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Magnesium,  12Mg
Magnesium ribbon.jpg
Magnesium ribbon stored in mineral oil to prevent oxidation
General properties
Name, symbol Magnesium, Mg
Appearance Silvery-white
Magnesium in the periodic table


Atomic number 12
Standard atomic weight (Ar) 24.305
Group, block 2 (alkaline earth metals); s-block
Period period 3
Electron configuration [Ne] 3s2
per shell
2, 8, 2
Physical properties
Phase 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  1.584 g/cm3
Heat of fusion 8.48 kJ/mol
Heat of 128 kJ/mol
Molar heat capacity 24.869 J/(mol·K)
Atomic properties
Oxidation states +2, +1 ​(a strongly basic oxide)
Electronegativity Pauling scale: 1.31
energies 1st: 737.7 kJ/mol
2nd: 1450.7 kJ/mol
3rd: 7732.7 kJ/mol
Atomic radius empirical: 160 pm
Covalent radius 141±7 pm
Van der Waals radius 173 pm
Crystal structure ​Hexagonal close-packed (hcp)
Speed of sound thin rod 4940 m/s (at ) (annealed)
Thermal expansion 24.8 µm/(m·K) (at 25 °C)
Thermal conductivity 156 W/(m·K)
Electrical resistivity 4.39·10-8 Ω·m (at 20 °C)
Magnetic ordering Paramagnetic
Young's modulus 45 GPa
Shear modulus 17 GPa
Bulk modulus 45 GPa
Poisson ratio 0.29
Mohs hardness 1–2.5
Brinell hardness 44–260 MPa
CAS Registry Number 7439-95-4
Naming After Magnesia, Greece
Discovery Joseph Black (1755)
First isolation Humphry Davy (1808)
· references

Magnesium is an alkaline earth metal with the symbol Mg and atomic number 12. It is one of the most powerful reducing agents available to the amateur chemist, even more so than zinc. However, it is impossible to plate magnesium out of an aqueous solution, and it is generally not feasible to recover the metal.


Physical properties

Magnesium is a light, grayish metal. Oxidized pieces are a darker shade of gray, and tend to have a white powder of magnesium oxide on the surface. It has a low melting point of 650 °C, though the metal will ignite in air before it reaches that temperature.

Chemical properties

Magnesium is an extremely powerful reducing agent, though it is relatively stable in air due to the formation of a partial passivation layer. In air, magnesium will ignite with a very hot white flame to form a mixture that consists mostly of magnesium oxide, but also contains some magnesium nitride.

2 Mg + O2 → 2 MgO
3 Mg + N2 → Mg3N2

The color of the flame is a noticeably purer white than that of titanium or zirconium flames, which appear slightly yellowish. Dangerous amounts of UV light is produced when magnesium burns, necessitating eye protection if one were to stare directly at the flame.

Magnesium reacts with carbon dioxide exothermically to form magnesium oxide and carbon:

2 Mg + CO2 → 2 MgO + C

Hence, carbon dioxide fuels rather than extinguishes magnesium fires.

Powdered Mg reacts with hydrogen at high pressure and temperature (200 atmospheres, 500 °C), in the presence of magnesium iodide catalyst, to yield magnesium hydride.

In water, magnesium pieces react only slowly to form magnesium hydroxide, due to the build-up of magnesium hydroxide on the surface of the pure magnesium metal, which slows down any further reaction. Powdered magnesium, however, reacts much more vigorously with water.

Mg + 2 H2O → Mg(OH)2 + H2

Magnesium metal does not react with alkali solutions. However, the metal will react vigorously in dilute acids to form corresponding magnesium salts.

Mg + 2 HX → MgX2 + H2

Most of these are soluble except for the hydroxide, fluoride, and carbonate.

Magnesium is the choice reducing agent when extracting lanthanide metals from their salts, as well as many transition metals that cannot be extracted via carbon thermoreduction. Magnesium metal can also be used to obtain alkali metals, like sodium, potassium and even rubidium and caesium (but not lithium) by reducing their respective hydroxide, either in a thermoreduction process or by stirring the Mg-MOH mixture in an inert solvent with a tertiary alcohol as catalyst at high temperature for several hours or days.

Availability and sources

A good and readily available source of magnesium is the sacrificial anodes used in many water heaters. They can be cheaply found at most plumbing stores. One rod generally has around 200 g of magnesium metal and costs around 8-12 $.

Fire starting kits often contain magnesium of 95% purity, which is sufficient for most simple reductions.

Some pencil sharpeners, such as those manufactured by KUM and Staedtler are made of magnesium, in case of the latter, 95% pure. A simple test to see if these are made of magnesium or not involves heating one in a blowtorch flame for about 30 seconds (outside!); if it is magnesium, it will catch fire and give off intense white light. A less destructive method involves adding a few drops of aqueous NaOH on the sharpener surface. Magnesium does not react with sodium hydroxide, but aluminium will. These can be found at University of California campuses. They are easily identifiable by their light weight.

Magnesium products of higher purity can be bought from GalliumSource. It is sold as turnings (coarse and fine), ribbon, ingots, rods, and foil.

Magnesium strips can also be bought from United Nuclear if you live in the US.

Magnesium powder can be purchased from eBay, at varying prices, depending on the particle size. In the EU, there are some restrictions regarding the sale and ownership of very fine magnesium powders, though there aren't any restrictions for coarse powders and turnings.


Elemental magnesium is difficult to prepare, due to its high reactivity. The industrial method involves the electrolysis of molten magnesium chloride or an eutectic mixture of MgCl2 and KCl (melting point 450 °C), in a Downs cell. This process requires the use of corrosion-resistant alloy crucibles, as molten magnesium chloride is very corrosive. The process takes place in an inert atmosphere, either argon or more often sulfur hexafluoride.




Magnesium and its compounds are not particularly toxic. Bulk magnesium is not prone to ignition, but magnesium powder and turnings are. Water and Carbon Dioxide extinguishers must NEVER be used to put out magnesium fires, as this accelerates the burning and can produce toxic and/or explosive gasses as a byproduct. Dry sand can be used to fight burning magnesium. Likewise, powdered magnesium oxide aka magnesia is also suitable for putting out magnesium fires.

Never consume magnesium or its compounds, when produced in the laboratory, as a supplement.


Magnesium metal will slowly corrode in air and turn dark gray. To prevent this, storage under mineral oil is sufficient. For long-term storage, ampouling is a viable solution, though rarely necessary. Keep it away from acidic vapors and corrosive gases.

In general, magnesium doesn't require special storage conditions, as long as the air from the storage space has low humidity and is not polluted. A simple zip lock bag is good enough to keep magnesium unaffected for several years.


No special disposal procedures are required for magnesium and magnesium compounds. Discard them as you wish.


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