Calcium carbonate#Industrial applications

Calcium carbonate shares the typical properties of other carbonates. Notably, it:

• reacts with acids, 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 °C in the case of {{chem2|CaCO3}}), to form calcium oxide, CaO, commonly called quicklime, with reaction enthalpy 178 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 caverns, 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 °C.

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:{{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}}

:{{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}}.{{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}} 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.

The thermodynamically stable form of {{chem2|CaCO3}} under normal conditions is hexagonal β-{{chem2|CaCO3}} (the mineral calcite). Other forms can be prepared, the denser (2.83 g/cm³) 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 °C; the vaterite form can be prepared by precipitation at 60 °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.{{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}} 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 ionic radii.{{citation needed|date=September 2023}}

Calcium carbonate crystallizes in three anhydrous polymorphs,{{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}}{{Cite book |last=Lippmann|first= Friedrich |title=Sedimentary carbonate minerals |date=1973 |publisher=Springer |isbn=3-540-06011-1 |oclc=715109304}} of which calcite is the thermodynamically most stable at room temperature, aragonite is only slightly less so, and vaterite is the least stable.{{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 |page=2203759 |doi=10.1002/advs.202203759 |pmid=36403251 |pmc=9811428 |s2cid=253707446 |issn=2198-3844}}

The calcite crystal structure is trigonal, with space group R{{overline|3}}c (No. 167 in the International Tables for Crystallography{{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}}), and Pearson symbol hR10.{{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}} Aragonite is orthorhombic, with space group Pmcn (No 62), and Pearson Symbol oP20.{{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}} Vaterite is composed of at least two different coexisting crystallographic structures. The major structure exhibits hexagonal symmetry in space group P6₃/mmc, the minor structure is still unknown.{{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}}

File:Calcite+Aragonite.png

All three polymorphs crystallize simultaneously from aqueous solutions under ambient conditions. 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.{{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 |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}} 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.{{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 |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}} 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.

File:Calcite+Vaterite.png

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,{{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 |url=https://onlinelibrary.wiley.com/doi/10.1002/adfm.202201394 |journal=Advanced Functional Materials |language=en |volume=32 |issue=25 |page=2201394 |doi=10.1002/adfm.202201394 |s2cid=247587883 |issn=1616-301X |access-date=15 December 2022 |archive-date=15 December 2022 |archive-url=https://web.archive.org/web/20221215200840/https://onlinelibrary.wiley.com/doi/10.1002/adfm.202201394 |url-status=live}} or by using proteins and peptides derived from biological calcium carbonate.{{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 |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}} Some polyamines such as cadaverine and Poly(ethylene imine) have been shown to facilitate the formation of aragonite over calcite.

Organisms, such as molluscs and arthropods, have shown the ability to grow all three crystal polymorphs of calcium carbonate, mainly as protection (shells) and muscle attachments.{{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}} 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.{{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}}{{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}}{{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 |page=107583 |doi=10.1016/j.jsb.2020.107583|pmid=32721585 |s2cid=220850117|doi-access=free}}

File:Silfurberg.jpg is the most stable 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.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]]

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

File:Calcium carbonate chunks.JPGshell]]

Eggshells, snail shells and most seashells are predominantly calcium carbonate and can be used as industrial sources of that chemical.{{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}} Oyster shells have enjoyed recent recognition as a source of dietary calcium, but are also a practical industrial source.{{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}}{{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}} Dark green vegetables such as broccoli and kale contain dietarily significant amounts of calcium carbonate, but they are not practical as an industrial source.{{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}}

Annelids 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 into calcium carbonate, which is later excreted into the dirt.{{Cite web |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 |website=The company of biologists |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}} The function of these glands is unknown but is believed to serve as a {{chem2|CO2}} regulation mechanism within the animals' tissues.{{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 }} This process is ecologically significant, stabilizing the pH of acid soils.{{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}}

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 and Huygens craters). This provides some evidence for the past presence of liquid water.{{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}}{{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}}

File:Rubaksa tufa plug.jpg in Rubaksa, Ethiopia]]

Carbonate is found frequently in geologic settings and constitutes an enormous carbon reservoir. Calcium carbonate occurs as aragonite, calcite and dolomite as significant constituents of the calcium cycle. The carbonate minerals form the rock types: limestone, chalk, marble, travertine, tufa, and others.

