Calcium carbonate
44731
226170257
2008-07-17T03:42:45Z
Wilson44691
4349401
/* Preparation */
{{Chembox new
| Name = Calcium carbonate
| ImageFile = Calcium carbonate.jpg
| OtherNames = [[Limestone]]; [[calcite]]; [[aragonite]]; [[chalk]]; [[marble]]
| Section1 = {{Chembox Identifiers
| CASNo = 471-34-1
| RTECS =
}}
| Section2 = {{Chembox Properties
| Formula = CaCO<sub>3</sub>
| MolarMass = 100.087 g/mol
| Appearance = White powder.
| Density = 2.71 g/cm³ ([[calcite]]); 2.83 g/cm³ ([[aragonite]])
| Solubility = Insoluble
| MeltingPt = 825 °C Decomposes
| BoilingPt = Decomposes
| pKa =
| pKb =
}}
| Section3 = {{Chembox Structure
| MolShape = Linear
| Dipole =
}}
| Section7 = {{Chembox Hazards
| MainHazards = Not hazardous.
| NFPA-H = 0 | NFPA-F = 0 | NFPA-R = 0
| FlashPt = Non-flammable.
| RPhrases = {{R36}}, {{R37}}, {{R38}}
| SPhrases = {{S26}}, {{S36}}
}}
}}
'''Calcium carbonate''' is a [[chemical compound]] with the [[chemical formula]] [[Calcium|Ca]][[Carbon|C]][[Oxygen|O]]<sub>3</sub>. It is a common substance found as [[Rock (geology)|rock]] in all parts of the world, and is the main component of [[seashells|shells of marine organisms]], [[snail]]s, and [[eggshell]]s. Calcium carbonate is the active ingredient in [[agricultural lime]], and is usually the principal cause of [[hard water]]. It is commonly used medicinally as a [[calcium]] supplement or as an [[antacid]], but high consumption can be hazardous.
==Occurrence==
Calcium carbonate is found naturally as the following [[mineral]]s and rocks:
* [[Aragonite]]
* [[Calcite]]
* [[Vaterite]] or (μ-CaCO<sub>3</sub>)
* [[Chalk]] (Blackboard chalk is [[calcium sulfate]], CaSO<sub>4</sub>)
* [[Limestone]]
* [[Marble]]
* [[Travertine]]
To test whether a mineral or rock contains carbonate, strong acids, such as [[hydrochloric acid]] or [[sulfuric acid]], can be added to it. If the sample does contain carbonate, it will fizz and produce carbon dioxide and water. Weak acids such as acetic acid will react, albeit less vigorously. All of the rocks/minerals mentioned above will react with acid.
To test for calcium, prepared a platinum or nichrome wire and dip it into some hydrochloric acid. Then, dip the wire into some crushed sample to be tested. Place the wire in a Bunsen Flame, if calcium is presented in the sample, brick-red flame will be produced.
If a sample gives positive results for both of the two tests above, it is calcium carbonate.
==Chemical properties==
:''See also: [[Carbonate]]''
Calcium carbonate shares the typical properties of other carbonates. Notably:
#it reacts with strong acids, releasing carbon dioxide:<br/>CaCO<sub>3(s)</sub> + 2HCl<sub>(aq)</sub> → CaCl<sub>2(aq)</sub> + CO<sub>2(g)</sub> + H<sub>2</sub>O<sub>(l)</sub>
#it releases carbon dioxide on heating (to above 840 °C in the case of CaCO<sub>3</sub>), to form [[calcium oxide]], commonly called [[quicklime]], with reaction [[enthalpy]] 178 kJ / mole:<br />CaCO<sub>3</sub> → CaO + CO<sub>2</sub>
Calcium carbonate will react with water that is saturated with carbon dioxide to form the soluble [[calcium bicarbonate]].
:CaCO<sub>3</sub> + CO<sub>2</sub> + H<sub>2</sub>O → Ca(HCO<sub>3</sub>)<sub>2</sub>
This reaction is important in the [[erosion]] of [[carbonate rock]]s, forming [[cavern]]s, and leads to [[hard water]] in many regions.
==Preparation==
[[Image:MississippianMarbleUT.JPG|thumb|right|[[Mississippian]] marble in Big Cottonwood Canyon, [[Wasatch Mountains]], [[Utah]].]]
The vast majority of calcium carbonate used in industry is extracted by mining or quarrying. Pure calcium carbonate (e.g. for food or pharmaceutical use), can be produced from a pure quarried source (usually marble).
