{
  "id": "magnesium",
  "title": "Magnesium",
  "lang": "en",
  "source_lang": "en",
  "summary": "Magnesium (Mg, atomic number 12) is a light alkaline-earth metal essential to photosynthesis and cell energy, and widely used in alloys and industry.",
  "category": "science",
  "subcategory": "chemistry",
  "type": "concept",
  "aliases": [
    "Mg",
    "magnesium (element)"
  ],
  "generated_at": "2026-10-08",
  "model": "pro/claude-haiku-5-5",
  "url": "https://aiwiki.page/en/magnesium/",
  "translations": {
    "zh-Hans": {
      "url": "https://aiwiki.page/zh/magnesium/",
      "json": "https://aiwiki.page/zh/magnesium.json"
    }
  },
  "sections": [
    {
      "heading": "",
      "text": "Magnesium is a chemical element with the symbol Mg and atomic number 12, a member of the alkaline-earth metals in group 2 of the periodic table. It is a silvery-white, lightweight metal that is highly reactive and almost never found uncombined in nature. It is the eighth most abundant element in Earth's crust and is largely present in minerals such as dolomite and magnesite and dissolved in seawater. Because its ion carries a stable +2 charge, it behaves very differently from most metals in both chemistry and biology.\n\nMagnesium matters for two broad reasons. In living things, the ion is a cofactor for hundreds of enzymes and is central to the chlorophyll molecule that lets plants capture light, as well as to the handling of ATP, DNA and RNA. In human industry, it is the lightest structural metal in common use and is employed in alloys, refractory materials, flame retardants, fertilizers and medicines. The element's history also runs from ancient salts such as Epsom salt, through its isolation by Humphry Davy in 1808, to modern mass production by electrolysis.\n\nThis entry first covers the element's atomic and physical properties, its chemistry and compounds, and its occurrence in rocks, oceans and the cosmos. It then traces discovery and industrial extraction, followed by magnesium's biological roles in plants and cells, its nutritional and clinical significance in humans, and its uses in metals and everyday materials. It closes with unresolved questions and active research areas."
    },
    {
      "heading": "Atomic and physical properties",
      "text": "Magnesium is a chemical element with atomic number 12, so each atom has 12 protons. It sits in group 2 of the periodic table, among the alkaline earth metals, in period 3. Its electrons form a neon-like core plus two outer electrons in the 3s subshell, which makes it an s-block element. The first two ionisation energies, about 7.6 eV and 15.0 eV, are low compared with the third, about 80 eV, so the two outer electrons are removed readily while the third is not. This pattern accounts for the near-universal +2 oxidation state of the element in its compounds.\n\nIsotopes and atomic mass\n\nNatural magnesium is a mixture of three stable isotopes. ²⁴Mg accounts for 79.0%, ²⁵Mg for 10.0%, and ²⁶Mg for 11.0%, according to standard tables; another compilation gives 78.99%, 10.00% and 11.01%. The three share 12 protons but carry 12, 13 and 14 neutrons respectively. The element is therefore not truly monoisotopic. However, the variation is small enough that the standard atomic weight is quoted as a narrow range, [24.304, 24.307], and routine chemistry uses a single value of about 24.305. Isotopic differences matter mainly in geochemical tracing and in specialised isotope-ratio studies, not in ordinary stoichiometry.\n\nMacroscopic physical traits\n\nA freshly cut magnesium surface is bright, but the metal tarnishes in air to form a thin oxide coating, which dulls its appearance. Its density at room temperature is about 1.74 g/cm³, a value listed as 1.738 in common reference tables. This is roughly two-thirds of the density of aluminium, which is why magnesium is one of the lightest structural metals in practical use. The melting point is about 650 °C (923 K) and the boiling point about 1090 °C, both far below those of iron or titanium, reflecting weak metallic bonding between the atoms. The metal is also soft, with a Mohs hardness near 2.5. It crystallises in a hexagonal close-packed lattice, in which each magnesium atom is surrounded by twelve nearest neighbors at equivalent distances. Because this structure offers few easy slip planes at room temperature, the metal is brittle relative to cubic metals when cold. Heating activates additional slip systems, so the metal becomes markedly more ductile and can be rolled, extruded or forged at elevated temperatures.\n\nSize and charge of the Mg²⁺ ion\n\nLosing its two valence electrons gives the Mg²⁺ ion, which has the same electron arrangement as neon. Its six-coordinate ionic radius is about 72 pm, noticeably smaller than that of Ca²⁺ (about 100 pm). Combining a +2 charge with this small radius produces a high charge density, which has two main consequences. In water, Mg²⁺ attracts and binds its neighbouring water molecules strongly, forming a tightly held hydration shell, typically six waters in an octahedral arrangement. Stripping this shell away is energetically costly, so the hydrated ion remains large and slow to exchange water. In crystals, the high charge density produces strong electrostatic attraction to anions. Magnesium and oxygen ions combine in a strong ionic bond, making it insoluble in water, and this strong ionic bonding underlies the very high melting point of magnesium oxide. The small cation also favours octahedral coordination in minerals such as brucite and periclase, a pattern that recurs throughout crystal chemistry of the element."
