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| {{chembox
| | My name: Seth Maitland<br>My age: 34<br>Country: Netherlands<br>Home town: Hoogwoud <br>Post code: 1718 Bk<br>Street: Burgemeester Hoogenboomlaan 175<br><br>My weblog: [http://salejobsinmyanmar.wordpress.com/ http://salejobsinmyanmar.wordpress.com/] |
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| | Watchedfields = changed
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| | verifiedrevid = 477003420
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| |Name = Calcium carbonate
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| |ImageFileL1 = calcium carbonate.png
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| |ImageSizeL1 = 120px
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| |ImageFileR1 = Calcium-carbonate-xtal-3D-vdW.png
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| |ImageSizeR1 = 150px
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| |ImageFile2 = Calcium carbonate.jpg
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| |ImageSize2 = 150px
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| |IUPACName = Calcium carbonate
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| |OtherNames = [[Limestone]]; [[calcite]]; [[aragonite]]; [[chalk]]; [[marble]]; [[pearl]]; [[oyster]]
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| |Section1 = {{Chembox Identifiers
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| | UNII_Ref = {{fdacite|correct|FDA}}
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| | UNII = H0G9379FGK
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| | ChEMBL_Ref = {{ebicite|changed|EBI}}
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| | ChEMBL = 1200539
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| | KEGG_Ref = {{keggcite|correct|kegg}}
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| | KEGG = D00932
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| | InChI = 1/CH2O3.Ca/c2-1(3)4;/h(H2,2,3,4);/q;+2/p-2
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| | ChEBI_Ref = {{ebicite|correct|EBI}}
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| | ChEBI = 3311
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| | SMILES = [Ca+2].[O-]C([O-])=O
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| | InChIKey = VTYYLEPIZMXCLO-NUQVWONBAS
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| | SMILES1 = C(=O)([O-])[O-].[Ca+2]
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| | StdInChI_Ref = {{stdinchicite|correct|chemspider}}
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| | StdInChI = 1S/CH2O3.Ca/c2-1(3)4;/h(H2,2,3,4);/q;+2/p-2
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| | StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
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| | StdInChIKey = VTYYLEPIZMXCLO-UHFFFAOYSA-L
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| | CASNo = 471-34-1
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| | CASNo_Ref = {{cascite|correct|CAS}}
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| | ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
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| | ChemSpiderID = 9708
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| | EINECS = 207-439-9
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| | PubChem = 10112
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| | RTECS = FF9335000
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| | ATCCode_prefix = A02
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| | ATCCode_suffix = AC01
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| | ATC_Supplemental = {{ATC|A12|AA04}}
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| }}
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| |Section2 = {{Chembox Properties
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| |Formula = CaCO<sub>3</sub>
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| |MolarMass = 100.0869 g/mol
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| |Appearance = Fine white powder; chalky taste
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| |Odor = odorless
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| |Density = 2.711 g/cm<sup>3</sup> ([[calcite]])<br/>2.83 g/cm<sup>3</sup> ([[aragonite]])
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| |Solubility = 0.0013 g/100 mL (25°C)<ref>{{cite book|title=SI Chemical Data Book (4th ed.) |publisher=John Wiley & Sons Australia, Ltd. |author= Gordon Aylward, Tristan Findlay|isbn=978-0-470-81638-7}}</ref>
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| |SolubilityProduct = 4.8{{e|-9}}<ref>{{cite book|last =Patnaik|first=Pradyot|year=2003|title=Handbook of Inorganic Chemical Compounds |publisher=McGraw-Hill|isbn =0-07-049439-8|url=http://books.google.com/?id=Xqj-TTzkvTEC|accessdate=2009-06-06}}</ref>
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| |Solvent = dilute acids
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| |SolubleOther = soluble
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| |MeltingPt = 825 °C (aragonite) <br> 1339 °C (calcite)<ref>{{cite web|url=http://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|accessdate=31 March 2011}}</ref>
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| |BoilingPt = decomposes
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| |RefractIndex = 1.59
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| |pKa = 9.0
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| |pKb =
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| }}
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| |Section3 = {{Chembox Structure
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| |CrystalStruct = Trigonal
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| |SpaceGroup = <span style="text-decoration: overline">3</span>2/m
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| }}
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| | Section4 = {{Chembox Thermochemistry
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| | DeltaHf = −1207 kJ·mol<sup>−1</sup><ref name=b1>{{cite book| author = Zumdahl, Steven S.|title =Chemical Principles 6th Ed.| publisher = Houghton Mifflin Company| year = 2009| isbn = 0-618-94690-X|page=A21}}</ref>
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| | Entropy = 93 J·mol<sup>−1</sup>·K<sup>−1</sup><ref name=b1 />
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| }}
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| |Section7 = {{Chembox Hazards
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| |ExternalMSDS = [http://www.inchem.org/documents/icsc/icsc/eics1193.htm ICSC 1193]
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| |EUIndex = Not listed
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| |MainHazards =
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| |NFPA-H = 1
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| |NFPA-F = 0
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| |NFPA-R = 0
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| |NFPA-O =
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| |RPhrases =
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| |SPhrases =
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| |FlashPtC = 825
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| |LD50 = 6450 mg/kg (oral, rat)
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| }}
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| |Section8 = {{Chembox Related
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| |OtherAnions = [[Calcium bicarbonate]]
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| |OtherCations = [[Magnesium carbonate]]<br/>[[Strontium carbonate]]<br/>[[Barium carbonate]]
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| |OtherCpds = [[Calcium sulfate]]
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| }}
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| }}
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| [[File:Calcite.GIF|thumb|right|Crystal structure of calcite]]
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| '''Calcium carbonate''' is a [[chemical compound]] with the [[chemical formula|formula]] [[Calcium|Ca]][[Carbon|C]][[Oxygen|O]]<sub>3</sub>. It is a common substance found in [[Rock (geology)|rock]]s in all parts of the world, and is the main component of [[seashells|shells of marine organisms]], [[snail]]s, [[coal ball]]s, [[pearls]], and [[eggshell]]s. Calcium carbonate is the active ingredient in [[agricultural lime]], and is created when Ca ions in [[hard water]] react with carbonate ions creating [[limescale]]. It is commonly used medicinally as a [[calcium]] supplement or as an [[antacid]], but excessive consumption can be hazardous.
