Stable isotopes of magnesium available from ISOFLEX
|Isotope||Z(p)||N(n)||Atomic Mass||Natural Abundance||Enrichment Level||Chemical Form|
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Magnesium was discovered in 1755 by Sir Humphrey Davy. It is a silvery-white metal, and its name originates from the Greek name Magnesia, a district of Thessaly. It occurs in all plants — its porphyrin complex, chlorophyll, is essential for photosynthesis. It is also an essential nutrient for humans.
Magnesium is a silvery-white metal with a close-packed hexagonal structure. It is soluble in dilute acids and reacts very slowly with water at ordinary temperatures. At room temperature, it is not attacked by air; however, when heated, it burns with a dazzling white light, forming the oxide MgO and the nitride Mg3N2. At ordinary temperatures, magnesium is stable in alkalis, both dilute and concentrated; however, hot solutions of alkalis above 60 ºC attack the metal. Magnesium combines with halogens at elevated temperatures, forming halides; and with nitrogen phosphorus, sulfur, and selenium at elevated temperatures, forming their binary compounds. Magnesium exhibits single displacement reactions, replacing lower metals in electrochemical series from their salt solutions or melt. Magnesium also reduces nonmetallic oxides such as carbon dioxide, carbon monoxide, sulfur dioxide and nitrous oxide, burning at elevated temperatures.
According to anecdotal history, in 1618, a farmer in Epsom, England, attempted to feed his cows water from a well. Because of the bitter taste, the cows refused to drink the water. However, the farmer noticed that the same water seemed to heal scratches and rashes. The fame of “Epsom salts” spread. Eventually the compound magnesium sulfate was recognized.
Today, magnesium has many practical applications, including the production of titanium by Kroll process and obtaining uranium from its fluoride. Magnesium alloys with aluminum, zinc, copper, nickel, lead, zirconium and other metals are used in automobile parts, aircraft, missiles, space vehicles, ship hulls, underground pipelines, memory discs, machine tools, furniture, lawnmowers, ladders, toys and sporting goods. It is also used in making small and lightweight dry cell batteries. Chemical applications of magnesium include its uses as a reducing agent, in preparation of Grignard reagents for organic syntheses, and for purifying gases. Magnesium is also used in blasting compositions, explosive sensitizers, incendiaries, signal flares and pyrotechnics.
Properties of Magnesium
|Standard state||Solid at 298 °K|
|CAS Registry ID||7439-95-4|
|Group in periodic table||2|
|Group name||Alkaline earth metal|
|Period in periodic table||3|
|Block in periodic table||s-block|
|Melting point||650 °C|
|Boiling point||1090 °C|
|Vaporization point||1090 ºC|
|Thermal conductivity||1.56 W/(m·K) at 298.2 °K|
|Electrical resistivity||4.46 µΩ·cm at 20 °C|
|Specific heat||1.05 kJ/kg K|
|Heat of vaporization||128 kJ·mol-1 at 1090 °C|
|Heat of fusion||8.7 kJ·mol-1|
|Density of liquid||1.57 g/cm3 at 650 ºC|
|Density of solid||1.74 g/cm3 at 20 °C|
|Atomic radius||1.60 Å|
|Ionic radius||Mg2+: 0.72 Å|
- Isotopic evidence for primordial molecular cloud material in metal-rich carbonaceous chondrites
- Structure of exotic Mg isotopes and temperature dependence of the symmetry energy of finite nuclei
- Magnesium isotope geochemistry in arc volcanism
- Stable isotope amount ratio analysis by using high-resolution continuum source graphite furnace molecular absorption spectrometry
- Magnesium isotope evidence that accretional vapour loss shapes planetary compositions