View source for Calcium carbonate

{{short description|Chemical compound}}
{{Use dmy dates|date=February 2015}}
{{chembox
| Verifiedfields = changed
| Watchedfields  = changed
| verifiedrevid  = 477003420
| Name           = Calcium carbonate
| ImageFile1     = Calcium carbonate structure.svg
| ImageFile3     = Calcium carbonate.jpg
| IUPACName      = Calcium carbonate
| OtherNames     = {{Ubl|[[Aragonite]]|[[Calcite]]|[[Chalk]]|[[Lime (material)|Lime]]|[[Limestone]]|[[Marble]]|[[Oyster]]shell|[[Pearl]]}}
| Section1       = {{Chembox Identifiers
|UNII_Ref = {{fdacite|correct|FDA}}
|UNII = H0G9379FGK
|ChEMBL_Ref = {{ebicite|changed|EBI}}
|ChEMBL = 1200539
|DrugBank = DB06724
|KEGG_Ref = {{keggcite|correct|kegg}}
|KEGG = D00932
|InChI = 1/CH2O3.Ca/c2-1(3)4;/h(H2,2,3,4);/q;+2/p-2
|ChEBI_Ref = {{ebicite|correct|EBI}}
|ChEBI = 3311
|SMILES = [Ca+2].[O-]C([O-])=O
|InChIKey = VTYYLEPIZMXCLO-NUQVWONBAS
|SMILES1 = C(=O)([O-])[O-].[Ca+2]
|StdInChI_Ref = {{stdinchicite|correct|chemspider}}
|StdInChI = 1S/CH2O3.Ca/c2-1(3)4;/h(H2,2,3,4);/q;+2/p-2
|StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
|StdInChIKey = VTYYLEPIZMXCLO-UHFFFAOYSA-L
|CASNo = 471-34-1
|CASNo_Ref = {{cascite|correct|CAS}}
|ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
|ChemSpiderID = 9708
|EINECS = 207-439-9
|PubChem = 10112
|RTECS = FF9335000
}}
| Section2       = {{Chembox Properties
|Formula = {{chem2|CaCO3}}
|MolarMass = 100.0869 g/mol
|Appearance = Fine white powder or colorless crystals; chalky taste
|Odor = odorless
|Density = 2.711 g/cm<sup>3</sup> ([[calcite]])<br />2.83 g/cm<sup>3</sup> ([[aragonite]])
|Solubility = 0.013 g/L (25 °C)<ref>{{cite book|title=SI Chemical Data Book |edition=4th |publisher=John Wiley & Sons Australia |last1=Aylward |first1=Gordon |last2=Findlay |first2=Tristan |isbn=978-0-470-81638-7|year=2008}}</ref><ref>{{cite book|title=Calcium Carbonate: From the Cretaceous Period Into the 21st Century|publisher=Springer Science & Business Media|year=2001|url=https://books.google.com/books?id=pbkKGa19k5QC&pg=RA1-PR2|last1=Rohleder |first1=J. |last2=Kroker |first2=E. |isbn=978-3-7643-6425-0}}</ref>
|SolubilityProduct = {{val|3.3e-9}}<ref>{{cite book|last =Benjamin|first=Mark M.|year=2002|title=Water Chemistry |publisher=McGraw-Hill|isbn =978-0-07-238390-4|url=https://books.google.com/books?id=67anQgAACAAJ}}</ref>
|Solvent = dilute acids
|SolubleOther = soluble
|MeltingPtC = 1339
|MeltingPt_notes = (calcite) <br /> {{cvt|825|°C|°F K}} (aragonite)<ref>{{cite web|url=https://www.cdc.gov/niosh/docs/81-123/pdfs/0090.pdf|title=Occupational safety and health guideline for calcium carbonate|publisher=US Dept. of Health and Human Services|access-date=31 March 2011|archive-date=30 April 2011|archive-url=https://web.archive.org/web/20110430015252/http://www.cdc.gov/niosh/docs/81-123/pdfs/0090.pdf|url-status=live}}</ref><ref>{{Cite web |url=http://diyhpl.us/~nmz787/mems/unorganized/CRC%20Handbook%20of%20Chemistry%20and%20Physics%2085th%20edition.pdf |title=CRC Handbook of Chemistry and Physics |access-date=29 October 2018 |archive-url=https://web.archive.org/web/20181029195852/http://diyhpl.us/~nmz787/mems/unorganized/CRC%20Handbook%20of%20Chemistry%20and%20Physics%2085th%20edition.pdf |archive-date=29 October 2018}}</ref>
|BoilingPt = [[Chemical decomposition|decomposes]]
|RefractIndex = 1.59
|MagSus = {{val|-3.82e-5|u=cm<sup>3</sup>/mol}}
}}
| Section3       = {{Chembox Structure
|CrystalStruct = Trigonal
|SpaceGroup = <span style="text-decoration: overline">3</span>2/m
}}
| Section4       = {{Chembox Thermochemistry
|DeltaHf = −1207&nbsp;kJ/mol<ref name=b1>{{cite book| author = Zumdahl, Steven S.|title =Chemical Principles 6th Ed| publisher = Houghton Mifflin Company| year = 2009| isbn = 978-0-618-94690-7|page=A21}}</ref>
|Entropy = 93&nbsp;J/(mol·K)<ref name=b1 />
}}
| Section5       = {{Chembox Pharmacology
|ATCCode_prefix = A02
|ATCCode_suffix = AC01
|ATC_Supplemental = {{ATC|A12|AA04}}
}}
| Section6       = {{Chembox Hazards
|ExternalSDS = [http://www.inchem.org/documents/icsc/icsc/eics1193.htm ICSC 1193]
|NFPA-H = 0
|NFPA-F = 0
|NFPA-R = 0
|LD50 = 6450 mg/kg (oral, rat)
|PEL = TWA 15 mg/m<sup>3</sup> (total) TWA 5 mg/m<sup>3</sup> (resp)<ref>{{PGCH|0090}}</ref>
}}
| Section7       = {{Chembox Related
|OtherAnions = [[Calcium bicarbonate]]
|OtherCations = {{ubl|[[Beryllium carbonate]]|[[Magnesium carbonate]]|[[Strontium carbonate]]|[[Barium carbonate]]|[[Radium carbonate]]|[[Zinc carbonate]]|[[Otavite|Cadmium carbonate]]|[[Lead(II) carbonate]]}}
|OtherCompounds = [[Calcium sulfate]]
}}
}}
[[File:Calcite.png|thumb|right|Crystal structure of calcite]]

'''Calcium carbonate''' is a [[chemical compound]] with the [[chemical formula]] {{chem2|[[Calcium|Ca]][[Carbonate|CO3]]}}. It is a common substance found in [[Rock (geology)|rocks]] as the [[mineral]]s [[calcite]] and [[aragonite]], most notably in [[chalk]] and [[limestone]], [[eggshell]]s, [[gastropod shell]]s, [[shellfish]] skeletons and [[pearl]]s. Materials containing much calcium carbonate or resembling it are described as [[calcareous]]. Calcium carbonate is the active ingredient in [[agricultural lime]] and is produced when calcium [[ion]]s in [[hard water]] react with [[carbonate ion]]s to form [[limescale]]. It has medical use as a [[calcium]] supplement or as an [[antacid]], but excessive consumption can be hazardous and cause [[hypercalcemia]] and digestive issues.<ref name=jer15>{{cite journal | title = Determination of <sup>210</sup>Po in calcium supplements and the possible related dose assessment to the consumers | last = Strumińska-Parulska | first = DI | journal = Journal of Environmental Radioactivity | volume = 150 | date = 2015 | pages = 121–125 | doi = 10.1016/j.jenvrad.2015.08.006| pmid = 26318774}}</ref>

==Chemistry==
Calcium carbonate shares the typical properties of other [[carbonate]]s. Notably, it:
*reacts with [[acid]]s, releasing [[carbonic acid]] which quickly disintegrates into [[carbon dioxide]] and [[water]]:
:{{chem2|CaCO3(s) + 2 H+(aq) → Ca(2+)(aq) + CO2(g) + H2O(l)}}
*releases carbon dioxide upon heating, called a [[thermal decomposition]] reaction, or [[calcination]] (to above 840&nbsp;°C in the case of {{chem2|CaCO3}}), to form [[calcium oxide]], CaO, commonly called [[quicklime]], with reaction [[enthalpy]] 178&nbsp;kJ/mol:
:{{chem2|CaCO3(s) → CaO(s) + CO2(g)}}

* reacts with gaseous hydrogen to form [[methane]] and water vapor plus solid [[calcium oxide]] or [[calcium hydroxide]] depending on temperature and product gas composition. Various metals including [[palladium]] and [[nickel]] are [[catalysts]] for the reaction.

Calcium carbonate reacts with water that is saturated with carbon dioxide to form the soluble [[calcium bicarbonate]].
:{{chem2|CaCO3(s) + CO2(g) + H2O(l) → Ca(HCO3)2(aq)}}

This reaction is important in the [[erosion]] of [[carbonate rock]], forming [[cavern]]s, and leads to [[hard water]] in many regions.

An unusual form of calcium carbonate is the hexahydrate [[ikaite]], {{chem2|CaCO3*6H2O}}. Ikaite is stable only below 8&nbsp;°C.

==Preparation==
The vast majority of calcium carbonate used in industry is extracted by mining or quarrying. Pure calcium carbonate (such as for food or pharmaceutical use), can be produced from a pure quarried source (usually [[marble]]).

Alternatively, calcium carbonate is prepared from [[calcium oxide]]. Water is added to give [[calcium hydroxide]] then [[carbon dioxide]] is passed through this solution to [[precipitate]] the desired calcium carbonate, referred to in the industry as precipitated calcium carbonate (PCC). This process is called [[carbonatation]]:<ref name="PCC">{{cite web|title = Precipitated Calcium Carbonate|access-date = 11 January 2014|url = http://www.lime.org/uses_of_lime/other_uses/precip_cc.asp|archive-url = https://web.archive.org/web/20140111165543/http://www.lime.org/uses_of_lime/other_uses/precip_cc.asp|archive-date = 11 January 2014}}</ref>

:{{chem2|CaO + H2O → Ca(OH)2}}
:{{chem2|Ca(OH)2 + CO2 → CaCO3 + H2O}}
In a laboratory, calcium carbonate can easily be crystallized from [[calcium chloride]] ({{chem2|CaCl2}}), by placing an [[aqueous solution]] of {{chem2|CaCl2}} in a [[desiccator]] alongside [[ammonium carbonate]] {{chem2|[NH4]2CO3}}.<ref>{{Cite journal |last1=Kim |first1=Yi-Yeoun |last2=Schenk |first2=Anna S. |last3=Ihli |first3=Johannes |last4=Kulak |first4=Alex N. |last5=Hetherington |first5=Nicola B. J. |last6=Tang |first6=Chiu C. |last7=Schmahl |first7=Wolfgang W. |last8=Griesshaber |first8=Erika |last9=Hyett |first9=Geoffrey |last10=Meldrum |first10=Fiona C. |date=September 2014 |title=A critical analysis of calcium carbonate mesocrystals |journal=Nature Communications |language=en |volume=5 |issue=1 |page=4341 |doi=10.1038/ncomms5341 |issn=2041-1723 |pmc=4104461 |pmid=25014563|bibcode=2014NatCo...5.4341K}}</ref> In the desiccator, ammonium carbonate is exposed to air and decomposes into [[ammonia]], carbon dioxide, and [[water]]. The carbon dioxide then diffuses into the aqueous solution of calcium chloride, reacts with the calcium ions and the water, and forms calcium carbonate.

