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Magnesium carbonate
Identifiers
CAS number 546-93-0 Yes check.svgY
13717-00-5 (monohydrate)
5145-48-2 (dihydrate)
14457-83-1 (trihydrate)
61042-72-6 (pentahydrate)
PubChem 11029
RTECS number OM2470000
Properties
Molecular formula MgCO3
Molar mass 84.3139 g/mol
Appearance white solid
hygroscopic
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.

Solubility in water 0.0012 mol/L (25 °C, anhydrous)
0.375 g/100 mL (20 °C, pentahydrate)
Solubility product, Ksp 1.0 x 10-5 [1]
Refractive index (nD) 1.717 (anhydrous)
1.458 (dihydrate)
1.412 (trihydrate)
Structure
Crystal structure Trigonal
Thermochemistry
Std enthalpy of
formation
ΔfHo298
−1111.69 kJ/mol
Standard molar
entropy
So298
65.84 J K−1 mol−1
Hazards
MSDS ICSC 0969
EU Index Not listed
Flash point Non-flammable
Related compounds
Other anions Magnesium bicarbonate
Other cations Beryllium carbonate
Calcium carbonate
Strontium carbonate
Barium carbonate
Related compounds Artinite
Hydromagnesite
Dypingite
 Yes check.svgY (what is this?)  (verify)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Magnesium carbonate, MgCO3, is a white solid that occurs in nature as a mineral. Several hydrated and basic forms of magnesium carbonate also exist as minerals. In addition, MgCO3 has a variety of uses.

Contents

Forms

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), hydromagnestite (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.

References to 'light' and 'heavy' magnesium carbonates actually refer to the magnesium hydroxy carbonates hydromagnesite and dypingite (respectively)[2].

Reactions

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 by exposing a magnesium hydroxide slurry to carbon dioxide under pressure (3.5 to 5 atm) below 50 °C, which gives 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

When dissolved with acid, magnesium carbonate decomposes with release of carbon dioxide:

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

At a temperature range between (250 °C - 800 °C), MgCO3 decomposes to magnesium oxide and carbon dioxide with reaction enthalpy 118 kJ / mole, this process is called calcining:

MgCO3 -{250-800 °C}→ MgO + CO2

Above 500 °C the process reaches a decomposition rate of 100% MgO, the publicated decomposition temperature in cases of material safety is 662 °C.

Uses

Magnesite and dolomite minerals are used to produce magnesium metal and basic 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 wet weather.[3]

Magnesium carbonate, most often referred to as 'chalk', is used as a drying agent for hands in rock climbing, gymnastics, weight lifting and powerlifting.

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 Hydroxide is used as a clay in face masks, it has mild astringent properties and helps to smooth and soften skin. Recommended for use on normal to dry skins.

Food additive

As a food additive magnesium carbonate is known as E504[citation needed], for which the only known side effect is that it may work as a laxative in high concentrations.[4]

Toxicology

Magnesium carbonate iself is not toxic. However, its excessive use may cause central nervous system depression and cardiac disturbances. [5] 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.

Compendial status

Notes & References

  1. ^ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0070494398
  2. ^ A. BOTHA and C. A. STRYDOM; “Preparation of a magnesium hydroxy carbonate from magnesium hydroxide;” Hydrometallurgy; Elsevier Science; December 2001; 62 (3): pp. 175–183.
  3. ^ "Morton Salt FAQ". http://www.mortonsalt.com/faqs/index.html#q3. Retrieved 2007-05-14. 
  4. ^ "Food-Info.net : E-numbers : E504: Magnesium carbonates". http://www.food-info.net/uk/e/e504.htm.  080419 food-info.net
  5. ^ https://fscimage.fishersci.com/msds/13340.htm Fischer Scientific
  6. ^ British Pharmacopoeia Commission Secretariat (2009). "Index, BP 2009". http://www.pharmacopoeia.co.uk/pdf/2009_index.pdf. Retrieved 31 January 2010. 
  7. ^ "Japanese Pharmacopoeia, Fifteenth Edition". 2006. http://jpdb.nihs.go.jp/jp15e/JP15.pdf. Retrieved 31 January 2010. 

See also

External links

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