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| ==Chemistry==
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| {{see also|Carbonate}}
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| Calcium carbonate shares the typical properties of other carbonates. Notably:
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| *it reacts with strong acids, releasing [[carbon dioxide]]:
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| :CaCO<sub>3(s)</sub> + 2 HCl<sub>(aq)</sub> → CaCl<sub>2(aq)</sub> + CO<sub>2(g)</sub> + H<sub>2</sub>O<sub>(l)</sub>
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| *it releases carbon dioxide on heating, called a [[thermal decomposition]] reaction, or [[calcination]], (to above 840 °C in the case of CaCO<sub>3</sub>), to form [[calcium oxide]], commonly called [[quicklime]], with reaction [[enthalpy]] 178 kJ / mole:
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| :CaCO<sub>3(s)</sub> → CaO<sub>(s)</sub> + CO<sub>2(g)</sub>
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| Calcium carbonate will react with water that is saturated with carbon dioxide to form the soluble [[calcium bicarbonate]].
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| :CaCO<sub>3</sub> + CO<sub>2</sub> + H<sub>2</sub>O → Ca(HCO<sub>3</sub>)<sub>2</sub>
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| This reaction is important in the [[erosion]] of [[carbonate rock]]s, forming [[cavern]]s, and leads to hard water in many regions.
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| An unusual form of calcium carbonate is [[ikaite]] with crystal water, CaCO<sub>3</sub>·6H<sub>2</sub>O. Ikaite is stable only below 6 °C.
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| ==Preparation==
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| The vast majority of calcium carbonate used in industry is extracted by mining or quarrying. Pure calcium carbonate (e.g. for food or pharmaceutical use), can be produced from a pure quarried source (usually marble).
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| Alternatively, calcium carbonate is prepared from [[calcium oxide]]. Water is added to give [[calcium hydroxide]], and [[carbon dioxide]] is passed through this solution to precipitate the desired calcium carbonate, referred to in the industry as precipitated calcium carbonate (PCC):<ref name="PCC">{{cite web|title = Precipitated Calcium Carbonate |accessdate = 2014-01-11|url = http://www.lime.org/uses_of_lime/other_uses/precip_cc.asp}}</ref>
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| : CaCO<sub>3</sub> → CaO + CO<sub>2</sub>
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| : CaO + H<sub>2</sub>O → Ca(OH)<sub>2</sub>
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| :Ca(OH)<sub>2</sub> + CO<sub>2</sub> → CaCO<sub>3</sub> + H<sub>2</sub>O
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| ==Occurrence==
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| ===Geological sources===
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| [[Calcite]], [[aragonite]] and [[vaterite]] are pure calcium carbonate minerals. Industrially important source rocks which are predominantly calcium carbonate include [[limestone]], [[chalk]], [[marble]] and [[travertine]].
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| ===Biological sources===
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| Eggshells, snail shells and most seashells are predominantly calcium carbonate and can be used as industrial sources of that chemical.
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| <ref>{{cite web |title=How are seashells created? |author=Francis Horne |date=23 October 2006 |work=Scientific American |accessdate=25 April 2012 |url=http://www.scientificamerican.com/article.cfm?id=how-are-seashells-created}}</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=WebMD: Oyster shell calcium |publisher=WebMD| accessdate=25 April 2012}}</ref><ref>{{cite web |title=Oyster Shell Calcium Carbonate|publisher=Caltron Clays & Chemicals|url=http://caltronclays.in/Oyster_CC.html}}</ref> While not practical as an industrial source, dark green vegetables such as Broccoli and Kale contain dietarily significant amounts of calcium carbonate.<ref>{{cite journal|year=1993 |title=Absorbability of Calcium from Brassica Vegetables: Broccoli, Bok Choy, and Kale |journal=Journal of Food Science |volume=58 |issue=6 |pages=1378–1380|doi=10.1111/j.1365-2621.1993.tb06187.x|last1=Heaney|first1=R.P.|last2=Weaver|first2=C.M.|last3=Hinders|first3=SM.|last4=Martin|first4=B.|last5=Packard|first5=P.T.}}</ref>
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| ==Geology==
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| Carbonate is found frequently in geologic settings and constitute an enormous [[carbon cycle|carbon reservoir]]. Calcium carbonate occurs as [[aragonite]] and [[calcite]]. The carbonate minerals form the rock types: [[limestone]], [[chalk]], [[marble]], [[travertine]], [[tufa]], and others. Purity of Calcium Carbonate has been found in some mines to be 99.2% pure.<ref>{{cite web|title=CongCal, McGraths Limestone Works Ltd, Cong, Co. Mayo, Ireland|url=http://www.congcal.com/|work=congcal.com|accessdate=5 August 2013}}</ref>
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| In tropic settings, the waters are warm and clear. [[Coral]]s are more abundant in this environment 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|sponge]]s, [[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 the [[ocean acidification]].
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| Where the oceanic crust is subducted under a [[continental plate]] sediments will be carried down to warmer zones in the astenosphere and [[mesosphere]] where the calcium carbonate is decomposed to [[carbon dioxide]] which will give rise to explosive [[volcano|volcanic eruptions]].
