Magnesium carbonate

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Magnesium carbonate
Magnesium carbonate.png
Uhličitan hořečnatý.PNG
Other names
Barringtonite (dihydrate)
Nesequehonite (trihydrate)
Lansfordite (pentahydrate)
546-93-0 (anhydrous) YesY
13717-00-5 (monohydrate) N
5145-48-2 (dihydrate) N
14457-83-1 (trihydrate) N
61042-72-6 (pentahydrate) N
ChEBI CHEBI:31793 YesY
ChEMBL ChEMBL1200736 N
ChemSpider 10563 YesY
Jmol interactive 3D Image
PubChem 11029
RTECS number OM2470000
Molar mass 84.3139 g/mol (anhydrous)
Appearance white solid
Odor odorless
Density 2.958 g/cm3 (anhydrous)
2.825 g/cm3 (dihydrate)
1.837 g/cm3 (trihydrate)
1.73 g/cm3 (pentahydrate)
Melting point 350 °C (662 °F; 623 K)
decomposes (anydrous)
165 °C (329 °F; 438 K)
0.0106 g/100ml (25 °C)
0.0063 g/100ml (100 °C)[1]
Solubility soluble in acid, aqueous CO2
insoluble in acetone, ammonia
1.717 (anhydrous)
1.458 (dihydrate)
1.412 (trihydrate)
75.6 J/mol·K[1]
65.7 J/mol·K[1][3]
-1113 kJ/mol[3]
-1029.3 kJ/mol[1]
ATC code A02AA01
Safety data sheet ICSC 0969
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 1: Exposure would cause irritation but only minor residual injury. E.g., turpentine Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
Flash point Non-flammable
US health exposure limits (NIOSH):
TWA 15 mg/m3 (total) TWA 5 mg/m3 (resp)[4]
Related compounds
Other anions
Magnesium bicarbonate
Other cations
Beryllium carbonate
Calcium carbonate
Strontium carbonate
Barium carbonate
Related compounds
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YesYN ?)
Infobox references

Magnesium carbonate, MgCO3 (archaic name magnesia alba), is an inorganic salt that is a white solid. Several hydrated and basic forms of magnesium carbonate also exist as minerals.


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.[5] 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.

References to 'light' and 'heavy' magnesium carbonates actually refer to the magnesium hydroxy carbonates hydromagnesite and dypingite (respectively).[6]


  • 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)
Mg(OH)2 + 2 CO2 → Mg(HCO3)2
Mg(HCO3)2 → MgCO3 + CO2 + H2O


With acids[edit]

Like many common group 2 metal carbonates, magnesium carbonate reacts with aqueous acids to release carbon dioxide and water:

MgCO3 + 2 HCl → MgCl2 + CO2 + H2O
MgCO3 + H2SO4 → MgSO4 + CO2 + H2O


At high temperatures MgCO3 decomposes to magnesium oxide and carbon dioxide. This process is important in the production of magnesium oxide.[5] This process is called calcining:

MgCO3 → MgO + CO2 (ΔH = +118 kJ/mol)

The decomposition temperature is given as 350 °C (662 °F).[7][8] However, calcination to the oxide is generally not considered complete below 900 °C due to interfering readsorption of liberated carbon dioxide.

It is also interesting to note that the hydrates of the salts lose water at different temperatures during decomposition.[9] 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:[10]

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. Magnesite and dolomite minerals are used to produce refractory bricks.[5] 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.[11] Magnesium carbonate, most often referred to as 'chalk', is used as a drying agent for hands in rock climbing, gymnastics, and weight lifting.

As a food additive magnesium carbonate is known as E504, for which the only known side effect is that it may work as a laxative in high concentrations.[12]

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 is non-toxic.

Compendial status[edit]

Notes and references[edit]

  1. ^ a b c d
  2. ^ Bénézeth, Pascale, et al. "Experimental determination of the solubility product of magnesite at 50 to 200 C." Chemical Geology 286.1 (2011): 21-31.
  3. ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. ISBN 0-618-94690-X. 
  4. ^ "NIOSH Pocket Guide to Chemical Hazards #0373". National Institute for Occupational Safety and Health (NIOSH). 
  5. ^ a b c d 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 
  6. ^ 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. 
  7. ^ "IAState MSDS". 
  8. ^ Weast, Robert C.; et al. (1978). CRC Handbook of Chemistry and Physics (59th ed.). West Palm Beach, FL: CRC Press. p. B-133. ISBN 0-8493-0549-8. 
  9. ^ "Conventional and Controlled Rate Thermal analysis of nesquehonite Mg(HCO3)(OH)·2(H2O)" (PDF). 
  10. ^ "Conventional and Controlled Rate Thermal analysis of nesquehonite Mg(HCO3)(OH)•2(H2O)" (PDF). 
  11. ^ "Morton Salt FAQ". Retrieved 2007-05-14. 
  12. ^ " : E-numbers : E504: Magnesium carbonates".  080419
  13. ^ British Pharmacopoeia Commission Secretariat (2009). "Index, BP 2009" (PDF). Retrieved 31 January 2010. 
  14. ^ "Japanese Pharmacopoeia, Fifteenth Edition" (PDF). 2006. Retrieved 31 January 2010. 

See also[edit]

External links[edit]

H2CO3 He
Li2CO3 BeCO3 B C (NH4)2CO3,
O F Ne
Al2(CO3)3 Si P S Cl Ar
Sc Ti V Cr MnCO3 FeCO3 CoCO3 NiCO3 CuCO3 ZnCO3 Ga Ge As Se Br Kr
Rb2CO3 SrCO3 Y Zr Nb Mo Tc Ru Rh Pd Ag2CO3 CdCO3 In Sn Sb Te I Xe
BaCO3   Hf Ta W Re Os Ir Pt Au Hg Tl2CO3 PbCO3 (BiO)2CO3 Po At Rn
Fr Ra   Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Fl Uup Lv Uus Uuo
La2(CO3)3 Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Ac Th Pa UO2CO3 Np Pu Am Cm Bk Cf Es Fm Md No Lr