|
|
| (40 intermediate revisions by the same user not shown) |
| Line 1: |
Line 1: |
| − | {{Use dmy dates|date=February 2015}}
| + | ==Top title on a page is surrounded by a double 'equals' sign== |
| − | {{chembox
| + | ===First subtitle is surrounded by a triple 'equals' sign=== |
| − | | Verifiedfields = changed
| + | ====Second subtitle is surrounded by a quadruple 'equals' sign==== |
| − | | Watchedfields = changed
| |
| − | | verifiedrevid = 477003420
| |
| − | | Name = Calcium carbonate
| |
| − | | ImageFileL1 = calcium carbonate.png
| |
| − | | ImageFileR1 = Calcium-carbonate-xtal-3D-SF.png
| |
| − | | ImageFile2 = Calcium carbonate.jpg
| |
| − | | IUPACName = Calcium carbonate
| |
| − | | OtherNames = [[calcite]]; [[aragonite]]; [[chalk]]; [[Lime (material)]]; [[Limestone]]; [[marble]]; [[oyster]]; [[pearl]];
| |
| − | |Section1={{Chembox Identifiers
| |
| − | | UNII_Ref = {{fdacite|correct|FDA}}
| |
| − | | UNII = H0G9379FGK
| |
| − | | ChEMBL_Ref = {{ebicite|changed|EBI}}
| |
| − | | ChEMBL = 1200539
| |
| − | | KEGG_Ref = {{keggcite|correct|kegg}}
| |
| − | | KEGG = D00932
| |
| − | | InChI = 1/CH2O3.Ca/c2-1(3)4;/h(H2,2,3,4);/q;+2/p-2
| |
| − | | ChEBI_Ref = {{ebicite|correct|EBI}}
| |
| − | | ChEBI = 3311
| |
| − | | SMILES = [Ca+2].[O-]C([O-])=O
| |
| − | | InChIKey = VTYYLEPIZMXCLO-NUQVWONBAS
| |
| − | | SMILES1 = C(=O)([O-])[O-].[Ca+2]
| |
| − | | StdInChI_Ref = {{stdinchicite|correct|chemspider}}
| |
| − | | StdInChI = 1S/CH2O3.Ca/c2-1(3)4;/h(H2,2,3,4);/q;+2/p-2
| |
| − | | StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
| |
| − | | StdInChIKey = VTYYLEPIZMXCLO-UHFFFAOYSA-L
| |
| − | | CASNo = 471-34-1
| |
| − | | CASNo_Ref = {{cascite|correct|CAS}}
| |
| − | | ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
| |
| − | | ChemSpiderID = 9708
| |
| − | | EINECS = 207-439-9
| |
| − | | PubChem = 10112
| |
| − | | RTECS = FF9335000
| |
| − | }}
| |
| − | |Section2={{Chembox Properties
| |
| − | | Formula = CaCO<sub>3</sub>
| |
| − | | MolarMass = 100.0869 g/mol
| |
| − | | Appearance = Fine white powder; chalky taste
| |
| − | | Odor = odorless
| |
| − | | Density = 2.711 g/cm<sup>3</sup> ([[calcite]])<br />2.83 g/cm<sup>3</sup> ([[aragonite]])
| |
| − | | Solubility = 0.013 g/L (25 °C)<ref>{{cite book|title=SI Chemical Data Book (4th ed.) |publisher=John Wiley & Sons Australia, Ltd. |author1=Aylward, Gordon |author2=Findlay, Tristan |isbn=978-0-470-81638-7|year=2008}}</ref><ref>{{cite book|title=Calcium Carbonate: From the Cretaceous Period Into the 21st Century|publisher=Springer Science & Business Media|year=2001|url=https://books.google.com/books?id=pbkKGa19k5QC&pg=RA1-PR2|author=Rohleder, J. |author2=Kroker, E. |isbn=3-7643-6425-4}}</ref>
| |
| − | | SolubilityProduct = 3.3{{e|-9}}<ref>{{cite book|last =Benjamin|first=Mark M.|year=2002|title=Water Chemistry |publisher=McGraw-Hill|isbn =0-07-238390-9|url=https://books.google.com/?id=67anQgAACAAJ}}</ref>
| |
| − | | Solvent = dilute acids
| |
| − | | SolubleOther = soluble
| |