File:Water-of-Five-colored-Pond Huanglong Sichuan China.jpg at Huanglong, Sichuan]]

In warm, clear tropical waters corals are more abundant than towards the poles where the waters are cold. Calcium carbonate contributors, including plankton (such as coccoliths and planktic foraminifera), coralline algae, sponges, brachiopods, echinoderms, bryozoa and 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 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 volcanic eruptions.

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.{{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}} 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.{{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 |pages=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}}

Calcium carbonate can 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. 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.{{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}}

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,{{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}} which had purely chitinous shells.

The main use of calcium carbonate is in the construction industry, either as a building material, or limestone aggregate for road building, as an ingredient of cement, or as the starting material for the preparation of builders' lime by burning in a kiln. However, because of weathering mainly caused by acid rain,{{cite web|title = Effects of Acid Rain|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|url-status = live}} 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 calcined in situ to give calcium oxide, which forms a slag with various impurities present, and separates from the purified iron.{{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}}

In the oil industry, calcium carbonate is added to drilling fluids 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 pools, as a pH corrector for maintaining alkalinity and offsetting the acidic properties of the disinfectant agent.{{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}}

It is also used as a raw material in the 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 for the precipitation of impurities in raw juice during carbonatation.{{cite book|last1=McGinnis|first1=R. A.|title=Beet-Sugar Technology|publisher=Beet Sugar Development Foundation|page=178|edition=2nd}}

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.{{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}}

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 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 in paints,{{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}} in particular matte emulsion paint where typically 30% by weight of the paint is either chalk or marble. It is also a popular filler in plastics. Some typical examples include around 15–20% loading of chalk in unplasticized polyvinyl chloride (uPVC) drainpipes, 5–15% loading of stearate-coated chalk or marble in uPVC window profile. PVC cables can use calcium carbonate at loadings of up to 70 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.{{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}} Here the percentage is often 20–40%. It also routinely used as a filler in thermosetting resins (sheet and bulk molding compounds) and has also been mixed with ABS, and other ingredients, to form some types of compression molded "clay" poker chips.{{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}} Precipitated calcium carbonate, made by dropping calcium oxide into water, is used by itself or with additives as a white paint, known as whitewashing.{{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}}{{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}}

Calcium carbonate is added to a wide range of trade and do it yourself adhesives, sealants, and decorating fillers. 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.{{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}}{{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}}{{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}}{{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}}

In ceramic glaze applications, calcium carbonate is known as whiting, 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 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, and will therefore not scratch glass and most other ceramics, 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.{{cite web|title = Ohio Historical Society Blog: Make It Shine|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}}

File:500 mg calcium supplements with vitamin D.jpg

Calcium carbonate is widely used medicinally as an inexpensive dietary calcium supplement or gastric antacid{{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.nlm.nih.gov/medlineplus/druginfo/medmaster/a601032.html |archive-url = https://web.archive.org/web/20071017031324/http://www.nlm.nih.gov/medlineplus/druginfo/medmaster/a601032.html |archive-date = 17 October 2007}} (such as Tums and 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 filler for tablets and other pharmaceuticals.{{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}}

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.{{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}}

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 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 grams daily, for prevention and treatment of osteoporosis,{{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}}{{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}} 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.{{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}}

As a food additive it is designated E170,{{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 and it has an INS number of 170. Used as an acidity regulator, anticaking agent, stabilizer or color it is approved for usage in the EU,{{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}} US{{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}} and Australia and New Zealand.{{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}} It is "added by law to all UK milled bread flour except wholemeal".{{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}}{{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}} 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.{{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}} 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,{{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}} posing a public health concern.{{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}} Lead is commonly found in natural sources of calcium.