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| ===Carbonate compensation depth===
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| 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. Calcium carbonate is unusual in that its solubility increases with decreasing temperature. Increasing pressure also increases the solubility of calcium carbonate. The CCD can range from 4–6 km below sea level.
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| ===Taphonomy===
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| 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 [[CaCO3|CaCO<sub>3</sub>]] [[permineralization]].<ref name="twoturn" /> This type of preservation preserves high levels of detail, even down to the microscopic level.<ref name="twoturn" /> However, it also leaves specimens vulnerable to [[weathering]] when exposed to the surface.<ref name="twoturn">Trexler, D., 2001, [http://books.google.com/books?id=mgc6CS4EUPsC&pg=PA98 Two Medicine Formation, Montana: geology and fauna]: In: Mesozoic Vertebrate Life, edited by Tanke, D. H., and Carpenter, K., Indiana University Press, pp. 298–309. ISBN 0-253-33907-3</ref>
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| ==Uses==
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| ===Industrial applications===
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| The main use of calcium carbonate is in the construction industry, either as a building material or limestone aggregate for roadbuilding or as an ingredient of cement or as the starting material for the preparation of builder's lime by burning in a kiln. However, due to weathering mainly caused by [[acid rain]], calcium carbonate (in limestone form) is no longer used for building purposes on its own, and only as a raw/primary substance for building materials.
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| 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.<ref>{{cite web|title = Blast Furnace|publisher = Science Aid|accessdate = 2007-12-30|url = http://www.scienceaid.co.uk/chemistry/industrial/blastfurnace.html}}</ref>
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| 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 pools, as a [[pH]] corrector for maintaining [[alkalinity]] and offsetting the acidic properties of the disinfectant agent.
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| 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 + carbon dioxide. This burnt lime is then slaked in sweet water to produce a calcium hydroxide suspension for the precipitation of impurities in raw juice during [[carbonatation]].
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| Calcium carbonate has traditionally been a major component of blackboard chalk. However, modern manufactured chalk is mostly [[gypsum]], hydrated [[calcium sulfate]] CaSO<sub>4</sub>·2H<sub>2</sub>O. Calcium carbonate is a main source for growing [[Seacrete]], or [[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}}</ref>{{Dead link|date=March 2013}}
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| Fine ground calcium carbonate (GCC) is an essential ingredient in the microporous film used in babies' [[diapers]] and some building films as the pores are nucleated around the calcium carbonate particles during the manufacture of the film by biaxial stretching. GCC or PCC is used as a filler in paper because they are cheaper than wood fiber. In terms of market volume, GCC are the most important types of fillers currently used.<ref>[http://www.ceresana.com/en/market-studies/additives/fillers/ Market Study Fillers, 2nd ed., published by Ceresana, September 2011]</ref> 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 has a very fine and controlled particle size, on the order of 2 micrometres in diameter, useful in [[Coated paper|coatings for paper]].
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| Calcium carbonate is widely used as an extender in paints,<ref name = reade>{{cite web|title = Calcium Carbonate Powder|publisher = Reade Advanced Materials |date=2006-02-04|accessdate = 2007-12-30|url = http://www.reade.com/Products/Minerals_and_Ores/calcium_carbonate.html}}</ref> in particular matte emulsion paint where typically 30% by weight of the paint is either chalk or marble. It is also a popular filler in plastics.<ref name = reade/> Some typical examples include around 15 to 20% loading of chalk in [[Polyvinyl chloride|unplasticized polyvinyl chloride]] (uPVC) drain pipe, 5 to 15% loading of stearate coated chalk or marble in uPVC window profile. [[Polyvinyl chloride|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). [[Polypropylene]] compounds are often filled with calcium carbonate to increase rigidity, a requirement that becomes important at high use 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 |accessdate=2008-08-01 |publisher=Imerys Performance Minerals}}</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 chips. Precipitated calcium carbonate, made by dropping [[calcium oxide]] into water, is used by itself or with additives as a white paint, known as [[whitewashing]].
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| 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.
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| In [[ceramics (art)|ceramics]]/glazing 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]], 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|publisher = Ohio Historical Society |url = http://ohiohistory.wordpress.com/2011/06/02/making-it-shine/}}</ref>
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| ===Health and dietary applications===
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| [[File:500 mg calcium supplements with vitamin D.jpg|thumb|500-milligram calcium supplements made from calcium carbonate]]
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| Calcium carbonate is widely used medicinally as an inexpensive dietary calcium supplement or [[antacid|gastric antacid]].<ref name = medline>{{cite web|publisher = [[National Institutes of Health]]|work = Medline Plus|title = Calcium Carbonate |date=2005-10-01|accessdate = 2007-12-30|url = http://www.nlm.nih.gov/medlineplus/druginfo/medmaster/a601032.html |archiveurl = http://web.archive.org/web/20071017031324/http://www.nlm.nih.gov/medlineplus/druginfo/medmaster/a601032.html <!-- Bot retrieved archive --> |archivedate = 2007-10-17}}</ref> It may be used as a [[phosphate binder]] for the treatment of [[hyperphosphatemia]] (primarily in patients with [[chronic renal failure]]). It is also used in the pharmaceutical industry as an inert [[Excipients#Fillers and diluents|filler]] for [[Tablet (pharmacy)|tablets]] and other [[pharmaceuticals]].<ref>{{cite book|author = Herbert A. Lieberman, Leon Lachman, Joseph B. Schwartz|title = Pharmaceutical Dosage Forms: Tablets|year = 1990|isbn = 0-8247-8044-2|page=153|publisher = Dekker|location = New York}}</ref>
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| Calcium carbonate is known among [[Irritable Bowel Syndrome|IBS]] sufferers to help reduce diarrhea{{Citation needed|date=May 2012}}. Some individuals report being symptom-free since starting supplementation. The process in which calcium carbonate reduces diarrhea is by binding water in the bowel, which creates a stool that is firmer and better formed. Calcium carbonate supplements are often combined with magnesium in various proportions. This should be taken into account as magnesium is known to cause [[diarrhea]].