Alternatively, [[calcium oxide]] is prepared by [[calcining]] crude calcium carbonate. Water is added to give [[calcium hydroxide]], and [[carbon dioxide]] is passed through this solution to precipitate the desired calcium carbonate, referred to in the industry as precipitated calcium carbonate (PCC):<ref>{{cite web | title = Solvay Precipitated Calcium Carbonate: Production | publisher = Solvay S. A. |date=2007-03-09 | accessdate = 2007-12-30 | url = http://www.solvaypcc.com/safety_environment/0,0,1000044-_EN,00.html}}</ref>
: CaCO<sub>3</sub> → CaO + CO<sub>2</sub>
: CaO + H<sub>2</sub>O → Ca(OH)<sub>2</sub>
:Ca(OH)<sub>2</sub> + CO<sub>2</sub> → CaCO<sub>3</sub> + H<sub>2</sub>O
==Uses==
=== Industrial applications ===
The main use of calcium carbonate is in the construction industry, either as a building material in its own right (e.g. [[marble]]) or limestone aggregate for roadbuilding or as an ingredient of [[cement]] or as the starting material for the preparation of builder's lime by burning in a kiln.
Calcium carbonate is also used in the purification of [[iron]] from [[iron ore]] in a [[blast furnace]]. Calcium carbonate is calcined ''in situ'' to give calcium oxide, which forms a slag with various impurities present, and separates from the purified iron.<ref>{{cite web | title = Blast Furnace | publisher = Science Aid | accessdate = 2007-12-30 | url = http://www.scienceaid.co.uk/chemistry/industrial/blastfurnace.html}}</ref>
Calcium carbonate is also used in the [[oil industry]] in [[drilling fluid]]s as a formation bridging and filtercake sealing agent and may also be used as a weighting material to increase the density of [[drilling fluid]]s to control downhole pressures.
Calcium carbonate is also one of the main sources used in growing [[Seacrete]], or [[Biorock]].
Calcium carbonate is widely used as an extender in paints,<ref name = reade>{{cite web | title = Calcium Carbonate Powder | publisher = Reade Advanced Materials |date=2006-02-04 | accessdate = 2007-12-30 | url = http://www.reade.com/Products/Minerals_and_Ores/calcium_carbonate.html}}</ref> in particular matte emulsion paint where typically 30% by weight of the paint is either chalk or marble.
Calcium carbonate is also widely used as a filler in plastics.<ref name = reade/> Some typical examples include around 15 to 20% loading of chalk in uPVC drain pipe, 5 to 15% loading of stearate coated chalk or marble in uPVC window profile. Fine ground calcium carbonate is an essential ingredient in the microporous film used in babies' [[diapers]] and some building films as the pores are nucleated around the calcium carbonate particles during the manufacture of the film by biaxial stretching. It has also been mixed with [[acrylonitrile butadiene styrene|ABS]], and other ingredients, to form some types of compression molded "clay" Poker chips.
Calcium carbonate is also used in a wide range of trade and DIY adhesives, sealants, and decorating fillers.<ref name = reade/> Ceramic tile adhesives typically contain 70 to 80% limestone. Decorating crack fillers contain similar levels of marble or dolomite. It is also mixed with putty in setting [[stained glass]] windows, and as a resist to prevent glass from sticking to kiln shelves when firing glazes and paints at high temperature.
Calcium carbonate is known as ''whiting'' in [[ceramics (art)|ceramics]]/glazing applications,<ref name = reade/> where it is used as a common ingredient for many glazes in its white powdered form. When a glaze containing this material is fired in a kiln, the whiting acts as a [[flux]] material in the glaze.
In North America, calcium carbonate has begun to replace [[Kaolinite|kaolin]] in the production of glossy paper. Europe has been practicing this as alkaline papermaking or acid-free papermaking for some decades. Carbonates are available in forms: ground calcium carbonate (GCC) or precipitated calcium carbonate (PCC). The latter has a very fine and controlled particle size, on the order of 2 micrometres in diameter, useful in coatings for paper.
Used in swimming pools as a pH corrector for maintaining [[alkalinity]] "buffer" to offset the acidic properties of the disinfectant agent.
It is commonly called [[chalk]] as it has been a major component of blackboard chalk. Chalk may consist of either calcium carbonate or [[gypsum]], hydrated [[calcium sulfate]] CaSO<sub>4</sub>·2H<sub>2</sub>O.
=== Health and dietary applications ===
[[Image:500 mg calcium supplements with vitamin D.jpg|thumb|500 milligram calcium supplements made from calcium carbonate]]
Calcium carbonate is widely used medicinally as an inexpensive dietary calcium supplement or [[antacid]].<ref name = medline>{{cite web | publisher = [[National Institutes of Health]] | work = Medline Plus | title = Calcium Carbonate |date=2005-10-01 | accessdate = 2007-12-30 | url = http://www.nlm.nih.gov/medlineplus/druginfo/medmaster/a601032.html}}</ref> It may be used as a [[phosphate binder]] for the treatment of [[hyperphosphatemia]] (primarily in patients with [[chronic renal failure]]). It is also used in the pharmaceutical industry as an inert filler for [[tablet]]s and other pharmaceuticals.<ref>{{cite book | author = Herbert A. Lieberman, Leon Lachman, Joseph B. Schwartz | title = Pharmaceutical Dosage Forms: Tablets | year = 1990 | isbn = 0824780442 | pages = 153}}</ref>
Calcium carbonate is also used in homeopathy as one of the constitutional remedies.
