Magnesium From Wikipedia, the free encyclopedia Not to be confused with Manganese. Magnesium 12Mg Be ↑ Mg ↓ Ca sodium ← magnesium → aluminium Magnesium in the periodic table Appearance shiny grey solid Spectral lines of Magnesium General properties Name, symbol, number magnesium, Mg, 12 Pronunciation /mæɡˈniːziəm/ mag-nee-zee-əm Element category alkaline earth metal Group, period, block 2 (alkaline earth metals), 3, s Standard atomic weight 24.305(1) Electron configuration [Ne] 3s2 2, 8, 2 Physical properties Phase solid Density (near r.t.) 1.738 g·cm−3 Liquid density at m.p. 1.584 g·cm−3 Melting point 923 K, 650 °C, 1202 °F Boiling point 1363 K, 1091 °C, 1994 °F Heat of fusion 8.48 kJ·mol−1 Heat of vaporization 128 kJ·mol−1 Molar heat capacity 24.869 J·mol−1·K−1 Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 701 773 861 971 1132 1361 Atomic properties Oxidation states +2, +1[1] (strongly basic oxide) Electronegativity 1.31 (Pauling scale) Ionization energies (more) 1st: 737.7 kJ·mol−1 2nd: 1450.7 kJ·mol−1 3rd: 7732.7 kJ·mol−1 Atomic radius 160 pm Covalent radius 141±7 pm Van der Waals radius 173 pm Miscellanea Crystal structure hexagonal close-packed Magnesium has a hexagonal close packed crystal structure Magnetic ordering paramagnetic Electrical resistivity (20 °C) 43.9 nΩ·m Thermal conductivity 156 W·m−1·K−1 Thermal expansion (25 °C) 24.8 µm·m−1·K−1 Speed of sound (thin rod) (r.t.) (annealed) 4940 m·s−1 Young's modulus 45 GPa Shear modulus 17 GPa Bulk modulus 45 GPa Poisson ratio 0.290 Mohs hardness 2.5 Brinell hardness 260 MPa CAS registry number 7439-95-4 History Naming after Magnesia, Greece Discovery Joseph Black (1755) First isolation Humphry Davy (1808) Most stable isotopes Main article: Isotopes of magnesium iso NA half-life DM DE (MeV) DP 24Mg 78.99% 24Mg is stable with 12 neutrons 25Mg 10.00% 25Mg is stable with 13 neutrons 26Mg 11.01% 26Mg is stable with 14 neutrons v t e · references Magnesium is a chemical element with the symbol Mg and atomic number 12. Its common oxidation number is +2. It is an alkaline earth metal and the eighth-most-abundant element in the Earth's crust[2] and ninth in the known universe as a whole.[3][4] Magnesium is the fourth-most-common element in the Earth as a whole (behind iron, oxygen and silicon), making up 13% of the planet's mass and a large fraction of the planet's mantle. The relative abundance of magnesium is related to the fact that it easily builds up in supernova stars from a sequential addition of three helium nuclei to carbon (which in turn is made from three helium nuclei).[citation needed] Due to magnesium ion's high solubility in water, it is the third-most-abundant element dissolved in seawater.[5] Magnesium is produced in stars larger than 3 solar masses by fusing helium and neon in the alpha process at temperatures above 600 megakelvins.[citation needed] The free element (metal) is not found naturally on Earth, as it is highly reactive (though once produced, it is coated in a thin layer of oxide (see passivation), which partly masks this reactivity). The free metal burns with a characteristic brilliant-white light, making it a useful ingredient in flares. The metal is now obtained mainly by electrolysis of magnesium salts obtained from brine. In commerce, the chief use for the metal is as an alloying agent to make aluminium-magnesium alloys, sometimes called magnalium or magnelium. Since magnesium is less dense than aluminium, these alloys are prized for their relative lightness and strength. In human biology, magnesium is the eleventh-most-abundant element by mass in the human body. Its ions are essential to all living cells, where they play a major role in manipulating important biological polyphosphate compounds like ATP, DNA, and RNA. Hundreds of enzymes, thus, require magnesium ions to function. Magnesium compounds are used medicinally as common laxatives, antacids (e.g., milk of magnesia), and in a number of situations where stabilization of abnormal nerve excitation and blood vessel spasm is required (e.g., to treat eclampsia). Magnesium ions are sour to the taste, and in low concentrations they help to impart a natural tartness to fresh mineral waters. In vegetation, magnesium is the metallic ion at the center of chlorophyll, and is, thus, a common additive to fertilizers.[6] Contents [hide] 1 Characteristics 1.1 Physical properties 1.2 Chemical properties 1.3 Occurrence 2 Forms 2.1 Alloy 2.2 Compounds 2.3 Isotopes 3 Production 4 History 5 Applications 5.1 As metal 5.2 In compounds 6 Biological roles 6.1 Detection in biological fluids 6.2 Disease 6.3 Magnesium overdose 7 Safety precautions for magnesium metal 8 See also 9 References 10 External links Characteristics[edit] Physical properties[edit] Elemental magnesium is a rather strong, silvery-white, light-weight metal (two-thirds the density of aluminium). It tarnishes slightly when exposed to air, although, unlike the alkali metals, an oxygen-free environment is unnecessary for storage because magnesium is protected by a thin layer of oxide that is fairly impermeable and difficult to remove. Like its lower periodic table group neighbor calcium, magnesium reacts with water at room temperature, though it reacts much more slowly than calcium. When submerged in water, hydrogen bubbles almost unnoticeably begin to form on the surface of the metal—though, if powdered, it reacts much more rapidly. The reaction occurs faster with higher temperatures (see precautions). Magnesium's ability to react with water can be harnessed to produce energy and run a magnesium-based engine. Magnesium also reacts exothermically with most acids, such as hydrochloric acid (HCl). As with aluminium, zinc, and many other metals, the reaction with HCl produces the chloride of the metal and releases hydrogen gas. Chemical properties[edit] Magnesium is a highly flammable metal, but, while it is easy to ignite when powdered or shaved into thin strips, it is difficult to ignite in mass or bulk. Once ignited, it is difficult to extinguish, being able to burn in nitrogen (forming magnesium nitride), carbon dioxide (forming magnesium oxide, and carbon) and water (forming magnesium oxide and hydrogen). This property was used in incendiary weapons used in the firebombing of cities in World War II, the only practical civil defense being to smother a burning flare under dry sand to exclude the atmosphere. On burning in air, magnesium produces a brilliant-white light that includes strong ultraviolet. Thus, magnesium powder (flash powder) was used as a source of illumination in the early days of photography. Later, magnesium ribbon was used in electrically ignited flashbulbs. Magnesium powder is used in the manufacture of fireworks and marine flares where a brilliant white light is required. Flame temperatures of magnesium and magnesium alloys can reach 3,100 °C (3,370 K; 5,610 °F),[7] although flame