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| Calcium carbonate is used in the production of toothpaste and has seen a resurgence as a food preservative and color retainer, when used in or with products such as organic apples or food.<ref>[http://chemistry.about.com/od/foodcookingchemistry/a/cadditives.htm Food Additives – Names Starting with C]. Chemistry.about.com (2012-04-10). Retrieved on 2012-05-24.</ref>
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| Excess calcium from supplements, fortified food and high-calcium diets, can cause the [[milk-alkali syndrome]], which has serious toxicity and can be fatal. In 1915, Bertram Sippy introduced the "Sippy regimen" of hourly ingestion of milk and cream, and the gradual addition of eggs and cooked cereal, for 10 days, combined with alkaline powders, which provided symptomatic relief for peptic ulcer disease. Over the next several decades, the Sippy regimen resulted in [[renal failure]], [[alkalosis]], and [[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. During the past 15 years, it has been reported in women taking calcium supplements above the recommended range of 1.2 to 1.5 g daily, for prevention and treatment of osteoporosis, and is exacerbated by [[dehydration]]. Calcium has been added to over-the-counter products, which contributes to inadvertent excessive intake. Excessive calcium intake can lead to [[hypercalcemia]], complications of which include vomiting, abdominal pain and altered mental status.<ref>{{cite journal|title=Clinical problem-solving, back to basics|author=Ilan Gabriely|journal=New England Journal of Medicine|year=2008|volume=358|pmid=18450607|doi=10.1056/NEJMcps0706188|issue=18|author-separator=,|display-authors=1|last2=Leu|first2=James P.|last3=Barzel|first3=Uriel S.|pages=1952–6}}</ref>
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| As a [[food additive]] it is designated E170;<ref>{{cite web|title=Food-Info.net : E-numbers : E170 Calcium carbonate|url=http://www.food-info.net/uk/e/e170.htm}} 080419 food-info.net</ref> INS number 170. Used as an acidity regulator, anticaking agent, stabiliser or colour it is approved for usage in the EU,<ref>UK Food Standards Agency: {{cite web |url=http://www.food.gov.uk/safereating/chemsafe/additivesbranch/enumberlist |title=Current EU approved additives and their E Numbers |accessdate=2011-10-27}}</ref> USA<ref>US Food and Drug Administration: {{cite web |url=http://www.fda.gov/Food/FoodIngredientsPackaging/FoodAdditives/FoodAdditiveListings/ucm091048.htm |title=Listing of Food Additives Status Part I |accessdate=2011-10-27}}</ref> and [[Australia]] and [[New Zealand]].<ref>Australia New Zealand Food Standards Code{{cite web |url=http://www.comlaw.gov.au/Details/F2011C00827 |title=Standard 1.2.4 – Labelling of ingredients |accessdate=2011-10-27}}</ref> It is used in some [[soy milk]] and [[almond milk]] products as a source of dietary calcium; one study suggests that calcium carbonate might be as [[bioavailable]] as the calcium in cow's milk.<ref>{{cite journal|author = Y. Zhao, B. R. Martin and C. M. Weaver|title = Calcium Bioavailability of Calcium Carbonate Fortified Soymilk Is Equivalent to Cow's Milk in Young Women|year = 2005|journal = J. Nutr.|volume = 135|issue = 10|pages = 2379–2382|pmid = 16177199}}</ref> Calcium carbonate is also used as a [[firming agent]] in many canned or bottled vegetable products.
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| ===Environmental applications===
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| In 1989, a researcher, Ken Simmons, introduced CaCO<sub>3</sub> into the Whetstone Brook in [[Massachusetts]].<ref>{{cite news|author = [[Associated Press]]|title =
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| Limestone Dispenser Fights Acid Rain in Stream |date=1989-06-13|url = http://query.nytimes.com/gst/fullpage.html?res=950DEFD9173FF930A25755C0A96F948260|publisher = New York Times}}</ref> His hope was that the calcium carbonate would counter the acid in the stream from acid rain and save the trout that had ceased to spawn. Although his experiment was a success, it did increase the amount of aluminium ions in the area of the brook that was not treated with the limestone. This shows that CaCO<sub>3</sub> can be added to neutralize the effects of acid rain in [[river]] ecosystems. Currently calcium carbonate is used to neutralize acidic conditions in both soil and water.<ref>{{cite web|title=Environmental Uses for Calcium Carbonate|url=http://www.congcal.com/markets/environmental/|work=http://www.congcal.com/markets/environmental/|publisher=Congcal|accessdate=5 August 2013}}</ref><ref>{{cite journal|author = R. K. Schreiber|title = Cooperative federal-state liming research on surface waters impacted by acidic deposition|year = 1988|journal =Water, Air, & Soil Pollution|volume = 41|issue = 1|pages = 53–73|doi=10.1007/BF00160344}}</ref><ref>{{cite web|title = Effects of low pH and high aluminum on Atlantic salmon smolts in Eastern Maine and liming project feasibility analysis|year = 2006|author = Dan Kircheis; Richard Dill|publisher = National Marine Fisheries Service and Maine Atlantic Salmon Commission|url = http://www.mainesalmonrivers.org/pages/Liming%20Project%20Rpt.pdf|format = reprinted at Downeast Salmon Federation}}</ref> 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>{{cite journal|author = M. Guhren, C. Bigler and I. Renberg|title = Liming placed in a long-term perspective: A paleolimnological study of 12 lakes in the Swedish liming program|year = 2007|journal =Journal of Paleolimnology|volume = 37|pages = 247–258|doi=10.1007/s10933-006-9014-9|issue = 2}}</ref>
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| Calcium carbonate is also used in [[flue gas desulfurisation]] applications eliminating harmful SO<sub>2</sub> and NO<sub>2</sub> emissions from coal and other fossil fuels burnt in large fossil fuel power stations.<ref>{{cite web|title=Calcium Carbonate in Flue Gas Desulphurisation|url=http://www.congcal.com/markets/environmental/|work=http://www.congcal.com/markets/environmental/|publisher=Congcal|accessdate=5 August 2013}}</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 °C, but stating an absolute threshold is misleading. Calcium carbonate exists in equilibrium with calcium oxide and [[carbon dioxide]] at any temperature. At each temperature there is a [[partial pressure]] of carbon dioxide that is in equilibrium with calcium carbonate. At room temperature the equilibrium overwhelmingly favors calcium carbonate, because the equilibrium CO<sub>2</sub> pressure is only a tiny fraction of the partial CO<sub>2</sub> pressure in air, which is about 0.035 kPa.