Excess calcium from supplements, fortified food and high-calcium diets, can cause the "milk alkali syndrome," which has serious toxicity and can be fatal. In 1915, Bertram Sippy introduced the "Sippy regimen" of hourly ingestion of milk and cream, and the gradual addition of eggs and cooked cereal, for 10 days, combined with alkaline powders, which provided symptomatic relief for peptic ulcer disease. Over the next several decades, the Sippy regimen resulted in renal failure, alkalosis, and hypercalemia, mostly in men with peptic ulcer disease. These adverse effects were reversed when the regimen stopped, but it was fatal in some patients with protracted vomiting. Milk alkali syndrome declined in men after effective treatments for peptic ulcer disease. But during the past 15 years, it has been reported in women taking calcium supplements above the recommended range of 1200 to 1500 mg daily, for prevention and treatment of osteoporosis, and is exacerbated by dehydration. Calcium has been added to over-the-counter products, which contributes to inadvertent excessive intake. Excessive calcium intake can lead to [[hypercalcemia]], complications of which include vomiting, abdominal pain and altered mental status.<ref>[http://content.nejm.org/cgi/content/full/358/18/1952 Clinical problem-solving, back to basics,] Ilan Gabriely et al, N Engl J Med, 1 May 2008, 358(18):1952.</ref>
A form of [[food additive]] is designated as E170.<ref>{{cite web|title=Food-Info.net : E-numbers : E170 Calcium carbonate|url=http://www.food-info.net/uk/e/e170.htm}} 080419 food-info.net</ref> It is used in some [[soy milk]] products as a source of dietary calcium; one study suggests that calcium carbonate might be [[bioavailable]] as the calcium in cow's milk.<ref>{{cite journal | author = Y. Zhao, B. R. Martin and C. M. Weaver | title = Calcium Bioavailability of Calcium Carbonate Fortified Soymilk Is Equivalent to Cow's Milk in Young Women | year = 2005 | journal = [[J. Nutr.]] | volume = 135 | issue = 10 | pages = 2379–2382 | doi = | url = http://jn.nutrition.org/cgi/content/abstract/135/10/2379 | pmid = 16177199 }}</ref>
=== Ecological applications ===
In 1989, a researcher, Ken Simmons, introduced CaCO<sub>3</sub> into the [[Whetstone Brook]] in Massachusetts.<ref>{{cite news | author = [[Associated Press]] | title =
Limestone Dispenser Fights Acid Rain in Stream |date=1989-06-13 | url = http://query.nytimes.com/gst/fullpage.html?res=950DEFD9173FF930A25755C0A96F948260 | publisher = [[New York Times]]}}</ref> His hope was that the calcium carbonate would counter the acid in the stream from acid rain and save the trout that had ceased to spawn. Although his experiment was a success, it did increase the amounts of aluminium ions in the area of the brook that was not treated with the limestone. This shows that CaCO<sub>3</sub> can be added to neutralize the effects of acid rain in [[river]] ecosystems. Currently calcium carbonate is used to neutralize acidic conditions in both soil and water.<ref>{{cite journal | author = R. K. Schreiber | title = Cooperative federal-state liming research on surface waters impacted by acidic deposition | year = 1988 | journal = [[Water, Air, & Soil Pollution]] | volume = 41 | issue = 1 | pages = 53–73 | doi = 10.1007/BF00160344 | doi_brokendate = 2008-06-25}}</ref><ref>{{cite web | title = Effects of low pH and high aluminum on Atlantic salmon smolts in Eastern Maine and liming project feasibility analysis | year = 2006 | author = Dan Kircheis; Richard Dill | publisher = National Marine Fisheries Service and Maine Atlantic Salmon Commission | url = http://www.mainesalmonrivers.org/pages/Liming%20Project%20Rpt.pdf | format = reprinted at Downeast Salmon Federation}}</ref>
==Calcination equilibrium==
{| border="1" cellspacing="0" cellpadding="2" style="margin: 0 0 0 0.5em; background: white; border-collapse: collapse; border-color: #C0C090;" align="right"
! {{chembox header}} | Equilibrium Pressure of CO<sub>2</sub> over CaCO<sub>3</sub><ref name=crc>''CRC Handbook of Chemistry and Physics'' 44th ed.</ref>
