Magnesium carbonate
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3D model (JSmol)
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PubChem CID
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| Properties | |
| Page Module:Chem2/styles.css has no content.MgCO3 | |
| Molar mass | 84.3139 g/mol (anhydrous) |
| Appearance | Colourless crystals or 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/100Template:Nnbspml (25 °C) 0.0063 g/100Template:Nnbspml (100 °C)[1] | |
Solubility product (Ksp)
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10−7.8[2] |
| Solubility | Soluble in acid, aqueous Page Module:Chem2/styles.css has no content.CO2 Insoluble in acetone, ammonia |
| −32.4·10−6 cm3/mol | |
Refractive index (nD)
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1.717 (anhydrous) 1.458 (dihydrate) 1.412 (trihydrate) |
| Structure | |
| Trigonal | |
| R3c, No. 167[3] | |
| Thermochemistry | |
Heat capacity (C)
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75.6 J/mol·K[1] |
Std molar
entropy (S⦵298) |
65.7 J/mol·K[1][4] |
Std enthalpy of
formation (ΔfH⦵298) |
−1113 kJ/mol[4] |
Gibbs free energy (ΔfG⦵)
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−1029.3 kJ/mol[1] |
| Pharmacology | |
| A02AA01 (WHO) Template:ATC | |
| Hazards | |
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| Flash point | Non-flammable |
| NIOSH (US health exposure limits): | |
PEL (Permissible)
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| Safety data sheet (SDS) | ICSC 0969 |
| Related compounds | |
Other anions
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Magnesium bicarbonate |
Other cations
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Beryllium carbonate Calcium carbonate Strontium carbonate Barium carbonate Radium carbonate |
Related compounds
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Artinite Hydromagnesite Dypingite |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Magnesium carbonate, Page Module:Chem2/styles.css has no content.MgCO3 (archaic name magnesia alba), is an inorganic salt that is a colourless or 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 (Page Module:Chem2/styles.css has no content.MgCO3), and the di, tri, and pentahydrates known as barringtonite (Page Module:Chem2/styles.css has no content.MgCO3·2H2O), nesquehonite (Page Module:Chem2/styles.css has no content.MgCO3·3H2O), and lansfordite (Page Module:Chem2/styles.css has no content.MgCO3·5H2O), respectively.[6] Some basic forms such as artinite (Page Module:Chem2/styles.css has no content.Mg2CO3(OH)2·3H2O), hydromagnesite (Page Module:Chem2/styles.css has no content.Mg5(CO3)4(OH)2·4H2O), and dypingite (Page Module:Chem2/styles.css has no content.Mg5(CO3)4(OH)2·5H2O) also occur as minerals. All of those minerals are colourless or white.
Magnesite consists of colourless or white trigonal crystals. The anhydrous salt is practically insoluble in water, acetone, and ammonia. All forms of magnesium carbonate react with acids. Magnesite crystallizes in the calcite structure wherein Page Module:Chem2/styles.css has no content.Mg2+ is surrounded by six oxygen atoms.[3]
| Carbonate coordination | Magnesium coordination | Unit cell |
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| File:Carbonate-coordination-in-magnesite-3D-bs-17.png | File:Magnesium-coordination-in-magnesite-3D-bs-17.png | File:Magnesite-unit-cell-3D-bs-17.png |
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.[7] The "light" form is precipitated from magnesium solutions using alkali carbonate at "normal temperatures" while the "heavy" may be produced from boiling concentrated solutions followed by precipitation to dryness, washing of the precipitate, and drying at 100 C.[8]
Preparation
Magnesium carbonate is ordinarily obtained by mining the mineral magnesite. Seventy percent of the world's supply is mined and prepared in China.[9]
Magnesium carbonate can be prepared in laboratory by reaction between any soluble magnesium salt and sodium bicarbonate:
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:
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.[6] The bicarbonate is then vacuum dried, causing it to lose carbon dioxide and a molecule of water:
- Page Module:Chem2/styles.css has no content.Mg(OH)2 + 2 CO2 → Mg(HCO3)2
- Page Module:Chem2/styles.css has no content.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:
- Page Module:Chem2/styles.css has no content.MgCO3 + 2 HCl → MgCl2 + CO2 + H2O
- Page Module:Chem2/styles.css has no content.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.[6] This process is called calcining:
- Page Module:Chem2/styles.css has no content.MgCO3 → MgO + CO2 (ΔH = +118 kJ/mol)
The decomposition temperature is given as 350 °C (662 °F).[10][11] 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.[12] For example, in the trihydrate Page Module:Chem2/styles.css has no content.MgCO3·3H2O, which molecular formula may be written as Page Module:Chem2/styles.css has no content.Mg(HCO3)(OH)·2H2O, the dehydration steps occur at 157 °C and 179 °C as follows:[12]
- Page Module:Chem2/styles.css has no content.Mg(HCO3)(OH)·2(H2O) → Mg(HCO3)(OH)·(H2O) + H2O at 157 °C
- Page Module:Chem2/styles.css has no content.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.[6] Page Module:Chem2/styles.css has no content.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, and colour retention in foods.
Because of its low solubility in water and hygroscopic properties, Page Module:Chem2/styles.css has no content.MgCO3 was first added to table salt (Page Module:Chem2/styles.css has no content.NaCl) in 1911 to make it flow more freely. The Morton Salt company adopted the slogan "When it rains it pours", highlighting that its salt, which contained Page Module:Chem2/styles.css has no content.MgCO3, would not stick together in humid weather.[13]
Powdered magnesium carbonate, known as climbing chalk or gym chalk is also used as a drying agent on athletes' hands in rock climbing, gymnastics, powerlifting, weightlifting and other sports in which a firm grip is necessary.[9] A variant is liquid chalk and another is mesoporous magnesium carbonate.
