| Discovery date | 1755 |
| Discovered by | Joseph Black |
| Origin of the name | The name is derived from Magnesia, a district of Eastern Thessaly in Greece. |
| Allotropes |
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Mg
Magnesium
12
24.305
| Group | 2 | Melting point | 650°C, 1202°F, 923 K |
| Period | 3 | Boiling point | 1090°C, 1994°F, 1363 K |
| Block | s | Density (g cm−3) | 1.74 |
| Atomic number | 12 | Relative atomic mass | 24.305 |
| State at 20°C | Solid | Key isotopes | 24Mg |
| Electron configuration | [Ne] 3s2 | CAS number | 7439-95-4 |
| ChemSpider ID | 4575328 | ChemSpider is a free chemical structure database | |
Image explanation
The image is inspired by chlorophyll, the molecule contained in green plants that enables them to photosynthesise. Chlorophyll contains a single atom of magnesium at its centre.
Appearance
A silvery-white metal that ignites easily in air and burns with a bright light.
Uses
Magnesium is one-third less dense than aluminium. It improves the mechanical, fabrication and welding characteristics of aluminium when used as an alloying agent. These alloys are useful in aeroplane and car construction.
Magnesium is used in products that benefit from being lightweight, such as car seats, luggage, laptops, cameras and power tools. It is also added to molten iron and steel to remove sulfur.
As magnesium ignites easily in air and burns with a bright light, it’s used in flares, fireworks and sparklers.
Magnesium sulfate is sometimes used as a mordant for dyes. Magnesium hydroxide is added to plastics to make them fire retardant. Magnesium oxide is used to make heat-resistant bricks for fireplaces and furnaces. It is also added to cattle feed and fertilisers. Magnesium hydroxide (milk of magnesia), sulfate (Epsom salts), chloride and citrate are all used in medicine.
Grignard reagents are organic magnesium compounds that are important for the chemical industry.
Biological role
Magnesium is an essential element in both plant and animal life. Chlorophyll is the chemical that allows plants to capture sunlight, and photosynthesis to take place. Chlorophyll is a magnesium-centred porphyrin complex. Without magnesium photosynthesis could not take place, and life as we know it would not exist.
In humans, magnesium is essential to the working of hundreds of enzymes. Humans take in about 250–350 milligrams of magnesium each day. We each store about 20 grams in our bodies, mainly in the bones.
Natural abundance
Magnesium is the eighth most abundant element in the Earth’s crust, but does not occur uncombined in nature. It is found in large deposits in minerals such as magnesite and dolomite. The sea contains trillions of tonnes of magnesium, and this is the source of much of the 850,000 tonnes now produced each year. It is prepared by reducing magnesium oxide with silicon, or by the electrolysis of molten magnesium chloride.
The first person to recognise that magnesium was an element was Joseph Black at Edinburgh in 1755. He distinguished magnesia (magnesium oxide, MgO) from lime (calcium oxide, CaO) although both were produced by heating similar kinds of carbonate rocks, magnesite and limestone respectively. Another magnesium mineral called meerschaum (magnesium silicate) was reported by Thomas Henry in 1789, who said that it was much used in Turkey to make pipes for smoking tobacco.
An impure form of metallic magnesium was first produced in 1792 by Anton Rupprecht who heated magnesia with charcoal. A pure, but tiny, amount of the metal was isolated in 1808 by Humphry Davy by the electrolysis of magnesium oxide. However, it was the French scientist, Antoine-Alexandre-Brutus Bussy who made a sizeable amount of the metal in 1831 by reacting magnesium chloride with potassium, and he then studied its properties.