|-
| 550 °C
| 0.055 k[[Pascal (unit)|Pa]]
|-
| 587 °C
| 0.13 k[[Pascal (unit)|Pa]]
|-
| 605 °C
| 0.31 k[[Pascal (unit)|Pa]]
|-
| 680 °C
| 1.80 k[[Pascal (unit)|Pa]]
|-
| 727 °C
| 5.9 k[[Pascal (unit)|Pa]]
|-
| 748 °C
| 9.3 k[[Pascal (unit)|Pa]]
|-
| 777 °C
| 14 k[[Pascal (unit)|Pa]]
|-
| 800 °C
| 24 k[[Pascal (unit)|Pa]]
|-
| 830 °C
| 34 k[[Pascal (unit)|Pa]]
|-
| 852 °C
| 51 k[[Pascal (unit)|Pa]]
|-
| 871 °C
| 72 k[[Pascal (unit)|Pa]]
|-
| 881 °C
| 80 k[[Pascal (unit)|Pa]]
|-
| 891 °C
| 91 k[[Pascal (unit)|Pa]]
|-
| 898 °C
| 101 k[[Pascal (unit)|Pa]]
|-
| 937 °C
| 179 k[[Pascal (unit)|Pa]]
|-
| 1082 °C
| 901 k[[Pascal (unit)|Pa]]
|-
| 1241 °C
| 3961 k[[Pascal (unit)|Pa]]
|-
|}
[[Calcination]] of limestone using charcoal fires to produce [[calcium oxide|quicklime]] has been practiced since antiquity by cultures all over the world. The temperature at which limestone yields calcium oxide is usually given as 825 °C, but stating an absolute threshold is misleading. Calcium carbonate exists in equilibrium with calcium oxide and carbon dioxide at any temperature. At each temperature there is a [[partial pressure]] of carbon dioxide that is in equilibrium with calcium carbonate. At room temperature the equilibrium overwhelmingly favors calcium carbonate, because the equilibrium CO<sub>2</sub> pressure is only a tiny fraction of the partial CO<sub>2</sub> pressure in air, which is about 0.035 k[[Pascal (unit)|Pa]].
At temperatures above 550 °C the equilibrium CO<sub>2</sub> pressure begins to exceed the CO<sub>2</sub> pressure in air. So above 550 °C, calcium carbonate begins to outgas CO<sub>2</sub> into air. But in a charcoal fired kiln, the concentration of CO<sub>2</sub> will be much higher than it is in air. Indeed if all the oxygen in the kiln is consumed in the fire, then the partial pressure of CO<sub>2</sub> in the kiln can be as high as 20 k[[Pascal (unit)|Pa]].
The table shows that this equilibrium pressure is not achieved until the temperature is nearly 800 °C. For the outgassing of CO<sub>2</sub> from calcium carbonate to happen at an economically useful rate, the equilibrium pressure must significantly exceed the ambient pressure of CO<sub>2</sub>. And for it to happen rapidly, the equilibrium pressure must exceed total atmospheric pressure of 101 k[[Pascal (unit)|Pa]], which happens at 898 °C.<br clear="right">
==Solubility==
===With varying CO<sub>2</sub> pressure===
{| border="1" cellspacing="0" cellpadding="2" style="margin: 0 0 0 0.5em; background: white; border-collapse: collapse; border-color: #C0C090;" align="right"
! {{chembox header}} colspan="3" | Calcium ion solubility<br> as a function of [[carbon dioxide|CO<sub>2</sub>]] [[partial pressure]] at 25 °C
|-
| width="80" align="center" bgcolor="E0E0E0"| <math>\scriptstyle P_{\text{CO}_2}</math> (atm)
| width="80" align="center" bgcolor="E0E0E0"| [[pH|p''H'']]
| width="120" align="center" bgcolor="E0E0E0"| [Ca<sup>2+</sup>] (mol/L)
|-
|10<sup>−12</sup> ||align="right"|12.0|| align="right"|5.19 × 10<sup>−3</sup>
|-
|10<sup>−10</sup> ||align="right"|11.3|| align="right"|1.12 × 10<sup>−3</sup>
|-
|10<sup>−8</sup> ||align="right"|10.7|| align="right"|2.55 × 10<sup>−4</sup>
|-
|10<sup>−6</sup> ||align="right"|9.83|| align="right"|1.20 × 10<sup>−4</sup>
|-
|10<sup>−4</sup> ||align="right"|8.62|| align="right"|3.16 × 10<sup>−4</sup>
|-
|'''3.5 × 10<sup>−4</sup>''' ||align="right"|'''8.27'''|| align="right"|'''4.70 × 10<sup>−4'''</sup>
|-
|10<sup>−3</sup> ||align="right"|7.96|| align="right"|6.62 × 10<sup>−4</sup>
|-
|10<sup>−2</sup> ||align="right"|7.30|| align="right"|1.42 × 10<sup>−3</sup>
|-
|10<sup>−1</sup> ||align="right"|6.63|| align="right"|3.05 × 10<sup>−3</sup>
|-
|1 ||align="right"|5.96|| align="right"|6.58 × 10<sup>−3</sup>
|-
|10 ||align="right"|5.30|| align="right"|1.42 × 10<sup>−2</sup>
|}
Calcium carbonate is poorly soluble in pure water (47 mg/L at normal atmospheric CO<sub>2</sub> partial pressure as shown below).