==Structure==
The thermodynamically stable form of {{chem2|CaCO3}} under normal conditions is [[Hexagonal crystal family|hexagonal]] β-{{chem2|CaCO3}} (the mineral [[calcite]]). Other forms can be prepared, the denser (2.83&nbsp;g/cm<sup>3</sup>) [[orthorhombic]] λ-{{chem2|CaCO3}} (the mineral [[aragonite]]) and hexagonal μ-{{chem2|CaCO3}}, occurring as the mineral [[vaterite]]. The aragonite form can be prepared by precipitation at temperatures above 85&nbsp;°C; the vaterite form can be prepared by precipitation at 60&nbsp;°C. Calcite contains calcium atoms coordinated by six oxygen atoms; in aragonite they are coordinated by nine oxygen atoms.{{citation needed|date=September 2023}} The vaterite structure is not fully understood.<ref name="DemichelisRaiteri2013">{{cite journal|last1=Demichelis|first1=Raffaella|last2=Raiteri|first2=Paolo|last3=Gale|first3=Julian D.|last4=Dovesi|first4=Roberto|title=The Multiple Structures of Vaterite|journal=Crystal Growth & Design|volume=13|issue=6|year=2013|pages=2247–2251|issn=1528-7483|doi=10.1021/cg4002972 |bibcode=2013CrGrD..13.2247D }}</ref> [[Magnesium carbonate]] ({{chem2|MgCO3}}) has the calcite structure, whereas [[strontium carbonate]] ({{chem2|SrCO3}}) and [[barium carbonate]] ({{chem2|BaCO3}}) adopt the aragonite structure, reflecting their larger [[atomic radius|ionic radii]].{{citation needed|date=September 2023}}

== Polymorphs ==
Calcium carbonate crystallizes in three [[anhydrous]] [[Polymorphism (materials science)|polymorphs]],<ref>{{Cite journal |last1=Morse |first1=John W. |last2=Arvidson |first2=Rolf S. |last3=Lüttge |first3=Andreas |date=2007-02-01 |title=Calcium Carbonate Formation and Dissolution |url=https://pubs.acs.org/doi/10.1021/cr050358j |journal=Chemical Reviews |language=en |volume=107 |issue=2 |pages=342–381 |doi=10.1021/cr050358j |pmid=17261071 |issn=0009-2665 |access-date=15 December 2022 |archive-date=1 December 2022 |archive-url=https://web.archive.org/web/20221201024832/https://pubs.acs.org/doi/10.1021/cr050358j |url-status=live|url-access=subscription }}</ref><ref>{{Cite book |last=Lippmann|first= Friedrich |title=Sedimentary carbonate minerals |date=1973 |publisher=Springer |isbn=3-540-06011-1 |oclc=715109304}}</ref> of which [[calcite]] is the thermodynamically most stable at room temperature, [[aragonite]] is only slightly less so, and [[vaterite]] is the least stable.<ref name="Nahi-2022">{{Cite journal |last1=Nahi |first1=Ouassef |last2=Kulak |first2=Alexander N. |last3=Zhang |first3=Shuheng |last4=He |first4=Xuefeng |last5=Aslam |first5=Zabeada |last6=Ilett |first6=Martha A. |last7=Ford |first7=Ian J. |last8=Darkins |first8=Robert |last9=Meldrum |first9=Fiona C. |date=2022-11-20 |title=Polyamines Promote Aragonite Nucleation and Generate Biomimetic Structures |journal=Advanced Science |volume=10 |issue=1 |article-number=2203759 |doi=10.1002/advs.202203759 |pmid=36403251 |pmc=9811428 |s2cid=253707446 |issn=2198-3844}}</ref>

=== Crystal structure ===
The [[calcite]] crystal structure is [[trigonal]], with [[space group]] R{{overline|3}}c (No. 167 in the International Tables for Crystallography<ref>{{cite book | title=International tables for crystallography. | publisher=International Union of Crystallography | publication-place=Chester, England | date=2006 | isbn=978-0-7923-6590-7 | oclc=166325528 | doi = 10.1107/97809553602060000001| s2cid=146060934 | editor-last1=Welberry | editor-first1=T. R}}</ref>), and [[Pearson symbol]] hR10.<ref>{{Cite journal |last1=Chessin |first1=H. |last2=Hamilton |first2=W. C. |last3=Post |first3=B. |date=1965-04-01 |title=Position and thermal parameters of oxygen atoms in calcite |url=https://scripts.iucr.org/cgi-bin/paper?S0365110X65001585 |journal=Acta Crystallographica |volume=18 |issue=4 |pages=689–693 |doi=10.1107/S0365110X65001585 |bibcode=1965AcCry..18..689C |issn=0365-110X |access-date=15 December 2022 |archive-date=15 December 2022 |archive-url=https://web.archive.org/web/20221215201234/https://scripts.iucr.org/cgi-bin/paper?S0365110X65001585 |url-status=live|url-access=subscription }}</ref> [[Aragonite]] is [[Orthorhombic crystal system|orthorhombic]], with space group Pmcn (No 62), and Pearson Symbol oP20.<ref>{{Cite journal |last=Negro |first=AD |date=1971 |title=Refinement of the crystal structure of aragonite |url=http://www.minsocam.org/ammin/AM56/AM56_768.pdf |journal=American Mineralogist: Journal of Earth and Planetary Materials |volume=56 |pages=768–772 |via=GeoScienceWorld |access-date=15 December 2022 |archive-date=15 December 2022 |archive-url=https://web.archive.org/web/20221215200845/http://www.minsocam.org/ammin/AM56/AM56_768.pdf |url-status=live}}</ref> [[Vaterite]] is composed of at least two different coexisting crystallographic structures. The major structure exhibits [[Hexagonal crystal family|hexagonal]] symmetry in space group P6<sub>3</sub>/mmc, the minor structure is still unknown.<ref>{{Cite journal |last1=Kabalah-Amitai |first1=Lee |last2=Mayzel |first2=Boaz |last3=Kauffmann |first3=Yaron |last4=Fitch |first4=Andrew N. |last5=Bloch |first5=Leonid |last6=Gilbert |first6=Pupa U. P. A. |last7=Pokroy |first7=Boaz |date=2013-04-26 |title=Vaterite Crystals Contain Two Interspersed Crystal Structures |journal=Science |volume=340 |issue=6131 |pages=454–457 |doi=10.1126/science.1232139 |pmid=23620047 |bibcode=2013Sci...340..454K |s2cid=206546317 |issn=0036-8075}}</ref>

=== Crystallization ===
[[File:Calcite+Aragonite.png|thumb|234x234px|Crystal structure of calcite and aragonite]]
All three polymorphs crystallize simultaneously from aqueous solutions under ambient conditions.<ref name="Nahi-2022" /> In additive-free aqueous solutions, calcite forms easily as the major product, while aragonite appears only as a minor product.

At high saturation, vaterite is typically the first phase precipitated, which is followed by a transformation of the vaterite to calcite.<ref>{{Cite journal |last1=Bots |first1=Pieter |last2=Benning |first2=Liane G. |last3=Rodriguez-Blanco |first3=Juan-Diego |last4=Roncal-Herrero |first4=Teresa |last5=Shaw |first5=Samuel |date=2012-07-03 |title=Mechanistic Insights into the Crystallization of Amorphous Calcium Carbonate (ACC) |url=https://pubs.acs.org/doi/10.1021/cg300676b |journal=Crystal Growth & Design |language=en |volume=12 |issue=7 |pages=3806–3814 |doi=10.1021/cg300676b |bibcode=2012CrGrD..12.3806B |issn=1528-7483 |access-date=15 December 2022 |archive-date=15 December 2022 |archive-url=https://web.archive.org/web/20221215200842/https://pubs.acs.org/doi/10.1021/cg300676b |url-status=live|url-access=subscription }}</ref> This behavior seems to follow [[Ostwald's rule]], in which the least stable polymorph crystallizes first, followed by the crystallization of different polymorphs via a sequence of increasingly stable phases.<ref>{{Cite journal |last1=Cardew |first1=Peter T. |last2=Davey |first2=Roger J. |date=2019-10-02 |title=The Ostwald Ratio, Kinetic Phase Diagrams, and Polymorph Maps |url=https://pubs.acs.org/doi/10.1021/acs.cgd.9b00815 |journal=Crystal Growth & Design |language=en |volume=19 |issue=10 |pages=5798–5810 |doi=10.1021/acs.cgd.9b00815 |bibcode=2019CrGrD..19.5798C |s2cid=202885778 |issn=1528-7483 |access-date=15 December 2022 |archive-date=15 December 2022 |archive-url=https://web.archive.org/web/20221215201316/https://pubs.acs.org/doi/10.1021/acs.cgd.9b00815 |url-status=live|url-access=subscription }}</ref> However, aragonite, whose stability lies between those of vaterite and calcite, seems to be the exception to this rule, as aragonite does not form as a precursor to calcite under ambient conditions.<ref name="Nahi-2022"/>
[[File:Calcite+Vaterite.png|thumb|Microscopic calcite and vaterite]]
Aragonite occurs in majority when the reaction conditions inhibit the formation of calcite and/or promote the nucleation of aragonite. For example, the formation of aragonite is promoted by the presence of magnesium ions,<ref>{{Cite journal |last1=Zhang |first1=Shuheng |last2=Nahi |first2=Ouassef |last3=Chen |first3=Li |last4=Aslam |first4=Zabeada |last5=Kapur |first5=Nikil |last6=Kim |first6=Yi-Yeoun |last7=Meldrum |first7=Fiona C. |date=June 2022 |title=Magnesium Ions Direct the Solid-State Transformation of Amorphous Calcium Carbonate Thin Films to Aragonite, Magnesium-Calcite, or Dolomite |journal=Advanced Functional Materials |language=en |volume=32 |issue=25 |article-number=2201394 |doi=10.1002/adfm.202201394 |s2cid=247587883 |issn=1616-301X |doi-access=free }}</ref> or by using proteins and peptides derived from biological calcium carbonate.<ref>{{Cite journal |last1=Metzler |first1=Rebecca A. |last2=Evans |first2=John Spencer |last3=Killian |first3=Christopher E. |last4=Zhou |first4=Dong |last5=Churchill |first5=Tyler H. |last6=Appathurai |first6=Narayana P. |last7=Coppersmith |first7=Susan N. |last8=Gilbert |first8=P. U. P. A. |date=2010-05-12 |title=Nacre Protein Fragment Templates Lamellar Aragonite Growth |url=https://pubs.acs.org/doi/10.1021/ja909735y |journal=Journal of the American Chemical Society |language=en |volume=132 |issue=18 |pages=6329–6334 |doi=10.1021/ja909735y |pmid=20397648 |bibcode=2010JAChS.132.6329M |issn=0002-7863 |access-date=15 December 2022 |archive-date=15 December 2022 |archive-url=https://web.archive.org/web/20221215200841/https://pubs.acs.org/doi/10.1021/ja909735y |url-status=live|url-access=subscription }}</ref> Some polyamines such as [[cadaverine]] and [[Polyethylenimine|Poly(ethylene imine)]] have been shown to facilitate the formation of aragonite over calcite.<ref name="Nahi-2022"/> Solid‑state NMR analysis has revealed that poly‑aspartate-stabilized ACC contains water molecules that undergo millisecond-timescale flips, illustrating dynamic hydration as a key factor in delaying crystallization.<ref>{{cite journal
 | last1 = Vinod-Kumar
 | last2 = Smith-Penzel
 | last3 = Redrouthu
 | last4 = Gindele
 | last5 = Gebauer
 | last6 = Mathies
 | title = Investigation of the structure and dynamics of amorphous calcium carbonate by NMR: stabilization by poly-aspartate and comparison to monohydrocalcite
 | journal = Physical Chemistry Chemical Physics
 | volume = 27
 | issue = 19
 | year = 2025
 | pages = 13291–13300
 | doi = 10.1039/d5cp01002k
 | pmid = 40326785
 | pmc = 
 | bibcode = 2025PCCP...2710348V
 | doi-access = free
 }}</ref>

