Homemade calcium carbonate
| IUPAC name
| Other names
|Molar mass||100.0869 g/mol|
|Density|| 2.83 g/cm3 (aragonite)|
2.711 g/cm3 (calcite)
|Melting point|| aragonite|
825 °C (1517 °F; 1,098 K)
1,339 °C (2,442 °F; 1,612 K)
|0.0013 g/100 ml (25 °C)|
|Solubility|| Reacts with acids|
Insoluble in organic solvents
Std enthalpy of
|Safety data sheet||Sigma-Aldrich|
|Lethal dose or concentration (LD, LC):|
LD50 (Median dose)
|6,450 mg/kg (oral, rat)|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Calcium carbonate is the inorganic compound with chemical formula CaCO3. It is found in limestone and most chalk, both of which it is sometimes referred to as, and many other sources in nature. It is often used in the production of other calcium compounds due to being relatively easy to obtain.
Among bases containing calcium, calcium carbonate is the least reactive and the safest to handle, making it ideal for the production of acid salts of calcium. It can be thermally decomposed into calcium oxide, or quicklime, at temperatures upwards of 600 °C, in a process often called calcination, though this reaction does not proceed with reliable speed until temperatures of 825 °C.
Calcium carbonate is a useful base with regards to neutralizing sulfuric acid in a mixture, such as those used in Fischer esterifications. The calcium sulfate produced is nearly insoluble in most solvents, and the water molecule produced is quickly absorbed by the anhydrous calcium sulfate, which is useful for water-sensitive mixtures that may hydrolyze.
Calcium carbonate most often appears as a fine (chalky) white powder, and has a chalky taste. It is insoluble in water (0.013 g/L at 25°C), making it easy to produce in double replacement reactions with a soluble carbonate. It decomposes if heated over 825 °C
Calcium carbonate is commonly available in a very wide range of stores. Powdered limestone, which is mostly calcium carbonate, can be found in many gardening or agricultural stores as a pH raiser for soil; this can vary in purity. Pure calcium carbonate may be found as a whiting agent in ceramics stores. Many calcium supplements or antacids found in pharmacies are also made of calcium carbonate, but also frequently contain glucose or other agents to improve taste. Chalk for blackboards, rock-climbing, and building were once made of calcium carbonate, but today they may be replaced with calcium sulfate or magnesium carbonate. Eggshells are also a cheap source of calcium carbonate, though they also have an organic layer, which can be removed by roasting them in an oven/kiln.
Calcium carbonate is a moderately high secondary component of plaster of Paris (which is mostly calcium sulfate. Reacting Plaster of Paris with some baking soda (this process takes time and agitation of the mixture) can give sufficiently high concentrations of calcium carbonate, which can then be used to make soluble calcium compounds. The excess calcium sulfate and other Plaster of Paris impurities can be filtered out, being insoluble.
Because many of the OTC sources here are impure but contain water-insoluble impurities, a decent way to purify them is simply by immersion in boiling water, filtration, and washing. Antacid tablets, however, contain sugar, which will caramelize if this is done, ruining the product. At the cost of some extra expense, the calcium carbonate within the product can be dissolved using hydrochloric acid, insoluble impurites can be filtered out, and the filtrate can be treated with sodium bicarbonate to precipitate a purer form.
Calcium carbonate can be precipitated by mixing solutions of calcium chloride and any water-soluble carbonate, or with sodium bicarbonate. If pure calcium chloride is not available, limestone can be dissolved in hydrochloric acid, filtered, and the filtrate can be combined with a carbonate solution, though this may contain impurities of magnesium.
The cheapest way to make calcium carbonate is to bubble carbon dioxide through an aqueous suspension of calcium hydroxide, aka slaked (or hydrated) lime, which can be cheaply bought from any construction store. While store-grade hydrated lime may also contain magnesium in the form of magnesium hydroxide/carbonate, this can be separated by simply bubbling an excess of carbon dioxide through the solution to form the bicarbonates of the two metals. As calcium bicarbonate is much more soluble (16.6 g/100 mL at 20°C) in water than magnesium bicarbonate (0.077 g / 100 mL), the soluble calcium can be filtered. Boiling the filtrate causes the calcium bicarbonate to decompose to calcium carbonate, which can be dried and used.
If impure calcium carbonate is sufficient, leaving wet calcium hydroxide in open air will slowly result in calcium carbonate. However, if the air is polluted, calcium sulfate and nitrate may also form.
- Make calcium oxide, or quicklime, used in making cement
- Make calcium nitrate for coloring flames or producing other nitrates
- Make calcium acetate for gelling alcohols and production of acetone
- Make calcium chloride if unavailable, used as a desiccant
- Make calcium bicarbonate
- Make carbon monoxide, by reacting an equimolar amount of calcium carbonate and zinc powder
- Make chemical scrubbers
- Acid neutralization
Pure calcium carbonate has low toxicity and is safe to handle without precaution. It is often used as a calcium supplement and food additive, but can be hazardous if large amounts are consumed. When finely divided it may be an irritant to eyes, nose and lungs.
Calcium carbonate is best stored in closed bottles. No special storage is required, however it's best to store it away from any acidic fumes.
Calcium carbonate occurs naturally and poses no threat to the environment.