The equilibrium of its solution is given by the equation (with dissolved calcium carbonate on the right):
::{| width="450"
| width="50%" height="30"| CaCO<sub>3</sub> {{Unicode|⇋}} Ca<sup>2+</sup> + CO<sub>3</sub><sup>2–</sup>
| ''K''<sub>sp</sub> = 3.7×10<sup>–9</sup> to 8.7×10<sup>–9</sup> at 25 °C
|}
where the [[solubility product]] for [Ca<sup>2+</sup>][CO<sub>3</sub><sup>2–</sup>] is given as anywhere from ''K''<sub>sp</sub> = 3.7×10<sup>–9</sup> to ''K''<sub>sp</sub> = 8.7×10<sup>–9</sup> at 25 °C, depending upon the data source.<ref>{{cite web | title = Selected Solubility Products and Formation Constants at 25 °C | publisher = [[California State University, Dominguez Hills]] | url = http://www.csudh.edu/oliver/chemdata/data-ksp.htm}}</ref><ref name = crc/> What the equation means is that the product of molar concentration of calcium ions ([[mole (unit)|moles]] of dissolved Ca<sup>2+</sup> per liter of solution) with the molar concentration of dissolved CO<sub>3</sub><sup>2–</sup> cannot exceed the value of ''K''<sub>sp</sub>. This seemingly simple solubility equation, however, must be taken along with the more complicated equilibrium of [[carbon dioxide]] with [[water]] (see [[carbonic acid]]). Some of the CO<sub>3</sub><sup>2–</sup> combines with H<sup>+</sup> in the solution according to:
::{| width="450"
| width="50%" height="25"| HCO<sub>3</sub><sup>–</sup> {{Unicode|⇋}} H<sup>+</sup> + CO<sub>3</sub><sup>2–</sup>
| ''K''<sub>a2</sub> = 5.61×10<sup>–11</sup> at 25 °C
|}
HCO<sub>3</sub><sup>–</sup> is known as the [[bicarbonate]] ion. [[Calcium bicarbonate]] is many times more soluble in water than calcium carbonate -- indeed it exists ''only'' in solution.
Some of the HCO<sub>3</sub><sup>–</sup> combines with H<sup>+</sup> in solution according to:
::{| width="450"
| width="50%" height="25"|H<sub>2</sub>CO<sub>3</sub> {{Unicode|⇋}} H<sup>+</sup> + HCO<sub>3</sub><sup>–</sup>
| ''K''<sub>a1</sub> = 2.5×10<sup>–4</sup> at 25 °C
|}
Some of the H<sub>2</sub>CO<sub>3</sub> breaks up into water and dissolved carbon dioxide according to:
::{| width="450"
| width="50%" height="25"| H<sub>2</sub>O + CO<sub>2</sub>(dissolved) {{Unicode|⇋}} H<sub>2</sub>CO<sub>3</sub>
| ''K''<sub>h</sub> = 1.70×10<sup>–3</sup> at 25 °C
|}
And dissolved carbon dioxide is in equilibrium with atmospheric carbon dioxide according to:
::{| width="500"
| width="45%" |<math>\frac{P_{\text{CO}_2}}{[\text{CO}_2]}\ =\ k_\text{H}</math>
| where ''k''<sub>H</sub> = 29.76 atm/(mol/L) at 25 °C ([[Henry's law|Henry constant]]), <math>\scriptstyle P_{\text{CO}_2}</math> being the CO<sub>2</sub> partial pressure.
|}
For ambient air, <math>\scriptstyle P_{\text{CO}_2}</math> is around 3.5×10<sup>–4</sup> atmospheres (or equivalently 35 [[Pascal (unit)|Pa]]). The last equation above fixes the concentration of dissolved CO<sub>2</sub> as a function of <math>\scriptstyle P_{\text{CO}_2}</math>, independent of the concentration of dissolved CaCO<sub>3</sub>. At atmospheric partial pressure of CO<sub>2</sub>, dissolved CO<sub>2</sub> concentration is 1.2×10<sup>–5</sub> moles/liter. The equation before that fixes the concentration of H<sub>2</sub>CO<sub>3</sub> as a function of [CO<sub>2</sub>]. For [CO<sub>2</sub>]=1.2×10<sup>–5</sub>, it results in [H<sub>2</sub>CO<sub>3</sub>]=2.0×10<sup>–8</sup> moles per liter. When [H<sub>2</sub>CO<sub>3</sub>] is known, the remaining three equations together with
::{| width="450"
| width="50%" height="25"| H<sub>2</sub>O {{Unicode|⇋}} H<sup>+</sup> + OH<sup>–</sup>
| ''K'' = 10<sup>–14</sup> at 25 °C
|}
(which is true for all aqueous solutions), and the fact that the solution must be electrically neutral,
::2[Ca<sup>2+</sup>] + [H<sup>+</sup>] = [HCO<sub>3</sub><sup>–</sup>] + 2[CO<sub>3</sub><sup>2–</sup>] + [OH<sup>–</sup>]
make it possible to solve simultaneously for the remaining five unknown concentrations (note that the above form of the neutrality equation is valid only if calcium carbonate has been put in contact with '''pure water''' or with a neutral pH solution; in the case where the origin water solvent pH is not neutral, the equation is modified).