    },
    {
      "heading": "Chemical reactivity and principal compounds",
      "text": "Reaction with oxygen and air\n\nMagnesium is a strong reducing agent: each atom readily gives up its two outer electrons, and most of its chemistry follows from that. Freshly exposed magnesium quickly reacts with air and forms a thin, dull grey oxide layer that hides its shiny finish. At ordinary pressure, air produces only the normal oxide, MgO, rather than peroxides or superoxides, which require more extreme conditions. Direct reaction of magnesium with air or oxygen at ambient pressure forms only the \"normal\" oxide MgO.\n\nIgnition needs an input of energy, but once started the burning is intensely exothermic. The flame supplies the heat needed for magnesium atoms to overcome their activation energy. In the combustion, magnesium gives up two electrons to oxygen atoms to form the white powdery product. The light is very bright and includes ultraviolet components, which is why welding-type eye protection is used when working with powdered magnesium. Powder and thin ribbon are the most hazardous forms, since they are highly flammable in their pure form when molten or in powder or ribbon form.\n\nThe heat of burning is high enough to drive reactions with gases that are normally unreactive. Burning magnesium can combine with nitrogen to form magnesium nitride, Mg₃N₂, and can also reduce carbon dioxide. The heat released can be so great that magnesium reacts with nitrogen and carbon dioxide, both of which are usually very unreactive. This is why magnesium fires cannot be put out with water, and why specialised extinguishing agents are needed.\n\nReaction with water, steam and acids\n\nMagnesium's behaviour in water depends strongly on temperature and on the physical form of the metal. In cold water the attack is slight. The reaction soon stops because the magnesium hydroxide formed is almost insoluble and forms a barrier on the metal that prevents further reaction. Sources differ on how to describe room-temperature behaviour: some report that hydrogen bubbles form slowly on the surface of the metal when it is submerged in water, and this happens much more rapidly with powdered magnesium. Others state that the metal reacts only with hot water and does not react with cold water.\n\nHot water speeds up the process because the protective hydroxide film dissolves more readily. Steam causes a much more violent reaction. At high temperatures the protective layer cannot keep pace with the rate at which water molecules attack the metal. Under these conditions the product is the oxide rather than the hydroxide, because metal hydroxides split on heating to give the metal oxide. The overall steam reaction is Mg + H₂O → MgO + H₂.\n\nMagnesium is oxidised and water is reduced in these reactions, which makes them redox processes in which the metal acts as the reductant. Magnesium is oxidised to its hydroxide, and water is reduced to hydrogen gas. Mineral acids such as hydrochloric acid dissolve the metal readily, releasing hydrogen and forming the corresponding magnesium salt (for example, Mg + 2HCl → MgCl₂ + H₂). Within Group 2, magnesium is less reactive toward water than calcium, strontium and barium, a trend that orders the metals as barium, strontium, calcium, magnesium and beryllium, in decreasing reactivity with water.\n\nPrincipal compounds and their solubility\n\nOxide (MgO). Magnesium oxide is a white solid, formed in combustion and by direct oxidation. It is only sparingly soluble in water, and this stability underlies its use as a high-temperature material (discussed in the applications sections).\n\nHydroxide (Mg(OH)₂). Magnesium hydroxide occurs naturally as the mineral brucite. Its solubility is low: one source gives a solubility of 0.00122 g per 100 ml in water, with a solubility product of 5.61 × 10⁻¹², while another reports about 0.00064 g per 100 mL at 25 °C, so values vary with the source and conditions. Because it is poorly soluble, it behaves as a weak base, and it dissolves in dilute acids and in solutions of ammonium salts. A water suspension is the familiar 8 percent suspension of solid magnesium hydroxide in water, sold as milk of magnesia. In the stomach, it reacts with excess acid, which reduces heartburn and upset stomach.\n\nSulfate (MgSO₄). Magnesium sulfate is highly soluble. The anhydrous salt dissolves at 35.1 g per 100 mL at 20 °C, and the heptahydrate is considerably more soluble, at 113 g per 100 mL at 20 °C.\n\nChloride (MgCl₂). Magnesium chloride is also very soluble in water and is the salt from which the metal is commonly produced by electrolysis, a point covered in the extraction section.\n\nCarbonate (MgCO₃). Magnesium carbonate is much less soluble, with a solubility of 0.0139 g per 100 mL at 25 °C. It dissolves in acid and in aqueous carbon dioxide, which is why carbonate rocks dissolve in acidic waters.\n\nThe broad pattern is that the sulfate and chloride are highly soluble, while the hydroxide, carbonate and oxide are sparingly soluble. This difference is exploited in extraction from seawater. Magnesium ions in seawater can be precipitated by adding calcium hydroxide, because calcium hydroxide is far more soluble in water than magnesium hydroxide, so the magnesium compound separates as a solid.\n\nOrganomagnesium chemistry and Grignard reagents\n\nMagnesium also forms compounds with direct carbon–metal bonds. When an organic halide reacts with magnesium metal in a dry ether solvent, it forms an organomagnesium halide, RMgX, known as a Grignard reagent. In this compound the carbon is polarised negatively relative to magnesium, so it acts as a nucleophile. Grignard reagents add to carbonyl groups and to many other electrophiles, producing new carbon–carbon bonds. This reaction was developed by Victor Grignard around 1900, and the work was recognised with the Nobel Prize in Chemistry in 1912. Because the reagents react readily with water and other protic compounds, they must be prepared and used under strictly anhydrous conditions. The same reactivity that makes magnesium a strong reductant in water also makes its organometallic derivatives powerful synthetic tools."
