Calcium bicarbonate
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| Calcium bicarbonate | |
|---|---|
| IUPAC name | Calcium bicarbonate |
| Other names | Calcium hydrogencarbonate |
| Properties | |
| Molecular formula | Ca(HCO3)2 |
| Molar mass | 162.11464 g/mol |
| Solubility in water | 16.6 g/100 ml (20°C) |
| Hazards | |
| Flash point | ?°C |
| Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox disclaimer and references |
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Calcium bicarbonate (Ca(HCO3)2), also called calcium hydrogen carbonate, does not refer to a known solid compound; it “exists” only in a solution containing the ions calcium (Ca2+), dissolved carbon dioxide (CO2), bicarbonate (HCO3–), and carbonate (CO32–). The relative concentrations of these carbon-containing species depend on the pH; bicarbonate predominates within the range 6.36-10.25 in fresh water.
All waters in contact with the atmosphere absorb carbon dioxide, and as these waters come into contact with rocks and sediments they acquire metal ions, most commonly calcium and magnesium, so most natural waters that come from streams, lakes, and especially wells, can be regarded as dilute solutions of these bicarbonates. These hard waters tend to form carbonate scale in pipes and boilers and they react with soaps to form an undesirable scum.
Attempts to prepare compounds such as calcium bicarbonate by evaporating its solution to dryness invariably yield the solid carbonate instead: Ca(HCO3)2(aq) → CO2(g) + H2O(l) + CaCO3(s). Very few solid bicarbonates other than those of the alkali metals are known to exist.
The above reaction is very important to the formation of stalactites, stalagmites, columns, and other speleothems within caves and, for that matter, in the formation of the caves themselves. As water containing carbon dioxide (including extra CO2 acquired from soil organisms) passes through limestone or other calcium carbonate containing minerals, it dissolves part of the calcium carbonate and hence becomes richer in bicarbonate. As the groundwater enters the cave, the excess carbon dioxide is released from the solution of the bicarbonate, causing the much less soluble calcium carbonate to be deposited.
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