| |
| | |
| At temperatures above 550 °C the equilibrium CO<sub>2</sub> pressure begins to exceed the CO<sub>2</sub> pressure in air. So above 550 °C, calcium carbonate begins to outgas CO<sub>2</sub> into air. However, in a charcoal fired kiln, the concentration of CO<sub>2</sub> will be much higher than it is in air. Indeed if all the [[oxygen]] in the kiln is consumed in the fire, then the partial pressure of CO<sub>2</sub> in the kiln can be as high as 20 kPa.<ref name="solvaypcc2007">{{cite web|title = Solvay Precipitated Calcium Carbonate: Production|publisher = Solvay S. A. |date=2007-03-09|accessdate = 2007-12-30|url = http://www.solvaypcc.com/safety_environment/0,0,1000044-_EN,00.html}}</ref>
| |
| | |
| The table shows that this equilibrium pressure is not achieved until the temperature is nearly 800 °C. For the outgassing of CO<sub>2</sub> from calcium carbonate to happen at an economically useful rate, the equilibrium pressure must significantly exceed the ambient pressure of CO<sub>2</sub>. And for it to happen rapidly, the equilibrium pressure must exceed total atmospheric pressure of 101 kPa, which happens at 898 °C.{{clear right}}
| |
| | |
| {| class="wikitable"
| |
| ! {{chembox header}} colspan=18|Equilibrium pressure of CO<sub>2</sub> over CaCO<sub>3</sub> (P) vs. 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 CO<sub>2</sub> pressure===
| |
| [[File:CanarySpring.jpg|thumb|right|Travertine calcium carbonate deposits from a [[hot spring]]]]
| |
| Calcium carbonate is poorly soluble in pure water (47 mg/L at normal atmospheric CO<sub>2</sub> partial pressure as shown below).
| |
| | |
| The equilibrium of its solution is given by the equation (with dissolved calcium carbonate on the right):
| |
| :{| width="450"
| |
| | style="width:50%; height:30px;"| CaCO<sub>3</sub> {{eqm}} Ca<sup>2+</sup> + CO<sub>3</sub><sup>2–</sup>
| |
| | ''K''<sub>sp</sub> = 3.7×10<sup>−9</sup> to 8.7×10<sup>−9</sup> at 25 °C
| |
| |}
| |
| | |
| where the [[solubility product]] for [Ca<sup>2+</sup>][CO<sub>3</sub><sup>2–</sup>] is given as anywhere from ''K''<sub>sp</sub> = 3.7×10<sup>−9</sup> to ''K''<sub>sp</sub> = 8.7×10<sup>−9</sup> at 25 °C, depending upon the data source.<ref 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}}</ref> What the equation means is that the product of molar concentration of calcium ions ([[mole (unit)|moles]] of dissolved Ca<sup>2+</sup> per liter of solution) with the molar concentration of dissolved CO<sub>3</sub><sup>2–</sup> cannot exceed the value of ''K''<sub>sp</sub>. This seemingly simple solubility equation, however, must be taken along with the more complicated equilibrium of [[carbon dioxide]] with [[water]] (see [[carbonic acid]]). Some of the CO<sub>3</sub><sup>2–</sup> combines with H<sup>+</sup> in the solution according to:
| |
| | |
| :{| width="450"
| |
| | style="width:50%; height:25px;"| HCO<sub>3</sub><sup>–</sup> {{eqm}} H<sup>+</sup> + CO<sub>3</sub><sup>2–</sup>
| |
| | ''K''<sub>a2</sub> = 5.61×10<sup>−11</sup> at 25 °C
| |
| |}
| |
| | |
| HCO<sub>3</sub><sup>–</sup> is known as the [[bicarbonate]] ion. [[Calcium bicarbonate]] is many times more soluble in water than calcium carbonate—indeed it exists ''only'' in solution.
