(anhydrous) , |
|Jmol-3D images||Image 1|
|Molar mass||84.31 g mol−1|
|Density||2.958 g/cm3 (anhydrous)
2.825 g/cm3 (dihydrate)
1.837 g/cm3 (trihydrate)
1.73 g/cm3 (pentahydrate)
|Melting point||540 °C (1,004 °F; 813 K)
165 °C (329 °F; 438 K)
|Solubility in water||anhydrous:
0.0106 g/100ml (25 °C)
0.0063 g/100ml (100 °C)
0.375 g/100ml (20 °C)
Solubility product (Ksp)
|Solubility||soluble in acid, aqueous CO2
insoluble in acetone, ammonia
Refractive index (nD)
Std enthalpy of
Gibbs free energy (ΔG)
|EU Index||Not listed|
|Other anions||Magnesium bicarbonate|
|Other cations||Beryllium carbonate
|Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)|
|(what is: / ?)|
The most common magnesium carbonate forms are the anhydrous salt called magnesite (MgCO3) and the di, tri, and pentahydrates known as barringtonite (MgCO3·2 H2O), nesquehonite (MgCO3·3 H2O), and lansfordite (MgCO3·5 H2O), respectively. Some basic forms such as artinite (MgCO3·Mg(OH)2·3 H2O), hydromagnesite (4 MgCO3·Mg(OH)2·4 H2O), and dypingite (4 MgCO3· Mg(OH)2·5 H2O) also occur as minerals.
Magnesite consists of white trigonal crystals. The anhydrous salt is practically insoluble in water, acetone, and ammonia. All forms of magnesium carbonate react in acids. Magnesium carbonate crystallizes in the calcite structure where in Mg2+ is surrounded by six oxygen atoms. The dihydrate one has a triclinic structure, while the trihydrate has a monoclinic structure.
- Magnesium carbonate is ordinarily obtained by mining the mineral magnesite.
- Magnesium carbonate can be prepared in laboratory by reaction between any soluble magnesium salt and sodium bicarbonate:
- MgCl2(aq) + 2NaHCO3(aq) → MgCO3(s) + 2NaCl(aq) + H2O(l) + CO2(g)
- Note that when the solution of magnesium chloride (or sulfate) is treated with aqueous sodium carbonate, a precipitate of basic magnesium carbonate is formed:
- 5MgCl2(aq) + 5Na2CO3(aq) + 5H2O(l) → Mg(OH)2·3MgCO3·3H2O(s) + Mg(HCO3)2(aq) + 10NaCl(aq)
- High purity industrial routes include a path through magnesium bicarbonate: combining magnesium hydroxide and carbon dioxide. A slurry of magnesium hydroxide is treated with 3.5 to 5 atm of carbon dioxide below 50 °C, giving the soluble bicarbonate, then vacuum drying the filtrate, which returns half of the carbon dioxide as well as water.
- Mg(OH)2 + 2 CO2 → Mg(HCO3)2
- Mg(HCO3)2 → MgCO3 + CO2 + H2O
- MgCO3 + 2 HCl → MgCl2 + CO2 + H2O
- MgCO3 + H2SO4 → MgSO4 + CO2 + H2O
- MgCO3 → MgO + CO2 (ΔH = +118 kJ/mol)
There is a lot of conflicting information about the decomposition temperature.
- Iowa State University has an MSDS which shows 350 °C
- Carnegie Mellon University offers a Gibbs free energy model which gives 575K (302 °C)
- ScienceLab.com has an MSDS which shows 662 °C
It is also interesting to note that the hydrates of the salts lose water at different temperatures during decomposition. For example in the trihydrate, which molecular formula may be written as Mg(HCO3)(OH)•2(H2O), the dehydration steps occur at 157 °C and 179 °C as follows:
- Mg(HCO3)(OH)•2(H2O) → Mg(HCO3)(OH)•(H2O) + H2O at 157 °C
- Mg(HCO3)(OH)•(H2O) → Mg(HCO3)(OH) + H2O at 179 °C
The primary use of magnesium carbonate is the production of magnesium oxide by calcining (see #Reactions). Magnesite and dolomite minerals are used to produce refractory bricks. MgCO3 is also used in flooring, fireproofing, fire extinguishing compositions, cosmetics, dusting powder, and toothpaste. Other applications are as filler material, smoke suppressant in plastics, a reinforcing agent in neoprene rubber, a drying agent, a laxative to loosen the bowels, and color retention in foods. In addition, high purity magnesium carbonate is used as antacid and as an additive in table salt to keep it free flowing.
Because of its water-insoluble, hygroscopic properties MgCO3 was first added to salt in 1911 to make the salt flow more freely. The Morton Salt company adopted the slogan "When it rains it pours" in reference to the fact that its MgCO3-containing salt would not stick together in humid weather. Magnesium carbonate, most often referred to as 'chalk', is used as a drying agent for hands in rock climbing, gymnastics, and weight lifting.
Magnesium carbonate is also used in taxidermy for whitening skulls. It can be mixed with hydrogen peroxide to create a paste, which is then spread on the skull to give it a white finish.
Magnesium carbonate itself is not toxic. It is slightly hazardous in case of skin and eye contact and may cause respiratory and digestive tract irritation in case of ingestion or inhalation.
Notes and references
- Bénézeth, Pascale, et al. "Experimental determination of the solubility product of magnesite at 50 to 200 C." Chemical Geology 286.1 (2011): 21-31.
- Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. ISBN 0-618-94690-X.
- Margarete Seeger; Walter Otto; Wilhelm Flick; Friedrich Bickelhaupt; Otto S. Akkerman (2005), "Magnesium Compounds", Ullmann's Encyclopedia of Industrial Chemistry, Weinheim: Wiley-VCH, doi:10.1002/14356007.a15_595.pub2
- Botha, A.; Strydom, C.A. (2001). "Preparation of a magnesium hydroxy carbonate from magnesium hydroxide". Hydrometallurgy 62 (3): 175. doi:10.1016/S0304-386X(01)00197-9.
- "IAState MSDS".
- "CMU Gibbs model".
- "Science Lab MSDS".
- "Conventional and Controlled Rate Thermal analysis of nesquehonite Mg(HCO3)(OH)·2(H2O)".
- "Conventional and Controlled Rate Thermal analysis of nesquehonite Mg(HCO3)(OH)•2(H2O)".
- "Morton Salt FAQ". Retrieved 2007-05-14.
- "Food-Info.net : E-numbers : E504: Magnesium carbonates". 080419 food-info.net
- British Pharmacopoeia Commission Secretariat (2009). "Index, BP 2009". Retrieved 31 January 2010.
- "Japanese Pharmacopoeia, Fifteenth Edition". 2006. Retrieved 31 January 2010.
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