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'''Calcium carbonateText in bold is surrounded by triple aporstrophe's''' ''Italicised text 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)|rocks]] as the [[mineral]]surrounded by double aporstrophe's [[calcite]] and [[aragonite]] (most notably as [[limestone]], which is a type of sedimentary rock build mainly of calcite) and is the main component of [[pearl]]s and the [[Seashell|shells of marine organisms]], [[snail]]s, and eggs. Calcium carbonate is the active ingredient in [[agricultural lime]] and is created when calcium ions in [[hard water]] react with [[carbonate ion]]s to create [[limescale]]. It is medicinally used as a [[calcium]] supplement or as an [[antacid]], but excessive consumption can be hazardous.''
==Occurrence== [[File:Calcium carbonate chunks.JPG|thumb|Calcium carbonate chunks from clamshell]] ===Geological sources===[[Calcite]], [[aragonite]] and [[vaterite]] are pure calcium carbonate minerals. Industrially important source rocks which are predominantly calcium carbonate include [[limestone]], [[chalk]], [[marble]] and [[travertine]]. [[File:Silfurberg.jpg|thumb|Calcite is the most stable polymorph of calcium carbonate. It is transparent to opaque. A transparent variety called [[Iceland spar]] (shown here) is used for optical purposes.{{clarify|date=January 2018}}]] ===Biological sources===Eggshells, snail shells and most seashells are predominantly calcium carbonate and can be used as industrial sources of that chemical.<ref>{{cite web |title=How are seashells created? |author=Horne, Francis |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> Dark green vegetables such as broccoli and kale contain dietarily significant amounts of calcium carbonate, however, they are not practical as an industrial source.<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> ===Extraterrestrial===Beyond Earth, strong evidence suggests the presence of calcium carbonate on [[Mars]]. Signs of calcium carbonate have been detected at more than one location (notably at [[Gusev crater|Gusev]] and [[Huygens (crater)|Huygens]] craters). This provides some evidence for the past presence of liquid water.<ref>{{cite journal| last1=Boynton |first1=WV| last2=Ming |first2=DW| last3=Kounaves |first3=SP| last4=Young |first4=SM| last5=Arvidson |first5=RE| last6=Hecht |first6=MH| last7=Hoffman |first7=J| last8=Niles |first8=PB| last9=Hamara |first9=DK| 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 |issueKey Stage 1=5936 |pages= 61–64| year=2009 |pmid=19574384 |bibcode=2009Sci...325...61B| display-authors=3| doi=10.1126/science.1172768| doi-broken-date=2017-01-31 }}</ref><ref name=Clark2007>{{cite journal| author1=Clark| year=2007| title=Evidence for montmorillonite or its compositional equivalent in Columbia Hills, Mars| journal=[[Journal of Geophysical Research]]| volume=112 |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| displayauthors=3 | url=http://dspace.stir.ac.uk/bitstream/1893/17119/1/Clark2007_Evidence_for_montmorillonite_or_its_compositional_equivalent_in_Columbia_Hills_Mars.pdf}}</ref> ==Geology==Carbonate is found frequently in geologic settings and constitutes an enormous [[carbon cycle|carbon reservoir]]. Calcium carbonate occurs as [[aragonite]], [[calcite]] and [[dolomite]]. The [[carbonate mineral]]s form the rock types: [[limestone]], [[chalk]], [[marble]], [[travertine]], [[tufa]], and others. In warm, clear tropical waters [[coral]]s are more abundant than towards the poles where the waters are cold. Calcium carbonate contributors, including [[plankton]] (such as [[coccolith]]s and planktic [[foraminifera]]), [[coralline algae]], [[sea sponge|sponges]], [[brachiopod]]s, [[echinoderm]]s, [[bryozoa]] and [[Mollusc shell|mollusks]], are typically found in shallow water environments where sunlight and filterable food are more abundant. Cold-water carbonates do exist at higher latitudes but have a very slow growth rate. The [[calcification]] processes are changed by [[ocean acidification]]. Where the [[oceanic crust]] is [[Subduction|subducted]] under a [[continental plate]] sediments will be carried down to warmer zones in the [[asthenosphere]] and [[lithosphere]]. Under these conditions calcium carbonate decomposes to produce [[carbon dioxide]] which, along with other gases, give rise to explosive [[volcano|volcanic