Agricultural lime, powdered chalk or limestone, is used as a cheap method of neutralising acidic soil, making it suitable for planting, also used in aquaculture industry for pH regulation of pond soil before initiating culture.{{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}} There is interest in understanding whether or not it can affect pesticide adsorption and desorption in calcareous soil.{{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}}

Calcium carbonate is a key ingredient in many household cleaning powders like Comet and is used as a scrubbing agent.

In 1989, a researcher, Ken Simmons, introduced {{chem2|CaCO3}} into the Whetstone Brook in Massachusetts.{{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}} 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.{{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/|url-status=dead}}{{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}} 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.{{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}}

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.

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.{{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= |date= |website= |publisher=EuroPlas |access-date=12 July 2024 |quote=}}

Calcination of limestone using charcoal fires to produce 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 {{chem2|CO2}} pressure is only a tiny fraction of the partial {{chem2|CO2}} pressure in air, which is about 0.035 kPa.

At temperatures above 550 °C the equilibrium {{chem2|CO2}} pressure begins to exceed the {{chem2|CO2}} pressure in air. So above 550 °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 kPa.{{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}}

The table shows that this partial pressure is not achieved until the temperature is nearly 800 °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 kPa, which happens at 898 °C.{{clear right}}

:

File:CanarySpring.jpg calcium carbonate deposits from a hot spring]]

Calcium carbonate is poorly soluble in pure water (47 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):

:

where the solubility product for {{chem2|[Ca(2+)][CO3(2−)]}} is given as anywhere from K_sp = {{val|3.7e-9}} to K_sp = {{val|8.7e-9}} at 25 °C, depending upon the data source.{{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}} What the equation means is that the product of molar concentration of calcium ions (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_sp. 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

:

{{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

:

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

:

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

:

For ambient air, P_{{{chem2|CO2}}} is around {{val|3.5e-4}} atm (or equivalently 35 Pa). The last equation above fixes the concentration of dissolved {{chem2|CO2}} as a function of P_{{{chem2|CO2}}}, 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

:

(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_sp = {{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 mg/L.
• As ambient {{chem2|CO2}} partial pressure is reduced below atmospheric levels, the solution becomes more and more alkaline. At extremely low P_{{{chem2|CO2}}}, 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_{{{chem2|CO2}}} = 10⁻¹² 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 stalagmites 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.

File:CaCO3-pH.gif

File:CaCO3-Temp.gif

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 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|H2CO3 ⇌ CO2(aq) + H2O}};{{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}}

|limited aeration in a deep water column; and

|periodic replenishment of bicarbonate to maintain buffer capacity (often estimated through measurement of 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 7 to pH 8 at 25 °C in fresh water).{{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}} 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_sp|[{{chem2|CO3(2−)}}]}}, and [{{chem2|CO3(2−)}}] = {{sfrac|K_a2 [{{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+)}}]_max = {{sfrac|K_sp|K_a2}} × {{sfrac|[{{chem2|H+}}]|[{{chem2|HCO3-}}]}}

The solubility product for {{chem2|CaCO3}} (K_sp) and the dissociation constants for the dissolved inorganic carbon species (including K_a2) are all substantially affected by temperature and salinity, with the overall effect that [{{chem2|Ca(2+)}}]_max 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|Mg(2+)}}, Tetrahydroxyborate and other ions in the pool, as well as supersaturation effects.{{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}}{{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}} 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.{{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}}

Solutions of strong (HCl), moderately strong (sulfamic) or weak (acetic, citric, sorbic, lactic, phosphoric) acids are commercially available. They are commonly used as descaling agents 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 g/mol):

::

: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 pK_a = 4.76) with decreasing total acid concentration [A] = [{{chem2|A−}}] + [AH], we obtain:

::

:For the same total acid concentration, the initial pH of the weak acid is less acid than the one 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 pK_a, 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

::

: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).

File:Vodní kámen.jpg

File:Calcium-48 carbonate.png

{{div col|colwidth=22em}}

• Cuttlebone
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{{reflist}}

• {{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]

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