The table on the right shows the result for [Ca<sup>2+</sup>] and [H<sup>+</sup>] (in the form of p''H'') as a function of ambient partial pressure of CO<sub>2</sub> (''K''<sub>sp</sub> = 4.47×10<sup>−9</sup> has been taken for the calculation).
:1) At atmospheric levels of ambient CO<sub>2</sub> the table indicates the solution will be slightly alkaline with a maximum CaCO<sub>3</sub> solubility of 47 mg/L.
:2) As ambient CO<sub>2</sub> partial pressure is reduced below atmospheric levels, the solution becomes more and more alkaline. At extremely low <math>\scriptstyle P_{\text{CO}_2}</math>, dissolved CO<sub>2</sub>, bicarbonate ion, and carbonate ion largely evaporate from the solution, leaving a highly alkaline solution of [[calcium hydroxide]], which is more soluble than CaCO<sub>3</sub>. Note that for <math>\scriptstyle P_{\text{CO}_2}</math> = 10<sup>−12</sup> atm, the [Ca<sup>2+</sup>][OH<sup>-</sup>]<sup>2</sup> product is still below the solubility product of Ca(OH)<sub>2</sub> (8×10<sup>−6</sup>). For still lower CO<sub>2</sub> pressure, Ca(OH)<sub>2</sub> precipitation will occur before CaCO<sub>3</sub> precipitation.
:3) As ambient CO<sub>2</sub> partial pressure increases to levels above atmospheric, pH drops, and much of the carbonate ion is converted to bicarbonate ion, which results in higher solubility of Ca<sup>2+</sup>.
The effect of the latter is especially evident in day to day life of people who have [[hard water]]. Water in aquifers underground can be exposed to levels of CO<sub>2</sub> much higher than atmospheric. As such water percolates through calcium carbonate rock, the CaCO<sub>3</sub> dissolves according to the second trend. When that same water then emerges from the tap, in time it comes into equilibrium with CO<sub>2</sub> levels in the air by outgassing its excess CO<sub>2</sub>. The calcium carbonate becomes less soluble as a result and the excess precipitates as lime scale. This same process is responsible for the formation of [[stalactites]] and [[stalagmite]]s in limestone caves.
Two hydrated phases of calcium carbonate, [[monohydrocalcite]], CaCO<sub>3</sub>·H<sub>2</sub>O and [[ikaite]], CaCO<sub>3</sub>·6H<sub>2</sub>O]], may [[precipitate]] from water at ambient conditions and persist as metastable phases.
===With varying p''H''===
We now consider the problem of the maximum solubility of calcium carbonate in normal atmospheric conditions (<math>\scriptstyle P_{\mathrm{CO}_2}</math> = 3.5 × 10<sup>−4</sup> atm) when the p''H'' of the solution is adjusted. This is for example the case in a swimming pool where the p''H'' is maintained between 7 and 8 (by addition of NaHSO<sub>4</sub> to decrease the p''H'' or of NaHCO<sub>3</sub> to increase it). From the above equations for the solubility product, the hydratation reaction and the two acid reactions, the following expression for the maximum [Ca<sup>2+</sup>] can be easily deduced:
:<math>[\text{Ca}^{2+}]_\text{max} = \frac{K_\text{sp}k_\text{H}} {K_\text{h}K_\text{a1}K_\text{a2}} \frac{[\text{H}^+]^2}{P_{\text{CO}_2}}</math>
showing a quadratic dependence in [H<sup>+</sup>]. The numerical application with the above values of the constants gives
{| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: {{{bgc|white}}}; border-collapse: collapse; border-color: {{{bc|#C0C090}}};"
|-
|{{chembox header}} width="170" align="center" |'''p''H'''''
|width="40"|7.0
|width="40"|7.2
|width="40"|7.4
|width="40"|7.6
|width="40"|7.8
|width="40"|8.0
|width="40"|8.2
|width="40"|8.27
|width="40"|8.4
|-
|{{chembox header}} width="170" align="center" |'''[Ca<sup>2+</sup>]<sub>max</sub> (10<sup>-4</sup>mol/L or °f)'''
|width="40"|1590
|width="40"|635
|width="40"|253
|width="40"|101
|width="40"|40.0
|width="40"|15.9
|width="40"|6.35
|width="40"|4.70
|width="40"|2.53
|-
|{{chembox header}} width="170" align="center"|'''[Ca<sup>2+</sup>]<sub>max</sub> (mg/L)'''
|width="40"|6390
|width="40"|2540
|width="40"|1010
|width="40"|403
|width="40"|160
|width="40"|63.9
|width="40"|25.4
|width="40"|18.9
|width="40"|10.1
|}
Comments:
*decreasing the p''H'' from 8 to 7 increases the maximum Ca<sup>2+</sup> concentration by a factor 100. Water with a p''H'' maintained to 7 can dissolve up to 15.9 g/L of CaCO<sub>3</sub>. This explains the high Ca<sup>2+</sup> concentration in some mineral waters with p''H'' close to 7.