    },
    {
      "heading": "Occurrence and geochemical cycle",
      "text": "Abundance and the absence of the free metal\n\nEstimates of magnesium's share of the crust cluster around 2 to 2.5 percent by mass, and rankings vary by source. Magnesium (Mg) is the eighth most abundant element and constitutes about 2% of the Earth's crust, while one review puts it sixth among the metals in the earth's crust. Its ion is far more prominent in the ocean: it is the second-most abundant metal ion in seawater. Seawater holds roughly 0.13 percent magnesium, mostly as the dissolved chloride, which imparts its characteristic bitter taste. Because the metal reacts readily with oxygen and water, it is never encountered as a free element; geologists find it only inside compounds, a point noted in standard references.\n\nPrincipal minerals\n\nMagnesium appears in more than 60 minerals, but only a few are mined for it. Only brucite, dolomite, magnesite and carnallite are commercially important for their magnesium content. Olivine is also on the commercial list. Magnesite is a carbonate, \\(\\mathrm{MgCO_3}\\), and dolomite is a mixed calcium-magnesium carbonate, \\(\\mathrm{CaMg(CO_3)_2}\\), while brucite is magnesium hydroxide, Mg(OH)2. Carnallite is a potassium-magnesium chloride hydrate, \\(\\mathrm{KMgCl_3\\cdot 6H_2O}\\). Dolomite typically forms in a different way: it is formed mainly from the replacement of limestone through chemical processes. Its ordering is also distinctive, since calcium and magnesium ions alternate in equal number in an ordered structure among carbonate ions.\n\nRetention during weathering\n\nThe primary route by which magnesium leaves rock is chemical weathering. Rainwater carrying dissolved carbon dioxide forms a weak acid that attacks silicate rocks on land, and this weak acid is what releases magnesium ions into solution. Through a process termed \"weathering\", this weak acid dissolves silicate rocks on the continents, releasing Ca++, Mg++, HCO3- (bicarbonate), and SiO2 (dissolved silica) into solution. Much of the magnesium, however, is not lost from the land. It is recycled into new solids: during weathering, magnesium is deposited as either carbonate minerals via dolomitization or silicate minerals through marine clay formation and alteration of oceanic crust. Under humid conditions the pattern differs at the surface. In this climate setting, dolomite and other carbonate minerals are susceptible to weathering, so Mg is typically depleted in topsoil compared to deeper soil. Long periods of intense weathering can strip much more, as in some Piedmont soils where intense, long–term weathering has destroyed much of the original mineralogy, and leached Mg from the soil profile.\n\nWater and carbon cycles\n\nRivers are the main carriers. Rivers then transport these products to the ocean, where organisms and inorganic reactions form carbonate minerals such as calcite and aragonite. In the sea, magnesium leaves the water through two sinks: Mg is removed from seawater via the formation of Mg-bearing carbonates and Mg-bearing silicates. Magnesium carbonates are harder to form than calcium carbonates. A study of more than 850 lake and water samples from British Columbia found that for HCO3-dominated waters, the data indicate maximum alkalinity thresholds of ~5 mEq/kg for Ca-carbonates and ~40 mEq/kg for Mg-carbonates. Magnesium-rich carbonates therefore need much stronger chemical conditions before they precipitate, which helps explain why dolomite forms through more specific pathways. Carbonate sediments eventually return to the deep Earth, since at subduction zones, the carbonate sediments are buried and forced back into the mantle. This sequence links the magnesium cycle to the carbon cycle and to the long-term regulation of atmospheric CO2.\n\nThe isotopic composition of magnesium helps trace these pathways. The majority of igneous silicate rocks have similar δ26Mg values that cluster around −0.25 ± 0.07‰, whereas carbonates have much lower values of δ26Mg, commonly between −1 and −5‰. Because the two sources differ so strongly, changes in the relative contribution of carbonate to silicate sources could lead to different Mg isotope compositions of Mg input into the oceans.\n\nSettings of major deposits\n\nEvaporite basins are the classic setting for soluble magnesium salts. As seawater or brine evaporates, sylvite precipitates first from mixed solutions of K+, Mg2+ and Cl−, leaving a brine enriched in magnesium from which the mixed halide carnallite then precipitates. Lake brines and bitterns, the concentrated residues of evaporation, are therefore important sources alongside seawater. Magnesium and its compounds also are recovered from seawater, brines found in lakes and wells, and bitterns (salts).\n\nUltramafic rocks, which are rich in iron and magnesium and poor in silica, are the other major setting. Brucite is most often associated with specialists studying contact metamorphic and ultramafic rocks, and in addition, serpentine (including asbestos tailings) and olivine are possible raw materials. Extracting magnesium directly from silicates has proven difficult; the extraction of magnesium from silicates is technologically feasible but economically challenging. Magnesite deposits are concentrated in a few countries, and China is now the biggest producer of crude magnesite in the world for refractory and agricultural uses."
    },
    {
      "heading": "Discovery, naming and early history",
      "text": "Magnesium's compounds were in use long before the metal was recognised as an element. The sulfate, the oxide and the carbonate were known as Epsom salts, magnesia and magnesite, while the silvery metal itself never appears free in nature. 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. The name goes back to a place: Magnesia, a district of Thessaly (Greece) where the mineral magnesia alba was first found. Epsom salts, a hydrated magnesium sulfate, owe their name to a spring in Epsom, England, and were credited with healing effects.\n\nBefore the 1750s, magnesia was regarded as a kind of lime. Magnesia, until then generally held to be a variety of lime, formed with the same acids wholly different salts, and was consequently to be regarded as a distinct substance. Joseph Black's work on this question is the turning point. Black, a Scottish physician, studied magnesia alba as part of his medical thesis, and his fuller account was read to the Edinburgh society in June 1755 and printed in 1756. With acids, magnesia alba behaved much like chalk, giving off a gas. Heating it produced magnesia usta, which, like quicklime, did not effervesce with acids. Unlike quicklime, however, it was not caustic or water-soluble. Black attributed the weight lost on heating to the gas released, which he identified as a distinct substance, the \"fixed air\" now known as carbon dioxide. Because the two alkaline earths had been confused for so long, magnesium and calcium were once thought to be the same substance. Black's experiments showed they were not.\n\nHumphry Davy's 1808 attempt was the first to produce the metal. Working at the Royal Institution in London, he applied electrolysis to a mixture of moist magnesia and mercuric oxide, then evaporated the mercury from the resulting magnesium amalgam. He used a mercury cathode, and his early experiments produced only tiny specks of an amalgam that seemed to contain a new metal. The product was impure for two reasons. Magnesium is highly reactive and has a strong affinity for oxygen, so the metal's high reactivity and strong affinity for oxygen made separation from mercury extremely difficult. Davy himself was never able to completely separate magnesium. He first proposed the name \"magnium\", which was later changed to magnesium.\n\nThe nineteenth century produced the first reasonably pure metal and the methods that would later support industry. In 1831, the French pharmacist Antoine-Alexandre Brutus Bussy isolated purer samples of magnesium metal from magnesium chloride with the help of highly reactive potassium metal. Bussy heated anhydrous magnesium chloride with potassium, and the potassium chloride formed in the reaction could then be dissolved away, leaving the metal behind. Bussy had to prepare perfectly anhydrous magnesium chloride, because any water would ruin the reaction, and he also needed to exclude oxygen so the product would not oxidise immediately. His process yielded a usable quantity of metal for the first systematic studies of its properties.\n\nElectrolysis returned to the picture in 1833, when Michael Faraday obtained magnesium metal by the electrolysis of fused anhydrous magnesium chloride. In 1852, Robert Bunsen developed an electrolytic cell that optimized the isolation of magnesium via the Faraday route. Bunsen's design, which used Davy's underlying concept in a bath of molten chloride, became the basis for later production. Commercial output followed only decades later, when commercial-level production first began in 1863, using the reduction techniques of Deville and Caron."