| − | | MeltingPt = 1,339 °C (2,442 °F; 1,612 K) (calcite) <br> 825 °C (1517 °F; 1,098 K) (aragonite) <ref>{{cite web|url=https://www.cdc.gov/niosh/docs/81-123/pdfs/0090.pdf|title=Occupational safety and health guideline for calcium carbonate|publisher=US Dept. of Health and Human Services|accessdate=31 March 2011}}</ref>
| |
| − | | BoilingPt = decomposes
| |
| − | | RefractIndex = 1.59
| |
| − | | pKa = 9.0
| |
| − | | pKb =
| |
| − | | MagSus = -38.2·10<sup>−6</sup> cm<sup>3</sup>/mol
| |
| − | }}
| |
| − | |Section3={{Chembox Structure
| |
| − | | CrystalStruct = Trigonal
| |
| − | | SpaceGroup = <span style="text-decoration: overline">3</span>2/m
| |
| − | }}
| |
| − | |Section5={{Chembox Thermochemistry
| |
| − | | 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>
| |
| − | | Entropy = 93 J·mol<sup>−1</sup>·K<sup>−1</sup><ref name=b1 />
| |
| − | }}
| |
| − | |Section6={{Chembox Pharmacology
| |
| − | | ATCCode_prefix = A02
| |
| − | | ATCCode_suffix = AC01
| |
| − | | ATC_Supplemental = {{ATC|A12|AA04}}
| |
| − | }}
| |
| − | |Section7={{Chembox Hazards
| |
| − | | ExternalSDS = [http://www.inchem.org/documents/icsc/icsc/eics1193.htm ICSC 1193]
| |
| − | | MainHazards =
| |
| − | | NFPA-H = 0
| |
| − | | NFPA-F = 0
| |
| − | | NFPA-R = 0
| |
| − | | NFPA-S =
| |
| − | | RPhrases =
| |
| − | | SPhrases =
| |
| − | | LD50 = 6450 mg/kg (oral, rat)
| |
| − | | PEL = TWA 15 mg/m<sup>3</sup> (total) TWA 5 mg/m<sup>3</sup> (resp)<ref>{{PGCH|0090}}</ref>
| |
| − | }}
| |
| − | |Section8={{Chembox Related
| |
| − | | OtherAnions = [[Calcium bicarbonate]]
| |
| − | | OtherCations = [[Magnesium carbonate]]<br />[[Strontium carbonate]]<br />[[Barium carbonate]]
| |
| − | | OtherCompounds = [[Calcium sulfate]]
| |
| − | }}
| |
| − | }}
| |
| − | [[File:Calcite.png|thumb|right|Crystal structure of calcite]]
| |
| | | | |
| − | '''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)|rocks]] as the [[mineral]]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. | + | '''Text in bold is surrounded by triple aporstrophe's''' |
| | + | ''Italicised text is surrounded by double aporstrophe's'' |
| | | | |
| − | ==Chemistry==
| + | To turn a word into a link use [[ this double bracket at the start and end. |
| − | Calcium carbonate shares the typical properties of other carbonates. Notably,
| + | To turn a word into a link, but have it link to a different page then use the | symbol to separate the word to be displayed from the title of the page. |
| − | * it reacts with [[acid]]s, releasing [[carbon dioxide]]:
| + | For example (I've added spaces to demonstrate what it would look like) [ [ evaporation | evaporating ] ]. Removing the spaces would make a link to the page called 'evaporation' while displaying the word 'evaporating' |
| − | :CaCO<sub>3</sub>(s) + 2H<sup>+</sup>(aq) → Ca<sup>2+</sup>(aq) + CO<sub>2</sub>(g) + H<sub>2</sub>O (l)
| + | [[Evaporation|evaporating]] |
| − | * it releases carbon dioxide upon 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:
| |
| − | :CaCO<sub>3</sub> (s) → CaO (s) + CO<sub>2</sub> (g)
| |
| | | | |
| − | Calcium carbonate will react with water that is saturated with carbon dioxide to form the soluble [[calcium bicarbonate]].