| |
| | |
| Some of the HCO<sub>3</sub><sup>–</sup> combines with H<sup>+</sup> in solution according to:
| |
| | |
| :{| width="450"
| |
| | style="width:50%; height:25px;"|H<sub>2</sub>CO<sub>3</sub> {{eqm}} H<sup>+</sup> + HCO<sub>3</sub><sup>–</sup>
| |
| | ''K''<sub>a1</sub> = 2.5×10<sup>−4</sup> at 25 °C
| |
| |}
| |
| | |
| Some of the H<sub>2</sub>CO<sub>3</sub> breaks up into water and dissolved carbon dioxide according to:
| |
| | |
| :{| width="450"
| |
| | style="width:50%; height:25px;"| H<sub>2</sub>O + CO<sub>2</sub>(dissolved) {{eqm}} H<sub>2</sub>CO<sub>3</sub>
| |
| | ''K''<sub>h</sub> = 1.70×10<sup>−3</sup> at 25 °C
| |
| |}
| |
| | |
| And dissolved carbon dioxide is in equilibrium with atmospheric carbon dioxide according to:
| |
| | |
| :{| width="500"
| |
| | style="width:45%;"|<math>\frac{P_{\text{CO}_2}}{[\text{CO}_2]}\ =\ k_\text{H}</math>
| |
| | where ''k''<sub>H</sub> = 29.76 atm/(mol/L) at 25 °C ([[Henry's law|Henry constant]]), <math>\scriptstyle P_{\text{CO}_2}</math> being the CO<sub>2</sub> partial pressure.
| |
| |}
| |
| | |
| For ambient air, <math>\scriptstyle P_{\text{CO}_2}</math> is around 3.5×10<sup>−4</sup> atmospheres (or equivalently 35 [[Pascal (unit)|Pa]]). The last equation above fixes the concentration of dissolved CO<sub>2</sub> as a function of <math>\scriptstyle P_{\text{CO}_2}</math>, independent of the concentration of dissolved CaCO<sub>3</sub>. At atmospheric partial pressure of CO<sub>2</sub>, dissolved CO<sub>2</sub> concentration is 1.2×10<sup>−5</sub> moles/liter. The equation before that fixes the concentration of H<sub>2</sub>CO<sub>3</sub> as a function of [CO<sub>2</sub>]. For [CO<sub>2</sub>]=1.2×10<sup>−5</sub>, it results in [H<sub>2</sub>CO<sub>3</sub>]=2.0×10<sup>−8</sup> moles per liter. When [H<sub>2</sub>CO<sub>3</sub>] is known, the remaining three equations together with
| |
| | |
| :{| width="450"
| |
| | style="width:50%; height:25px;"| H<sub>2</sub>O {{eqm}} H<sup>+</sup> + OH<sup>–</sup>
| |
| | ''K'' = 10<sup>−14</sup> at 25 °C
| |
| |}
| |
| | |
| (which is true for all aqueous solutions), and the fact that the solution must be electrically neutral,
| |
| | |
| :2[Ca<sup>2+</sup>] + [H<sup>+</sup>] = [HCO<sub>3</sub><sup>–</sup>] + 2[CO<sub>3</sub><sup>2–</sup>] + [OH<sup>–</sup>]
| |
| | |
| make it possible to solve simultaneously for the remaining five unknown concentrations (note that the above form of the neutrality equation is valid only if calcium carbonate has been put in contact with pure water or with a neutral pH solution; in the case where the origin water solvent pH is not neutral, the equation is modified).
| |
| | |
| {| class="wikitable" style="float:right;"
| |
| ! {{chembox header}} colspan="3"|Calcium ion solubility as a function of<br> [[carbon dioxide|CO<sub>2</sub>]] [[partial pressure]] at 25 °C (''K''<sub>sp</sub> = 4.47×10<sup>−9</sup>)
| |
| |-
| |
| !<math>\scriptstyle P_{\text{CO}_2}</math> (atm)
| |
| ![[pH]]
| |
| ![Ca<sup>2+</sup>] (mol/L)
| |
| |-
| |
| |10<sup>−12</sup> ||12.0||5.19 × 10<sup>−3</sup>
| |
| |-
| |
| |10<sup>−10</sup> ||11.3||1.12 × 10<sup>−3</sup>
| |
| |-
| |
| |10<sup>−8</sup> ||10.7||2.55 × 10<sup>−4</sup>
| |
| |-
| |
| |10<sup>−6</sup> ||9.83||1.20 × 10<sup>−4</sup>
| |
| |-
| |
| |10<sup>−4</sup> ||8.62||3.16 × 10<sup>−4</sup>
| |
| |-
| |
| |3.5 × 10<sup>−4</sup>||8.27||4.70 × 10<sup>−4</sup>
| |
| |-
| |
| |10<sup>−3</sup> ||7.96||6.62 × 10<sup>−4</sup>
| |
| |-
| |
| |10<sup>−2</sup> ||7.30||1.42 × 10<sup>−3</sup>
| |
| |-
| |
| |10<sup>−1</sup> ||6.63||3.05 × 10<sup>−3</sup>
| |
| |-
| |
| |1 ||5.96||6.58 × 10<sup>−3</sup>
| |
| |-
| |
| |10 ||5.30||1.42 × 10<sup>−2</sup>
| |
| |}
| |
| | |
| The table on the right shows the result for [Ca<sup>2+</sup>] and [H<sup>+</sup>] (in the form of pH) as a function of ambient partial pressure of CO<sub>2</sub> (''K''<sub>sp</sub> = 4.47×10<sup>−9</sup> has been taken for the calculation).
| |
| | |
| *At atmospheric levels of ambient CO<sub>2</sub> the table indicates the solution will be slightly alkaline with a maximum CaCO<sub>3</sub> solubility of 47 mg/L.