| Atomic radius, non-bonded (Å) | 1.73 | Covalent radius (Å) | 1.40 |
| Electron affinity (kJ mol−1) | Not stable |
Electronegativity (Pauling scale) |
1.31 |
|
Ionisation energies (kJ mol−1) |
1st
737.75
2nd
1450.683
3rd
7732.692
4th
10542.519
5th
13630.48
6th
18019.6
7th
21711.13
8th
25661.24
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| Common oxidation states | 2 | ||||
| Isotopes | Isotope | Atomic mass | Natural abundance (%) | Half life | Mode of decay |
| 24Mg | 23.985 | 78.99 | - | - | |
| 25Mg | 24.986 | 10 | - | - | |
| 26Mg | 25.983 | 11.01 | - | - | |
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Specific heat capacity (J kg−1 K−1) |
1023 | Young's modulus (GPa) | 44.7 | |||||||||||
| Shear modulus (GPa) | 17.3 | Bulk modulus (GPa) | 44.7 | |||||||||||
| Vapour pressure | ||||||||||||||
| Temperature (K) |
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| Pressure (Pa) |
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| Listen to Magnesium Podcast |
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Transcript :
Chemistry in its element: magnesium (Promo) You're listening to Chemistry in its element brought to you by Chemistry World, the magazine of the Royal Society of Chemistry (End promo) Chris Smith Hello, this week we meet the substance whose chemical claim to fame is that its quite literally hit a bum note in the past as a cure for constipation. But its explosive role isn't just confined to the colon because it's also the basis of incendriary bombs and even the existence of life on earth. And to tell the story of Magnesium, here's John Emsley. John Emsley It was once the destroyer of cities - now it's a saver of energy The summer of 1618 saw England gripped by drought, but as Henry Wicker walked across Epsom Common he was came across a pool of water from which thirsty cattle refused to drink. He found that the water tasted bitter and on evaporation it yielded a salt which had a remarkable effect: it acted as a laxative. This became the famous Epsom's salt (magnesium sulfate, MgSO4) and became a treatment for constipation for the next 350 years. The first person to propose that magnesium was an element was Joseph Black of Edinburgh in 1755, and an impure form of metallic magnesium was produced in 1792 by Anton Rupprecht who heated magnesia (magnesium oxide, MgO) with charcoal. He named the element austrium after his native Austria. A small sample of the pure metal was isolated by Humphry Davy in 1808, by the electrolysis of moist MgO, and he proposed the name magnium based on the mineral magnesite (MgCO3) which came from Magnesia in Greece. Neither name survived and eventually it was called magnesium. Magnesium is essential to almost all life on Earth - it is at the heart of the chlorophyll molecule, which plants use to convert carbon dioxide into glucose, and then to cellulose, starch, and many other molecules which pass along the food chain. Humans take in around 300 mg of magnesium per day and we need at least 200 mg, but the body has a store of around 25 g of this element in its skeleton so there is rarely a deficiency. Almonds, brazil nuts, cashew nuts, soybeans, parsnips, bran, and even chocolate are all rich in magnesium. Some brands of beer contain a lot, such as Webster's Yorkshire Bitter - it may owe some of its flavour to the high levels of magnesium sulfate in the water used to brew it. Magnesium is the seventh most abundant element in the Earth's crust, and third most abundant if the Earth's mantle is also taken into consideration because this consists largely of olivine and pyroxene, which are magnesium silicates. It is also abundant in sea water (1200 p.p.m.) so much so that this was the source of magnesium for bombs in World War II. The metal itself was produced by the electrolysis of the molten chloride. Once magnesium starts to burn it is almost impossible to extinguish, because it reacts exothermically with oxygen, nitrogen and water. It burns with a bright light and was used for photographic flash bulbs It made an ideal incendiary agent and in some air raids during World War II as many as half a million 2 kg magnesium bombs would be scattered over a city in the space of an hour. The result was massive conflagrations and firestorms. Bulk magnesium metal is not easily ignited so this had to be done by a thermite reaction at the heart of the bomb. The thermite reaction, between aluminium powder and iron oxide, releases more than enough heat to cause the magnesium casing of the bomb to burn fiercely. Many minerals are known which contain magnesium; but the main ones are dolomite (calcium magnesium carbonate, CaMg(CO3)2) and magnesite which are mined to the extent of 10 million tonnes per year. Magnesite is heated to convert it to magnesia (MgO), and this has several applications: fertilizers; cattle feed supplement; a bulking agent in plastics; and