Difference between revisions of "Magnesium carbonate"
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==Preparation== | ==Preparation== | ||
| − | Magnesium carbonate can be easily prepared by bubbling carbon dioxide in an aqueous suspension of [[magnesium hydroxide]] then boil off the excess water. Another route involves adding [[sodium bicarbonate]] or [[ammonium bicarbonate]] to a soluble magnesium salt, like [[magnesium chloride]] or [[magnesium sulfate|sulfate]], which causes magnesium carbonate to precipitate out of the solution. The precipitate is filtered and | + | Magnesium carbonate can be easily prepared by bubbling carbon dioxide in an aqueous suspension of [[magnesium hydroxide]] then boil off the excess water. If an excess of carbon dioxide is used (which is recommended, given that both the hydroxide and the carbonate are insoluble in water), [[magnesium bicarbonate]] is produced. Unlike [[calcium bicarbonate]], magnesium bicarbonate is insoluble in water, so separating the bicarbonate from the carbonate/hydroxide cannot be done by solubilizing the bicarbonate like in the case of calcium. |
| + | |||
| + | : Mg(OH)<sub>2</sub> + H<sub>2</sub>O + CO<sub>2</sub> → Mg(HCO<sub>3</sub>)<sub>2</sub> + H<sub>2</sub>O | ||
| + | |||
| + | Boiling off the water will cause the magnesium bicarbonate to decompose to magnesium carbonate. Small amounts of magnesium oxide will be produced if the heating is too strong or uneven, which is a fairly common occurence. | ||
| + | |||
| + | : Mg(HCO<sub>3</sub>)<sub>2</sub> → MgCO<sub>3</sub> + H<sub>2</sub>O + CO<sub>2</sub> | ||
| + | |||
| + | Another route involves adding [[sodium bicarbonate]] or [[ammonium bicarbonate]] to a soluble magnesium salt, like [[magnesium chloride]] or [[magnesium sulfate|sulfate]], which causes magnesium carbonate to precipitate out of the solution. The resulting precipitate is filtered and dried to remove all the water. | ||
| + | |||
| + | : MgCl<sub>2</sub> + Na<sub>2</sub>CO<sub>3</sub> → MgCO<sub>3</sub> + 2 NaCl | ||
If the magnesium salt 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. | If the magnesium salt 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. | ||
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===Storage=== | ===Storage=== | ||
| − | + | Magnesium carbonate should be kept in closed plastic or glass bottles, away from acids and moisture. | |
===Disposal=== | ===Disposal=== | ||
Latest revision as of 18:03, 17 November 2019
| Names | |
|---|---|
| IUPAC name
Magnesium carbonate
| |
| Systematic IUPAC name
Magnesium carbonate | |
| Other names
Magnesia alba
Magnesite Barringtonite (dihydrate) Nesequehonite (trihydrate) Lansfordite (pentahydrate) Upsalite | |
| Properties | |
| MgCO3 | |
| Molar mass | 84.3139 g/mol (anhydrous) |
| Appearance | White solid |
| 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) (anydrous) (decomposes) 165 °C (329 °F; 438 K) (trihydrate) |
| Boiling point | Complete decomposition around 8-900 °C |
| anhydrous 0.0139 g/100 ml (25 °C) 0.00603 g/100 ml (100 °C) | |
| Solubility | Reacts with acids, aq. CO2 Insoluble in organic solvents |
| Vapor pressure | ~0 mmHg |
| Thermochemistry | |
| Std molar
entropy (S |
65.7 J·mol-1·K-1 |
| Std enthalpy of
formation (ΔfH |
-1,113 kJ/mol |
| Hazards | |
| Safety data sheet | Sigma-Aldrich (hydrate) |
| Flash point | Non-flammable |
| Lethal dose or concentration (LD, LC): | |
| LD50 (Median dose)
|
8,000 mg/kg (rat, oral) |
| Related compounds | |
| Related compounds
|
Calcium carbonate Strontium carbonate Barium carbonate |
| Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). | |
| Infobox references | |
Magnesium carbonate is an insoluble inorganic salt of magnesium. It has the chemical formula MgCO3, but it is more commonly encountered as hydrated and basic forms, which can also be found in nature.
Contents
Properties
Chemical
Magnesium carbonate reacts with acids to form magnesium salts of said acids and releasing carbon dioxide.
- MgCO3 + H2SO4 → MgSO4 + H2O + CO2
Physical
Magnesium carbonate is a white hygroscopic solid, insoluble in solvents, but reacts with acids.
Availability
Magnesium carbonate is sold as soil amendment or agricultural lime, in the form of dolomite. It is more available in the basic carbonate form as climbing chalk and can be bought from climbing/mountain gear stores.
Higher purity magnesium carbonate can be bought from chemical suppliers, as hydrate or in the basic form.
Preparation
Magnesium carbonate can be easily prepared by bubbling carbon dioxide in an aqueous suspension of magnesium hydroxide then boil off the excess water. If an excess of carbon dioxide is used (which is recommended, given that both the hydroxide and the carbonate are insoluble in water), magnesium bicarbonate is produced. Unlike calcium bicarbonate, magnesium bicarbonate is insoluble in water, so separating the bicarbonate from the carbonate/hydroxide cannot be done by solubilizing the bicarbonate like in the case of calcium.
- Mg(OH)2 + H2O + CO2 → Mg(HCO3)2 + H2O
Boiling off the water will cause the magnesium bicarbonate to decompose to magnesium carbonate. Small amounts of magnesium oxide will be produced if the heating is too strong or uneven, which is a fairly common occurence.
- Mg(HCO3)2 → MgCO3 + H2O + CO2
Another route involves adding sodium bicarbonate or ammonium bicarbonate to a soluble magnesium salt, like magnesium chloride or sulfate, which causes magnesium carbonate to precipitate out of the solution. The resulting precipitate is filtered and dried to remove all the water.
- MgCl2 + Na2CO3 → MgCO3 + 2 NaCl
If the magnesium salt 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.
Higher purity magnesium carbonate can be produced by converting the less pure MgCO3 to magnesium bicarbonate, which is done by combining a slurry of magnesium hydroxide and carbon dioxide at high pressure and moderate temperature. The bicarbonate is filtered, then vacuum dried, causing it to lose carbon dioxide and a molecule of water, leaving behind very pure magnesium bicarbonate. This route however, may not be economical for the amateur chemist, and very pure MgCO3 is not always necessary.
Completely anhydrous magnesium carbonate (upsalite) was considered to be impossible to synthesize, until 2013 when Swedish chemists were able to prepare it from a slurry of magnesium oxide, carbon dioxide and methanol, at a pressure of 3 bar, at 50 °C for 3 h. Cooling down results in a gel which after depressurization is solidified in a furnace at 70 °C for 2 days. This material can be calcined at 300 °C, which results in pure dry MgCO3.[1]
Projects
- Make antiacid
- Preparation of magnesium salts
- Make magnesium oxide
Handling
Safety
Magnesium carbonate is practically non-toxic, though powdered magnesium carbonate is irritant.
Storage
Magnesium carbonate should be kept in closed plastic or glass bottles, away from acids and moisture.
Disposal
No special disposal is required. Discard it as you wish.