| + | To start a new paragraph you can simply press enter twice |
| − | :CaCO<sub>3</sub> + CO<sub>2</sub> + H<sub>2</sub>O → Ca(HCO<sub>3</sub>)<sub>2</sub>
| |
| | | | |
| − | This reaction is important in the [[erosion]] of [[carbonate rock]], forming [[cavern]]s, and leads to [[hard water]] in many regions.
| + | : Or you can use : at the start of each new paragraph which will also indent it. |
| | | | |
| − | An unusual form of calcium carbonate is the hexahydrate, [[ikaite]], CaCO<sub>3</sub>·6H<sub>2</sub>O. Ikaite is stable only below 6 °C.
| + | : The * symbol can be used as a bullet point at the start of a line. |
| | | | |
| − | ==Preparation==
| + | * Like this |
| − | 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]]).
| + | **or this for a subclause of the first bullet point |
| | + | **and again |
| | + | **and again |
| | + | ***or this for a subclause of the third subclause of the original bullet point |
| | | | |
| − | 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):<ref name="PCC">{{cite web|title = Precipitated Calcium Carbonate |accessdate = 11 January 2014|url = http://www.lime.org/uses_of_lime/other_uses/precip_cc.asp}}</ref>
| + | The # symbol can be used to number things |
| | + | #like this |
| | + | #or this |
| | + | #or this |
| | + | ##or a subclause |
| | + | ##like this |
| | + | ##and this |
| | | | |
| − | : CaO + H<sub>2</sub>O → Ca(OH)<sub>2</sub>
| + | Finally you can add all of these together |
| − | :<chem>Ca(OH)2 + CO2 -> CaCO3(v) + H2O</chem>
| |
| | | | |
| − | ==Structure==
| + | :*# like this |
| − | The thermodynamically stable form of CaCO<sub>3</sub> under normal conditions is hexagonal β-CaCO<sub>3</sub>, (the mineral [[calcite]]).<ref name ="Ropp">{{cite book|last=R C Ropp Elsevier|title=Encyclopedia of the alkaline earth compounds|publisher=Elsevier|isbn=9780444595508|pages=359–370}}</ref> Other forms can be prepared, the denser,(2.83 g/cc) orthorhombic λ-CaCO<sub>3</sub> ( the mineral [[aragonite]]) and μ-CaCO<sub>3</sub>, occurring as the mineral [[vaterite]].<ref name ="Ropp"/> The aragonite form can be prepared by precipitation at temperatures above 85 °C, the vaterite form can be prepared by precipitation at 60 °C.<ref name ="Ropp"/> Calcite contains calcium atoms coordinated by 6 oxygen atoms, in aragonite they are coordinated by 9 oxygen atoms.<ref name ="Ropp"/> The vaterite structure is not fully understood.<ref name="DemichelisRaiteri2013">{{cite journal|last1=Demichelis|first1=Raffaella|last2=Raiteri|first2=Paolo|last3=Gale|first3=Julian D.|last4=Dovesi|first4=Roberto|title=The Multiple Structures of Vaterite|journal=Crystal Growth & Design|volume=13|issue=6|year=2013|pages=2247–2251|issn=1528-7483|doi=10.1021/cg4002972}}</ref> Magnesium carbonate MgCO<sub>3</sub> has the calcite structure, whereas strontium and barium carbonate (SrCO<sub>3</sub> and BaCO<sub>3</sub>) adopt the aragonite structure, reflecting their larger ionic radii.<ref name ="Ropp"/>
| + | :*# and this. |
| | | | |
| − | ==Occurrence== | + | ==Key Stage 1== |
| − | | + | {| border="0" style="border-collapse:collapse" |
| − | [[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===
| + | |[[File:HydrogenSymbol1.png|center|31px|link=Hydrogen]] |
| − | 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
| + | |[[File:HeliumSymbol1.png|center|31px|link=Helium]] |