for heat-resistant bricks for fireplaces and furnaces. The metal itself is being produced in increasing amounts. It was originally introduced for racing bicycles which were the first vehicles to use pure magnesium frames, giving a better combination of strength and lightness than other metals. (A steel frame is nearly five times heavier than a magnesium one.) For use as a metal, magnesium is alloyed with a few percent of aluminium, plus traces of zinc and manganese, to improve strength, corrosion resistance and welding qualities, and this alloy is used to save energy by making things lighter. It is found in car and aircraft seats, lightweight luggage, lawn mowers, power tools, disc drives and cameras. At the end of its useful life the magnesium in all these products can be recycled at very little cost. Because it is an electropositive metal, magnesium can be act as a 'sacrificial' electrode to protect iron and steel structures because it corrodes away preferentially when they are exposed to water which otherwise would cause rusting Chris Smith So better bikes, better bombs and better bums. Thank you very much to science writer John Emsley for telling the tale of Magnesium. Next week the illuminating story of the element that spawned a light bulb but really needs to work on its image. Quentin Cooper If any element needs a change of PR this is the one. It's brittle, prone to ponginess and arguably the dunce of the periodic table. Even the man who discovered osmium treated it rather sniffily. It reeked - or at least some of its compounds did. Tennant described the "pungent and penetrating smell" as one of the new element's "most distinguishing characters". So he called it osmium - osme being the Greek for odour. Chris Smith That's Quentin Cooper who will be undressing osmium for us in next week's Chemistry in its element, I hope you can join us. I'm Chris Smith, thank you for listening, see you next time. (Promo) Chemistry in its element is brought to you by the Royal Society of Chemistry and produced by thenakedscientists.com. There's more information and other episodes of Chemistry in its element on our website at chemistryworld.org/elements. (End promo)
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Visual Elements images and videos
© Murray Robertson 1998-2017.
W. M. Haynes, ed., CRC Handbook of Chemistry and Physics, CRC Press/Taylor and Francis, Boca Raton, FL, 95th Edition, Internet Version 2015, accessed December 2014.
Tables of Physical & Chemical Constants, Kaye & Laby Online, 16th edition, 1995. Version 1.0 (2005), accessed December 2014.
J. S. Coursey, D. J. Schwab, J. J. Tsai, and R. A. Dragoset, Atomic Weights and Isotopic Compositions (version 4.1), 2015, National Institute of Standards and Technology, Gaithersburg, MD, accessed November 2016.
T. L. Cottrell, The Strengths of Chemical Bonds, Butterworth, London, 1954.
John Emsley, Nature’s Building Blocks: An A-Z Guide to the Elements, Oxford University Press, New York, 2nd Edition, 2011.
Thomas Jefferson National Accelerator Facility - Office of Science Education, It’s Elemental - The Periodic Table of Elements, accessed December 2014.
Periodic Table of Videos, accessed December 2014.
Derived in part from material provided by the British Geological Survey © NERC.
Elements 1-112, 114, 116 and 117 © John Emsley 2012. Elements 113, 115, 117 and 118 © Royal Society of Chemistry 2017.
Produced by The Naked Scientists.
Created by video journalist Brady Haran working with chemists at The University of Nottingham.
© Murray Robertson 1998-2017.
Data
W. M. Haynes, ed., CRC Handbook of Chemistry and Physics, CRC Press/Taylor and Francis, Boca Raton, FL, 95th Edition, Internet Version 2015, accessed December 2014.
Tables of Physical & Chemical Constants, Kaye & Laby Online, 16th edition, 1995. Version 1.0 (2005), accessed December 2014.
J. S. Coursey, D. J. Schwab, J. J. Tsai, and R. A. Dragoset, Atomic Weights and Isotopic Compositions (version 4.1), 2015, National Institute of Standards and Technology, Gaithersburg, MD, accessed November 2016.
T. L. Cottrell, The Strengths of Chemical Bonds, Butterworth, London, 1954.
Uses and properties
John Emsley, Nature’s Building Blocks: An A-Z Guide to the Elements, Oxford University Press, New York, 2nd Edition, 2011.
Thomas Jefferson National Accelerator Facility - Office of Science Education, It’s Elemental - The Periodic Table of Elements, accessed December 2014.
Periodic Table of Videos, accessed December 2014.
Supply risk data
Derived in part from material provided by the British Geological Survey © NERC.
History text
Elements 1-112, 114, 116 and 117 © John Emsley 2012. Elements 113, 115, 117 and 118 © Royal Society of Chemistry 2017.
Podcasts
Produced by The Naked Scientists.
Periodic Table of Videos
Created by video journalist Brady Haran working with chemists at The University of Nottingham.