eruptions]]. ===Carbonate compensation depth===The [[carbonate compensation depth]] (CCD) is the point in the ocean where the rate of precipitation of calcium carbonate is balanced by the rate of dissolution due to the conditions present. Deep in the ocean, the temperature drops and pressure increases. Calcium carbonate is unusual in that its solubility increases with decreasing temperature. Increasing pressure also increases the solubility of calcium carbonate. The carbonate compensation depth can range from 4–6 km below sea level. ===Role in taphonomy===Calcium carbonate can [[taphonomy|preserve fossils]] through [[permineralization]]. Most of the vertebrate fossils of the [[Two Medicine Formation]]—a [[geologic formation]] known for its [[duck-billed dinosaur]] eggs—are preserved by CaCO<sub>3</sub> permineralization.<ref nameborder="twoturn0" /> This type of preservation conserves high levels of detail, even down to the microscopic level. However, it also leaves specimens vulnerable to [[weathering]] when exposed to the surface.<ref namestyle="twoturn">Trexler, D. (2001) [https://books.google.com/books?id=mgc6CS4EUPsC&pg=PA98 "Two Medicine Formation, Montanaborder-collapse: geology and faunacollapse"], pp. 298–309 in ''Mesozoic Vertebrate Life'', Tanke, D. H., and Carpenter, K. (eds), Indiana University Press. {{ISBN|0-253-33907-3}}</ref> [[Trilobite]] populations were once thought to have composed the majority of aquatic life during the [[Cambrian]], due to the fact that their calcium carbonate-rich shells were more easily preserved than those of other species,<ref>{{Cite book|url=https://www.nap.edu/catalog/11630/out-of-thin-air-dinosaurs-birds-and-earths-ancient-atmosphere|title=Out of Thin Air: Dinosaurs, Birds, and Earth's Ancient Atmosphere|last=Ward|first=Peter|date=|publisher=|year=|isbn=9780309666121|location=|pages=|language=en|doi=10.17226/11630}}</ref> which had purely chitinous shells. ==Uses== ===Industrial applications=== The main use of calcium carbonate is in the construction industry, either as a building material or limestone aggregate for road building or as an ingredient of cement or as the starting material for the preparation of builder's lime by burning in a kiln. However, because of weathering mainly caused by [[acid rain]],<ref>{{cite web|title = Effects of Acid Rain|publisher = US Environmental Protection Agency|accessdate = 14 March 2015|url = http://www.epa.gov/acidrain/effects/materials.html}}</ref> calcium carbonate (in limestone form) is no longer used for building purposes on its own, but only as a raw/primary substance for building materials. Calcium carbonate is also used in the purification of [[iron]] from [[iron ore]] in a [[blast furnace]]. The carbonate is 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 = 30 December 2007|url = http://www.scienceaid.co.uk/chemistry/industrial/blastfurnace.html}}</ref> In the [[oil industry]], calcium carbonate is added to [[drilling fluid]]s as a formation-bridging and filtercake-sealing agent; it is also a weighting material which increases the density of drilling fluids to control the downhole pressure. Calcium carbonate is added to swimming pools, as a [[pH]] corrector for maintaining [[alkalinity]] and offsetting the acidic properties of the disinfectant agent.{{citation needed|date=June 2015}} 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 sweet water to produce a calcium hydroxide suspension for the precipitation of impurities in raw juice during [[carbonatation]].<ref>{{cite book|last1=McGinnis|first1=R.A.|title=Beet-Sugar Technology|publisher=Beet Sugar Development Foundation|page=178|edition=2nd}}</ref> Calcium carbonate 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 |deadurl=yes |archiveurl=https://web.archive.org/web/20140725032803/http://www.aristocratholding.com/calris-5.html |archivedate=25 July 2014 }}</ref> Fine ground calcium carbonate (GCC) is an essential ingredient in the microporous film used in [[diapers]] and some building films as the pores are nucleated around the calcium carbonate particles during the manufacture of the film by biaxial stretching. GCC 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 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 micrometres.