Wikipedia

Calcium sulfate: Difference between revisions

Browse history interactively
← Previous edit
Content deleted Content added
VisualWikitext
Citation bot (talk | contribs)
Bots
5,220,041 edits
Altered title. Added bibcode. | Use this bot. Report bugs. | Suggested by Abductive | Category:E-number additives | #UCB_Category 235/313
 
(29 intermediate revisions by 18 users not shown)
Line 5:
| verifiedrevid = 443496462
| Name = Calcium sulfate
| ImageFile1= CaSO4simple.svg
| ImageFile = CaSO4.tif
| ImageSize =
| ImageName = Calcium sulphate anhydrous
Line 11:
| ImageSize2 =
| ImageName2 = Calcium sulfate hemihydrate
| OtherNames = Sulfate of lime<br/>[[Plaster of Paris]]<br/>[[Drierite]]<br/>[[Gypsum]]
|Section1={{Chembox Identifiers
| index3_label = (dihydrate)
Line 52:
|Section2={{Chembox Properties
| Density = 2.96 g/cm<sup>3</sup> (anhydrous) <br> 2.32 g/cm<sup>3</sup> (dihydrate)
| Solubility = 0.26 g/100ml at 25&nbsp;°C (dihydrate)<ref>{{cite journal |last1=Lebedev |first1=A. L. |last2=Kosorukov |first2=V. L. |date=2017 |title= Gypsum Solubility in Water at 25°C |url=https://link.springer.com/content/pdf/10.1134/S0016702917010062.pdf |journal=Geochemistry International |volume=55 |issue=2 |pages=171–177 |doi=10.1134/S0016702917010062|bibcode=2017GeocI..55..205L |s2cid=132916752 }}</ref>
| Solvent = [[glycerol]]
| SolubleOther = slightly soluble (dihydrate)
Line 92:
}}
 
'''Calcium sulfate''' (or '''calcium sulphate''') is the inorganic compound with the formula CaSO<sub>4</sub> and related [[hydrate]]s. In the form of γ-[[anhydrite]] (the [[anhydrous]] form), it is used as a [[desiccant]]. One particular hydrate is better known as [[plaster of Paris]], and another occurs naturally as the mineral [[gypsum]]. It has many uses in industry. All forms are white solids that are poorly soluble in water.<ref name=Ullmanns>Franz Wirsching "Calcium Sulfate" in Ullmann's Encyclopedia of Industrial Chemistry, 2012 Wiley-VCH, Weinheim. {{DOIdoi|10.1002/14356007.a04_555}}</ref> Calcium sulfate causes [[permanent hardness]] in water.
 
==Hydration states and crystallographic structures==
[[Image:CaSO4.tif|thumb|left|Structure of the hemihydrate of calcium sulfate reveals a dense network of Ca-O-S bonds. Color code: red (O), green (Ca), orange (S).]]
The compound exists in three levels of hydration corresponding to different crystallographic structures and to minerals:
* {{chem|CaSO|4}} ([[anhydrite]]): anhydrous state.<ref>{{cite journal| doi = 10.1107/S0567740875007145| title = Anhydrite: A refinement| journal = Acta Crystallographica Section B| volume = 31| issue = 8| pages = 2164| year = 1975| last1 = Morikawa| first1 = H.| last2 = Minato| first2 = I.| last3 = Tomita| first3 = T.| last4 = Iwai| first4 = S.| bibcode = 1975AcCrB..31.2164M}}</ref> The structure is related to that of [[zirconium orthosilicate]] (zircon): {{chem|Ca||2+}} is 8-coordinate, {{chem|SO|4|2-}} is tetrahedral, O is 3-coordinate.
* {{chem|CaSO|4|·2H|2|O}} ([[gypsum]] and [[selenite (mineral)]]): dihydrate.<ref>{{cite journal| doi = 10.1107/S0567740874004055| title = A refinement of the crystal structure of gypsum {{chem|CaSO|4|·2H|2|O}}| journal = Acta Crystallographica Section B| volume = 30| issue = 4| pages = 921| year = 1974| last1 = Cole| first1 = W.F.| last2 = Lancucki| first2 = C.J.| doi-access = free}}</ref>
* {{chem|CaSO|4|·{{sfrac|2}}H|2|O}} ([[bassanite]]): hemihydrate, also known as [[plaster of Paris]]. Specific hemihydrates are sometimes distinguished: α-hemihydrate and β-hemihydrate.<ref name=t>Taylor H.F.W. (1990) ''Cement Chemistry''. Academic Press, {{ISBN|0-12-683900-X}}, pp. 186-187.</ref>
 
