Difference between revisions of "Calcium carbonate"
| Line 61: | Line 61: | ||
| MeltingPtC = | | MeltingPtC = | ||
| MeltingPt_ref = | | MeltingPt_ref = | ||
| − | | MeltingPt_notes = ''aragonite''<br>825 °C (1517 °F; 1,098 K)<br>''calcite''<br>1,339 °C (2,442 °F; 1,612 K) | + | | MeltingPt_notes = ''aragonite''<br>825 °C (1517 °F; 1,098 K) (decomposition)<br>''calcite''<br>1,339 °C (2,442 °F; 1,612 K) (decomposition) |
| Odor = Odorless | | Odor = Odorless | ||
| pKa = 9.0 | | pKa = 9.0 | ||
Revision as of 20:10, 19 October 2020
 Homemade calcium carbonate
| |
| Names | |
|---|---|
| IUPAC name
Calcium carbonate
| |
| Other names
Aragonite
Calcite Chalk Limestone Marble Oyster Pearl | |
| Properties | |
| CaCO3 | |
| Molar mass | 100.0869 g/mol |
| Appearance | White solid |
| Odor | Odorless |
| Density | 2.83 g/cm3 (aragonite) 2.711 g/cm3 (calcite) |
| Melting point | aragonite 825 °C (1517 °F; 1,098 K) (decomposition) calcite 1,339 °C (2,442 °F; 1,612 K) (decomposition) |
| Boiling point | Decomposes |
| 0.0013 g/100 ml (25 °C) | |
| Solubility | Reacts with acids Insoluble in organic solvents |
| Vapor pressure | ~0 mmHg |
| Acidity (pKa) | 9.0 |
| Thermochemistry | |
| Std molar
entropy (S |
93 J·mol−1·K−1 |
| Std enthalpy of
formation (ΔfH |
−1,207 kJ/mol |
| Hazards | |
| Safety data sheet | Sigma-Aldrich |
| Flash point | Non-flammable |
| Lethal dose or concentration (LD, LC): | |
| LD50 (Median dose)
|
6,450 mg/kg (oral, rat) |
| Related compounds | |
| Related compounds
|
Magnesium carbonate Strontium sulfate Barium carbonate |
| Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). | |
| Infobox references | |
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.
Contents
Properties
Chemical
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.
- CaCO3 + 2 HX → CaX2 + H2O + CO2
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, as the hot calcium oxide will quickly reabsorb any calcium dioxide from air to reform the carbonate.
- CaCO3 → CaO + CO2
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. However, if a slight excess of calcium carbonate isn't used, minute amounts of unreacted sulfuric acid may be trapped in the resulting CaSO4, which could cause trouble later. Alternatively, one could simply stir the mixture for long periods of time, to make sure all the acid reacted, though this might take a while, depending on the amount used.
Bubbling carbon dioxide through an aq. suspension of calcium carbonate will cause it to dissolve, by converting it into calcium bicarbonate, which is only soluble as solution.
- CaCO3(s) + CO2 → Ca(HCO3)2(aq)
Evaporating the solution to dryness causes the bicarbonate to release its carbon dioxide, leaving behind calcium carbonate.
Physical
Calcium carbonate most often appears as a fine (chalky) white powder, and has a chalky taste. It is practically 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 to CaO if heated over 825 °C for long periods of time.
Availability
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 and snail shells are also a cheap source of calcium carbonate, though they also have organic layers (proteins or polysaccharides), which can be removed by roasting them in an oven/kiln, preferably in an oxidizing atmosphere to eliminate all carbon traces.
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 impurities can be filtered out, and the filtrate can be treated with sodium bicarbonate to precipitate a purer form.
Preparation
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.
Projects
- 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
Handling
Safety
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.
Storage
Calcium carbonate is best stored in any clean plastic bottles. No special storage is required, however it's best to store it away from any acidic fumes.
Disposal
Calcium carbonate, unless contaminated from chemical reactions, is practically non-toxic and doesn't require special disposal. You can simply dump it in the trash or ground or wherever you want. Since calcium carbonate occurs naturally in both mineral and biotic sources, it poses no threat to the environment.