=== Selection by organisms ===
Organisms, such as [[mollusca|molluscs]] and [[arthropod]]s, have shown the ability to grow all three crystal polymorphs of calcium carbonate, mainly as protection (shells) and muscle attachments.<ref>{{Cite book |last1=Lowenstam |first1=H.A. |title=On Biomineralization |last2=Weiner |first2=S. |publisher=Oxford University Press |year=1989 |isbn=978-0-19-504977-0}}</ref> Moreover, they exhibit a remarkable capability of phase selection over calcite and aragonite, and some organisms can switch between the two polymorphs. The ability of phase selection is usually attributed to the use of specific macromolecules or combinations of macromolecules by such organisms.<ref>{{Cite journal |last1=Belcher |first1=A. M. |last2=Wu |first2=X. H. |last3=Christensen |first3=R. J. |last4=Hansma |first4=P. K. |last5=Stucky |first5=G. D. |last6=Morse |first6=D. E. |date=May 1996 |title=Control of crystal phase switching and orientation by soluble mollusc-shell proteins |url=https://www.nature.com/articles/381056a0 |journal=Nature |language=en |volume=381 |issue=6577 |pages=56–58 |doi=10.1038/381056a0 |bibcode=1996Natur.381...56B |s2cid=4285912 |issn=1476-4687 |access-date=15 December 2022 |archive-date=15 December 2022 |archive-url=https://web.archive.org/web/20221215212206/https://www.nature.com/articles/381056a0 |url-status=live|url-access=subscription }}</ref><ref>{{Cite journal |last1=Falini |first1=Giuseppe |last2=Albeck |first2=Shira |last3=Weiner |first3=Steve |last4=Addadi |first4=Lia |date=1996-01-05 |title=Control of Aragonite or Calcite Polymorphism by Mollusk Shell Macromolecules |url=https://www.science.org/doi/10.1126/science.271.5245.67 |journal=Science |language=en |volume=271 |issue=5245 |pages=67–69 |doi=10.1126/science.271.5245.67 |bibcode=1996Sci...271...67F |s2cid=95357556 |issn=0036-8075 |access-date=15 December 2022 |archive-date=15 December 2022 |archive-url=https://web.archive.org/web/20221215212206/https://www.science.org/doi/10.1126/science.271.5245.67 |url-status=live|url-access=subscription }}</ref><ref>{{Cite journal |last=Marin |first=Frédéric |date=October 2020 |title=Mollusc shellomes: Past, present and future |journal=Journal of Structural Biology |language=en |volume=212 |issue=1 |article-number=107583 |doi=10.1016/j.jsb.2020.107583|pmid=32721585 |s2cid=220850117|doi-access=free}}</ref>

==Occurrence==
[[File:Silfurberg.jpg|thumb|[[Calcite]] is the most stable [[polymorphism (materials science)|polymorph]] of calcium carbonate. It is transparent to opaque. A transparent variety called [[Iceland spar]] (shown here) was used to create [[polarized light]] in the 19th century.<ref name="Russell2008">Russell, Daniel E . 17 February 2008. Retrieved December 31, 2010. "[http://www.mindat.org/article.php/190/Helgustadir+Iceland+Spar+Mine Helgustadir Iceland Spar Mine] {{Webarchive|url=https://web.archive.org/web/20190508192834/https://www.mindat.org/article.php/190/Helgustadir+Iceland+Spar+Mine |date=8 May 2019}}" ''[[mindat.org]]''</ref>]]

===Geological sources===
[[Calcite]], [[aragonite]] and [[vaterite]] are pure calcium carbonate minerals. Industrially important source rocks which are predominantly calcium carbonate include [[limestone]], [[chalk]], [[marble]] and [[travertine]].

===Biological sources===
[[File:Calcium carbonate chunks.JPG|thumb|Calcium carbonate chunks from [[clam]]shell]]
[[Eggshell]]s, [[snail]] shells and most [[seashell]]s are predominantly calcium carbonate and can be used as industrial sources of that chemical.<ref>{{cite web |title=How are seashells created? |last=Horne |first=Francis |date=23 October 2006 |website=Scientific American |access-date=25 April 2012 |url=http://www.scientificamerican.com/article.cfm?id=how-are-seashells-created |archive-date=19 March 2011 |archive-url=https://web.archive.org/web/20110319230625/http://www.scientificamerican.com/article.cfm?id=how-are-seashells-created |url-status=live}}</ref> [[Oyster]] shells have enjoyed recent recognition as a source of dietary calcium, but are also a practical industrial source.<ref>{{cite web |url=http://www.webmd.com/drugs/drug-16642-Natural+Oyster+Shell+Calcium+Oral.aspx?drugid=16642&drugname=Natural+Oyster+Shell+Calcium+Oral| title=Oyster shell calcium |website=WebMD| access-date=25 April 2012}}</ref><ref>{{cite web|title=Oyster Shell Calcium Carbonate|publisher=Caltron Clays & Chemicals|url=http://caltronclays.in/Oyster_CC.html|access-date=25 April 2012|archive-url=https://web.archive.org/web/20130910033017/http://caltronclays.in/Oyster_CC.html|archive-date=10 September 2013}}</ref> Dark [[green vegetables]] such as [[broccoli]] and [[kale]] contain dietarily significant amounts of calcium carbonate, but they are not practical as an industrial source.<ref>{{cite journal|last=Mangels|first=Ann Reed|date=June 4, 2014|title=Bone nutrients for vegetarians|journal=The American Journal of Clinical Nutrition|volume=100|issue=1|pages=469S–475S|doi=10.3945/ajcn.113.071423|pmid=24898231|doi-access=free}}</ref>

[[Annelid]]s in the family [[Lumbricidae]], earthworms, possess a regionalization of the digestive track called [[calciferous glands]], Kalkdrüsen, or glandes de Morren, that processes calcium and [[Carbon dioxide|{{chem2|CO2}}]] into calcium carbonate, which is later excreted into the dirt.<ref>{{Cite journal |title=The Function of the Calciferous Glands of Earthworms |url=https://journals.biologists.com/jeb/article/13/3/279/4425/The-Function-of-the-Calciferous-Glands-of |access-date=2024-02-05 |journal=Journal of Experimental Biology |date=1936 |doi=10.1242/jeb.13.3.279 |archive-date=5 February 2024 |archive-url=https://web.archive.org/web/20240205140813/https://journals.biologists.com/jeb/article/13/3/279/4425/The-Function-of-the-Calciferous-Glands-of |url-status=live |last1=Robertson |first1=James D. |volume=13 |issue=3 |pages=279–297 |bibcode=1936JExpB..13..279R }}</ref> The function of these glands is unknown but is believed to serve as a {{chem2|CO2}} regulation mechanism within the animals' tissues.<ref>{{Cite journal |title=Stable isotopes reveal that the calciferous gland of earthworms is a CO2-fixing organ |url=https://www.sciencedirect.com/science/article/pii/S0038071707003872 |access-date=2024-02-05 |journal=Soil Biology and Biochemistry |date=2008 |doi=10.1016/j.soilbio.2007.09.012 |archive-date=29 January 2012 |archive-url=https://web.archive.org/web/20120129144311/http://www.sciencedirect.com/science/article/pii/S0038071707003872 |url-status=live |last1=Briones |first1=María Jesús Iglesias |last2=Ostle |first2=Nicholas J. |last3=Piearce |first3=Trevor G. |volume=40 |issue=2 |pages=554–557 |url-access=subscription }}</ref> This process is ecologically significant, stabilizing the [[Soil pH|pH of acid soils]].<ref>{{Cite web |title=Ecological functions of earthworms in soil |url=https://edepot.wur.nl/352649 |access-date=2024-02-05 |website=eDepot |archive-date=5 February 2024 |archive-url=https://web.archive.org/web/20240205140817/https://edepot.wur.nl/352649 |url-status=live}}</ref>