| |
| | |
| *As ambient CO<sub>2</sub> partial pressure is reduced below atmospheric levels, the solution becomes more and more alkaline. At extremely low <math>\scriptstyle P_{\text{CO}_2}</math>, dissolved CO<sub>2</sub>, bicarbonate ion, and carbonate ion largely evaporate from the solution, leaving a highly alkaline solution of [[calcium hydroxide]], which is more soluble than CaCO<sub>3</sub>. Note that for <math>\scriptstyle P_{\text{CO}_2}</math> = 10<sup>−12</sup> atm, the [Ca<sup>2+</sup>][OH<sup>−</sup>]<sup>2</sup> product is still below the solubility product of Ca(OH)<sub>2</sub> (8×10<sup>−6</sup>). For still lower CO<sub>2</sub> pressure, Ca(OH)<sub>2</sub> precipitation will occur before CaCO<sub>3</sub> precipitation.
| |
| | |
| *As ambient CO<sub>2</sub> partial pressure increases to levels above atmospheric, pH drops, and much of the carbonate ion is converted to bicarbonate ion, which results in higher solubility of Ca<sup>2+</sup>.
| |
| | |
| The effect of the latter is especially evident in day-to-day life of people who have hard water. Water in aquifers underground can be exposed to levels of CO<sub>2</sub> much higher than atmospheric. As such water percolates through calcium carbonate rock, the CaCO<sub>3</sub> dissolves according to the second trend. When that same water then emerges from the tap, in time it comes into equilibrium with CO<sub>2</sub> levels in the air by outgassing its excess CO<sub>2</sub>. The calcium carbonate becomes less soluble as a result and the excess precipitates as lime scale. This same process is responsible for the formation of [[stalactites]] and [[stalagmite]]s in limestone caves.
| |
| | |
| Two hydrated phases of calcium carbonate, [[monohydrocalcite]], CaCO<sub>3</sub>·H<sub>2</sub>O and [[ikaite]], CaCO<sub>3</sub>·6H<sub>2</sub>O, may [[precipitate]] from water at ambient conditions and persist as metastable phases.
| |
| | |
| ===With varying pH===
| |
| The maximum solubility of calcium carbonate in normal atmospheric conditions (<math>\scriptstyle P_{\mathrm{CO}_2}</math> = 3.5 × 10<sup>−4</sup> atm) varies as the pH of the solution is adjusted. This is for example the case in a swimming pool where the pH is maintained between 7 and 8 (by addition of [[sodium bisulfate]] NaHSO<sub>4</sub> to decrease the pH or of [[sodium bicarbonate]] NaHCO<sub>3</sub> to increase it). From the above equations for the solubility product, the hydration reaction and the two acid reactions, the following expression for the maximum [Ca<sup>2+</sup>] can be easily deduced:
| |
| :<math>[\text{Ca}^{2+}]_\text{max} = \frac{K_\text{sp}k_\text{H}} {K_\text{h}K_\text{a1}K_\text{a2}} \frac{[\text{H}^+]^2}{P_{\text{CO}_2}}</math>
| |
| showing a quadratic dependence in [H<sup>+</sup>]. The numerical application with the above values of the constants gives{{Citation needed|date=September 2009}}
| |
| | |
| {| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: white; border-collapse: collapse; border-color: #C0C090;"
| |
| |-
| |
| |{{chembox header}} width="170" align="center" |'''pH'''
| |
| |7.0
| |
| |7.2
| |
| |7.4
| |
| |7.6
| |
| |7.8
| |
| |8.0
| |
| |8.2
| |
| |8.27
| |
| |8.4
| |
| |-
| |
| |{{chembox header}} width="170" align="center" |'''[Ca<sup>2+</sup>]<sub>max</sub> (10<sup>−6</sup>mol/L)'''
| |
| |180
| |
| |71.7
| |
| |28.5
| |
| |11.4
| |
| |4.52
| |
| |1.80
| |
| |0.717
| |
| |0.519
| |
| |0.285
| |
| |-
| |
| |{{chembox header}} width="170" align="center"|'''[Ca<sup>2+</sup>]<sub>max</sub> (mg/L)'''
| |
| |7.21
| |
| |2.87
| |
| |1.14
| |
| |0.455
| |
| |0.181
| |
| |0.0721
| |
| |0.0287
| |
| |0.0208
| |
| |0.0114
| |
| |}
| |
| Comments:
| |
| *decreasing the pH from 8 to 7 increases the maximum Ca<sup>2+</sup> concentration by a factor 100. Water with a pH maintained to 7 can dissolve up to 15.9 mg/L of CaCO<sub>3</sub>. This explains the high Ca<sup>2+</sup> concentration in some mineral waters with pH close to 7.
| |
| *note that the Ca<sup>2+</sup> concentration of the previous table is recovered for pH = 8.27
| |
| *keeping the pH to 7.4 in a swimming pool (which gives optimum HClO/ClO<sup>−</sup> [[hypochlorous acid|ratio]] in the case of "chlorine" maintenance) results in a maximum Ca<sup>2+</sup> concentration of 1010 mg/L. This means that successive cycles of water evaporation and partial renewing may result in a very [[hard water]] before CaCO<sub>3</sub> precipitates (water with a Ca<sup>2+</sup> concentration above 120 mg/L is considered very hard). Addition of a calcium [[chelation|sequestering agent]] or complete renewing of the water will solve the problem.
| |
| | |
| ===Solubility in a strong or weak acid solution===
| |
| Solutions of [[strong acid|strong]] ([[hydrochloric acid|HCl]]), 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 CaCO<sub>3</sub> that can be "dissolved" by one liter of an acid solution can be calculated using the above equilibrium equations.