    },
    {
      "heading": "Extraction and industrial production",
      "text": "Magnesium is almost never mined as metal. Producers start from a chloride or an oxide-bearing rock and reduce the magnesium ion to metal, using either electrolysis or a thermal reduction step. Primary production is defined as metal smelted from ores and brines, which excludes recycled metal and compounds such as magnesia. Primary production covers magnesium metal smelted from ores and brines, by electrolysis of magnesium chloride or by silicothermic reduction of dolomite.\n\nThe electrolytic route\n\nElectrolytic recovery needs a chloride feed, and that feed is usually prepared from seawater or brine (electrolytic recovery of magnesium requires a magnesium chloride feedstock that normally is prepared from seawater or brines). In the Dow process, partly dehydrated magnesium chloride can be obtained by mixing seawater with lightly burned reactive dolomite in a flocculator. The insoluble hydroxide that forms is collected, reacted with hydrochloric acid, and dried by evaporation to about 25 percent water (converted to magnesium chloride by reaction with hydrochloric acid, and dried in a series of evaporation steps to 25 percent water content). Some plants instead dry the chloride to an anhydrous form, which Norsk Hydro favoured.\n\nThe Dow route began in Michigan. Dow's first pure metal was produced in 1916 at a pilot plant in Midland, using brine that also contained far more sodium chloride than magnesium chloride; lime was needed to precipitate the hydroxide first (While MgCl_2 was present in the brine, so was NaCl at a much higher concentration, so chemical processing was required to isolate MgCl_2. Calcium oxide was used to react and result in solid Mg(OH)_2). Because brine supply was limited, the company moved to seawater, and its largest plant was built near Freeport, Texas (When Dow set out to construct the U.S.'s largest magnesium plant, it found a perfect site near Freeport, Texas.). Dow closed its magnesium operation in November 1998 (Dow closed its magnesium operation in November 1998.).\n\nIn the cell, the molten chloride is split into metal and chlorine. The bath is a salt mixture of sodium, calcium and magnesium chlorides, with a graphite anode and a steel cathode pot. The metal rises to the surface and is skimmed off, while chlorine is released at the anode (The composition of the bath is 54% NaCl, 23% CaCl_2, and 22% MgCl_2. The newly formed magnesium is scooped out after the accumulation at the surface of the bath.). Chlorine is also fed back into the process: it converts magnesium oxide into the chloride needed by the cells (Chlorine gas introduced through portholes in the furnace reacts with the magnesium oxide to produce molten magnesium chloride).\n\nThe thermal (silicothermic) route\n\nThe main alternative is the Pidgeon process, a vacuum reduction of calcined dolomite. Calcined dolomite is pelleted with ferrosilicon as the reductant and a little fluorite as catalyst, then heated in an externally fired tank. Under the catalysis of fluorite, silicon atoms in ferrosilicon reduce the magnesium ions of MgO to metallic magnesium. Metallic magnesium sublimates at high temperature into magnesium vapor. The vapour is then condensed to solid metal. Estimates for the thermal route vary; one conference study puts the raw material demand at about 11 tonnes of dolomite and 1 tonne of ferrosilicon per tonne of magnesium (Pidgeon process consumes 11 tons of dolomite and 1 ton of ferrosilicon per ton of magnesium.). The process is widely used: over 80 % of global Mg metal production relies on the Pidgeon process.\n\nEnergy and cost\n\nBoth routes are energy-intensive. One U.S. Bureau of Mines-related estimate put the Dow process at about 358 million BTU per short ton, with roughly 40 percent of that consumed before the electrolysis step (the Dow process consumes 358.times.10.sup.6 BTU per short ton of magnesium produced. Of this energy consumption, about 40% is used up in the process stages leading up to electrolysis.). A patent's background section claims that the preparation of the chloride feed accounts for about half of the cost and energy of electrolytic production (Fifty percent of the cost and energy consumption for the production of magnesium is reported to come from the preparation of magnesium chloride.). The Pidgeon route is described as energy-intensive but cost-effective, and Chinese producers are said to reduce costs by using waste heat from coal gas production (Chinese magnesium producers can achieve a low cost because they take advantage of waste heat energy associated with coal gas production to drive the process).\n\nGrowth of output and recent producers\n\nIndustrial production began in Germany in 1886 and dominated it until the Second World War. Wartime demand then pushed production elsewhere: during the Second World War, Dow Chemical and Magnesium Elektron in the UK began electrolytic reduction using seawater from Galveston Bay and the North Sea. After the war, Dow broadened civilian markets through wrought products and surface treatment systems (Dow Chemical broadened civilian markets by developing wrought products, photoengraving technology, and surface treatment systems.). Later refinements included internal heating of retorts from 1961 and improved cell technology from about 1970 (improvements in electrolytic cell technology from about 1970).\n\nIn recent years output has been highly concentrated. In 2024 world primary production was about 1,080,000 tonnes, of which China produced 953,000 tonnes. The remaining production came chiefly from Russia, Israel, Brazil, Kazakhstan, Turkey and Iran. Estimates of China's share differ by source and year, from about 85 percent to over 90 percent. One market analysis reports that US primary magnesium production was zero from 2023 to 2025, and the sole producer US Magnesium shut down completely after bankruptcy in 2025. Within China, output is concentrated in a few provinces, notably Liaoning, Shaanxi, and Shanxi.\n\nSecondary metal from scrap is treated separately from primary output. Scrap is remelted and refined, and this recycled metal is not counted in primary production statistics."