{{citation needed|date=June 2015}} Calcium carbonate is widely used as an extender in paints,<ref name = reade>{{cite web|title = Calcium Carbonate Powder|publisher = Reade Advanced Materials |date=4 February 2006|accessdate = 30 December 2007|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).{{citation needed|date=June 2015}} [[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=1 August 2008 |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.{{citation needed|date=June 2015}} Precipitated calcium carbonate, made by dropping [[calcium oxide]] into water, is used by itself or with additives as a white paint, known as [[whitewashing]].{{citation needed|date=June 2015}} 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.{{citation needed|date=June 2015}} 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> ===Health and dietary applications===[[File:500 mg calcium supplements with vitamin DHydrogenSymbol1.jpgpng|thumbcenter|500-milligram calcium supplements made from calcium carbonate]]Calcium carbonate is widely used medicinally as an inexpensive dietary calcium supplement for [[antacid31px|gastric antacid]]<ref name = medline>{{cite web|work = Medline Plus|publisher = [[National Institutes of Health]]|title = Calcium Carbonate |date=1 October 2005|accessdate = 30 December 2007|url = https://www.nlm.nih.gov/medlineplus/druginfo/medmaster/a601032.html |archiveurl link= https://web.archive.org/web/20071017031324/http://www.nlm.nih.gov/medlineplus/druginfo/medmaster/a601032.html <!-- Bot retrieved archive --> |archivedate = 17 October 2007}}</ref> (e.g., [[Tums]]). It may be used as a [[phosphate binder]] for the treatment of [[hyperphosphatemiaHydrogen]] (primarily in patients with [[chronic renal failure]]). It is also used in the pharmaceutical industry as an inert [[Excipient|filler]] for [[Tablet (pharmacy)|tablets]] and other [[pharmaceuticals]].<ref>{{cite book|author1=Lieberman, Herbert A. |author2=Lachman, Leon |author3=Schwartz, Joseph B. |title = Pharmaceutical Dosage Forms: Tablets|year = 1990|isbn = 0-8247-8044-2|page=153|publisher = Dekker|location = New York}}</ref> Calcium carbonate is used in the production of calcium oxide as well as toothpaste and has seen a resurgence as a food preservative and color retainer, when used in or with products such as organic apples.<ref>[http://chemistry.about.com/od/foodcookingchemistry/a/cadditives.htm Food Additives – Names Starting with C]. Chemistry.about.com (10 April 2012). Retrieved 2012-05-24.</ref> 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 [[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=Gabriely, Ilan |journal=New England Journal of Medicine|year=2008|volume=358|pmid=18450607|doi=10.1056/NEJMcps0706188|issue=18|last2=Leu|first2=James P.|last3=Barzel|first3=Uriel S.|pages=1952–6}}</ref> As a [[food additive]] it is designated 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> and it has an INS number of 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=27 October 2011}}</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=27 October 2011 |deadurl=yes |archiveurl=https://web.archive.org/web/20130314104055/http://www.fda.gov/Food/FoodIngredientsPackaging/FoodAdditives/FoodAdditiveListings/ucm091048.htm |archivedate=14 March 2013 |df=dmy }}</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=27 October 2011}}</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| pmid = 16177199| year = 2005| author1 = 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 = 2379–82| last2 = Martin| first2 = B. R.| last3 = Weaver| first3 = C. M.}}</ref> Calcium carbonate is also used as a [[firming agent]] in many canned or bottled vegetable products. ===Agricultural use===[[Agricultural lime]], powdered chalk or limestone, is used as a cheap method for neutralising acidic soil, making it suitable for planting.<ref name="Oates2008">{{cite book|first=J. A. H.|last=Oates|title=Lime and LimestoneFile: Chemistry and Technology, Production and Uses|url=https://booksHeliumSymbol1.google.com/books?id=MVoEMNI5Vb0C&pg=PA111png|date=11 July 2008center|publisher=John Wiley & Sons31px|isbn=978-3-527-61201-7|pages=111–3}}</ref> ===Household use===Calcium carbonate is a key ingredient in many household cleaning powders like [[Comet (cleanser)]] and is used as a scrubbing agent. ===Environmental applications=== In 1989, a researcher, Ken Simmons, introduced CaCO<sub>3</sub> into the Whetstone Brook in [[Massachusetts]].