*note that the Ca<sup>2+</sup> concentration of the previous table is recovered for p''H'' = 8.27
*keeping the p''H'' to 7.4 in a swimming pool (which gives optimum HClO/ClO<sup>−</sup> [[hypochlorous acid|ratio]] in the case of "chlorine" maintenance) results in a maximum Ca<sup>2+</sup> concentration of 1010 mg/L. This means that successive cycles of water evaporation and partial renewing may result in a very [[hard water]] before CaCO<sub>3</sub> precipitates (water with a Ca<sup>2+</sup> concentration above 120 mg/L or 30 °f is considered very hard). Addition of a calcium sequestrant or complete renewing of the water will solve the problem.
===Solubility in a strong or weak acid solution===
Solutions of [[strong acid|strong]] ([[hydrochloric acid|HCl]]) or [[weak acid|weak]] ([[acetic acid|acetic]], [[phosphoric acid|phosphoric]]) acids are commercially available. They are commonly used to remove [[limescale]] deposits. The maximum amount of CaCO<sub>3</sub> that can be "dissolved" by one liter of an acid solution can be calculated using the above equilibrium equations.
*In the case of a strong monoacid with decreasing acid concentration [A] = [A<sup>−</sup>], we obtain (with CaCO<sub>3</sub> molar mass = 100 g):
{| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: {{{bgc|white}}}; border-collapse: collapse; border-color: {{{bc|#C0C090}}};"
|-
|width="160" {{chembox header}} |'''[A] (mol/L)'''
|width="40"|1
|width="40"|10<sup>−1</sup>
|width="40"|10<sup>−2</sup>
|width="40"|10<sup>−3</sup>
|width="40"|10<sup>−4</sup>
|width="40"|10<sup>−5</sup>
|width="40"|10<sup>−6</sup>
|width="40"|10<sup>−7</sup>
|width="40"|10<sup>−10</sup>
|-
|width="160" {{chembox header}} |'''Initial p''H'''''
|0.00||1.00||2.00||3.00||4.00||5.00||6.00||6.79||7.00
|-
|width="160" {{chembox header}} |'''Final p''H'''''
|6.75||7.25||7.75||8.14||8.25||8.26||8.26||8.26||8.27
|-
|width="160" {{chembox header}} |'''Dissolved CaCO<sub>3</sub> (g per liter of acid)'''
|50.0||5.00||0.514||0.0849||0.0504||0.0474||0.0471||0.0470||0.0470
|}
where the initial state is the acid solution with no Ca<sup>2+</sup> (not taking into account possible CO<sub>2</sub> dissolution) and the final state is the solution with saturated Ca<sup>2+</sup>. For strong acid concentrations, all species have a negligible concentration in the final state with respect to Ca<sup>2+</sup> and A<sup>−</sup> so that the neutrality equation reduces approximately to 2[Ca<sup>2+</sup>] = [A<sup>−</sup>] yielding <math>\scriptstyle[\mathrm{Ca}^{2+}] \simeq \frac{[\mathrm{A}^-]}{2}</math>. When the concentration decreases, [HCO<sub>3</sub><sup>−</sup>] becomes non negligible so that the preceding expression is no longer valid. For vanishing acid concentrations, we recover the final p''H'' and the solubility of CaCO<sub>3</sub> in pure water.
*In the case of a weak monoacid (here we take acetic acid with p''K''<sub>A</sub> = 4.76) with decreasing total acid concentration [A] = [A<sup>−</sup>]+[AH], we obtain:
{| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: {{{bgc|white}}}; border-collapse: collapse; border-color: {{{bc|#C0C090}}};"
|-
|width="160" {{chembox header}} |'''[A] (mol/L)'''
|width="40"|1
|width="40"|10<sup>−1</sup>
|width="40"|10<sup>−2</sup>
|width="40"|10<sup>−3</sup>
|width="40"|10<sup>−4</sup>
|width="40"|10<sup>−5</sup>
|width="40"|10<sup>−6</sup>
|width="40"|10<sup>−7</sup>
|width="40"|10<sup>−10</sup>
|-
|width="160" {{chembox header}} |'''Initial p''H'''''
|2.38||2.88||3.39||3.91||4.47||5.15||6.02||6.79||7.00
|-
|width="160" {{chembox header}} |'''Final p''H'''''
|6.75||7.25||7.75||8.14||8.25||8.26||8.26||8.26||8.27
|-
|width="160" {{chembox header}} |'''Dissolved CaCO<sub>3</sub> (g per liter of acid)'''
|49.5||4.99||0.513||0.0848||0.0504||0.0474||0.0471||0.0470||0.0470
|}
We see that for the same total acid concentration, the initial p''H'' of the weak acid is less acid than the one of the strong acid; however, the maximum amount of CaCO<sub>3</sub> which can be dissolved is approximately the same. This is because in the final state, the p''H'' is larger that the p''K''<sub>A</sub>, so that the weak acid is almost completely dissociated, yielding in the end as many H<sup>+</sup> ions as the strong acid to "dissolve" the calcium carbonate.
