Magnesium carbonate

Magnesium carbonate
Names
Other names
Magnesite
Barringtonite (dihydrate)
Nesequehonite (trihydrate)
Lansfordite (pentahydrate)
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.008.106
E number E504 (acidity regulators, ...)
RTECS number OM2470000
Properties
MgCO3
Molar mass 84.3139 g/mol (anhydrous)
Appearance white solid
hygroscopic
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)
(trihydrate)
anhydrous:
0.0139 g/100ml (25 °C)
0.00603 g/100ml (100 °C)[1]
10−7.8[2]
Solubility soluble in acid, aqueous CO2
insoluble in acetone, ammonia
−32.4·10−6 cm3/mol
1.717 (anhydrous)
1.458 (dihydrate)
1.412 (trihydrate)
Structure
Trigonal
Thermochemistry
75.6 J/mol·K[1]
65.7 J/mol·K[1][3]
-1113 kJ/mol[3]
-1029.3 kJ/mol[1]
Pharmacology
A02AA01 (WHO) A06AD01 (WHO)
Hazards
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
0
1
0
Flash point Non-flammable
US health exposure limits (NIOSH):
PEL (Permissible)
 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
 Artinite
Hydromagnesite
Dypingite
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.

Forms

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]

Preparation

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)

If magnesium chloride (or sulfate) is treated with aqueous sodium carbonate, a precipitate of basic magnesium carbonate—a hydrated complex of magnesium carbonate and magnesium hydroxide—rather than magnesium carbonate itself 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, which can be formed by combining a slurry of magnesium hydroxide and carbon dioxide at high pressure and moderate temperature.[5] The bicarbonate is then vacuum dried, causing it to lose carbon dioxide and a molecule of water:

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

Reactions

With acids

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

Decomposition

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

Uses

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 low solubility in water and hygroscopic properties, MgCO3 was first added to salt in 1911 to make it flow more freely. The Morton Salt company adopted the slogan "When it rains it pours" with 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 also used as a drying agent on athletes' palms 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.

As a matte white coating for projection screens.[13]

Safety

Magnesium carbonate is non-toxic.

Compendial status

See also

Notes and references

  1. 1 2 3 4 http://chemister.ru/Database/properties-en.php?dbid=1&id=634
  2. ↑ Bénézeth, Pascale; Saldi, Giuseppe D.; Dandurand, Jean-Louis; Schott, Jacques (2011). "Experimental determination of the solubility product of magnesite at 50 to 200 °C". Chemical Geology. 286 (1–2): 21–31. doi:10.1016/j.chemgeo.2011.04.016.
  3. 1 2 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. 1 2 3 4 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. ↑ "Food-Info.net : E-numbers : E504: Magnesium carbonates". 080419 food-info.net
  13. ↑ Noronha, Shonan (2015). Certified Technology Specialist-Installation. McGraw Hill Education. p. 256. ISBN 978-0071835657.
  14. ↑ British Pharmacopoeia Commission Secretariat (2009). "Index, BP 2009" (PDF). Archived from the original (PDF) on 11 April 2009. Retrieved 31 January 2010.
  15. ↑ "Japanese Pharmacopoeia, Fifteenth Edition" (PDF). 2006. Archived from the original (PDF) on 22 July 2011. Retrieved 31 January 2010.
Carbonates
H2CO3 He
Li2CO3,
LiHCO3		 BeCO3 B C (NH4)2CO3,
NH4HCO3		 O F Ne
Na2CO3,
NaHCO3,
Na3H(CO3)2		 MgCO3,
Mg(HCO3)2		 Al2(CO3)3 Si P S Cl Ar
K2CO3,
KHCO3		 CaCO3,
Ca(HCO3)2		 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
Cs2CO3,
CsHCO3		 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 Nh Fl Mc Lv Ts Og
↓
La2(CO3)3 Ce2(CO3)3 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
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