    },
    {
      "heading": "Magnesium in plants and photosynthesis",
      "text": "Magnesium's most visible role in plants is inside chlorophyll. A magnesium ion sits at the centre of the tetrapyrrole core, and a long hydrocarbon chain called phytol is attached to the ring and anchors the pigment in membranes. The most common form of chlorophyll found in plants is chlorophyll a. The conjugated ring system soaks up red and blue-violet light, and it is able to absorb red, blue, and violet wavelengths of light while reflecting green light. Within the pigment, the central metal ion is usually described as an electron acceptor that helps pass on the energy of absorbed photons, which feeds into the photosynthetic light reactions.\n\nWhen magnesium runs short, the first visible sign is usually a loss of colour in older leaves. Since magnesium is mobile within the plant, symptoms often first appear on lower areas typically towards older leaves before spreading upwards. The characteristic pattern is interveinal chlorosis, where the areas between leaf veins turn yellow while the veins themselves remain green. The shortage also affects the pigment's supply chain. Mg deficiency restrains chlorophyll synthesis and promotes chlorophyll degradation, subsequently reducing the photosynthetic electron transport rate. In a 16-week experiment on citrus seedlings grown without supplemental magnesium, interveinal chlorosis, vein enlargement and corkiness, and alterations of gas exchange and pigments were observed in middle and lower leaves. Damage reached the photosystem II complex, and the structural damage to thylakoids and inactivation of the oxygen-evolving complex lowered energy transfer and the rate of CO2 assimilation. Species respond differently: in rice and cucumber grown in hydroponics, both Mg-deficient plants had lower biomass, leaf area, and chlorophyll content, but cucumber showed a more marked decline in chlorophyll and carotenoid content.\n\nInside the chloroplast, magnesium plays a second role as a regulatory signal. Its concentration is held within narrow limits by dedicated transporters, and three different transporters from two distinct families have been identified in the chloroplast inner envelope of the model plant Arabidopsis thaliana. Illumination changes the free magnesium pool: the light-induced increase in the pH and the Mg2+ concentration in the stroma is thought to be obligatory for light activation of Calvin cycle enzymes. Magnesium also contributes to the ribosomal machinery, where it stabilises rRNA folding and helps hold the two ribosomal subunits together, so protein synthesis in plastids as in the cytosol depends on an adequate supply.\n\nThe Calvin cycle illustrates magnesium's dual character. Magnesium is required for key CBB cycle enzymes, as well as rubisco activation. In Rubisco, the enzyme is activated when carbamylated (CO2 added to Lys201) and with Mg++ bound. Too much free magnesium can also be counterproductive. In isolated barley chloroplasts, two millimolar Mg(2+) blocked the light activation of phosphoribulokinase, NADP-linked glyceraldehyde-3-phosphate dehydrogenase, and fructose 1,6-diphosphatase, which shows that the light-driven regulation of these enzymes depends on tightly balanced magnesium levels.\n\nFor agriculture, magnesium has received less attention than nitrogen, phosphorus and potassium. Mg-deficiency, one of the most common physiological disorders influencing productivity and quality of crops, is being intensified in intensive agricultural systems, where N, P, and K fertilizers have been heavily applied. Diagnosis is also difficult, because it is difficult to detect latent Mg deficiency. Soil magnesium comes mainly from the weathering of silicate and carbonate minerals and from exchangeable ions held on clay particles. Sandy, acidic or heavily leached soils are most prone to shortage, and heavy potassium inputs can suppress magnesium uptake. Remedies include liming with magnesium-containing materials on acid soils and, in some cases, foliar treatment that promotes photosynthesis and combats magnesium deficiencies."
    },
    {
      "heading": "Biochemistry in living cells",
      "text": "Inside cells, magnesium acts mainly as the divalent cation Mg²⁺, and most of its biological roles follow from that charge and its small, tightly hydrated size. Its most widespread function is as a cofactor for enzymes. One widely cited count puts the number of enzymatic reactions that depend on it at more than 300, a critical cofactor in more than 300 enzymatic reactions. Kinases are a prominent example, since magnesium ions are required by a large number of enzymes referred to as kinases. Enzymes of the citric acid cycle and the electron transport chain also require or are regulated by magnesium, including 2-oxoglutarate dehydrogenase and the mitochondrial ATP synthase complex.\n\nMagnesium's closest partner in energy metabolism is ATP. Its phosphate groups carry several negative charges, and Mg²⁺ binds them, so that ATP must be complexed to magnesium in order to be biologically active. In practice, the species recognised by many enzymes is the Mg–ATP complex. Magnesium is also needed by the ATPases that keep intracellular electrolyte levels in order, as well as by many other ATP-driven pumps located in different cell compartments in maintaining adequate intracellular electrolyte content. Because ATP is central to metabolism, this coupling makes magnesium status a constraint on cellular energy use as a whole.\n\nMagnesium also shapes the chemistry of DNA and RNA. Nucleic acids are polyanions, and their phosphate backbones must be neutralised by positive counter-ions. Mg²⁺ does this effectively, and it can bind both diffusely, through electrostatic attraction to the surrounding negative field, and at specific sites. Experimental work has shown that the diffuse mode and the site-specific mode make different energetic contributions to RNA folding RNA stabilization arises from two energetically distinct modes of Mg2+ binding: diffuse- and site-binding. In transfer RNA, Mg²⁺ holds the native tertiary fold together even when monovalent ions are present in large excess Mg21 strongly stabilized the native tertiary structure of tRNA even in the presence of a large excess of monovalent ions. Magnesium is also involved in genetic processes. It helps keep ribosomes aggregated into polysomes able to initiate protein synthesis, and it is a required cofactor in several DNA repair pathways Mg2+ is required for the removal of DNA damage. It also supports DNA ligase and DNA polymerase activity during DNA synthesis. Some RNA catalysts, known as ribozymes, depend on it as well, and Mg²⁺ can serve as their cofactor when they cleave target mRNA Mg2+ can also act as a cofactor for ribonucleic acid enzymes (ribozymes) capable of specifically recognizing and cleaving the target mRNA.\n\nMagnesium's role in controlling ion movement across the cell-membrane is closely tied to its relationship with calcium, potassium and sodium. Among intracellular cations, magnesium is the second most abundant cellular cation after potassium. Its transport is handled partly by ion channels. Stretch-activated cation channels in mammalian cells are permeable for Mg2+ ions at different extracellular concentrations including physiological ones, and members of the TRP channel family, particularly TRPM6 and TRPM7, are regarded as key effectors of magnesium uptake. Magnesium also affects the flux of other ions. It inhibits Ca2+ influx by blocking the voltage-dependent calcium channels in some experimental systems, and it can modulate potassium channels, which in turn affect how cells repolarise. Because these ions compete for the same channels and transport systems, changes in one tend to ripple into the others, and this interplay is part of how cells maintain homeostasis. Even so, our understanding of how cells regulate Mg2+ homeostasis and transport still remains incomplete."
