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| Names | |
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| Other names | |
| Identifiers | |
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3D model (JSmol) | |
| ChEBI | |
| ChEMBL | |
| ChemSpider |
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| ECHA InfoCard | 100.008.106 |
| E number | E504(i)(acidity regulators, ...) |
| RTECS number |
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| UNII |
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| Properties | |
| 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/100 ml (25 °C) 0.0063 g/100 ml (100 °C)[1] | |
Solubility product (Ksp) | 10−7.8[2] |
| Solubility | Soluble in acid, aqueousCO2 Insoluble inacetone,ammonia |
| −32.4·10−6 cm3/mol | |
Refractive index (nD) | 1.717 (anhydrous) 1.458 (dihydrate) 1.412 (trihydrate) |
| Structure | |
| Trigonal | |
| R3c, No. 167[3] | |
| Thermochemistry | |
| 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⦵) | −1029.3 kJ/mol[1] |
| Pharmacology | |
| A02AA01 (WHO) A06AD01 (WHO) | |
| Hazards | |
| NFPA 704 (fire diamond) | |
| Flash point | Non-flammable |
| NIOSH (US health exposure limits): | |
PEL (Permissible) |
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| Safety data sheet (SDS) | ICSC 0969 |
| Related compounds | |
Otheranions | Magnesium bicarbonate |
Othercations | 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 theirstandard state (at 25 °C [77 °F], 100 kPa). | |
 | You can helpexpand this article with text translated fromthe corresponding article in German. (December 2018)Click [show] for important translation instructions.
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Magnesium carbonate,MgCO3 (archaic namemagnesia alba), is an inorganic salt that is a colourless or white solid. Severalhydrated andbasic forms of magnesium carbonate also exist asminerals.
The most common magnesium carbonate forms are theanhydrous salt calledmagnesite (MgCO3), and the di, tri, and pentahydrates known as barringtonite (MgCO3·2H2O), nesquehonite (MgCO3·3H2O), and lansfordite (MgCO3·5H2O), respectively.[6] Some basic forms such asartinite (Mg2CO3(OH)2·3H2O),hydromagnesite (Mg5(CO3)4(OH)2·4H2O), anddypingite (Mg5(CO3)4(OH)2·5H2O) also occur asminerals. All of those minerals are colourless or white.
Magnesite consists of colourless or whitetrigonalcrystals. The anhydrous salt is practicallyinsoluble inwater,acetone, andammonia. All forms of magnesium carbonate react withacids. Magnesite crystallizes in thecalcite structure whereinMg2+ issurrounded by sixoxygen atoms.[3]
| Carbonate coordination | Magnesium coordination | Unit cell |
|---|---|---|
 |  |  |
The dihydrate has atriclinic structure, while the trihydrate has amonoclinic structure.
References to "light" and "heavy" magnesium carbonates actually refer to the magnesium hydroxy carbonateshydromagnesite anddypingite, 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]
Magnesium carbonate is ordinarily obtained by mining the mineralmagnesite. 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 andsodium bicarbonate:
If magnesium chloride (or sulfate) is treated with aqueous sodium carbonate, a precipitate of basic magnesium carbonate – ahydrated complex of magnesium carbonate andmagnesium hydroxide – rather than magnesium carbonate itself is formed:
High purity industrial routes include a path throughmagnesium bicarbonate, which can be formed by combining aslurry of magnesium hydroxide andcarbon 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:
Like many common group 2 metal carbonates, magnesium carbonate reacts with aqueous acids to releasecarbon dioxide andwater:
At high temperatures MgCO3 decomposes tomagnesium oxide andcarbon dioxide. This process is important in the production of magnesium oxide.[6] This process is calledcalcining:
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 trihydrateMgCO3·3H2O, which molecular formula may be written asMg(HCO3)(OH)·2H2O, the dehydration steps occur at 157 °C and 179 °C as follows:[12]
The primary use of magnesium carbonate is the production ofmagnesium oxide bycalcining. Magnesite anddolomite minerals are used to producerefractory bricks.[6]MgCO3 is also used inflooring,fireproofing,fire extinguishing compositions,cosmetics,dusting powder, andtoothpaste. Other applications are asfiller material, smoke suppressant in plastics, a reinforcing agent inneoprene rubber, adrying agent, and colour retention in foods.
Because of its low solubility in water andhygroscopic properties,MgCO3 was first added totable salt (NaCl) in 1911 to make it flow more freely. TheMorton Salt company adopted the slogan "When it rains it pours", highlighting that its salt, which containedMgCO3, would not stick together in humid weather.[13]
Powdered magnesium carbonate, known asclimbing chalk orgym chalk is also used as a drying agent on athletes' hands inrock climbing,gymnastics, powerlifting,weightlifting and other sports in which a firm grip is necessary.[9] A variant isliquid chalk and another ismesoporous 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 afood 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 intaxidermy for whitening skulls. It can be mixed withhydrogen 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 forprojection screens.[18]
It is alaxative to loosen thebowels.
In addition, high purity magnesium carbonate is used as anantacid and as an additive intable salt to keep it free flowing. Magnesium carbonate can do this because it does not dissolve in water, only in acid, where it willeffervesce (bubble).[19]
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.
{{isbn}}: ignored ISBN errors (link).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.
Most climbing chalks consist of magnesium carbonate (MgCO₃), often derived from mined magnesite deposits in China.
Gym Blow describes its lab-synthesized magnesium carbonate as being produced from purified magnesium salts and carbonate compounds to remove impurities and improve consistency.