Line 105 ⟶ 106:
 
===Hydration and dehydration reactions===
<!-- The title is linked from elsewhere -->
With judicious heating, gypsum converts to the partially dehydrated mineral called [[bassanite]] or [[plaster of Paris]]. This material has the formula CaSO<sub>4</sub>·(''n''H<sub>2</sub>O), where 0.5 ≤ ''n'' ≤ 0.8.<ref name=t/> Temperatures between {{cvt|100|and(-)|150|°C}} are required to drive off the water within its structure. The details of the temperature and time depend on ambient humidity. Temperatures as high as {{cvt|170|°C}} are used in industrial calcination, but at these temperatures γ-anhydrite begins to form. The heat energy delivered to the gypsum at this time (the heat of hydration) tends to go into driving off water (as water vapor) rather than increasing the temperature of the mineral, which rises slowly until the water is gone, then increases more rapidly. The equation for the partial dehydration is:
 
Line 120 ⟶ 122:
 
===Food industry===
The calcium sulfate hydrates are used as a [[flocculation|coagulant]] in products such as [[tofu]].<ref>{{cite web |url = http://www.soymilkmaker.com/coagulant.html |title = About tofu coagulant|website = www.soymilkmaker.com|publisher = Sanlinx Inc. |date = 31 August 2015|access-date = 10 January 2008|archive-date = 14 March 2015|archive-url = https://web.archive.org/web/20150314030506/http://www.soymilkmaker.com/coagulant.html|url-status = dead}}</ref>
 
For the [[FDA]], it is permitted in Cheesecheese and Relatedrelated Cheesecheese Productsproducts; Cerealcereal Floursflours; Bakerybakery Productsproducts; Frozenfrozen Dessertsdesserts; Artificialartificial Sweetenerssweeteners for Jellyjelly & Preservespreserves; Condimentcondiment Vegetablesvegetables; and Condimentcondiment Tomatoestomatoes and some candies.<ref name=pc>{{cite web |title=Compound Summary for CID 24497 - Calcium Sulfate |url=https://pubchem.ncbi.nlm.nih.gov/compound/calcium_sulfate#section=Food-Additives-and-Ingredients |publisher=PubChem}}</ref>
 
It is known in the [[E number]] series as '''E516''', and the UN's [[ Food and Agriculture Organization | FAO ]] knows it as a firming agent, a flour treatment agent, a sequestrant, and a leavening agent.<ref name=pc/>
 
=== Dentistry ===
Calcium sulfate has a long history of use in dentistry.<ref>{{Cite journal|last1=Titus|first1=Harry W.|last2=McNally|first2=Edmund|last3=Hilberg|first3=Frank C.|date=1933-01-01|title=Effect of Calcium Carbonate and Calcium Sulphate on Bone Development|journal=Poultry Science|language=en|volume=12|issue=1|pages=5–8|doi=10.3382/ps.0120005|issn=0032-5791|doi-access=free}}</ref> It has been used in bone regeneration as a graft material and graft binder (or extender) and as a barrier in guided bone tissue regeneration. It is a biocompatible material and is completely resorbed following implantation.<ref>{{Cite journal |lastlast1=Thomas |firstfirst1=Mark V. |last2=Puleo |first2=David A. |last3=Al-Sabbagh |first3=Mohanad |date=2005 |title=Calcium sulfate: a review |url=https://pubmed.ncbi.nlm.nih.gov/16393128/ |journal=Journal of Long-Term Effects of Medical Implants |volume=15 |issue=6 |pages=599–607 |doi=10.1615/jlongtermeffmedimplants.v15.i6.30 |issn=1050-6934 |pmid=16393128}}</ref> It does not evoke a significant host response and creates a calcium-rich milieu in the area of implantation.<ref>{{Cite web|date=2020-03-25|title=Biphasic Calcium Sulfate - Overview|url=https://www.augmabio.com/abca/clinical-literature/biphasic-calcium-sulfate-overview/|access-date=2020-07-16|website=Augma Biomaterials}}</ref>
 
