View source for Calcium bicarbonate

{{Short description|Chemical compound existing only in aqueous solution (Ca(HCO₃)₂)}}
{{Chembox
|verifiedrevid=430733457
|Name=Calcium bicarbonate
|SystematicName=Calcium bicarbonate
|OtherNames=Cleansing lime<br/>Bicarbonate of lime<br/>Rain salt
|ImageFile=Calcium bicarbonate.png
|ImageName=Calcium bicarbonate
|IUPACName=Calcium hydrogencarbonate
|Section1={{Chembox Identifiers 
| SMILES = OC(=O)O[Ca]OC(=O)O
| SMILES1 = [Ca+2].[O-]C(=O)O.[O-]C(=O)O
| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
| ChemSpiderID = 8351767
| PubChem = 10176262
| UNII_Ref = {{fdacite|correct|FDA}}
| UNII = 7PRA4BLM2L
| InChI = 1/2CH2O3.Ca/c2*2-1(3)4;/h2*(H2,2,3,4);/q;;+2/p-2
| InChIKey = NKWPZUCBCARRDP-NUQVWONBAN
| StdInChI_Ref = {{stdinchicite|correct|chemspider}}
| StdInChI = 1S/2CH2O3.Ca/c2*2-1(3)4;/h2*(H2,2,3,4);/q;;+2/p-2
| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
| StdInChIKey = NKWPZUCBCARRDP-UHFFFAOYSA-L
| CASNo_Ref = {{cascite|correct|CAS}}
| CASNo = 3983-19-5
| RTECS = 
  }}
|Section2={{Chembox Properties
| Formula = Ca(HCO<sub>3</sub>)<sub>2</sub>
| MolarMass = 162.11464 g/mol
| Appearance = 
| Density =
| Solubility = 16.1 g/100 mL (0 °C) <br> 16.6 g/100 mL (20 °C) <br> 18.5 g/100 mL (100 °C) {{reference needed|reason=This data seems erratic.|date=July 2022}}
| MeltingPt = 
| BoilingPt = 
| pKa = 
| pKb = 
| Viscosity = 
  }}
|Section3={{Chembox Structure
| MolShape = 
| Coordination = 
| CrystalStruct = 
| Dipole =
  }}
|Section7={{Chembox Hazards
| ExternalSDS = 
| MainHazards = Irritant 
| FlashPt = Non-flammable
  }}
|Section8={{Chembox Related
| OtherAnions =
| OtherCations = [[Magnesium bicarbonate]]
| OtherCompounds = [[Sodium bicarbonate]]
  }}
}}

'''Calcium bicarbonate''', also called '''calcium hydrogencarbonate''', has the chemical formula Ca(HCO<sub>3</sub>)<sub>2</sub>. The term does not refer to a known solid compound; it exists only in [[aqueous]] solution containing [[calcium]] (Ca<sup>2+</sup>), [[bicarbonate]] ({{chem|HCO|3|−}}), and [[carbonate]] ({{chem|CO|3|2−}}) ions, together with dissolved [[carbon dioxide]] (CO<sub>2</sub>). 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 water]]s 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 solid calcium bicarbonate by evaporating its solution to dryness invariably yield instead the solid [[calcium carbonate]]:<ref>{{cite web |url=http://www.newton.dep.anl.gov/askasci/chem99/chem99492.htm |url-status=dead |archive-url=https://web.archive.org/web/20000929065617/http://www.newton.dep.anl.gov/askasci/chem99/chem99492.htm |archive-date=2000-09-29 |title=Baking Soda}}</ref>
:Ca(HCO<sub>3</sub>)<sub>2</sub>([[Aqueous solution|aq]]) → CO<sub>2</sub>(g)  +  H<sub>2</sub>O(l)  +  [[Calcium carbonate|CaCO<sub>3</sub>]](s).
Very few solid bicarbonates other than those of the [[alkali metal]]s (other than [[ammonium bicarbonate]]) are known to exist.{{clarify|reason=meaning is unclear and unsourced|date=June 2020}}

The above reaction is very important to the formation of [[stalactites]], [[stalagmite]]s, columns, and other [[speleothems]] within [[caves]], and for that matter, in the formation of the caves themselves.  As water containing carbon dioxide (including extra CO<sub>2</sub> acquired from soil organisms) passes through limestone or other calcium carbonate-containing minerals, it dissolves part of the calcium carbonate,  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.

In the reverse process, dissolved carbon dioxide (CO<sub>2</sub>) in rainwater (H<sub>2</sub>O) reacts with [[limestone]] [[calcium carbonate]] (CaCO<sub>3</sub>) to form soluble calcium bicarbonate (Ca(HCO<sub>3</sub>)<sub>2</sub>). This soluble compound is then washed away with the rainwater. This form of weathering is called [[carbonation]] and [[carbonatation]].

In medicine, calcium bicarbonate is sometimes administered intravenously to immediately correct the cardiac depressor effects of [[hyperkalemia]] by increasing calcium concentration in serum, and at the same time, correcting the acid usually present.

== References ==
<references/>

{{Calcium compounds}}
{{Carbonates}}

[[Category:Bicarbonates]]
[[Category:Acid salts]]
[[Category:Calcium compounds]]