Recent developments in sports chalk manufacturing have explored the differences between naturally mined and laboratory-synthesized forms of magnesium carbonate. Laboratory-produced MgCO3 is created through controlled precipitation of purified magnesium salts and carbonate compounds, producing higher-purity material with fewer mineral impurities than mined magnesite.[14][15] Manufacturers of specialty grip chalks have reported that such refined magnesium carbonate reduces skin irritation and provides a more consistent moisture-absorption profile during athletic use.[16]
As a food additive, magnesium carbonate is known as E504. Its only known side effect is that it may work as a laxative in high concentrations.[17]
Magnesium carbonate is used in taxidermy for whitening skulls. It can be mixed with hydrogen peroxide to create a paste, which is spread on the skull to give it a white finish.
Magnesium carbonate is used as a matte white coating for projection screens.[18]
Medical use
It is a laxative to loosen the bowels.
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 does not dissolve in water, only in acid, where it will effervesce (bubble).[19]
Compendial status
See also
- Calcium acetate/magnesium carbonate
- Upsalite, a reported amorphous form of magnesium carbonate
Notes and references
Page Template:Reflist/styles.css has no content.
- ^ a b c d Page Module:Citation/CS1/styles.css has no content."Magnesium carbonate".
- ^ Page Module:Citation/CS1/styles.css has no content.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.
- ^ a b Page Module:Citation/CS1/styles.css has no content.Ross, Nancy L. (1997). "The equation of state and high-pressure behavior of magnesite". Am. Mineral. 82 (7–8): 682–688. Bibcode:1997AmMin..82..682R. doi:10.2138/am-1997-7-805. S2CID 43668770.
- ^ a b Page Module:Citation/CS1/styles.css has no content.Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. ISBN 978-0-618-94690-7.
- ^ Page Module:Citation/CS1/styles.css has no content.NIOSH Pocket Guide to Chemical Hazards. "#0373". National Institute for Occupational Safety and Health (NIOSH).
- ^ a b c d Script error: No such module "Template wrapper".
- ^ Page Module:Citation/CS1/styles.css has no content.Botha, A.; Strydom, C.A. (2001). "Preparation of a magnesium hydroxy carbonate from magnesium hydroxide". Hydrometallurgy. 62 (3): 175. Bibcode:2001HydMe..62..175B. doi:10.1016/S0304-386X(01)00197-9.
- ^ Page Module:Citation/CS1/styles.css has no content.J.R. Partington (1951). General and inorganic chemistry, 2nd ed.
- ^ a b Page Module:Citation/CS1/styles.css has no content.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.
- ^ Page Module:Citation/CS1/styles.css has no content."IAState MSDS". Archived from the original on 25 April 2016. Retrieved 4 May 2013.
- ^ Page Module:Citation/CS1/styles.css has no content.Weast, Robert C.; et al. (1978). CRC Handbook of Chemistry and Physics (59th ed.). West Palm Beach, FL: CRC Press. p. B-133. Template:ISBN.
- ^ a b Page Module:Citation/CS1/styles.css has no content."Conventional and Controlled Rate Thermal analysis of nesquehonite Mg(HCO3)(OH)·2(H2O)" (PDF).
- ^ Page Module:Citation/CS1/styles.css has no content."Her Debut - Morton Salt". Retrieved 27 December 2017.
- ^ Page Module:Citation/CS1/styles.css has no content.Technical Evaluation Report: Magnesium Carbonate and Magnesium Carbonate Hydroxide Handling/Processing (PDF) (Report). U.S. Department of Agriculture – Agricultural Marketing Service. December 2023. Retrieved 7 November 2025.
The report notes that magnesium carbonate can be produced synthetically under controlled laboratory conditions to achieve high-purity grades used in food and pharmaceutical products.
- ^ Page Module:Citation/CS1/styles.css has no content.Allf, Bradley (21 May 2018). "The Hidden Environmental Cost of Climbing Chalk". Climbing Magazine. Cruz Bay Publishing. Retrieved 7 November 2025.
Most climbing chalks consist of magnesium carbonate (MgCO3), often derived from mined magnesite deposits in China.
- ^ Page Module:Citation/CS1/styles.css has no content."The Cold, Dusty Truth About CaCO3 vs. MgCO3". Gym Blow Blog. Gym Blow. 13 April 2025. Retrieved 7 November 2025.
Gym Blow describes its lab-synthesized magnesium carbonate as being produced from purified magnesium salts and carbonate compounds to remove impurities and improve consistency.
- ^ Page Module:Citation/CS1/styles.css has no content."Food-Info.net : E-numbers : E504: Magnesium carbonates". 080419 food-info.net
- ^ Page Module:Citation/CS1/styles.css has no content.Noronha, Shonan (2015). Certified Technology Specialist-Installation. McGraw Hill Education. p. 256. ISBN 978-0071835657.
- ^ Page Module:Citation/CS1/styles.css has no content."What Is Magnesium Carbonate?". Sciencing. 24 April 2017. Retrieved 15 April 2018.
- ^ Page Module:Citation/CS1/styles.css has no content.British Pharmacopoeia Commission Secretariat (2009). "Index, BP 2009" (PDF). Archived from the original (PDF) on 11 April 2009. Retrieved 31 January 2010.
- ^ Page Module:Citation/CS1/styles.css has no content."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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