<ref>{{cite news|agency = [[Associated Press]]|title = Limestone Dispenser Fights Acid Rain in Stream |date=13 June 1989|url = https://query.nytimes.com/gst/fullpage.html?res=950DEFD9173FF930A25755C0A96F948260|work = The 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 name=env>{{cite web|title=Environmental Uses for Calcium Carbonate|url=http://www.congcal.com/markets/environmental/|publisher=Congcal|accessdate=5 August 2013}}</ref><ref>{{cite journal|author = Schreiber, 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|url=https://link.springer.com/article/10.1007%2FBF00160344|doi-broken-date = 2017-01-31}}</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|author1=Kircheis, Dan |author2=Dill, Richard |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 |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 }}</ref> 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 name=env/> ==Calcination equilibrium==[[Calcination]] of [[limestone]] using [[charcoalHelium]] 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=9 March 2007|accessdate = 30 December 2007|url = http://www.solvaypcc.com/safety_environment/0,0,1000044-_EN,00.html}}</ref> The table shows that this partial 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}} |Equilibrium pressure of CO<sub>2</sub> over CaCO<sub>3</sub> (P) vs. temperature (T).<ref name=crc>{{RubberBible86th}}</ref>
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|'''P (kPa)'''[[File:LithiumSymbol1.png|center|031px|link=Lithium]]|[[File:BerylliumSymbol1.055png|center|0.1331px|link=Beryllium]]|||||0.31||1.80||5.9||9[[File:BoronSymbol1.3png|center|31px|14link=Boron]]|[[File:CarbonSymbol1.png|24center|31px|34link=Carbon]]|[[File:NitrogenSymbol1.png|51center|31px|72 link=Nitrogen]]|[[File:OxygenSymbol1.png|80center|31px|91link=Oxygen]]|[[File:FluorineSymbol1.png|101center|31px|179link=Fluorine]]|[[File:NeonSymbol1.png|901center|31px|3961link=Neon]]
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|'''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:CanarySpringSodiumSymbol1.jpgpng|thumbcenter|right31px|[[Travertinelink=Sodium]] 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="500"| 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> = 8File:MagnesiumSymbol1.7×10<sup>−9</sup> at 25 °C, depending upon the data source.<ref name = crc/><ref>{{cite webpng|title = Selected Solubility Products and Formation Constants at 25 °Ccenter|publisher = [[California State University, Dominguez Hills]]31px|url link= 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)|molesMagnesium]] 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="500"| 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="500"| 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="500"| style="width[[File: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> = 1AluminiumSymbol1.70×10<sup>−3</sup> at 25 °Cpng|} And dissolved carbon dioxide is in equilibrium with atmospheric carbon dioxide according to: :{center| width="500"31px| stylelink="width:45%;"|<math chem>\frac{P_{\ce{CO2}}}{[\ce{CO2}Aluminium]]}\ =\ k_\ce{H}</math>| where ''k''<sub>H</sub> = 29.76 atm/(mol/L) at 25 °C ([[Henry's lawFile:SiliconSymbol1.png|Henry constantcenter|31px|link=Silicon]]), <math chem>P_{\ce{CO2}}</math> being the CO<sub>2</sub> partial pressure.|} For ambient air, <math chem>P_{\ce{CO2}}</math> is around 3.5×10<sup>−4</sup> atmospheres (or equivalently 35 [[Pascal (unit)|Pa]])File:PhosphorusSymbol1. The last equation above fixes the concentration of dissolved CO<sub>2</sub> as a function of <math chem>P_{\ce{CO2}}</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</sup> 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</sup>, 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{png| class="wikitable floatright"center|+ {{chembox header}} 31px|Calcium ion solubility as a function of [[carbon dioxide|CO<sub>2</sub>link=Phosphorus]] [[partial pressure]] at 25 °C {{math|1=(''K''<sub>sp</sub> = 4.47×10<sup>−9</sup>)}}|-!<math chem>\scriptstyle P_\ce{CO2}</math> (atm)![[pH]]![Ca<sup>2+</sup>] (mol/L)|-| 10<sup>−12</sup> ||12File:SulphurSymbol1.0png|center|5.19 × 10<sup>−3</sup>31px|-link=Sulphur]]| 10<sup>−10</sup> ||11[[File:ChlorineSymbol1.3png|center|1.12 × 10<sup>−3</sup>31px|-link=Chlorine]]| 10<sup>−8</sup> ||10[[File:ArgonSymbol1.7png|center|2.55 × 10<sup>−4</sup>31px|-| 10<sup>−6</sup> ||9.83||1.20 × 10<sup>−4</sup>|-| 10<sup>−4</sup> ||8.62||3.16 × 10<sup>−4</sup>link=Argon]]