===Extraterrestrial===
Beyond Earth, strong evidence suggests the presence of calcium carbonate on [[Mars]]. Signs of calcium carbonate have been detected at more than one location (notably at [[Gusev crater|Gusev]] and [[Huygens (crater)|Huygens]] craters). This provides some evidence for the past presence of liquid water.<ref>{{cite journal | last1=Boynton | first1=W. V. | last2=Ming | first2=D. W. | last3=Kounaves | first3=S. P. | last4=Young | first4=S. M. | last5=Arvidson | first5=R. E. | last6=Hecht | first6=M. H. | last7=Hoffman | first7=J. | last8=Niles | first8=P. B. | last9=Hamara | first9=D. K. | last10=Quinn | first10=R. C. | last11=Smith | first11=P. H. | last12=Sutter | first12=B. | last13=Catling | first13=D. C. | last14=Morris | first14=R. V. | title=Evidence for Calcium Carbonate at the Mars Phoenix Landing Site | url=http://planetary.chem.tufts.edu/Boynton%20etal%20Science%202009v325p61.pdf | journal=Science | volume=325 | issue=5936 | pages=61–64 | year=2009 | pmid=19574384 | bibcode=2009Sci...325...61B | display-authors=3 | doi=10.1126/science.1172768 | s2cid=26740165 | access-date=7 January 2015 | archive-date=5 March 2016 | archive-url=https://web.archive.org/web/20160305022836/http://planetary.chem.tufts.edu/Boynton%20etal%20Science%202009v325p61.pdf | url-status=live}}</ref><ref name=Clark2007>{{cite journal |last1 = Clark |first1 = B. C. III |year = 2007 |title = Evidence for montmorillonite or its compositional equivalent in Columbia Hills, Mars |journal = [[Journal of Geophysical Research]] |volume = 112 |issue = E6 |pages = E06S01 |doi = 10.1029/2006JE002756 |last2 = Arvidson |first2 = R. E. |last3 = Gellert |first3 = R. |last4 = Morris |first4 = R. V. |last5 = Ming |first5 = D. W. |last6 = Richter |first6 = L. |last7 = Ruff |first7 = S. W. |last8 = Michalski |first8 = J. R. |last9 = Farrand |first9 = W. H. |last10 = Yen |first10 = A. |last11 = Herkenhoff |first11 = K. E. |last12 = Li |first12 = R. |last13 = Squyres |first13 = S. W. |last14 = Schröder |first14 = C. |last15 = Klingelhöfer |first15 = G. |last16 = Bell |first16 = J. F. |bibcode = 2007JGRE..112.6S01C |url = http://dspace.stir.ac.uk/bitstream/1893/17119/1/Clark2007_Evidence_for_montmorillonite_or_its_compositional_equivalent_in_Columbia_Hills_Mars.pdf |hdl = 1893/17119 |hdl-access = free |access-date = 20 April 2018 |archive-date = 29 July 2018 |archive-url = https://web.archive.org/web/20180729112107/https://dspace.stir.ac.uk/bitstream/1893/17119/1/Clark2007_Evidence_for_montmorillonite_or_its_compositional_equivalent_in_Columbia_Hills_Mars.pdf |url-status = live}}</ref>

==Geology==
[[File:Rubaksa tufa plug.jpg|thumb|left|Surface precipitation of {{chem2|CaCO3}} as [[tufa]] in [[Rubaksa]], Ethiopia]]
Carbonate is found frequently in geologic settings and constitutes an enormous [[carbon cycle|carbon reservoir]]. Calcium carbonate occurs as [[aragonite]], [[calcite]] and [[dolomite (mineral)|dolomite]] as significant constituents of the [[calcium cycle]]. The [[carbonate mineral]]s form the rock types: [[limestone]], [[chalk]], [[marble]], [[travertine]], [[tufa]], and others.
[[File:Water-of-Five-colored-Pond Huanglong Sichuan China.jpg|thumb|[[Tufa]] at [[Huanglong Scenic and Historic Interest Area|Huanglong]], [[Sichuan]]]]
In warm, clear tropical waters [[coral]]s are more abundant than towards the poles where the waters are cold. Calcium carbonate contributors, including [[plankton]] (such as [[coccolith]]s and planktic [[foraminifera]]), [[coralline algae]], [[sea sponge|sponges]], [[brachiopod]]s, [[echinoderm]]s, [[bryozoa]] and [[Mollusc shell|mollusks]], are typically found in shallow water environments where sunlight and filterable food are more abundant. Cold-water carbonates do exist at higher latitudes but have a very slow growth rate. The [[calcification]] processes are changed by [[ocean acidification]].

Where the [[oceanic crust]] is [[Subduction|subducted]] under a [[continental plate]] sediments will be carried down to warmer zones in the [[asthenosphere]] and [[lithosphere]]. Under these conditions calcium carbonate decomposes to produce [[carbon dioxide]] which, along with other gases, give rise to explosive [[volcano|volcanic eruptions]].

===Carbonate compensation depth===
The [[carbonate compensation depth]] (CCD) is the point in the ocean where the rate of precipitation of calcium carbonate is balanced by the rate of dissolution due to the conditions present. Deep in the ocean, the temperature drops and pressure increases. Increasing pressure also increases the solubility of calcium carbonate. Calcium carbonate is unusual in that its solubility increases with decreasing temperature.<ref name="weyl1959">{{cite journal |last1=Weyl |first1=P.K. |title=The change in solubility of calcium carbonate with temperature and carbon dioxide content |journal=Geochimica et Cosmochimica Acta |date=1959 |volume=17 |issue=3–4 |pages=214–225 |doi=10.1016/0016-7037(59)90096-1|bibcode=1959GeCoA..17..214W}}</ref> The carbonate compensation depth ranges from 4,000 to 6,000 meters below sea level in modern oceans, and the various polymorphs (calcite, aragonite) have different compensation depths based on their stability.<ref>{{Cite book |last=Burton |first=Elizabeth |date=1990 |title=Carbonate compensation depth |chapter=Carbonate carbonatescompensation depthcompensation depth |chapter-url=https://link.springer.com/referenceworkentry/10.1007/1-4020-4496-8_46  |page=73 |doi=10.1007/1-4020-4496-8_46 |isbn=978-1-4020-4496-0 |via=Elsevier |access-date=22 December 2023 |archive-date=22 December 2023 |archive-url=https://web.archive.org/web/20231222213347/https://link.springer.com/referenceworkentry/10.1007/1-4020-4496-8_46 |url-status=live}}</ref>

===Role in taphonomy===
Calcium carbonate can [[taphonomy|preserve fossils]] through [[permineralization]]. Most of the vertebrate fossils of the [[Two Medicine Formation]]—a [[geologic formation]] known for its [[duck-billed dinosaur]] eggs—are preserved by {{chem2|CaCO3}} permineralization.<ref name="twoturn" /> This type of preservation conserves high levels of detail, even down to the microscopic level. However, it also leaves specimens vulnerable to [[weathering]] when exposed to the surface.<ref name="twoturn">{{cite book |last=Trexler |first=D. |date=2001 |chapter-url=https://books.google.com/books?id=mgc6CS4EUPsC&pg=PA98 |chapter=Two Medicine Formation, Montana: geology and fauna |pages=[https://archive.org/details/mesozoicvertebra0000unse/page/298 298–309] |title=Mesozoic Vertebrate Life |editor1-last=Tanke |editor1-first=D. H. |editor2-last=Carpenter |editor2-first=K. |publisher=Indiana University Press |isbn=978-0-253-33907-2 |url=https://archive.org/details/mesozoicvertebra0000unse/page/298}}</ref>

[[Trilobite]] populations were once thought to have composed the majority of aquatic life during the [[Cambrian]], due to the fact that their calcium carbonate-rich shells were more easily preserved than those of other species,<ref>{{Cite book|url=https://www.nap.edu/catalog/11630/out-of-thin-air-dinosaurs-birds-and-earths-ancient-atmosphere|title=Out of Thin Air: Dinosaurs, Birds, and Earth's Ancient Atmosphere|last=Ward|first=Peter|year=2006|isbn=978-0-309-66612-1|language=en|doi=10.17226/11630|access-date=31 December 2017|archive-date=1 January 2018|archive-url=https://web.archive.org/web/20180101030500/https://www.nap.edu/catalog/11630/out-of-thin-air-dinosaurs-birds-and-earths-ancient-atmosphere|url-status=live}}</ref> which had purely [[chitin]]ous shells.

==Uses==
===Construction===
The main use of calcium carbonate is in the construction industry, either as a building material, or limestone [[Construction aggregate|aggregate]] for road building, as an ingredient of [[cement]], or as the starting material for the preparation of [[slaked lime|builders' lime]] by burning in a [[kiln]]. However, because of weathering mainly caused by [[acid rain]],<ref>{{cite web|title = Effects of Acid Rain| date=8 September 2006 |publisher = US Environmental Protection Agency|access-date = 14 March 2015|url = http://www.epa.gov/acidrain/effects/materials.html|archive-date = 2 March 2015|archive-url = https://web.archive.org/web/20150302173739/http://www.epa.gov/acidrain/effects/materials.html}}</ref> calcium carbonate (in limestone form) is no longer used for building purposes on its own, but only as a raw primary substance for building materials.

Calcium carbonate is also used in the purification of [[iron]] from [[iron ore]] in a [[blast furnace]]. The carbonate is [[Calcination|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|access-date = 30 December 2007|url = http://www.scienceaid.co.uk/chemistry/industrial/blastfurnace.html|archive-url = https://web.archive.org/web/20071217143213/http://www.scienceaid.co.uk/chemistry/industrial/blastfurnace.html|archive-date = 17 December 2007}}</ref>

In the [[oil industry]], calcium carbonate is added to [[drilling fluid]]s as a formation-bridging and filtercake-sealing agent; it is also a weighting material which increases the density of drilling fluids to control the downhole pressure. Calcium carbonate is added to [[swimming pool]]s, as a [[pH]] corrector for maintaining [[alkalinity]] and offsetting the acidic properties of the [[disinfectant]] agent.<ref>{{Cite book|url=https://books.google.com/books?id=8jF-AwAAQBAJ&q=Calcium+carbonate+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&pg=PT1601|title=Health & Drugs: Disease, Prescription & Medication|last=Sfetcu|first=Nicolae|date=2014-05-02|publisher=Nicolae Sfetcu|language=en}}</ref>

It is also used as a raw material in the [[Sugar refining|refining of sugar]] from [[sugar beet]]; it is calcined in a kiln with [[anthracite]] to produce calcium oxide and carbon dioxide. This burnt lime is then slaked in fresh water to produce a calcium hydroxide [[Suspension (chemistry)|suspension]] for the precipitation of impurities in raw juice during [[carbonatation]].<ref>{{cite book|last1=McGinnis|first1=R. A.|title=Beet-Sugar Technology|publisher=Beet Sugar Development Foundation|page=178|edition=2nd}}</ref>

Calcium carbonate in the form of [[chalk]] has traditionally been a major component of [[blackboard]] chalk. However, modern manufactured chalk is mostly [[gypsum]], hydrated [[calcium sulfate]] {{chem2|CaSO4*2H2O}}. Calcium carbonate is a main source for growing [[biorock]]. Precipitated calcium carbonate (PCC), pre-dispersed in [[slurry]] form, is a common filler material for [[latex gloves]] with the aim of achieving maximum saving in material and production costs.<ref name=precaco3>{{cite web|title=Precipitated Calcium Carbonate uses |url=http://www.aristocratholding.com/calris-5.html |archive-url=https://web.archive.org/web/20140725032803/http://www.aristocratholding.com/calris-5.html |archive-date=25 July 2014}}</ref>

Fine ground calcium carbonate (GCC) is an essential ingredient in the microporous film used in [[diapers]] and some building films, as the pores are nucleated around the calcium carbonate particles during the manufacture of the film by biaxial stretching. GCC and PCC are used as a filler in [[paper]] because they are cheaper than [[wood fiber]]. Printing and writing paper can contain 10–20% calcium carbonate. 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. PCC used for paper filling and paper coatings is precipitated and prepared in a variety of shapes and sizes having characteristic narrow particle size distributions and equivalent spherical diameters of 0.4 to 3 micrometers.{{citation needed|date=June 2015}}