===Other usesDesiccant===
[[File:Drierite.jpg|thumb|left|upright|150px|The desiccant Drierite|right]]
When sold at the anhydrous state as a desiccant with a color-indicating agent under the name [[Drierite]], it appears blue (anhydrous) or pink (hydrated) due to impregnation with [[cobalt(II) chloride]], which functions as a moisture indicator.
 
===Sulfuric acid production===
Up to the 1970s, commercial quantities of [[sulfuric acid]] were produced in [[Whitehaven]] ([[Cumbria]], UK) from anhydrous calcium sulfate. Upon being mixed with [[shale]] or [[marl]], and roasted, the sulfate liberates [[sulfur dioxide]] gas, a precursor in [[sulfuric acid]] production, the reaction also produces [[calcium silicate]], a mineral phase essential in [[cement]] [[Clinker (cement)|clinker]] production.<ref>[http://www.lakestay.co.uk/whitehavenmininghistory.html Whitehaven Coast Archeological Survey]</ref>
Up to the 1970s, commercial quantities of [[sulfuric acid]] were produced from anhydrous calcium sulfate.<ref>[https://www.cementkilns.co.uk/cement_kiln_whitehaven.html Whitehaven Cement Plant]</ref> Upon being mixed with [[shale]] or [[marl]], and roasted at 1400°C, the sulfate liberates [[sulfur dioxide]] gas, a precursor to [[sulfuric acid]]. The reaction also produces [[calcium silicate]], used in [[cement]] [[Clinker (cement)|clinker]] production.<ref name="anhydrite process">[https://www.cementkilns.co.uk/cemkilndoc054.html Anhydrite Process]</ref><ref>[https://d28rz98at9flks.cloudfront.net/9554/Rec1949_044.pdf COMMONWEALTH OF AUSTRALIA. DEPARTMENT OF SUPPLY AND SHIPPING. BUREAU OF MINERAL RESOURCES GEOLOGY AND GEOPHYSICS. REPORT NO.1949/44 (Geol. Ser. No. 27) by E.K. Sturmfels THE PRODUCTION OF SULPHURIC ACID AND PORTLAND CEMENT FROM CALCIUM SULPHATE AND ALUMINIUM SILICATES]</ref>
:{{chem2|2 CaSO4 + 2 SiO2 + C → 2 CaSiO3 + 2 SO2 + CO2}}
 
Some component reactions pertaining to calcium sulfate:
: 2 CaSO<sub>4</sub> + 2 SiO<sub>2</sub> → 2 CaSiO<sub>3</sub> + 2 SO<sub>2</sub> + O<sub>2</sub> <ref>[https://d28rz98at9flks.cloudfront.net/9554/Rec1949_044.pdf COMMONWEALTH OF AUSTRALIA. DEPARTMENT OF SUPPLY AND SHIPPING. BUREAU OF MINERAL RESOURCES GEOLOGY AND GEOPHYSICS. REPORT NO.1949/44 (Geol. Ser. No. 27) by E.K. Sturmfels THE PRODUCTION OF SULPHURIC ACID AND PORTLAND CEMENT FROM CALCIUM SULPHATE AND ALUMINIUM SILICATES]</ref>
:{{chem2|CaSO4 + 2 C → CaS + 2 CO2}}
 
:{{chem2|3 CaSO4 + CaS + 2 SiO2 → 2 Ca2SiO4 + 4 SO2}}
The plant made sulfuric acid by the “Anhydrite Process” , in which [[cement clinker]] itself was a by-product. In this process, anhydrite (calcium sulfate) replaces limestone in a cement rawmix, and under reducing conditions, [[sulfur dioxide]] is evolved instead of [[carbon dioxide]]. The sulfur dioxide is converted to sulfuric acid by the [[Contact Process]] using a [[vanadium pentoxide]] [[catalyst]]. <ref>[https://www.cementkilns.co.uk/waste.html#anhydrite Whitehaven anhydrate process]</ref>
:{{chem2|3 CaSO4 + CaS → 4 CaO + 4 SO2}}
 