    },
    {
      "heading": "Human physiology, nutrition and clinical use",
      "text": "Intake, body stores and food sources\n\nAdult recommended intakes sit at roughly 400–420 mg per day for men and 310–320 mg for women, with the lower figures for younger adults. The RDAs for magnesium for adults between ages nineteen and thirty are 400 milligrams per day for males and 310 milligrams per day for females. For adults above age thirty, the RDA increases slightly to 420 milligrams per day for males and 320 milligrams for females. Actual consumption often falls short. In a US national survey covering 2017 to March 2020, average intake was 339 mg/day for men (ages ≥19 years) and 269 mg/day for women, and 57% of men and 52% of women had inadequate dietary intakes relative to the estimated average requirement.\n\nTotal body content is usually given as about 20–25 g. One common estimate puts an adult body at around 25 gram (g) of magnesium, 50–60% of which the skeletal system stores, with the remainder found in muscle, soft tissues, and bodily fluids. Skeletal muscle is itself a major reservoir. Good dietary sources include green leafy vegetables, unrefined grains, legumes, beans, and nuts. Absorption is mainly in the distal small bowel through paracellular pathways and a saturable transport system, and it becomes more efficient when intake is low. Vitamin D may influence uptake, although the relationship is not clear.\n\nRenal regulation of magnesium balance\n\nBody magnesium is controlled by three linked processes: intestinal absorption, renal reabsorption, and storage in bone. The kidney is the principal control point. Approximately 70% to 80% of plasma magnesium in ionized or complexed forms is ultrafilterable in the kidneys. Reabsorption happens in two segments. In the thick ascending limb it proceeds passively through the paracellular route, whereas in the distal convoluted tubule the final urinary concentration is set by an active transcellular process. Renal magnesium handling at the DCT occurs via the apical TRPM6, a channel whose delivery to the membrane depends on epidermal growth factor signalling. Serum magnesium itself is the main signal: the serum magnesium concentration is the major regulator of renal magnesium reabsorption, so a rapid rise after intravenous replacement reduces the stimulus for retention. Diagnostic work-up relies on serum levels and a 24-hour urine collection, where a 24-hour urine Mg2+ higher than 24 mg/day suggests renal Mg2+ wasting, while lower values point to poor intake or gastrointestinal loss.\n\nCauses of hypomagnesaemia\n\nHypomagnesaemia is defined by a low serum level; one clinical threshold is serum magnesium concentration < 1.8 mg/dL (< 0.70 mmol/L). Mechanistically it arises from decreased gastrointestinal absorption or increased renal Mg2+ excretion. Reduced intake is a common starting point, often through starvation or malabsorption. Gastrointestinal losses include prolonged diarrhoea, while heavy alcohol use lowers both food intake and renal retention. Loop and thiazide diuretics are a frequent medication cause: inhibiting the NKCC2 transporter in the thick ascending limb can produce excessive urinary magnesium loss. Hypercalcaemia and aminoglycoside antibiotics also promote renal wasting, and mutations in the genes that code for claudin-16 and claudin-19 cause a hereditary disease, familial hypomagnesemia with hypercalciuria and nephrocalcinosis. Because hypocalcemia is a common manifestation of moderate-to-severe magnesium depletion, calcium abnormalities often accompany the low magnesium.\n\nPhysiological roles and associations\n\nMagnesium is central to neuromuscular excitability. Deficiency is linked with muscle cramps, fatigue, appetite loss, and disruptions in calcium and potassium homeostasis, and with abnormal muscle excitability as well as convulsions, psychiatric disturbances. Cardiac rhythm is similarly sensitive, which is why intravenous magnesium is used against certain arrhythmias (see below). Observational studies associate low magnesium with increased incidence of diabetes mellitus, metabolic syndrome and coronary mortality, and with type 2 diabetes and hypertension more broadly. These are associations rather than established causes. A small crossover study of intravenous magnesium sulfate in healthy volunteers found that serum glucose concentration increased in the treated group, so the effect of magnesium on glycaemia is not simple. Its role in blood pressure regulation remains an area of ongoing research.\n\nClinical use\n\nOral magnesium is used to correct deficiency, but supplemental intake carries limits: the Tolerable Upper Intake Level (UL) is 350 mg/day, and this applies specifically to supplemental magnesium. Intravenous magnesium sulfate is the standard response to severe or symptomatic deficiency, and the goal of maintaining the serum magnesium concentration of more than 1.0 mg/dL guides treatment. Its established indications are narrower than its historical use. Guideline-level evidence supports magnesium for eclampsia and preeclampsia, and for torsade de pointes and managing rapid atrial fibrillation. In severe acute asthma, parenteral magnesium supplementation improves peak expiratory flow rate and forced expiratory volume in one second, but routine use in all acute asthma is not supported. Evidence for migraine is weaker, and for other settings it is contested: one review notes that the use and efficacy of the drug, however, are controversial, and a 2024 meta-analysis in spinal surgery described the evidence as partial and controversial. Safety depends on kidney function, since parenteral use in the presence of renal insufficiency may lead to magnesium intoxication."