| |
| *In the case of a strong monoacid with decreasing acid concentration [A] = [A<sup>−</sup>], we obtain (with CaCO<sub>3</sub> molar mass = 100 g):
| |
| | |
| {| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: white; border-collapse: collapse; border-color: #C0C090;"
| |
| |-
| |
| |width="160" {{chembox header}} |'''[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>
| |
| |-
| |
| |width="160" {{chembox header}} |'''Initial pH'''
| |
| |0.00||1.00||2.00||3.00||4.00||5.00||6.00||6.79||7.00
| |
| |-
| |
| |width="160" {{chembox header}} |'''Final pH'''
| |
| |6.75||7.25||7.75||8.14||8.25||8.26||8.26||8.26||8.27
| |
| |-
| |
| |width="160" {{chembox header}} |'''Dissolved CaCO<sub>3</sub> (g per liter of acid)'''
| |
| |50.0||5.00||0.514||0.0849||0.0504||0.0474||0.0471||0.0470||0.0470
| |
| |}
| |
| | |
| where the initial state is the acid solution with no Ca<sup>2+</sup> (not taking into account possible CO<sub>2</sub> dissolution) and the final state is the solution with saturated Ca<sup>2+</sup>. For strong acid concentrations, all species have a negligible concentration in the final state with respect to Ca<sup>2+</sup> and A<sup>−</sup> so that the neutrality equation reduces approximately to 2[Ca<sup>2+</sup>] = [A<sup>−</sup>] yielding <math>\scriptstyle[\mathrm{Ca}^{2+}] \simeq \frac{[\mathrm{A}^-]}{2}</math>. When the concentration decreases, [HCO<sub>3</sub><sup>−</sup>] becomes non-negligible so that the preceding expression is no longer valid. For vanishing acid concentrations, one can recover the final pH and the solubility of CaCO<sub>3</sub> in pure water.
| |
| | |
| *In the case of a weak monoacid (here we take acetic acid with p''K''<sub>A</sub> = 4.76) with decreasing total acid concentration [A] = [A<sup>−</sup>]+[AH], we obtain:
| |
| | |
| {| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: white; border-collapse: collapse; border-color: #C0C090;"
| |
| |-
| |
| |width="160" {{chembox header}} |'''[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>
| |
| |-
| |
| |width="160" {{chembox header}} |'''Initial pH'''
| |
| |2.38||2.88||3.39||3.91||4.47||5.15||6.02||6.79||7.00
| |
| |-
| |
| |width="160" {{chembox header}} |'''Final pH'''
| |
| |6.75||7.25||7.75||8.14||8.25||8.26||8.26||8.26||8.27
| |
| |-
| |
| |width="160" {{chembox header}} |'''Dissolved CaCO<sub>3</sub> (g per liter of acid)'''
| |
| |49.5||4.99||0.513||0.0848||0.0504||0.0474||0.0471||0.0470||0.0470
| |
| |}
| |
| 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 CaCO<sub>3</sub> 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 H<sup>+</sup> ions as the strong acid to "dissolve" the calcium carbonate.
| |
| | |
| *The calculation in the case of [[phosphoric acid]] (which is the most widely used for domestic applications) is more complicated since the concentrations of the four dissociation states corresponding to this acid must be calculated together with [HCO<sub>3</sub><sup>−</sup>], [CO<sub>3</sub><sup>2−</sup>], [Ca<sup>2+</sup>], [H<sup>+</sup>] and [OH<sup>−</sup>]. The system may be reduced to a seventh degree equation for [H<sup>+</sup>] the numerical solution of which gives
| |
| | |
| {| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: white; border-collapse: collapse; border-color: #C0C090;"
| |
| |-
| |
| |width="160" {{chembox header}} |'''[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>
| |
| |-
| |
| |width="160" {{chembox header}} |'''Initial pH'''
| |
| |1.08||1.62||2.25||3.05||4.01||5.00||5.97||6.74||7.00
| |
| |-
| |
| |width="160" {{chembox header}} |'''Final pH'''
| |
| |6.71||7.17||7.63||8.06||8.24||8.26||8.26||8.26||8.27
| |
| |-
| |
| |width="160" {{chembox header}} |'''Dissolved CaCO<sub>3</sub> (g per liter of acid)'''
| |
| |62.0||7.39||0.874||0.123||0.0536||0.0477||0.0471||0.0471||0.0470
| |
| |}
| |
| | |
| where [A] = [H<sub>3</sub>PO<sub>4</sub>] + [H<sub>2</sub>PO<sub>4</sub><sup>−</sup>] + [HPO<sub>4</sub><sup>2−</sup>] + [PO<sub>4</sub><sup>3−</sup>] is the total acid concentration. Thus phosphoric acid is more efficient than a monoacid since at the final almost neutral pH, the second dissociated state concentration [HPO<sub>4</sub><sup>2−</sup>] is not negligible (see [[phosphoric acid#pH and composition of a phosphoric acid aqueous solution|phosphoric acid]]).
| |
| | |
| ==See also==
| |
| {{colbegin|3}}
| |
| *[[Cuttlebone]]
| |
| *[[Cuttlefish]]
| |
| *[[Gesso]]
| |
| *[[Limescale]]
| |
| *[[Marble]]
| |
| *[[Ocean acidification]]
| |
| {{colend}}
| |
| | |
| ==References==
| |
| {{reflist|30em}}
| |
| | |
| ==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]
| |
| *[http://www.cdc.gov/niosh/npg/npgd0090.html CDC - NIOSH Pocket Guide to Chemical Hazards - Calcium Carbonate]
| |
| | |
| {{Calcium compounds}}
| |
| {{Antacids}}
| |
| {{Drugs for treatment of hyperkalemia and hyperphosphatemia}}
| |
| | |
| {{DEFAULTSORT:Calcium Carbonate}}
| |
| [[Category:Calcium compounds]]
| |
| [[Category:Carbonates]]
| |
| [[Category:Limestone]]
| |
| [[Category:Phosphate binders]]
| |
| [[Category:Excipients]]
| |
| [[Category:Antacids]]
| |