| − | | 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 |issue=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 name="twoturn" /> This type of preservation conserves high levels of detail, even down to the microscopic level. However, it also leaves specimens vulnerable to [[weathering]] when exposed to the surface.<ref name="twoturn">Trexler, D. (2001) [https://books.google.com/books?id=mgc6CS4EUPsC&pg=PA98 "Two Medicine Formation, Montana: geology and fauna"], 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 D.jpg|thumb|500-milligram calcium supplements made from calcium carbonate]] | |
| − | Calcium carbonate is widely used medicinally as an inexpensive dietary calcium supplement for [[antacid|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 = 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 [[hyperphosphatemia]] (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 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=111–3}}</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 [[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=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>
| |
| | |- | | |- |
| − | |'''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 | + | |[[File:LithiumSymbol1.png|center|31px|link=Lithium]] |
| | + | |[[File:BerylliumSymbol1.png|center|31px|link=Beryllium]] |
| | + | | |
| | + | | |
| | + | | |
| | + | | |
| | + | | |
| | + | | |
| | + | | |
| | + | | |
| | + | | |
| | + | | |
| | + | |[[File:BoronSymbol1.png|center|31px|link=Boron]] |
| | + | |[[File:CarbonSymbol1.png|center|31px|link=Carbon]] |
| | + | |[[File:NitrogenSymbol1.png|center|31px|link=Nitrogen]] |
| | + | |[[File:OxygenSymbol1.png|center|31px|link=Oxygen]] |
| | + | |[[File:FluorineSymbol1.png|center|31px|link=Fluorine]] |
| | + | |[[File:NeonSymbol1.png|center|31px|link=Neon]] |
| | |- | | |- |
| − | |'''T (°C)'''||550||587||605||680||727||748||777||800||830||852||871||881||891||898||937||1082||1241 | + | |[[File:SodiumSymbol1.png|center|31px|link=Sodium]] |
| − | |}
| + | |[[File:MagnesiumSymbol1.png|center|31px|link=Magnesium]] |
| − | | + | | |
| − | ==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="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
| + | |[[File:AluminiumSymbol1.png|center|31px|link=Aluminium]] |
| − | |}
| + | |[[File:SiliconSymbol1.png|center|31px|link=Silicon]] |
| − | | + | |[[File:PhosphorusSymbol1.png|center|31px|link=Phosphorus]] |
| − | 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:
| + | |[[File:SulphurSymbol1.png|center|31px|link=Sulphur]] |
| − | | + | |[[File:ChlorineSymbol1.png|center|31px|link=Chlorine]] |
| − | :{| width="500"
| + | |[[File:ArgonSymbol1.png|center|31px|link=Argon]] |
| − | | 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: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 chem>\frac{P_{\ce{CO2}}}{[\ce{CO2}]}\ =\ k_\ce{H}</math> | |
| − | | where ''k''<sub>H</sub> = 29.76 atm/(mol/L) at 25 °C ([[Henry's law|Henry constant]]), <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]]). 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
| |
| − | {| class="wikitable floatright"
| |
| − | |+ {{chembox header}} |Calcium ion solubility as a function of [[carbon dioxide|CO<sub>2</sub>]] [[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> ||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> | + | |[[File:PotassiumSymbol1.png|center|31px|link=Potassium]] |
| − | |- | + | |[[File:CalciumSymbol1.png|center|31px|link=Calcium]] |