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| 3[[File:PotassiumSymbol1.5 × 10<sup>−4</sup>png|center|31px|link=Potassium]]|8[[File:CalciumSymbol1.27png|center|31px|link=Calcium]]|4[[File:ScandiumSymbol1.70 × 10<sup>−4</sup>png|center|31px|link=Scandium]]|-[[File:TitaniumSymbol1.png|center|31px|link=Titanium]]| 10<sup>−3</sup> [[File:VanadiumSymbol1.png|center|731px|link=Vanadium]]|[[File:ChromiumSymbol1.96png|center|31px|link=Chromium]]|6[[File:ManganeseSymbol1.62 × 10<sup>−4</sup>png|center|31px|link=Manganese]]|-[[File:IronSymbol1.png|center|31px|link=Iron]]| 10<sup>−2</sup> [[File:CobaltSymbol1.png|center|731px|link=Cobalt]]|[[File:NickelSymbol1.30png|center|31px|link=Nickel]]|1[[File:CopperSymbol1.42 × 10<sup>−3</sup>png|center|31px|link=Copper]]|-[[File:ZincSymbol1.png|center| 10<sup>−1</sup> 31px|link=Zinc]]|6[[File:GalliumSymbol1.63png|center|3.05 × 10<sup>−3</sup>31px|-link=Gallium]]| 1 [[File:GermaniumSymbol1.png|center|5.9631px|link=Germanium]]|6[[File:ArsenicSymbol1.58 × 10<sup>−3</sup>png|-center| 10 31px|link=Arsenic]]|5[[File:SeleniumSymbol1.30png|center|1.42 × 10<sup>−2</sup>31px|link=Selenium]]|} [[File:{BromineSymbol1.png| widthcenter|31px|link="450"Bromine]]| style="width[[File:50%; height:25px;"KryptonSymbol1.png|center| H<sub>2</sub>O {{eqm}} H<sup>+</sup> + OH<sup>−</sup>31px| ''K'' link= 10<sup>−14</sup> at 25 °CKrypton]]
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=== With varying pH, temperature and salinity: CaCO<sub>3</sub> scaling in swimming pools Meaning===[[File:CaCO3-pH.gif|thumb|alt=Effects of salinity and pH on the maximum calcium ion level before scaling is anticipated at 25 C and 1 mM bicarbonate (e.g. in swimming pools)]][[File:CaCO3-Temp.gif|thumb|alt=Effects of temperature and bicarbonate concentration on the maximum calcium ion level before scaling '''Testpage''' is anticipated at pH 7 and 5,000 ppm salinity (e.g. in swimming pools)]]In contrast to show the open equilibrium scenario above, many swimming pools are managed by addition type of [[sodium bicarbonate]] (NaHCO<sub>3</sub>) to about 2 mM as a buffer, then control of pH through use of HCl, NaHSO<sub>4</sub>, Na<sub>2</sub>CO<sub>3</sub>, NaOH or chlorine formulations errors in this wiki that are acidic or basicneed correcting. In this situation, dissolved inorganic carbon ([[total inorganic carbon]]) is far from equilibrium with atmospheric CO<sub>2</sub>File:AppleFruit. Progress towards equilibrium through outgassing of CO<sub>2</sub> is slowed by (i) the slow reaction [[Carbonic acid|H<sub>2</sub>CO<sub>3</sub>]] ⇌ CO<sub>2</sub>(aq) + H<sub>2</sub>O;<ref>{{Cite journal png| doi = 10.1021/jp909019uright| pmid = 20039712100px| title = Comprehensive Study of the Hydration and Dehydration Reactions of Carbon Dioxide in Aqueous Solutionthumb| journal = The Journal A picture 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..114a dragon.1734W}}</ref> (ii) limited aeration in a deep water column and (iii) periodic replenishment of bicarbonate to maintain buffer capacity (often estimated through measurement of [[alkalinity|‘total alkalinity’]]).
:Ca<sup>2+</sup><sub>max</sub> = (K<sub>sp</sub> / K<sub>a2</sub>) × ('''Testpage is to show the type... *The last sentence should not have been completely in bold type. It happened because the the sentence began with three apostrophe's but more should have been added after the word '''testpage'''.*TestPage is the title of this page so it should never appear as a link to another page. This happens sometimes when brackets [H<sup>+</sup>[ are used because they can also make the page bold, but if the word is spelled wrong or there is a capital in the wrong place it becomes a link.*Some links may be 'red' when there is already a page with a similar name that they should be linked to. [[Pollinate]] / and [HCO<sub>3</sub><sup>−</sup>[Pollination]])can both be fit on the same page so if you know there is a page that it could link to, then it shouldn't be red.
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{| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: white; border-collapse: collapse; border-color: #C0C090;" class="'wikitable"|-! 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>'
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