* The calculation in the case of [[phosphoric acid]] (which is the most widely used for domestic applications) is more complicated since the concentrations of the four dissociation states corresponding to this acid must be calculated together with [HCO<sub>3</sub><sup>−</sup>], [CO<sub>3</sub><sup>2−</sup>], [Ca<sup>2+</sup>], [H<sup>+</sup>] and [OH<sup>−</sup>]. The system may be reduced to a seventh degree equation for [H<sup>+</sup>] the numerical solution of which gives
{| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: {{{bgc|white}}}; border-collapse: collapse; border-color: {{{bc|#C0C090}}};"
|-
|width="160" {{chembox header}} |'''[A] (mol/L)'''
|width="40"|1
|width="40"|10<sup>−1</sup>
|width="40"|10<sup>−2</sup>
|width="40"|10<sup>−3</sup>
|width="40"|10<sup>−4</sup>
|width="40"|10<sup>−5</sup>
|width="40"|10<sup>−6</sup>
|width="40"|10<sup>−7</sup>
|width="40"|10<sup>−10</sup>
|-
|width="160" {{chembox header}} |'''Initial p''H'''''
|1.08||1.62||2.25||3.05||4.01||5.00||5.97||6.74||7.00
|-
|width="160" {{chembox header}} |'''Final p''H'''''
|6.71||7.17||7.63||8.06||8.24||8.26||8.26||8.26||8.27
|-
|width="160" {{chembox header}} |'''Dissolved CaCO<sub>3</sub> (g per liter of acid)'''
|62.0||7.39||0.874||0.123||0.0536||0.0477||0.0471||0.0471||0.0470
|}
where [A] = [H<sub>3</sub>PO<sub>4</sub>] + [H<sub>2</sub>PO<sub>4</sub><sup>−</sup>] + [HPO<sub>4</sub><sup>2−</sup>] + [PO<sub>4</sub><sup>3−</sup>] is the total acid concentration. We see that phosphoric acid is more efficient than a monoacid since at the final almost neutral p''H'', the second dissociated state concentration [HPO<sub>4</sub><sup>2−</sup>] is not negligible (see [[phosphoric acid#pH and composition of a phosphoric acid solution|phosphoric acid ]]).
==See also==
[[Image:CanarySpring.jpg|thumb|Travertine calcium carbonate deposits from a [[hot spring]]]]
* [[Ocean acidification]]
* [[Gesso]]
* [[Cuttlefish]]
* [[Cuttlebone]]
*[[Marble]] Industrial use of (ground) marble
==References==
{{reflist}}
==External links==
* {{ICSC|1193|11}}
* {{PubChemLink|516889}}
* [[ATC codes]]: {{ATC|A02|AC01}} and {{ATC|A12|AA04}}
[[Category:Calcium compounds]]
[[Category:Carbonates]]
[[Category:Limestone]]
[[Category:Phosphate binders]]
[[ar:كربونات كالسيوم]]
[[bs:Kalcijum karbonat]]
[[bg:Калциев карбонат]]
[[ca:Carbonat de calci]]
[[cs:Uhličitan vápenatý]]
[[da:Calciumcarbonat]]
[[de:Calciumcarbonat]]
[[et:Kaltsiumkarbonaat]]
[[es:Carbonato de calcio]]
[[eu:Kaltzio karbonato]]
[[fa:کلسیم کربنات]]
[[fr:Carbonate de calcium]]
[[it:Carbonato di calcio]]
[[he:סידן פחמתי]]
[[lv:Kalcija karbonāts]]
[[hu:Kalcium-karbonát]]
[[nl:Calciumcarbonaat]]
[[ja:炭酸カルシウム]]
[[no:Kalsiumkarbonat]]
[[nn:Kalsiumkarbonat]]
[[pl:Węglan wapnia]]
[[pt:Carbonato de cálcio]]
[[ro:Carbonat de calciu]]
[[ru:Карбонат кальция]]
[[simple:Calcium carbonate]]
[[sk:Uhličitan vápenatý]]
[[sl:Kalcijev karbonat]]
[[sr:Калцијум карбонат]]
[[fi:Kalsiumkarbonaatti]]
[[sv:Kalciumkarbonat]]
[[th:แคลเซียมคาร์บอเนต]]
[[vi:Cacbonat canxi]]
[[uk:Карбонат кальцію]]
[[zh:碳酸鈣]]