| − | | 10<sup>−3</sup> ||7.96||6.62 × 10<sup>−4</sup> | + | |[[File:ScandiumSymbol1.png|center|31px|link=Scandium]] |
| − | |- | + | |[[File:TitaniumSymbol1.png|center|31px|link=Titanium]] |
| − | | 10<sup>−2</sup> ||7.30||1.42 × 10<sup>−3</sup> | + | |[[File:VanadiumSymbol1.png|center|31px|link=Vanadium]] |
| − | |- | + | |[[File:ChromiumSymbol1.png|center|31px|link=Chromium]] |
| − | | 10<sup>−1</sup> ||6.63||3.05 × 10<sup>−3</sup> | + | |[[File:ManganeseSymbol1.png|center|31px|link=Manganese]] |
| − | |- | + | |[[File:IronSymbol1.png|center|31px|link=Iron]] |
| − | | 1 ||5.96||6.58 × 10<sup>−3</sup> | + | |[[File:CobaltSymbol1.png|center|31px|link=Cobalt]] |
| − | |- | + | |[[File:NickelSymbol1.png|center|31px|link=Nickel]] |
| − | | 10 ||5.30||1.42 × 10<sup>−2</sup> | + | |[[File:CopperSymbol1.png|center|31px|link=Copper]] |
| − | |} | + | |[[File:ZincSymbol1.png|center|31px|link=Zinc]] |
| − | | + | |[[File:GalliumSymbol1.png|center|31px|link=Gallium]] |
| − | :{| width="450" | + | |[[File:GermaniumSymbol1.png|center|31px|link=Germanium]] |
| − | | style="width:50%; height:25px;"| H<sub>2</sub>O {{eqm}} H<sup>+</sup> + OH<sup>−</sup> | + | |[[File:ArsenicSymbol1.png|center|31px|link=Arsenic]] |
| − | | ''K'' = 10<sup>−14</sup> at 25 °C | + | |[[File:SeleniumSymbol1.png|center|31px|link=Selenium]] |
| | + | |[[File:BromineSymbol1.png|center|31px|link=Bromine]] |
| | + | |[[File:KryptonSymbol1.png|center|31px|link=Krypton]] |
| | |} | | |} |
| | | | |
| − | (which is true for all aqueous solutions), and the fact that the solution must be electrically neutral,
| + | [https://mediawiki.org MediaWiki] |
| − | | |
| − | :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 initial water solvent pH is not neutral, the equation is modified).
| |
| − | | |
| − | The table on the right shows the result for [Ca<sup>2+</sup>] and [H<sup>+</sup>] (in the form of 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 chem>P_{\ce{CO2}}</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 chem>P_{\ce{CO2}} = 10^{-12} \mathrm{atm}</math>, 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, 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)]]
| + | '''Testpage''' is to show the type of errors in this wiki that need correcting.[[File:AppleFruit.png|right|100px|thumb|A picture of a dragon.]] |
| − | [[File:CaCO3-Temp.gif|thumb|alt=Effects of temperature and bicarbonate concentration on the maximum calcium ion level before scaling is anticipated at pH 7 and 5,000 ppm salinity (e.g. in swimming pools)]]
| |
| − | In contrast to the open equilibrium scenario above, many swimming pools are managed by addition 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 that are acidic or basic. In this situation, dissolved inorganic carbon ([[total inorganic carbon]]) is far from equilibrium with atmospheric CO<sub>2</sub>. 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 | doi = 10.1021/jp909019u| pmid = 20039712| title = Comprehensive Study of the Hydration and Dehydration Reactions of Carbon Dioxide in Aqueous Solution| journal = The Journal of Physical Chemistry A| volume = 114| issue = 4| pages = 1734–40| year = 2010| last1 = Wang | first1 = X. | last2 = Conway | first2 = W. | last3 = Burns | first3 = R. | last4 = McCann | first4 = N. | last5 = Maeder | first5 = M. | bibcode = 2010JPCA..114.1734W}}</ref> (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’]]).
