Magnesium

Magnesium is a chemical element, represented by the symbol Mg and carrying atomic number 12. This shiny gray metal has a low density, a low melting point, and is highly reactive chemically. Like other alkaline earth metals in group 2 of the periodic table, it is never found free in nature, only combined with other elements, and it almost always carries a +2 oxidation state. It quickly reacts with air, forming a thin coating of magnesium oxide that protects the metal underneath from further corrosion. When the pure metal burns, it produces a brilliant white light. Most magnesium is obtained by electrolyzing magnesium salts extracted from brine. It is less dense than aluminum and is mainly used to create strong, lightweight alloys, often with aluminum.

In space, magnesium forms inside large, aging stars when three helium nuclei fuse with a carbon nucleus. When these stars explode as supernovas, much of the magnesium is released into the interstellar medium, where it can later be incorporated into new star systems. On Earth, magnesium is the eighth most abundant element in the crust and the fourth most common element overall (after iron, oxygen, and silicon), accounting for 13% of the planet’s mass and a large portion of its mantle. It is also the third most abundant element dissolved in seawater, after sodium and chlorine.

In the human body, magnesium is the eleventh most abundant element by mass. It is essential for all cells and for the function of about 300 enzymes. Magnesium ions interact with polyphosphate compounds like ATP, DNA, and RNA. Hundreds of enzymes require magnesium ions to work. Medicinally, magnesium compounds are used as common laxatives and antacids (such as milk of magnesia), and to calm abnormal nerve activity or blood vessel spasms in conditions like eclampsia.

**Characteristics**

**Physical properties** Elemental magnesium is a gray-white, lightweight metal, with a density two-thirds that of aluminum. It has the lowest melting point (923 K, or 650 °C) and the lowest boiling point (1,363 K, or 1,090 °C) of all the alkaline earth metals. Pure polycrystalline magnesium is brittle and breaks easily along shear bands. It becomes much more malleable when alloyed with small amounts of other metals, such as 1% aluminum. Its malleability can also be greatly improved by reducing its grain size to about 1 micrometer or less.

**Chemical properties**

**Oxidation** Magnesium is widely used as a reducing agent. Although it oxidizes in air, it does not need an inert atmosphere for storage because it forms a protective layer of magnesium oxide. Direct reaction with air or oxygen at normal pressure yields only the ordinary oxide, MgO. This oxide can be combined with hydrogen peroxide to form magnesium peroxide (MgO₂), and at low temperatures the peroxide can react further with ozone to form magnesium superoxide (Mg(O₂)₂). If powdered magnesium is heated to just below its melting point, it reacts with nitrogen in the solid state to form magnesium nitride (Mg₃N₂). Magnesium reacts with water at room temperature, but much more slowly than calcium. When submerged in water, hydrogen bubbles form slowly on the metal’s surface; this reaction is much faster with powdered magnesium and also speeds up at higher temperatures. This reversible reaction with water can be used to store energy and run a magnesium-based engine. Magnesium reacts exothermically with most acids, such as hydrochloric acid, producing magnesium chloride and hydrogen gas. Although difficult to ignite in bulk, magnesium metal will ignite. It can also be used as an igniter for thermite, a mixture of aluminum and iron oxide powder that ignites only at very high temperatures.

**Reaction with water** When finely powdered, magnesium reacts with water to produce magnesium hydroxide and hydrogen gas: Mg(s) + 2 H₂O(l) → Mg(OH)₂(aq) + H₂(g). This reaction is much less dramatic than those of alkali metals because the magnesium hydroxide builds up on the metal’s surface and slows further reaction. With steam, magnesium produces magnesium oxide and hydrogen: Mg(s) + H₂O(g) → MgO(aq) + H₂(g).

**Organic chemistry** Organomagnesium compounds are common in organic chemistry, most notably as Grignard reagents. These are formed by reacting magnesium with haloalkanes or aryl halides in diethyl ether. Examples include phenylmagnesium bromide and ethylmagnesium bromide. Grignard reagents act as nucleophiles, attacking electrophilic groups such as the carbon atom in the polar bond of a carbonyl group. Another important organomagnesium reagent is magnesium anthracene, or magnesocene, used as a source of highly active magnesium. First prepared in 1954 by independent groups led by Ernst Otto Fischer and Albert Wilkinson, magnesocene is a white to off-yellow pyrophoric powder that violently hydrolyzes in water. The related butadiene-magnesium adduct provides a source of the butadiene dianion. Complexes of dimagnesium(I) have also been observed.

**Detection in solution** Magnesium ions can be detected by adding ammonium chloride, ammonium hydroxide, and monosodium phosphate to an aqueous or dilute HCl solution of the salt. A white precipitate indicates magnesium ions. Azo violet dye can also be used, turning deep blue in an alkaline solution of a magnesium salt; the color comes from the dye being adsorbed by Mg(OH)₂.

**Forms**

**Alloys** As of 2013, consumption of magnesium alloys was less than one million tonnes per year, compared with 50 million tonnes of aluminum.