    },
    {
      "heading": "Alloys and structural applications",
      "text": "Why magnesium is alloyed\n\nMagnesium is the lightest of the structural metals in common use, and its low density (about 1.7 g/cm³) is the main reason it is chosen for components where mass matters. Pure magnesium is too soft and creeps too readily for most structural roles, so practical parts are almost always alloys. The dominant cast families are Mg–Al alloys of the AZ and AM types, rare-earth-bearing alloys such as WE43 and E21 that retain strength at higher temperatures, and zinc–zirconium ZK alloys. Aluminium improves strength and castability, zinc contributes to strength, manganese is added mainly to improve corrosion resistance, and rare earths are used for elevated-temperature performance. Machinability and good castability are the other reasons for its use: parts can be cast in complex shapes with fewer tooling steps than many steel assemblies require.\n\nThe limits are equally specific. Magnesium's creep resistance is weaker than aluminium's, so it is not used for the largest, hottest engine parts such as the block and cylinder head, and useful service temperatures are commonly capped near 175 °C. Transmission housings, which run cooler, are a more natural fit.\n\nEarly automotive and aerospace use\n\nMagnesium entered motor sport first. By the 1920s it was used in racing cars, where saving mass directly improved acceleration and speed. Its use in production followed. A die-cast magnesium piston was made in 1925 by Elektronmetal GmbH, a Mahle subsidiary, and the post-war period brought magnesium body and frame panels to racing designs such as the Mercedes 300SL and the Porsche 962 (the latter as a competition car rather than a road model).\n\nThe Volkswagen Beetle is the best-known production example, though sources differ on the details. Some histories date the first magnesium transmission housing for its air-cooled engine to the mid-1930s, while others credit the 1950s Beetle with a magnesium gearbox and engine blocks. Reported magnesium content also varies, from about 17 kg for crankcase and transmission housings to roughly 20 kg for the whole car. In the 1970s Volkswagen cast air-cooled engine blocks from Mg–Al alloys with silicon additions, but this practice was abandoned as engine operating temperatures rose and the metal's cost increased.\n\nAerospace has a similarly long record, with weight-critical civil and military airframe and engine components. Today magnesium is generally described as the third most used structural metal in automobiles after steel and aluminium. Cast magnesium is typically found in instrument-panel beams, seat frames, steering-column parts and gearbox housings, mostly in the front and upper parts of the vehicle where weight savings have the greatest effect on handling.\n\nCorrosion behaviour and surface protection\n\nCorrosion is the main obstacle to wider use. Magnesium is the most electrochemically active of the common structural metals, and its corrosion is dominated by hydrogen evolution at the cathodic sites. When a painted or otherwise coated part is scratched, the exposed magnesium becomes anodic relative to the surrounding material and can corrode locally at a rate that may be more damaging than if the part had never been coated. Unlike zinc on galvanised steel, no reliable self-sacrificing coating for magnesium parts had been identified in the patent literature examined, so protection relies on barrier layers, conversion treatments, anodising, decorative electroplating and organic paint systems.\n\nOne notable development is the magnesium-rich primer. Magnesium powder pigmented into organic coatings acts as a sacrificial pigment that protects aerospace aluminium alloys, offering an alternative to chromate-based pretreatments, which are carcinogenic and widely regulated.\n\nMagnesium as a sacrificial anode\n\nA sacrificial anode works because a more reactive metal corrodes preferentially and so protects a less reactive structure. Magnesium is used this way on buried and submerged steel such as pipelines, ships, bridges, drilling platforms and heat exchangers. Magnesium and zinc are the most common galvanic anodes for pipelines. Because magnesium's potential is roughly 1 V more negative than steel's, a single anode protects only a limited length of pipe. Alloying tunes the behaviour: adding aluminium or zinc reduces the anode's rate of self-dissolution and helps avoid over-protection, and aluminium-bearing alloys have been studied for steel reinforcement in chloride-contaminated concrete."
    },
    {
      "heading": "Other industrial, agricultural and everyday uses",
      "text": "Most magnesium compounds that find industrial use are not the metal itself but its oxide and hydroxide, and the largest volume is in high-temperature engineering. Brick containing roughly 80–95% MgO is the usual lining for the hottest zones of cement kilns, steelmaking furnaces and non-ferrous smelters (Magnesia bricks contain 80–95% MgO (magnesium oxide), manufactured from dead-burned magnesite or fused magnesia.). The oxide's value rests on its melting point, which is high for a moderately priced oxide (Magnesium oxide has the highest melting point of the moderately priced oxides and is therefore an important raw material for refractory bricks and other materials.). Basic slags attack acidic linings, so the chemical match also matters: at about 1700 °C, dense magnesia can dissolve several times its own weight of iron oxide while remaining solid (Hard-fired magnesia bricks have the desirable properties of high strength at high temperature, a reasonable spalling resistance and excellent slag resistance.). In steel plants, carbon-bearing magnesia bricks with up to about 30% carbon have displaced older magnesia–chrome designs in electric-arc furnace walls (Magnesia-carbon bricks containing up to 30% carbon have been developed and have successfully replaced magnesia-chrome bricks in the side walls of electric-arc steelmaking furnaces.). The chrome-based bricks fell out of favour partly because of their environmental and disposal problems, since their heated chromium can form a hazardous hexavalent species (Because of the formation of hexavalent chrome which is classified as a hazardous substance by the federal government and which consequently posed a number of problems).\n\nMagnesium hydroxide serves a different purpose in polymers, acting as a halogen-free flame retardant. Older additives such as organohalogen compounds and antimony trioxide can generate considerable smoke and toxic gases in a fire, which has pushed manufacturers toward alternatives (such additives can generate a significant amount of smoke and toxic gases upon exposure to flame. As a result, attempts have been made to minimize their use in favor of less toxic alternatives. More recently, the use of magnesium hydroxide particles or aluminum hydroxide particles as a safer alternative has become increasingly prevalent.). Its action is largely physical. Decomposition starts at roughly 330 °C, and the reaction is endothermic, so it draws heat out of the surrounding polymer (The endothermic decomposition commencing at about 330^o C for magnesium hydroxide (versus about 230^o C for ATH) withdraws heat from the substrate, slowing the rate of thermal degradation of the plastic.). The water vapour released dilutes the flammable gases (The release of water vapor upon decomposition of magnesium hydroxide dilutes the fuel supply present in the gas phase.), and the residue can form an insulating char. Its higher thermal stability also lets it be used in polymers such as polypropylene and nylon, which are difficult to filter with aluminium hydroxide (Higher stability allows the use of higher plastics processing temperatures and thus enables a filler based on magnesium hydroxide to be used in a wider range of polymers, especially polypropylene and thermoplastic polyamides such as nylon). The trade-off is loading: fillers can make up a large share of the plastic's mass, which can reduce tensile strength and other mechanical properties (Including large amounts of a particulate filler, however, can have a negative effect on the tensile strength and other mechanical properties of the plastic material.).