|Jmol-3D images||Image 1|
|Molar mass||84.3139 g/mol|
|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 decomp. (anydrous)
165 °C (trihydrate)
|Solubility in water||0.0106 g/100 mL (25 °C, anhydrous)
3.75 g/L (20 °C, pentahydrate)
|Solubility product, Ksp||1.0×10−5 |
|Solubility||soluble in acid, aqueous CO2
insoluble in acetone, ammonia
|Refractive index (nD)||1.717 (anhydrous)
|Std enthalpy of
|EU Index||Not listed|
|Other anions||Magnesium bicarbonate|
|Other cations||Beryllium carbonate
| (what is: / ?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
The most common magnesium carbonate forms are the anhydrous salt called magnesite (MgCO3) and the di, tri, and pentahydrates known as barringtonite (MgCO3·2H2O), nesquehonite (MgCO3·3H2O), and lansfordite (MgCO3·5H2O), respectively. Some basic forms such as artinite (MgCO3·Mg(OH)2·3H2O), hydromagnesite (4MgCO3·Mg(OH)2·4H2O), and dypingite (4MgCO3· Mg(OH)2·5H2O) 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.
Although magnesium carbonate is ordinarily obtained by mining the mineral magnesite, the trihydrate salt, MgCO3·3H2O, can be prepared by mixing solutions of magnesium and carbonate ions under an atmosphere of carbon dioxide. Magnesium carbonate can also be synthesized from 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 soluble magnesium bicarbonate:
- Mg(OH)2 + 2 CO2 → Mg(HCO3)2
Following the filtration of the solution, the filtrate is dried under vacuum to produce magnesium carbonate as a hydrated salt:
- Mg2+ + 2 HCO3– → MgCO3 + CO2 + H2O
Like many common carbonates, magnesium carbonate reacts with acids to release of carbon dioxide:
- 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 has an MSDS which shows 350°C
- The original text here had 400°C
- The original chembox here had 540°C
- CMU offers a Gibbs free energy model which gives 575K (302°C)
- Science Lab has an MSDS which shows 662°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
- Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
- 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".
- "prior Wikipedia "reactions" text".
- "prior Wikipedia "chembox" value".
- "CMU Gibbs model".
- "Science Lab MSDS".
- "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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