Calcium carbonate is widely used as an [[Extender (ink)|extender]] in [[paint]]s,<ref name = reade>{{cite web|title = Calcium Carbonate Powder|publisher = Reade Advanced Materials|date = 4 February 2006|access-date = 30 December 2007|url = http://www.reade.com/Products/Minerals_and_Ores/calcium_carbonate.html|archive-url = https://web.archive.org/web/20080222003757/http://www.reade.com/Products/Minerals_and_Ores/calcium_carbonate.html|archive-date = 22 February 2008}}</ref> in particular [[Gloss and matte paint|matte]] [[emulsion paint]] where typically 30% by weight of the paint is either chalk or marble. It is also a popular filler in plastics.<ref name = reade/> Some typical examples include around 15–20% loading of chalk in [[Polyvinyl chloride|unplasticized polyvinyl chloride]] (uPVC) [[Rain gutter|drainpipes]], 5–15% loading of [[stearic acid|stearate]]-coated chalk or marble in uPVC window profile. [[Polyvinyl chloride|PVC]] cables can use calcium carbonate at loadings of up to 70&nbsp;phr (parts per hundred parts of resin) to improve mechanical properties (tensile strength and elongation) and electrical properties (volume resistivity).{{citation needed|date=June 2015}} [[Polypropylene]] compounds are often filled with calcium carbonate to increase rigidity, a requirement that becomes important at high usage temperatures.<ref name= Imerys>{{cite web |url=http://www.imerys-perfmins.com/calcium-carbonate/eu/calcium-carbonate-plastic.htm |title=Calcium carbonate in plastic applications |access-date=1 August 2008 |publisher=Imerys Performance Minerals |archive-url=https://web.archive.org/web/20080804020529/http://www.imerys-perfmins.com/calcium-carbonate/eu/calcium-carbonate-plastic.htm |archive-date=4 August 2008}}</ref> Here the percentage is often 20–40%. It also routinely used as a filler in [[Thermosetting plastic|thermosetting resins]] (sheet and bulk molding compounds)<ref name = Imerys/> and has also been mixed with [[acrylonitrile butadiene styrene|ABS]], and other ingredients, to form some types of compression molded "clay" [[poker chip]]s.<ref>{{Cite web|url=http://www.xintuchemical.com/why-do-calcium-carbonate-play-an-important-part-in-industrial-a-59.html|title=Why do calcium carbonate play an important part in Industrial|website=www.xintuchemical.com|language=en|access-date=2018-10-07|archive-date=7 October 2018|archive-url=https://web.archive.org/web/20181007183653/http://www.xintuchemical.com/why-do-calcium-carbonate-play-an-important-part-in-industrial-a-59.html|url-status=live}}</ref> Precipitated calcium carbonate, made by dropping [[calcium oxide]] into water, is used by itself or with additives as a white paint, known as [[whitewash]]ing.<ref>{{Cite web|url=https://www.dgci.be/18122-11/precipitated_calcium_carbonate_commodity_price|title=precipitated calcium carbonate commodity price|website=www.dgci.be|access-date=2018-10-07|archive-url=https://web.archive.org/web/20181007223232/https://www.dgci.be/18122-11/precipitated_calcium_carbonate_commodity_price|archive-date=7 October 2018}}</ref><ref name="JimohEtAl2017">{{cite journal | url=http://www.scielo.org.za/pdf/sajc/v70/01.pdf | title=Understanding the Precipitated Calcium Carbonate (PCC) Production Mechanism and Its Characteristics in the Liquid–Gas System Using Milk of Lime (MOL) Suspension | author=Jimoh, O.A. | display-authors=et al | journal=South African Journal of Chemistry | year=2017 | volume=70 | pages=1–7 | doi=10.17159/0379-4350/2017/v70a1 | doi-access=free | access-date=7 October 2018 | archive-date=21 September 2018 | archive-url=https://web.archive.org/web/20180921161721/http://www.scielo.org.za/pdf/sajc/v70/01.pdf | url-status=live}}</ref>

Calcium carbonate is added to a wide range of trade and [[do it yourself]] 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.<ref>{{Cite web|url=https://www.chemicalprocessing.com/experts/environmental-protection/show/533|access-date=2021-02-03|website=www.chemicalprocessing.com | title = Topic: Re: Can our calcium carbonate "waste" be utilized in other industries so we can divert it from landfills? | date = 4 March 2010 | url-status = live|archive-url=https://web.archive.org/web/20170323143544/http://www.chemicalprocessing.com:80/experts/environmental-protection/show/533/ |archive-date=23 March 2017}}</ref><ref>{{Cite web|title=Why do calcium carbonate play an important part in Industry?|url=http://www.xintuchemical.com/why-do-calcium-carbonate-play-an-important-part-in-industrial-a-59.html|access-date=2021-02-03|website=www.xintuchemical.com|archive-date=7 October 2018|archive-url=https://web.archive.org/web/20181007183653/http://www.xintuchemical.com/why-do-calcium-carbonate-play-an-important-part-in-industrial-a-59.html|url-status=live}}</ref><ref>{{Cite web|title=Calcium Carbonates / Calcite/ Limestone. CaCO3 {{!}} Rajasthan Minerals & Chemicals|url=http://www.rmcl.co.in/calcium-carbonates_calcite.html|access-date=2021-02-03|website=www.rmcl.co.in|archive-date=15 April 2021|archive-url=https://web.archive.org/web/20210415090019/http://www.rmcl.co.in/calcium-carbonates_calcite.html}}</ref><ref>{{Cite web|url=http://kamceramics.com/portfolio/calcium-carbonate/|title=Calcium Carbonate|access-date=2021-02-03|website=kamceramics.com|archive-date=15 April 2021|archive-url=https://web.archive.org/web/20210415064720/http://kamceramics.com/portfolio/calcium-carbonate/|url-status=live}}</ref>

In [[ceramic glaze]] applications, calcium carbonate is known as ''whiting'',<ref name = reade/> and is 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 [[Ceramic flux|flux]] material in the glaze. Ground calcium carbonate is an [[abrasive]] (both as scouring powder and as an ingredient of household scouring creams), in particular in its calcite form, which has the relatively low hardness level of 3 on the [[Mohs scale of mineral hardness|Mohs scale]], and will therefore not scratch [[glass]] and most other [[ceramic]]s, [[Vitreous enamel|enamel]], [[bronze]], [[iron]], and [[steel]], and have a moderate effect on softer metals like [[aluminium]] and [[copper]]. A paste made from calcium carbonate and [[deionized water]] can be used to clean [[tarnish]] on [[silver]].<ref name="Make it Shine">{{cite web|title = Ohio Historical Society Blog: Make It Shine| work=Ohio History Connection Collections Blog | date=2 June 2011 |publisher = Ohio Historical Society|url = http://ohiohistory.wordpress.com/2011/06/02/making-it-shine/|access-date = 2 June 2011|archive-url = https://web.archive.org/web/20120323201653/http://ohiohistory.wordpress.com/2011/06/02/making-it-shine/|archive-date = 23 March 2012}}</ref>

===Health and diet===
[[File: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|gastric antacid]]<ref name = medline>{{cite web|website = Medline Plus|publisher = [[National Institutes of Health]]|title = Calcium Carbonate |date=1 October 2005|access-date = 30 December 2007|url = https://www.medlineplus.gov/druginfo/meds/a601032.html |url-status = live|archive-url = https://web.archive.org/web/20071017031324/http://www.nlm.nih.gov/medlineplus/druginfo/medmaster/a601032.html <!-- Bot retrieved archive --> |archive-date = 17 October 2007}}</ref> (such as [[Tums]] and [[Eno (drug)|Eno]]). It may be used as a [[phosphate binder]] for the treatment of [[hyperphosphatemia]] (primarily in patients with [[chronic kidney failure]]). It is used in the pharmaceutical industry as an inert [[Excipient|filler]] for [[Tablet (pharmacy)|tablets]] and other [[pharmaceuticals]].<ref>{{cite book|last1=Lieberman |first1=Herbert A. |last2=Lachman |first2=Leon |last3=Schwartz |first3=Joseph B. |title = Pharmaceutical Dosage Forms: Tablets|url=https://archive.org/details/pharmaceuticaldo03lach |url-access=limited |year = 1990|isbn = 978-0-8247-8044-9|page=[https://archive.org/details/pharmaceuticaldo03lach/page/n171 153]|publisher = Dekker|location = New York}}</ref>

Calcium carbonate is used in the production of calcium oxide as well as toothpaste and has seen a resurgence as a food preservative and color retainer, when used in or with products such as organic apples.<ref>{{cite web |url=http://chemistry.about.com/od/foodcookingchemistry/a/cadditives.htm |title=Food Additives – Names Starting with C |website=Chemistry.about.com |date=10 April 2012 |access-date=2012-05-24 |archive-url=https://web.archive.org/web/20061016112555/http://chemistry.about.com/od/foodcookingchemistry/a/cadditives.htm |archive-date=16 October 2006}}</ref>

Calcium carbonate is used therapeutically as phosphate binder in patients on maintenance [[haemodialysis]]. It is the most common form of phosphate binder prescribed, particularly in non-dialysis chronic kidney disease. Calcium carbonate is the most commonly used phosphate binder, but clinicians are increasingly prescribing the more expensive, non-calcium-based phosphate binders, particularly [[sevelamer]].

Excess calcium from supplements, fortified food, and high-calcium diets can cause [[milk-alkali syndrome]], which has serious toxicity and can be fatal. In 1915, [[Bertram Welton Sippy|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 [[kidney failure]], [[alkalosis]], and [[hypercalcaemia]], 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 arose. Since the 1990s it has been most frequently reported in women taking calcium supplements above the recommended range of 1.2 to 1.5&nbsp;grams daily, for prevention and treatment of osteoporosis,<ref name="pmid17483976">{{cite journal |vauthors=Caruso JB, Patel RM, Julka K, Parish DC |title=Health-behavior induced disease: return of the milk-alkali syndrome |journal=J Gen Intern Med |volume=22 |issue=7 |pages=1053–5 |date=July 2007 |pmid=17483976 |doi=10.1007/s11606-007-0226-0 |pmc=2219730}}</ref><ref name="pmid16702792">{{cite journal |vauthors=Beall DP, Henslee HB, Webb HR, Scofield RH |title=Milk-alkali syndrome: a historical review and description of the modern version of the syndrome |journal=Am. J. Med. Sci. |volume=331 |issue=5 |pages=233–42 |date=May 2006 |pmid=16702792 |doi= 10.1097/00000441-200605000-00001|s2cid=45802184}}</ref> 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>{{cite journal|title=Clinical problem-solving, back to basics|last1=Gabriely|first1=Ilan|journal=New England Journal of Medicine|year=2008|volume=358|pmid=18450607|doi=10.1056/NEJMcps0706188|issue=18|last2=Leu|first2=James P.|last3=Barzel|first3=Uriel S.|pages=1952–6}}</ref>

