applications.[61][62]
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.[72]
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.[73] It was used as a premium feature because of its light weight in some laptops.[74]
Magnesium is flammable, burning at a temperature of approximately 3,100 °C (3,370 K; 5,610 °F),[63] and the autoignition temperature of magnesium ribbon is approximately 473 °C (746 K; 883 °F).[75] Magnesium's high combustion temperature makes it a useful tool for starting emergency fires. 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.[76][77] Magnesium filament used in electrically ignited single-use photography flashbulbs replaced this usage eventually. Magnesium powder is used in fireworks and marine flares where a brilliant light is required, and in trick self-relighting birthday candles. It was also used for various theatrical effects,[78] such as lightning,[79] pistol flashes,[80] and supernatural appearances.[81]
Magnesium is 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.[82]
Flame temperatures of magnesium and magnesium alloys can reach 3,100 °C (5,610 °F),[63] although flame height above the burning metal is usually less than 300 mm (12 in).[83] 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),[84] in carbon dioxide (forming magnesium oxide and carbon),[85] and in water (forming magnesium oxide and hydrogen, which also combusts due to heat in the presence of additional oxygen). This property[which?] 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[why?] to exclude atmosphere from the combustion.[citation needed]
In the form of turnings or ribbons, to prepare Grignard reagents, which are useful in organic synthesis.[86]
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.[97]
Magnesium reacts with haloalkanes to give Grignard reagents, which are used for a wide variety of organic reactions forming carbon–carbon bonds.[98]
Magnesium salts are included in various foods,[99] fertilizers[100] (magnesium is a component of chlorophyll[101]), and microbe culture media.[102]
Magnesium sulfite is used in the manufacture of paper (sulfite process).[103]
Magnesium phosphate is used to fireproof wood used in construction.[104]
Magnesium hexafluorosilicate is used for moth-proofing textiles.[105]
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.[106]
Magnesium intake—especially from diet—may modestly lower blood pressure and reduce risks of stroke and sudden cardiac death, but evidence is mixed, effects are small, and more robust clinical trials are needed to clarify its role in cardiovascular disease prevention.[107]
Spices, nuts, cereals, cocoa and vegetables are good sources of magnesium.[15] Green leafy vegetables such as spinach are also rich in magnesium.[108]
In the UK, the recommended daily values by the Dietary Reference Intake for magnesium are 300 mg for men and 270 mg for women.[109]
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.[110]
Numerous pharmaceutical preparations of magnesium and dietary supplements are available. Most people get enough magnesium through a healthy diet, though supplements may help in specific conditions like magnesium deficiency, pregnancy complications, or certain chronic health issues.[111]
A literature review found limited evidence supporting oral magnesium supplementation for migraine prevention, suggesting that increasing dietary magnesium intake may be a more effective option for patients open to lifestyle changes.[112]
An adult body contains 22–26 grams of magnesium,[15][113] 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[114] and can either prevent hypercalcemia or cause hypocalcemia depending on the initial level.[114] 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.[115]
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.[116] A retention of 20% or more of the injected amount indicates deficiency.[117] As of , no biomarker has been established for magnesium.[118]
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. The newborn children of mothers who received parenteral magnesium sulfate during labor may exhibit toxicity with normal serum magnesium levels.[119]
Low plasma magnesium (hypomagnesemia) is common: it is found in 2.5–15% of the general population.[120] From to , 48 percent of the United States population consumed less magnesium than recommended in the Dietary Reference Intake.[121] 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.[120] Alcoholism is often associated with magnesium deficiency. Chronically low serum magnesium levels are associated with metabolic syndrome, diabetes mellitus type 2, fasciculation, and hypertension.[122]
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. Overdose is more likely in the presence of impaired renal function.[120] Overdose is likely in cases of excessive intake of supplements. Megavitamin therapy has caused death in a child,[128] and severe hypermagnesemia in a woman and a young girl who had healthy kidneys.[129][130] 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.[114]
Plants require magnesium to synthesize chlorophyll, essential for photosynthesis.[131] 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,[132] especially in older leaves, and can be corrected by either applying epsom salts (which is rapidly leached), or crushed dolomitic limestone, to the soil.
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.[135]
Magnesium is capable of reducing water and releasing highly flammable hydrogen gas:[136]
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:[85]
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.[137]
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