| |
| | | | |
| − | In this situation, the dissociation constants for the much faster reactions H<sub>2</sub>CO<sub>3</sub> ⇌ H<sup>+</sup> + HCO<sub>3</sub><sup>‾</sup> ⇌ 2 H<sup>+</sup> + CO<sub>3</sub><sup>2−</sup> allow the prediction of concentrations of each dissolved inorganic carbon species in solution, from the added concentration of HCO<sub>3</sub><sup>−</sup> (which constitutes more than 90% of [[Bjerrum plot]] species from pH 7 to pH 8 at 25 °C in fresh water).<ref name="Mook 2000">Mook, W. (2000) [http://www-naweb.iaea.org/napc/ih/documents/global_cycle/vol%20I/cht_i_09.pdf "Chemistry of carbonic acid in water"], pp. 143–165 in ''Environmental Isotopes in the Hydrological Cycle: Principles and Applications''. INEA/UNESCO: Paris.</ref> Addition of HCO<sub>3</sub><sup>−</sup> will increase CO<sub>3</sub><sup>2−</sup> concentration at any pH. Rearranging the equations given above, we can see that [Ca<sup>2+</sup>] = Ksp / [CO<sub>3</sub><sup>2−</sup>], and [CO<sub>3</sub><sup>2−</sup>] = K<sub>a2</sub> × [HCO<sub>3</sub><sup>−</sup>] / [H<sup>+</sup>]. Therefore, when HCO<sub>3</sub><sup>−</sup> concentration is known, the maximum concentration of Ca<sup>2+</sup> ions before scaling through CaCO<sub>3</sub> precipitation can be predicted from the formula:
| + | ===About Error checking in Testpage==== |
| | + | *There is an extra '=' in the title 'About Error checking in Testpage'. |
| | + | *In the last sentence '''testpage''' should have been in bold]] because it is the title of this page. |
| | + | *The ']]' show that the word 'bold' was supposed to be a link but the first brackets were not included. |
| | | | |
| − | :Ca<sup>2+</sup><sub>max</sub> = (K<sub>sp</sub> / K<sub>a2</sub>) × ([H<sup>+</sup>] / [HCO<sub>3</sub><sup>−</sup>]) | + | : '''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 [[ 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 [[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. |
| | | | |
| − | The solubility product for CaCO<sub>3</sub> (K<sub>sp</sub>) and the dissociation constants for the dissolved inorganic carbon species (including K<sub>a2</sub>) are all substantially affected by temperature and [[salinity]],<ref name="Mook 2000" /> with the overall effect that Ca<sup>2+</sup><sub>max</sub> increases from fresh to salt water, and decreases with rising temperature, pH, or added bicarbonate level, as illustrated in the accompanying graphs.
| + | ***Bullet points should only have one bullet, not two in a row. |
| | | | |
| − | The trends are illustrative for pool management, but whether scaling occurs also depends on other factors including interactions with Mg<sup>2+</sup>, B(OH)<sub>4</sub><sup>−</sup> and other ions in the pool, as well as supersaturation effects.<ref>{{cite journal|author=Wojtowicz, 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}}</ref><ref>{{cite journal|author=Wojtowicz, 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}}</ref> Scaling is commonly observed in electrolytic chlorine generators, where there is a high pH near the cathode surface and scale deposition further increases temperature. This is one reason that some pool operators prefer borate over bicarbonate as the primary pH buffer, and avoid the use of pool chemicals containing calcium.<ref>Birch, R. G. (2013) [http://members.iinet.net.au/~jorobbirch/BABES.pdf BABES: a better method than "BBB" for pools with a salt-water chlorine generator.] iinet.net.au</ref>
| + | Sometimes a sentence is far too long, or they have some parts to them which make it more complicated, or too many different ways of saying the same thing and this can make it difficult for people, especially young children, to read the sentence and get the meaning from that sentence because there's too much going on and maybe they're a little tired or have difficulty staring at a screen, which makes it harder to pay attention, so you lose track of what the sentence was originally trying to say before you reach the end and have to go back over it to make sure what you read makes sense. |
| − | | + | *That sentence is way too long and complicated. If you have to go back to re-read anything, then it is too complicated and needs correcting. |
| − | ===Solubility in a strong or weak acid solution===
| + | *Most pictures and diagrams will have a description. If the description doesn't match the image then it needs correcting. |
| − | 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.
