Magnesium - Simple English Wikipedia, the free encyclopedia

Magnesium, ₀₀Mg

Magnesium
Pronunciation	/mæɡˈniːziəm/ ​(mag-NEE-zee-əm)
Appearance	shiny grey solid
Standard atomic weight Ar°(Mg)
	[24.304, 24.307][1]
Magnesium in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson Be ↑ Mg ↓ Ca sodium ← magnesium → aluminium
Group	group 2 (alkaline earth metals)
Period	period 3
Block	s-block
Electron configuration	[Ne] 3s2
Electrons per shell	2, 8, 2
Physical properties
Phase at STP	solid
Melting point	923 K ​(650 °C, ​1202 °F)
Boiling point	1363 K ​(1091 °C, ​1994 °F)
Density (near r.t.)	1.738 g/cm3
when liquid (at m.p.)	1.584 g/cm3
Heat of fusion	8.48 kJ/mol
Heat of vaporization	128 kJ/mol
Molar heat capacity	24.869 J/(mol·K)
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	0,[2] +1,[3] +2 (a strongly basic oxide)
Electronegativity	Pauling scale: 1.31
Ionization energies	1st: 737.7 kJ/mol 2nd: 1450.7 kJ/mol 3rd: 7732.7 kJ/mol (more)
Atomic radius	empirical: 160 pm
Covalent radius	141±7 pm
Van der Waals radius	173 pm
Spectral lines of magnesium
Other properties
Natural occurrence	primordial
Crystal structure	​hexagonal close-packed (hcp)
Speed of sound thin rod	4940 m/s (at r.t.) (annealed)
Thermal expansion	24.8 µm/(m⋅K) (at 25 °C)
Thermal conductivity	156 W/(m⋅K)
Electrical resistivity	43.9 nΩ⋅m (at 20 °C)
Magnetic ordering	paramagnetic
Molar magnetic susceptibility	+13.1·10−6 cm3/mol (298 K)[4]
Young's modulus	45 GPa
Shear modulus	17 GPa
Bulk modulus	35.4[5] GPa
Poisson ratio	0.290
Mohs hardness	1–2.5
Brinell hardness	44–260 MPa
CAS Number	7439-95-4
History
Naming	after Magnesia, Greece
Discovery	Joseph Black (1755)
First isolation	Humphry Davy (1808)
Isotopes of magnesium v e
Main isotopes[6] Decay abun­dance half-life (t1/2) mode pro­duct 24Mg 79% stable 25Mg 10% stable 26Mg 11% stable
Category: Magnesium view talk edit | references

Magnesium ( /mæɡˈniːziəm/ mag-NEE-zee-əm) is a chemical element. It has the symbol Mg, atomic number 12 and common oxidation state +2. It is an alkaline earth metal and the eighth most abundant element in the Earth's crust. It makes up 2% of the mass of the crust. It is the ninth most common element in the known universe.^[8] This is because it is easily built up in supernova stars by addition of three helium nuclei to carbon (which in turn is made from three helium nuclei). Magnesium ion's high solubility in water helps ensure that it is the third most abundant element dissolved in seawater.^[9]

Magnesium is the 11th most abundant element by mass in the human body. Its ions are essential to all living cells. The ions 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 is also the metallic ion at the center of chlorophyll. It is a common additive to fertilizers.^[10] Magnesium ions are sour to the taste, and in low concentrations help to impart a natural tartness to fresh mineral waters.

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 brilliant white light, making it a useful ingredient in flares. The metal is now mainly obtained by electrolysis of magnesium salts obtained from brine.

Commercially, the main 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.

Magnesium is used in fireworks to make a brilliant bright light. Another use is to mix it with other metals to make it strong, lightweight alloys such as those used to make bicycle frames.

Magnesium compounds are used medicinally as common laxatives, antacids (i.e. milk of magnesia), and where stabilization of abnormal nerve excitation and blood vessel spasm is required (that is, to treat eclampsia).

Magnesium is used in electronic devices, including: mobile phones, laptop computers, cameras, and other electronic components. Magnesium's low weight, good mechanical and electrical properties are good for these uses.

