Magnesium carbonate

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Magnesium carbonate
Magnesium carbonate.png
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Names
Other names
Magnesite
Barringtonite (dihydrate)
Nesequehonite (trihydrate)
Lansfordite (pentahydrate)
Identifiers
  • 546-93-0 (anhydrous) checkY
  • 17968-26-2 (monohydrate) checkY
  • 5145-48-2 (dihydrate) ☒N
  • 14457-83-1 (trihydrate) ☒N
  • 61042-72-6 (pentahydrate) ☒N
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.008.106 Edit this at Wikidata
E number E504(i) (acidity regulators, ...)
RTECS number
  • OM2470000
UNII
  • InChI=1S/CH2O3.Mg/c2-1(3)4;/h(H2,2,3,4);/q;+2/p-2 checkY
    Key: ZLNQQNXFFQJAID-UHFFFAOYSA-L checkY
  • InChI=1/CH2O3.Mg/c2-1(3)4;/h(H2,2,3,4);/q;+2/p-2
    Key: ZLNQQNXFFQJAID-NUQVWONBAU
  • [Mg+2].[O-]C([O-])=O
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 (anhydrous)
165 °C (329 °F; 438 K)
(trihydrate)
Anhydrous:
0.0139 g/100 ml (25 °C)
0.0063 g/100 ml (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]
Std enthalpy of
formation (ΔfH⦵298)
 −1113 kJ/mol[3]
Gibbs free energy (ΔfG˚)
 −1029.3 kJ/mol[1]
Pharmacology
A02AA01 (WHO) A06AD01 (WHO)
Hazards
Safety data sheet ICSC 0969
NFPA 704 (fire diamond)
1
0
0
Flash point Non-flammable
NIOSH (US health exposure limits):
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
Radium 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  (what is checkY☒N ?)
Infobox references

Magnesium carbonate, Mg CO3 (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 with acids. Magnesium carbonate crystallizes in the calcite structure wherein Mg2+ is surrounded by six oxygen atoms. The dihydrate 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. Seventy percent of the world's supply is mined and prepared in China.[7]

Magnesium carbonate can be prepared in laboratory by reaction between any soluble magnesium salt and sodium bicarbonate:

MgCl2(aq) + 2 NaHCO3(aq) → MgCO3(s) + 2 NaCl(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:

5 MgCl2(aq) + 5 Na2CO3(aq) + 5 H2O(l) → Mg(OH)2·3 MgCO3·3 H2O(s) + Mg(HCO3)2(aq) + 10 NaCl(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

Chemical properties[]

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).[8][9] However, calcination to the oxide is generally not considered complete below 900 °C due to interfering readsorption of liberated carbon dioxide.

The hydrates of the salts lose water at different temperatures during decomposition.[10] 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:[11]

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 colour retention in foods. In addition, high purity magnesium carbonate is used as an antacid and as an additive in table salt to keep it free flowing. Magnesium carbonate can do this because it doesn't dissolve in water, only acid, where it will effervesce (bubble).[12]

Climber Jan Hojer blows surplus chalk from his hand. Boulder World Cup 2015

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.[13] Magnesium carbonate, most often referred to as "chalk", is also used as a drying agent on athletes' hands in rock climbing, gymnastics, weight lifting and other sports in which a firm grip is necessary.[14]

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.[15]

Magnesium carbonate is 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 used as a matte white coating for projection screens.[16]

Safety[]

Magnesium carbonate is non-toxic and non-flammable.

Compendial status[]

See also[]

Notes and references[]

  1. ^ a b c d "Magnesium carbonate".
  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. Bibcode:2011ChGeo.286...21B. doi:10.1016/j.chemgeo.2011.04.016.
  3. ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. ISBN 978-0-618-94690-7.
  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. "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. ^ Allf, Bradley (21 May 2018). "The Hidden Environmental Cost of Climbing Chalk". Climbing Magazine. Cruz Bay Publishing. Retrieved 22 May 2018. In fact, China produces 70 percent of the world’s magnesite. Most of that production—both mining and processing—is concentrated in a small corner of Liaoning, a hilly industrial province in northeast China between Beijing and North Korea.
  8. ^ "IAState MSDS".
  9. ^ 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.
  10. ^ "Conventional and Controlled Rate Thermal analysis of nesquehonite Mg(HCO3)(OH)·2(H2O)" (PDF).
  11. ^ "Conventional and Controlled Rate Thermal analysis of nesquehonite Mg(HCO3)(OH)•2(H2O)" (PDF).
  12. ^ "What Is Magnesium Carbonate?". Sciencing. Retrieved 15 April 2018.
  13. ^ "Her Debut - Morton Salt". Retrieved 27 December 2017.
  14. ^ Allf, Bradley (22 April 2021). "The Hidden Environmental Impact of Climbing Chalk". Climbing. Retrieved 23 April 2021.
  15. ^ "Food-Info.net : E-numbers : E504: Magnesium carbonates". 080419 food-info.net
  16. ^ Noronha, Shonan (2015). Certified Technology Specialist-Installation. McGraw Hill Education. p. 256. ISBN 978-0071835657.
  17. ^ British Pharmacopoeia Commission Secretariat (2009). "Index, BP 2009" (PDF). Archived from the original (PDF) on 11 April 2009. Retrieved 31 January 2010.
  18. ^ "Japanese Pharmacopoeia, Fifteenth Edition" (PDF). 2006. Archived from the original (PDF) on 22 July 2011. Retrieved 31 January 2010.

External links[]

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