As a [[food additive]] it is designated [[E numbers|E170]],<ref>{{cite web|website=Food-Info.net|title=E-numbers: E170 Calcium carbonate|url=http://www.food-info.net/uk/e/e170.htm|access-date=19 April 2008|archive-date=14 October 2022|archive-url=https://web.archive.org/web/20221014050719/http://www.food-info.net/uk/e/e170.htm|url-status=live}} 080419 food-info.net</ref> and it has an [[International Numbering System for Food Additives|INS]] number of 170. Used as an [[acidity regulator]], [[anticaking agent]], [[Stabilizer (food)|stabilizer]] or [[Food coloring|color]] it is approved for usage in the EU,<ref>{{cite web |publisher=UK Food Standards Agency |url=http://www.food.gov.uk/safereating/chemsafe/additivesbranch/enumberlist |title=Current EU approved additives and their E Numbers |access-date=27 October 2011 |archive-date=7 October 2010 |archive-url=https://web.archive.org/web/20101007124435/http://www.food.gov.uk/safereating/chemsafe/additivesbranch/enumberlist |url-status=live}}</ref> US<ref>{{cite web|url=https://www.fda.gov/Food/FoodIngredientsPackaging/FoodAdditives/FoodAdditiveListings/ucm091048.htm|publisher=US [[Food and Drug Administration]] |title=Listing of Food Additives Status Part I |access-date=27 October 2011 |archive-url=https://web.archive.org/web/20130314104055/https://www.fda.gov/Food/FoodIngredientsPackaging/FoodAdditives/FoodAdditiveListings/ucm091048.htm |archive-date=14 March 2013}}</ref> and [[Australia]] and [[New Zealand]].<ref>{{cite web |url=http://www.comlaw.gov.au/Details/F2011C00827 |title=Standard 1.2.4 – Labelling of ingredients |date=8 September 2011 |access-date=27 October 2011 |publisher=Australia New Zealand Food Standards Code |archive-date=2 September 2013 |archive-url=https://web.archive.org/web/20130902084805/http://www.comlaw.gov.au/Details/F2011C00827 |url-status=live}}</ref> It is "added by law to all UK milled bread flour except wholemeal".<ref name="lhbread">{{cite news |last1=Holdstock |first1=Lee |title=Why go organic? |url=https://www.sustainweb.org/realbread/why_go_organic/ |access-date=3 April 2021 |agency=Soil Association Certification Limited |publisher=Real Bread Campaign |archive-date=14 October 2022 |archive-url=https://web.archive.org/web/20221014050716/https://www.sustainweb.org/realbread/why_go_organic/ |url-status=live}}</ref><ref name="defra13">{{cite news |title=Bread and Flour Regulations 1998 A summary of responses to the consultation and Government Reply |url=https://assets.publishing.service.gov.uk/government/uploads/system/uploads/attachment_data/file/226553/bread-flour-sum-resp-130805.pdf |publisher=Department for Environment, Food and Rural Affairs |date=August 2013 |access-date=9 April 2021 |archive-date=19 September 2021 |archive-url=https://web.archive.org/web/20210919172916/https://assets.publishing.service.gov.uk/government/uploads/system/uploads/attachment_data/file/226553/bread-flour-sum-resp-130805.pdf |url-status=live}}</ref> It is used in some [[soy milk]] and [[almond milk]] products as a source of dietary calcium; at least one study suggests that calcium carbonate might be as [[bioavailable]] as the calcium in [[cow's milk]].<ref>{{Cite journal| pmid = 16177199| year = 2005| last1 = Zhao| first1 = Y.| title = Calcium bioavailability of calcium carbonate fortified soymilk is equivalent to cow's milk in young women| journal = The Journal of Nutrition| volume = 135| issue = 10| pages = 2379–2382| last2 = Martin| first2 = B. R.| last3 = Weaver| first3 = C. M.| doi = 10.1093/jn/135.10.2379| doi-access = free}}</ref> Calcium carbonate is also used as a [[firming agent]] in many canned and bottled vegetable products.

Several calcium supplement formulations have been documented to contain the chemical element [[lead]],<ref name="Kauffman-2007">{{Cite journal|date=2007-07-01|title=Lead in pharmaceutical products and dietary supplements|url=https://www.sciencedirect.com/science/article/abs/pii/S0273230007000360|journal=Regulatory Toxicology and Pharmacology|language=en|volume=48|issue=2|pages=128–134|doi=10.1016/j.yrtph.2007.03.001|issn=0273-2300|last1=Kauffman|first1=John F.|last2=Westenberger|first2=Benjamin J.|last3=Robertson|first3=J. David|last4=Guthrie|first4=James|last5=Jacobs|first5=Abigail|last6=Cummins|first6=Susan K.|pmid=17467129|access-date=11 July 2021|archive-date=11 July 2021|archive-url=https://web.archive.org/web/20210711074707/https://www.sciencedirect.com/science/article/abs/pii/S0273230007000360|url-status=live|url-access=subscription}}</ref> posing a [[public health]] concern.<ref name="ross20">{{cite journal |doi=10.1001/jama.284.11.1425|title=Lead Content of Calcium Supplements|year=2000|last1=Ross|first1=Edward A.|last2=Szabo|first2=N. J.|last3=Tebbett|first3=I. R.|journal=JAMA|volume=284|issue=11|pages=1425–1429|pmid=10989406}}</ref> Lead is commonly found in natural sources of calcium.<ref name="Kauffman-2007" />

===Agriculture and aquaculture===
[[Agricultural lime]], powdered chalk or limestone, is used as a cheap method of neutralising [[Soil pH|acidic soil]], making it suitable for planting, also used in aquaculture industry for pH regulation of pond soil before initiating culture.<ref name="Oates2008">{{cite book |first=J. A. H. |last=Oates |title=Lime and Limestone: Chemistry and Technology, Production and Uses |url=https://books.google.com/books?id=MVoEMNI5Vb0C&pg=PA111 |date=11 July 2008 |publisher=John Wiley & Sons |isbn=978-3-527-61201-7 |pages=111–113}}</ref> There is interest in understanding whether or not it can affect pesticide adsorption and desorption in calcareous soil.<ref>{{Cite journal |last1=El-Aswad |first1=Ahmed F. |last2=Fouad |first2=Mohamed R. |last3=Badawy |first3=Mohamed E. I. |last4=Aly |first4=Maher I. |date=2023-05-31 |title=Effect of Calcium Carbonate Content on Potential Pesticide Adsorption and Desorption in Calcareous Soil |url=https://www.tandfonline.com/doi/full/10.1080/00103624.2022.2146131 |journal=Communications in Soil Science and Plant Analysis |language=en |volume=54 |issue=10 |pages=1379–1387 |doi=10.1080/00103624.2022.2146131 |bibcode=2023CSSPA..54.1379E |s2cid=253559627 |issn=0010-3624 |access-date=18 August 2023 |archive-date=18 August 2023 |archive-url=https://web.archive.org/web/20230818130720/https://www.tandfonline.com/doi/full/10.1080/00103624.2022.2146131 |url-status=live|url-access=subscription }}</ref>

===Household cleaning===
Calcium carbonate is a key ingredient in many household cleaning powders like [[Comet (cleanser)|Comet]] and is used as a scrubbing agent.

===Pollution mitigation===
In 1989, a researcher, Ken Simmons, introduced {{chem2|CaCO3}} into the Whetstone Brook in [[Massachusetts]].<ref name="New York Times">{{cite news|agency = [[Associated Press]]|title = Limestone Dispenser Fights Acid Rain in Stream|date = 13 June 1989|url = https://www.nytimes.com/1989/06/13/science/limestone-dispenser-fights-acid-rain-in-stream.html|work = The New York Times|access-date = 27 July 2018|archive-date = 28 July 2018|archive-url = https://web.archive.org/web/20180728002911/https://www.nytimes.com/1989/06/13/science/limestone-dispenser-fights-acid-rain-in-stream.html|url-status = live}}</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 amount of aluminium ions in the area of the brook that was not treated with the limestone. This shows that {{chem2|CaCO3}} 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 name=env>{{cite web|title=Environmental Uses for Calcium Carbonate|date=6 September 2012|url=http://www.congcal.com/markets/environmental/|publisher=Congcal|access-date=5 August 2013|archive-date=4 January 2014|archive-url=https://web.archive.org/web/20140104055647/http://www.congcal.com/markets/environmental/}}</ref><ref name="Schreiber-1988">{{cite journal|last = Schreiber|first = R. K.|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|bibcode = 1988WASP...41...53S|s2cid = 98404326|url = https://link.springer.com/article/10.1007%2FBF00160344|access-date = 28 August 2017|archive-date = 10 January 2018|archive-url = https://web.archive.org/web/20180110054856/https://link.springer.com/article/10.1007%2FBF00160344|url-status = live|url-access = subscription}}</ref> Since the 1970s, such ''liming'' has been practiced on a large scale in Sweden to mitigate acidification and several thousand lakes and streams are limed repeatedly.<ref name="Guhrén-2006">{{Cite journal |doi= 10.1007/s10933-006-9014-9 |title= Liming placed in a long-term perspective: A paleolimnological study of 12 lakes in the Swedish liming program |journal= Journal of Paleolimnology |volume= 37 |issue= 2 |pages= 247–258 |year= 2006 |last1= Guhrén |first1= M. |last2= Bigler |first2= C. |last3= Renberg |first3= I. |bibcode= 2007JPall..37..247G |s2cid= 129439066}}</ref>

Calcium carbonate is also used in [[flue-gas desulfurization]] applications eliminating harmful {{chem2|SO2}} and {{chem2|NO2}} emissions from coal and other fossil fuels burnt in large fossil fuel power stations.<ref name=env/>

===Plastics===
Calcium carbonate is commonly used in the plastic industry as a filler. When it is incorporated in a plastic material, it can improve the hardness, stiffness, dimensional stability and processability of the material.<ref>{{cite web |url=https://europlas.com.vn/en-US/blog-1/why-calcium-carbonate-used-in-plastic-industry |title=Why calcium carbonate used in plastic industry |author=<!--Not stated--> |date= |website= |publisher=EuroPlas |access-date=12 July 2024 |quote=}}</ref>

==Calcination equilibrium==
[[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&nbsp;°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 {{chem2|CO2}} pressure is only a tiny fraction of the partial {{chem2|CO2}} pressure in air, which is about 0.035&nbsp;kPa.