\n\nIn textile work, magnesium sulfate, or Epsom salt, plays a supporting rather than primary role. A mordant is a substance that binds a dye to the fibre. An early US patent described fixing magnesium hydroxide within cotton by padding the fabric with a magnesium salt, preferring Epsom salt as the cheapest soluble option (By preference I employ magnesium sulfate (Epsom salts) as being the cheapest known soluble salt of magnesium.). In modern dyebaths the salt is mainly used as an electrolyte that helps dye uptake and stabilises pH (As a magnesium salt, it serves multiple functions in the dyeing process, from acting as an electrolyte to improving dye adherence and stabilizing the dye bath's pH.). In meta-mordant methods, it is added alongside aluminium sulfate, while aluminium and iron salts remain the main mordants (Zinc sulfate and magnesium sulfate were used as additives to provide an effective simultaneous mordanting and dyeing process (meta mordant) for wool yarns with madder dyes extracted by ultrasound. In the dyeing procedure, aluminum sulfate and ferrous sulfate were used as mordants.).\n\nMagnesium compounds also enter agriculture and animal nutrition. As a secondary macronutrient, magnesium sits at the centre of the chlorophyll molecule, so a shortage shows first as yellowing between the veins of older leaves. Soils that are sandy or acidic, or heavily dosed with potassium fertilizer, are most prone to deficiency, which is corrected with magnesium-bearing fertilizers such as dolomitic lime, kieserite or Epsom salt. In livestock, magnesium oxide and magnesium sulfate are added to rations to prevent low blood magnesium. Grazing ruminants on lush spring pasture are especially at risk of grass tetany, a condition of muscle spasm that can be fatal if untreated.\n\nMagnesium's most dramatic everyday use is light. Its combustion releases a large amount of energy and produces a brilliant white flame, which is why it is used as fuel in fireworks, signal flares and illumination flares. In the late nineteenth century, magnesium ribbon and powder gave photographers a short, intense light source, and magnesium flash powder, mixed with an oxidiser, was widely used until flashbulbs replaced it. Because the hot metal also reacts with water to release hydrogen, water cannot be used to extinguish a magnesium fire; specialised dry powders are needed instead."
    },
    {
      "heading": "Open questions and research frontiers",
      "text": "Several research areas that involve magnesium remain unsettled. Some are engineering problems, some are electrochemical, and some concern whether supplementation changes health outcomes. Each depends on measurements and models that are still being refined.\n\nBiodegradable magnesium implants\n\nMagnesium is attractive for temporary implants because the body can absorb it, which removes the need for a second operation to take out hardware. Its main drawback is speed. Magnesium reacts with water to form magnesium hydroxide and hydrogen gas, and the rate of hydrogen evolution is equivalent to the corrosion rate. A fixation device therefore has to lose strength at a pace that matches bone healing. If it corrodes too fast, it fails mechanically before the fracture is stable. Research has focused on three approaches: purifying the alloy to remove iron, nickel and copper, which accelerate corrosion; adding small amounts of zinc, manganese, calcium or rare earth elements; and applying surface coatings such as calcium phosphate, polymers or metal–organic frameworks. Reviews generally conclude that no single strategy fully solves the problem, and that the gas produced during degradation and the behaviour of vascular stents in vivo still need better understanding. Clinical adoption of magnesium bone implants has therefore remained limited.\n\nMagnesium-based batteries\n\nMagnesium offers a divalent metal anode that is abundant and, in principle, does not form dendrites in the way lithium can. Its practical limits, however, are electrochemical. In most non-aqueous electrolytes, the metal surface quickly forms a dense passivation film. This layer, which can contain magnesium oxide, hydroxide and products from salt decomposition such as magnesium fluoride, conducts Mg²⁺ poorly, so plating and stripping become irreversible. Early chloride-containing electrolytes improved reversibility but attacked cathode current collectors at higher voltages, which narrowed the usable electrochemical window. The cathode is a second problem. Because Mg²⁺ carries two charges in a small ion, it binds strongly to many host lattices, which slows its diffusion through the cathode and makes many materials that work with lithium ineffective with magnesium. Current research concentrates on halide-free electrolytes, artificial interphases, and cathode architectures designed for divalent ions. Commercially viable electrolytes with wide enough stability windows have not yet been established, so magnesium cells remain at the laboratory and prototype stage.\n\nSupplementation and health outcomes\n\nThe evidence on magnesium supplementation is most reliable for blood pressure. A 2025 meta-analysis of 38 randomized controlled trials with 2,709 participants found reductions of about 2.8 mm Hg systolic and 2.1 mm Hg diastolic compared with placebo, with larger reductions in people taking blood pressure medication and in people with low magnesium status. The authors also found high heterogeneity between studies and no clear dose–response pattern. For blood sugar, a meta-analysis of 23 trials in type 2 diabetes reported a modest fall in fasting glucose but only minimal change in glycated haemoglobin. A separate review of prediabetic adults did not find a statistically significant fasting glucose effect. Most sleep and cognition data are observational or come from small trials, and results are inconsistent. Many outcomes studied are surrogate measures rather than hard clinical endpoints such as cardiovascular events. Most researchers therefore regard the evidence for these uses as suggestive but incomplete, and better-designed, larger trials are still needed.\n\nMeasuring free intracellular magnesium\n\nMost intracellular magnesium is bound to proteins, nucleotides and membranes, and the free Mg²⁺ concentration in the cytosol is commonly estimated at around 0.5–1 mM, far below the total content. Measuring the free fraction is difficult. Fluorescent indicators such as mag-fura-2 and mag-indo-1 respond to Mg²⁺, but they also bind calcium, have limited selectivity, and are affected by pH, so calibration and controls are essential. Their low affinity for magnesium can also saturate or lose sensitivity in some compartments. Phosphorus-31 nuclear magnetic resonance can infer magnesium binding from shifts in ATP and related phosphate signals, but it lacks the spatial resolution to map individual cells. Open questions include how free magnesium changes within mitochondria, the endoplasmic reticulum and the nucleus, how quickly it responds to hormonal or metabolic signals, and whether small shifts in free concentration have measurable effects on enzymes and channels."
    }
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  ],
  "sources": [
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    },
    {
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    }
  ],
  "license": "https://creativecommons.org/licenses/by/4.0/"
}