Magnesium reacted with an alkyl halide gives a Gringard reagent, which is a very useful tool for preparing alcohols.

Magnesium is also used in incendiary bombs, which are bombs that blow up and spread fire everywhere.

As of 2013, magnesium alloys was used less than one million tonnes per year. This was a lot less than aluminium alloys. In the same year, 50 million tonnes of aluminium alloys were used. Using magnesium alloys have limited. This is because magnesium alloys can corrode,^[11] creep at high temperatures, and combust.^[12]

Magnesium can makes many different compound. They are important in industry and biology. Some common magnesium compounds are magnesium carbonate, magnesium chloride, magnesium citrate, magnesium hydroxide, magnesium oxide, magnesium sulfate, and magnesium sulfate heptahydrate (Epsom salts).^[13][14]

Magnesium has three stable isotopes. These isotopes are ²⁴
Mg, ²⁵
Mg and ²⁶
Mg. All of these isotopes are very common in nature. About 79% of Mg is ²⁴
Mg. The isotope ²⁸
Mg is radioactive. In the 1950s to 1970s, many nuclear power plants made ²⁸
Mg for experiments. This isotope has a fairly short half-life of 21 hours.

Magnesium burns at a very high temperature, and cannot be put out with water or an ordinary fire extinguisher. It must be extinguished with a class D fire extinguisher. Burning magnesium also produces ultraviolet radiation, which can harm the eyes if it is viewed directly.

• Periodic table
• List of common elements

1. ↑ "Standard Atomic Weights: Magnesium". CIAAW. 2011.
2. ↑ Mg(0) has been synthesized in a compound containing a Na₂Mg₂²⁺ cluster coordinated to a bulky organic ligand; see Rösch, B.; Gentner, T. X.; Eyselein, J.; Langer, J.; Elsen, H.; Li, W.; Harder, S. (2021). "Strongly reducing magnesium(0) complexes". Nature. 592 (7856): 717–721. Bibcode:2021Natur.592..717R. doi:10.1038/s41586-021-03401-w. PMID 33911274. S2CID 233447380
3. ↑ Bernath, P. F.; Black, J. H. & Brault, J. W. (1985). "The spectrum of magnesium hydride" (PDF). Astrophysical Journal. 298: 375. Bibcode:1985ApJ...298..375B. doi:10.1086/163620.. See also Low valent magnesium compounds.
4. ↑ Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN 0-8493-0464-4.
5. ↑ K. A. Gschneider, Solid State Phys. 16, 308 (1964)
6. ↑ Kondev, F. G.; Wang, M.; Huang, W. J.; Naimi, S.; Audi, G. (2021). "The NUBASE2020 evaluation of nuclear properties" (PDF). Chinese Physics C. 45 (3): 030001. doi:10.1088/1674-1137/abddae.
7. ↑ Bernath, P. F.; Black, J. H.; Brault, J. W. (1985). "The spectrum of magnesium hydride" (PDF). Astrophysical Journal. 298: 375. Bibcode:1985ApJ...298..375B. doi:10.1086/163620.
8. ↑ Ash, Russell (2005). The Top 10 of Everything 2006: The Ultimate Book of Lists. Dk Pub. ISBN 0756613213. Archived from the original on 2010-02-10. Retrieved 2011-09-26..
9. ↑ Anthoni, J Floor (2006). "The chemical composition of seawater".
10. ↑ "Magnesium in health".
11. ↑ Makar, G. L.; Kruger, J. (1993). "Corrosion of magnesium". International Materials Reviews. 38 (3): 138–153. Bibcode:1993IMRv...38..138M. doi:10.1179/imr.1993.38.3.138.
12. ↑ Dodson, Brian (29 August 2013). "Stainless magnesium breakthrough bodes well for manufacturing industries". Gizmag.com. Retrieved 29 August 2013.
13. ↑ "8 Types of magnesium and their benefits". www.medicalnewstoday.com. 2021-03-23. Retrieved 2024-05-04.
14. ↑ "Chemistry of Magnesium (Z=12)". Chemistry LibreTexts. 2013-10-02. Retrieved 2024-05-04.