At temperatures above 550&nbsp;°C the equilibrium {{chem2|CO2}} pressure begins to exceed the {{chem2|CO2}} pressure in air. So above 550&nbsp;°C, calcium carbonate begins to outgas {{chem2|CO2}} into air. However, in a charcoal fired kiln, the concentration of {{chem2|CO2}} 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 {{chem2|CO2}} in the kiln can be as high as 20&nbsp;kPa.<ref name="solvaypcc2007">{{cite web|title = Solvay Precipitated Calcium Carbonate: Production|publisher = Solvay|date = 9 March 2007|access-date = 30 December 2007|url = http://www.solvaypcc.com/safety_environment/0,0,1000044-_EN,00.html|archive-date = 19 October 2007|archive-url = https://web.archive.org/web/20071019031826/http://www.solvaypcc.com/safety_environment/0,0,1000044-_EN,00.html|url-status = live}}</ref>

The table shows that this partial pressure is not achieved until the temperature is nearly 800&nbsp;°C. For the outgassing of {{chem2|CO2}} from calcium carbonate to happen at an economically useful rate, the equilibrium pressure must significantly exceed the ambient pressure of {{chem2|CO2}}. And for it to happen rapidly, the equilibrium pressure must exceed total atmospheric pressure of 101&nbsp;kPa, which happens at 898&nbsp;°C.{{clear right}}

:{| class="wikitable"
|+ Equilibrium pressure of {{chem2|CO2}} over {{chem2|CaCO3}} (''P'') versus temperature (''T'').<ref name=crc>{{RubberBible86th}}</ref>
|-
!''P'' (kPa)
|0.055||0.13||0.31||1.80||5.9||9.3||14||24||34||51||72 ||80||91||101||179||901||3961
|-
!''T'' (°C)
|550||587||605||680||727||748||777||800||830||852||871||881||891||898||937||1082||1241
|}

==Solubility==
===With varying {{chem2|CO2}} pressure===
[[File:CanarySpring.jpg|thumb|right|[[Travertine]] calcium carbonate deposits from a [[hot spring]]]]
Calcium carbonate is poorly soluble in pure water (47&nbsp;mg/L at normal atmospheric {{chem2|CO2}} partial pressure as shown below).

The equilibrium of its solution is given by the equation (with dissolved calcium carbonate on the right):
:{| width="500"
| style="width:50%; height:30px;"| {{chem2|CaCO3 ⇌ Ca(2+) + CO3(2−)}}
| ''K''<sub>sp</sub> = {{val|3.7e-9}} to {{val|8.7e-9}} at 25&nbsp;°C
|}

where the [[solubility product]] for {{chem2|[Ca(2+)][CO3(2−)]}} is given as anywhere from ''K''<sub>sp</sub> = {{val|3.7e-9}} to ''K''<sub>sp</sub> = {{val|8.7e-9}} at 25&nbsp;°C, depending upon the data source.<ref name = crc/><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|access-date = 7 June 2007|archive-date = 25 May 2006|archive-url = https://web.archive.org/web/20060525044457/http://www.csudh.edu/oliver/chemdata/data-ksp.htm}}</ref> What the equation means is that the product of molar concentration of calcium ions ([[mole (unit)|moles]] of dissolved {{chem2|Ca(2+)}} per liter of solution) with the molar concentration of dissolved {{chem2|CO3(2−)}} 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 {{chem2|CO3(2−)}} combines with {{chem2|H+}} in the solution according to

:{| width="500"
| style="width:50%; height:25px;"| {{chem2|HCO3− ⇌ H+ + CO3(2−)}} &nbsp;&nbsp;
| ''K''<sub>a2</sub> = {{val|5.61e-11}} at 25&nbsp;°C
|}

{{chem2|HCO3−}} 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 {{chem2|HCO3−}} combines with {{chem2|H+}} in solution according to

:{| 
| width=250|{{chem2|H2CO3 ⇌ H+ + HCO3−}}
| ''K''<sub>a1</sub> = {{val|2.5e-4}} at 25&nbsp;°C
|}

Some of the {{chem2|H2CO3}} breaks up into water and dissolved carbon dioxide according to

:{| 
|width=250|{{chem2|H2O + CO2([[aqueous|aq]]) ⇌ H2CO3}}
| ''K''<sub>h</sub> = {{val|1.70e-3}} at 25&nbsp;°C
|}

And dissolved carbon dioxide is in equilibrium with atmospheric carbon dioxide according to

:{| 
| width=250|{{sfrac|''P''<sub>{{chem2|CO2}}</sub>|{{chem2|[CO2]}}}} = <math>H_{\rm v}</math>
| where <math>H_{\rm v}</math> = 29.76 atm/(mol/L) at 25&nbsp;°C ([[Henry's law|Henry volatility]]), and ''P''<sub>{{chem2|CO2}}</sub> is the {{chem2|CO2}} partial pressure.
|}

For ambient air, ''P''<sub>{{chem2|CO2}}</sub> is around {{val|3.5e-4}} [[standard atmosphere (unit)|atm]] (or equivalently 35&nbsp;[[Pascal (unit)|Pa]]). The last equation above fixes the concentration of dissolved {{chem2|CO2}} as a function of ''P''<sub>{{chem2|CO2}}</sub>, independent of the concentration of dissolved {{chem2|CaCO3}}. At atmospheric partial pressure of {{chem2|CO2}}, dissolved {{chem2|CO2}} concentration is {{val|1.2e-5}} moles per liter. The equation before that fixes the concentration of {{chem2|H2CO3}} as a function of {{chem2|CO2}} concentration. For [{{chem2|CO2}}] = {{val|1.2e-5}}, it results in {{chem2|[H2CO3]}} = {{val|2.0e-8}} moles per liter. When {{chem2|[H2CO3]}} is known, the remaining three equations together with
{| class="wikitable floatright"
|+ Calcium ion solubility as a function of [[carbon dioxide|{{chem2|CO2}}]] [[partial pressure]] at 25&nbsp;°C (''K''<sub>sp</sub>&nbsp;=&nbsp;{{val|4.47e-9}})
|-
!''P''<sub>{{chem2|CO2}}</sub> (atm)
![[pH]]
!{{chem2|[Ca(2+)]}} (mol/L)
|-
| 10<sup>−12</sup> ||12.0||{{val|5.19e-3}}
|-
| 10<sup>−10</sup> ||11.3||{{val|1.12e-3}}
|-
| 10<sup>−8</sup> ||10.7||{{val|2.55e-4}}
|-
| 10<sup>−6</sup> ||9.83||{{val|1.20e-4}}
|-
| 10<sup>−4</sup> ||8.62||{{val|3.16e-4}}
|-
| {{val|3.5e-4}}||8.27||{{val|4.70e-4}}
|-
| 10<sup>−3</sup> ||7.96||{{val|6.62e-4}}
|-
| 10<sup>−2</sup> ||7.30||{{val|1.42e-3}}
|-
| 10<sup>−1</sup> ||6.63||{{val|3.05e-3}}
|-
| 1 ||5.96||{{val|6.58e-3}}
|-
| 10 ||5.30||{{val|1.42e-2}}
|}

:{| width="450"
| style="width:50%; height:25px;"|{{chem2|H2O ⇌ H+ + OH−}}
| ''K'' = 10<sup>−14</sup> at 25&nbsp;°C
|}

(which is true for all aqueous solutions), and the constraint that the solution must be electrically neutral, i.e., the overall charge of dissolved positive ions {{chem2|[Ca(2+)] + 2 [H+]}} must be cancelled out by the overall charge of dissolved negative ions {{chem2|[HCO3−] + [CO3(2−)] + [OH−]}}, make it possible to solve simultaneously for the remaining five unknown concentrations (the previously mentioned form of the neutrality 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 initial water solvent pH is not neutral, the balance is not neutral).

The adjacent table shows the result for {{chem2|[Ca(2+)]}} and {{chem2|[H+]}} (in the form of pH) as a function of ambient partial pressure of {{chem2|CO2}} (''K''<sub>sp</sub> = {{val|4.47e-9}} has been taken for the calculation).
* At atmospheric levels of ambient {{chem2|CO2}} the table indicates that the solution will be slightly alkaline with a maximum {{chem2|CaCO3}} solubility of 47&nbsp;mg/L.
* As ambient {{chem2|CO2}} partial pressure is reduced below atmospheric levels, the solution becomes more and more alkaline. At extremely low ''P''<sub>{{chem2|CO2}}</sub>, dissolved {{chem2|CO2}}, bicarbonate ion, and carbonate ion largely evaporate from the solution, leaving a highly alkaline solution of [[calcium hydroxide]], which is more soluble than {{chem2|CaCO3}}. For ''P''<sub>{{chem2|CO2}}</sub> = 10<sup>−12</sup>&nbsp;atm, the {{chem2|[Ca(2+)][OH−]2}} product is still below the solubility product of {{chem2|Ca(OH)2}} ({{val|8e-6}}). For still lower {{chem2|CO2}} pressure, {{chem2|Ca(OH)2}} precipitation will occur before {{chem2|CaCO3}} precipitation.
* As ambient {{chem2|CO2}} 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 {{chem2|Ca(2+)}}.

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 {{chem2|CO2}} much higher than atmospheric. As such, water percolates through calcium carbonate rock, the {{chem2|CaCO3}} dissolves according to one of the trends above. When that same water then emerges from the tap, in time, it comes into equilibrium with {{chem2|CO2}} levels in the air by outgassing its excess {{chem2|CO2}}. 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]] {{chem2|CaCO3*H2O}} and [[ikaite]] {{chem2|CaCO3*6H2O}}, may [[precipitate]] from water at ambient conditions and persist as metastable phases.

=== With varying pH, temperature and salinity: {{chem2|CaCO3}} scaling in swimming pools ===
[[File:CaCO3-pH.gif|thumb|alt=Effects of salinity and pH on the maximum calcium ion level before scaling is anticipated at 25 °C and 1 mmol/L bicarbonate concentration (e.g. in swimming pools)]]
[[File:CaCO3-Temp.gif|thumb|alt=Effects of temperature and bicarbonate concentration on the maximum calcium ion level before scaling is anticipated at pH&nbsp;7 and 5,000&nbsp;ppm salinity (such as in swimming pools)]]
In contrast to the open equilibrium scenario above, many swimming pools are managed by addition of [[sodium bicarbonate]] ({{chem2|NaHCO3}}) to the concentration of about 2&nbsp;mmol/L as a buffer, then control of pH through use of HCl, {{chem2|NaHSO4}}, {{chem2|Na2CO3}}, NaOH or chlorine formulations that are acidic or basic. In this situation, dissolved inorganic carbon ([[total inorganic carbon]]) is far from equilibrium with atmospheric {{chem2|CO2}}. Progress towards equilibrium through outgassing of {{chem2|CO2}} is slowed by
{{ordered list
|the slow reaction
:{{chem2|[[Carbonic acid|H2CO3]] ⇌ CO2([[aqueous solution|aq]]) + H2O}};<ref>{{Cite journal | doi = 10.1021/jp909019u| pmid = 20039712| title = Comprehensive Study of the Hydration and Dehydration Reactions of Carbon Dioxide in Aqueous Solution| journal = The Journal of Physical Chemistry A| volume = 114| issue = 4| pages = 1734–40| year = 2010| last1 = Wang | first1 = X. | last2 = Conway | first2 = W. | last3 = Burns | first3 = R. | last4 = McCann | first4 = N. | last5 = Maeder | first5 = M. | bibcode = 2010JPCA..114.1734W}}</ref>
|limited aeration in a deep water column; and
|periodic replenishment of bicarbonate to maintain buffer capacity (often estimated through measurement of [[alkalinity|total alkalinity]]).}}