| + | *Some pictures are too small to see the important detail. Some may be too large to fit on the screen. These need to be resized. |
| − | * 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): | + | *Some tables with pictures in have small pictures but a large box: |
| − | | + | {| class='wikitable' |
| − | {| 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)
| + | |[[File:MusicalInstruments.png|center|200px]] |
| − | | 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
| + | |I did a '''test''' to find out which instrument makes the deepest sound. |
| − | | 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><br />(g/[[liter|L]] of acid)
| |
| − | | 50.0||5.00||0.514||0.0849||0.0504||0.0474||0.0471||0.0470||0.0470
| |
| | |} | | |} |
| | | | |
| − | where the initial state is the acid solution with no 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.
| + | The table should look like this: |
| − | * 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;" 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>
| |
| − | |-
| |
| − | ! 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><br />(g/[[liter|L]] of acid)
| |
| − | | 49.5||4.99||0.513||0.0848||0.0504||0.0474||0.0471||0.0470||0.0470
| |
| − | |}
| |
| − | For the same total acid concentration, the initial pH of the weak acid is less acid than 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;" class="wikitable" | + | {| 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>
| |
| | |- | | |- |
| − | ! width="160" {{chembox header}} |Initial pH
| + | |[[File:MusicalInstruments.png|center|200px]] |
| − | | 1.08||1.62||2.25||3.05||4.01||5.00||5.97||6.74||7.00
| |
| | |- | | |- |
| − | ! width="160" {{chembox header}} |Final pH
| + | | style="height:20px; width:200px; text-align:center;" |I did a '''test''' to find out which instrument makes the deepest sound. |
| − | | 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><br />(g/[[liter|L]] 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]]).
| + | <ul id="nav"><li><span class="plainlinks">'''[[Main_Page|Shortcuts]]'''</span><ul><li><span class="plainlinks">'''[[Carnivore]]'''</span></li><li><span class="plainlinks">'''[[Herbivore]]'''</span></li><li><span class="plainlinks">'''[[Reptile]]'''</span></li><li><span class="plainlinks">'''[[Amphibian]]'''</span></li><li><span class="plainlinks">'''[[Mammal]]'''</span></li></ul></li></ul></ul> |
| − | | |
| − | ==See also==
| |
| − | {{div col|colwidth=22em}}
| |
| − | * [[Cuttlebone]]
| |
| − | * [[Cuttlefish]]
| |
| − | * [[Gesso]]
| |
| − | * [[Limescale]]
| |
| − | * [[Marble]]
| |
| − | * [[Ocean acidification]]
| |
| − | {{div col end}}
| |
| − | | |
| − | ==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]
| |
| − | * [https://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}}
| |
| − | | |
| − | {{Authority control}}
| |
| − | | |
| − | {{DEFAULTSORT:Calcium Carbonate}}
| |
| − | [[Category:Calcium compounds]]
| |
| − | [[Category:Carbonates]]
| |
| − | [[Category:Limestone]]
| |
| − | [[Category:Phosphate binders]]
| |
| − | [[Category:Excipients]]
| |
| − | [[Category:Antacids]]
| |
| − | [[Category:Food stabilizers]]
| |