In this situation, the dissociation constants for the much faster reactions
:{{chem2|H2CO3 ⇌ H+ + HCO3− ⇌ 2 H+ + CO3(2−)}}
allow the prediction of concentrations of each dissolved inorganic carbon species in solution, from the added concentration of {{chem2|HCO3−}} (which constitutes more than 90% of [[Bjerrum plot]] species from pH&nbsp;7 to pH&nbsp;8 at 25&nbsp;°C in fresh water).<ref name="Mook 2000">{{cite book|last=Mook|first=W.|date=2000|url=http://www-naweb.iaea.org/napc/ih/documents/global_cycle/vol%20I/cht_i_09.pdf|contribution=Chemistry of carbonic acid in water|pages=143–165|title=Environmental Isotopes in the Hydrological Cycle: Principles and Applications|publisher=INEA/UNESCO|location=Paris|access-date=18 March 2014|archive-url=https://web.archive.org/web/20140318074927/http://www-naweb.iaea.org/napc/ih/documents/global_cycle/vol%20I/cht_i_09.pdf|archive-date=18 March 2014}}</ref> Addition of {{chem2|HCO3−}} will increase {{chem2|CO3(2−)}} concentration at any pH. Rearranging the equations given above, we can see that {{chem2|[Ca(2+)]}} = {{sfrac|''K''<sub>sp</sub>|[{{chem2|CO3(2−)}}]}}, and [{{chem2|CO3(2−)}}] = {{sfrac|''K''<sub>a2</sub> [{{chem2|HCO3−}}]|[{{chem2|H+}}]}}. Therefore, when {{chem2|HCO3−}} concentration is known, the maximum concentration of {{chem2|Ca(2+)}} ions before scaling through {{chem2|CaCO3}} precipitation can be predicted from the formula:

:[{{chem2|Ca(2+)}}]<sub>max</sub> = {{sfrac|''K''<sub>sp</sub>|''K''<sub>a2</sub>}} × {{sfrac|[{{chem2|H+}}]|[{{chem2|HCO3-}}]}}

The solubility product for {{chem2|CaCO3}} (''K''<sub>sp</sub>) and the dissociation constants for the dissolved inorganic carbon species (including ''K''<sub>a2</sub>) are all substantially affected by temperature and [[salinity]],<ref name="Mook 2000" /> with the overall effect that [{{chem2|Ca(2+)}}]<sub>max</sub> increases from freshwater to saltwater, and decreases with rising temperature, pH, or added bicarbonate level, as illustrated in the accompanying graphs.

The trends are illustrative for pool management, but whether scaling occurs also depends on other factors including interactions with {{chem2|[[magnesium|Mg]](2+)}}, [[Tetrahydroxyborate|{{chem2|[B(OH)4]−}}]] and other ions in the pool, as well as supersaturation effects.<ref>{{cite journal |last=Wojtowicz |first=J. A. |year=1998 |title=Factors affecting precipitation of calcium carbonate |journal=Journal of the Swimming Pool and Spa Industry |volume=3 |issue=1 |pages=18–23 |url=http://jspsi.poolhelp.com/ARTICLES/JSPSI_V3N1_pp18-23.pdf |access-date=18 March 2014 |archive-url=https://web.archive.org/web/20140318061934/http://jspsi.poolhelp.com/ARTICLES/JSPSI_V3N1_pp18-23.pdf |archive-date=18 March 2014}}</ref><ref>{{cite journal|last=Wojtowicz|first=J. A.|year=1998|title=Corrections, potential errors, and significance of the saturation index|journal=Journal of the Swimming Pool and Spa Industry|volume=3|issue=1|pages=37–40|url=http://jspsi.poolhelp.com/ARTICLES/JSPSI_V3N1_pp37-40.pdf|access-date=18 March 2014|archive-url=https://web.archive.org/web/20120824152654/http://jspsi.poolhelp.com/ARTICLES/JSPSI_V3N1_pp37-40.pdf|archive-date=24 August 2012}}</ref> Scaling is commonly observed in electrolytic chlorine generators, where there is a high pH near the cathode surface and scale deposition further increases temperature. This is one reason that some pool operators prefer borate over bicarbonate as the primary pH buffer, and avoid the use of pool chemicals containing calcium.<ref>{{cite web |last=Birch |first=R. G. |date=2013 |url=https://scithings.id.au/BABES.pdf |title=BABES: a better method than "BBB" for pools with a salt-water chlorine generator |website=scithings.id.au |access-date=11 October 2020 |archive-date=15 April 2021 |archive-url=https://web.archive.org/web/20210415065559/https://scithings.id.au/BABES.pdf |url-status=live}}</ref>

===Solubility in a strong or weak acid solution===
Solutions of [[strong acid|strong]] ([[hydrochloric acid|HCl]]), moderately strong ([[sulfamic acid|sulfamic]]) or [[weak acid|weak]] ([[acetic acid|acetic]], [[citric acid|citric]], [[sorbic acid|sorbic]], [[lactic acid|lactic]], [[phosphoric acid|phosphoric]]) acids are commercially available. They are commonly used as [[descaling agent]]s to remove [[limescale]] deposits. The maximum amount of {{chem2|CaCO3}} 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] = [{{chem2|A−}}], we obtain (with {{chem2|CaCO3}} molar mass = 100&nbsp;g/mol):
::{| class="wikitable"
! [A] (mol/L)
| 1
| 10<sup>−1</sup>
| 10<sup>−2</sup>
| 10<sup>−3</sup>
| 10<sup>−4</sup>
| 10<sup>−5</sup>
| 10<sup>−6</sup>
| 10<sup>−7</sup>
| 10<sup>−10</sup>
|-
! Initial pH
| 0.00||1.00||2.00||3.00||4.00||5.00||6.00||6.79||7.00
|-
!Final pH
| 6.75||7.25||7.75||8.14||8.25||8.26||8.26||8.26||8.27
|-
! Dissolved {{chem2|CaCO3}} (g/L 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 {{chem2|Ca(2+)}} (not taking into account possible {{chem2|CO2}} dissolution) and the final state is the solution with saturated {{chem2|Ca(2+)}}. For strong acid concentrations, all species have a negligible concentration in the final state with respect to {{chem2|Ca(2+)}} and {{chem2|A−}} so that the neutrality equation reduces approximately to 2[{{chem2|Ca(2+)}}] = [{{chem2|A−}}] yielding [{{chem2|Ca(2+)}}] ≈ 0.5 [{{chem2|A−}}]. When the concentration decreases, [{{chem2|HCO3−}}] becomes non-negligible so that the preceding expression is no longer valid. For vanishing acid concentrations, one can recover the final pH and the solubility of {{chem2|CaCO3}} in pure water.
* In the case of a weak monoacid (here we take acetic acid with [[Acid dissociation constant|p''K''<sub>a</sub>]] = 4.76) with decreasing total acid concentration [A] = [{{chem2|A−}}] + [AH], we obtain:

::{| class="wikitable"
![A] (mol/L)
| [{{chem2|Ca(2+)}}] ≈ 0.5 [{{chem2|A−}}]<br />{{clarify|date=May 2024}}
| 10<sup>−1</sup>
| 10<sup>−2</sup>
| 10<sup>−3</sup>
| 10<sup>−4</sup>
| 10<sup>−5</sup>
| 10<sup>−6</sup>
| 10<sup>−7</sup>
| 10<sup>−10</sup>
|-
!Initial pH
| 2.38||2.88||3.39||3.91||4.47||5.15||6.02||6.79||7.00
|-
!Final pH
| 6.75||7.25||7.75||8.14||8.25||8.26||8.26||8.26||8.27
|-
!|Dissolved {{chem2|CaCO3}} (g/L of acid)
| 49.5||4.99||0.513||0.0848||0.0504||0.0474||0.0471||0.0470||0.0470
|}
:For the same total acid concentration, the initial pH of the weak acid is less acid than that of the strong acid; however, the maximum amount of {{chem2|CaCO3}} which can be dissolved is approximately the same. This is because in the final state, the pH is larger than the p''K''<sub>a</sub>, so that the weak acid is almost completely dissociated, yielding in the end as many {{chem2|H+}} 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 [{{chem2|HCO3−}}], [{{chem2|CO3(2−)}}], [{{chem2|Ca(2+)}}], [{{chem2|H+}}] and [{{chem2|OH−}}]. The system may be reduced to a seventh degree equation for [{{chem2|H+}}] the numerical solution of which gives

::{| class="wikitable"
! [A] (mol/L)
| 1
| 10<sup>−1</sup>
| 10<sup>−2</sup>
| 10<sup>−3</sup>
| 10<sup>−4</sup>
| 10<sup>−5</sup>
| 10<sup>−6</sup>
| 10<sup>−7</sup>
| 10<sup>−10</sup>
|-
! Initial pH
| 1.08||1.62||2.25||3.05||4.01||5.00||5.97||6.74||7.00
|-
! Final pH
| 6.71||7.17||7.63||8.06||8.24||8.26||8.26||8.26||8.27
|-
! Dissolved {{chem2|CaCO3}} (g/L of acid)
| 62.0||7.39||0.874||0.123||0.0536||0.0477||0.0471||0.0471||0.0470
|}

:where [A] = {{chem2|[H3PO4] + [H2PO4−] + [HPO4(2−)] + [PO4(3−)]}} is the total acid concentration. Thus phosphoric acid is more efficient than a monoacid since at the final almost neutral pH, the second dissociated state concentration [{{chem2|HPO4(2−)}}] is not negligible (see [[phosphoric acid#pH and composition of a phosphoric acid aqueous solution|phosphoric acid]]).

== See also ==
[[File:Vodní kámen.jpg|thumb|Electron micrograph of needle-like calcium carbonate crystals formed as limescale in a kettle]]
[[File:Calcium-48 carbonate.png|thumb|Around 2&nbsp;g of calcium-48 carbonate]]
{{div col|colwidth=22em}}
* [[Cuttlebone]]
* [[Cuttlefish]]
* [[Gesso]]
* [[Limescale]]
* [[Marble]]
* [[Ocean acidification]]
* [[Whiting event]]
* [[List of climate engineering topics]]
* [[Lysocline]]{{div col end}}

==References==
{{reflist}}

==External links==
* {{ICSC|1193|11}}
* {{PubChemLink|516889}}
* [[ATC codes]]: {{ATC|A02|AC01}} and {{ATC|A12|AA04}}
* [http://calcium-carbonate.org.uk/calcium-carbonate.asp The British Calcium Carbonate Association – What is calcium carbonate] {{Webarchive|url=https://web.archive.org/web/20080524102528/http://www.calcium-carbonate.org.uk/calcium-carbonate.asp |date=24 May 2008}}
* [https://www.cdc.gov/niosh/npg/npgd0090.html CDC – NIOSH Pocket Guide to Chemical Hazards – Calcium Carbonate]

{{Calcium compounds}}
{{Carbonates}}
{{Antacids}}
{{Drugs for treatment of hyperkalemia and hyperphosphatemia}}
{{Authority control}}

[[Category:Calcium compounds]]
[[Category:Carbonates]]
[[Category:Limestone]]
[[Category:Phosphate binders]]
[[Category:Excipients]]
[[Category:Antacids]]
[[Category:Food stabilizers]]
[[Category:E-number additives]]
[[Category:Over-the-counter drugs in the United States]]