:{{chem2|Ca2SiO4 + CaO → Ca3OSiO4}}
CaSO<sub>4</sub> + 2 C → CaS + 2CO<sub>2</sub>
 
3 CaSO<sub>4</sub> + CaS + 2 SiO<sub>2</sub> → 2 Ca<sub>2</sub>SiO<sub>4</sub> ([[belite]]) + 4 SO<sub>2</sub>
 
3 CaSO<sub>4</sub> + CaS → 4 CaO + 4 SO<sub>2</sub>
 
Ca<sub>2</sub>SiO<sub>4</sub> + CaO → Ca<sub>3</sub>OSiO<sub>4</sub> ([[alite]])
 
2 SO<sub>2</sub> + O<sub>2</sub> → 2 SO<sub>3</sub>
(in the presence of the [[catalyst]] [[vanadium pentoxide]])
 
SO<sub>3</sub> + H2O → H<sub>2</sub>SO<sub>4</sub> <ref>[https://www.cementkilns.co.uk/waste.html#anhydrite Whitehaven anhydrate process]</ref>
 
Because of its use in an expanding niche market, the Whitehaven plant continued to expand in a manner not shared by the other Anhydrite Process plants. The anhydrite mine opened on 11/1/1955, and the acid plant started on 14/11/1955. For a while in the early 1970s, it became the largest sulfuric acid plant in the UK, making about 13% of national production, and it was by far the largest Anhydrite Process plant ever built. <ref>[https://www.cementkilns.co.uk/cement_kiln_whitehaven.htmlInternet website cement kilns whitehaven]</ref>
 
==Production and occurrence==
Line 158 ⟶ 149:
 
In addition to natural sources, calcium sulfate is produced as a by-product in a number of processes:
*In [[flue-gas desulfurization]], exhaust gases from [[fossil-fuel power station]]s and other processes (e.g. cement manufacture) are scrubbed to reduce their sulfur oxidedioxide content, by injecting finely ground [[limestone]]:<ref>{{cite book |doi=10.1002/0471238961.0701190519160509.a01|chapter=Fuels, Synthetic, Gaseous Fuels |title=Kirk‐OthmerKirk-Othmer Encyclopedia of Chemical Technology |year=2000 |last1=Speight |first1=James G. |isbn=9780471484943 }}</ref>
:{{chem2|SO2 + 0.5 O2 + CaCO3 -> CaSO4 + CO2}}
Related sulfur-trapping methods use [[calcium hydroxide|lime]] and some produces an impure [[calcium sulfite]], which oxidizes on storage to calcium sulfate.
Line 166 ⟶ 157:
*Calcium sulfate can also be recovered and re-used from scrap drywall at construction sites.
 
These precipitation processes tend to concentrate radioactive elements in the calcium sulfate product. This issue is particular with the phosphate by-product, since phosphate ores naturally contain [[uranium]] and its [[decay product]]s such as [[radium-226]], [[lead-210]] and [[polonium-210]]. Extraction of uranium from phosphorus ores can be economical on its own depending on prices on the [[uranium market]] or the separation of uranium can be mandated by environmental legislation and its sale is used to recover part of the cost of the process.<ref>{{citationcite web needed|date url=Aprilhttps://www.osti.gov/biblio/6654998 | osti=6654998 | title=Recovery of uranium from phosphate rocks | last1=Wang | first1=R. D. | last2=Field | first2=L. A. | last3=Gillet d'Auriac | first3=F. S. }}</ref><ref>{{cite web | url=https://world-nuclear.org/information-library/nuclear-fuel-cycle/uranium-resources/uranium-from-phosphates.aspx | title=Uranium from Phosphates &#124; Phosphorite Uranium - World Nuclear Association }}</ref><ref>{{cite web | url=https://www.world-nuclear-news.org/Articles/Agreement-on-Brazilian-uranium-phosphate-extractio | title=Brazil plans uranium-phosphate extraction plant in Santa Quitéria : Uranium & Fuel - World Nuclear News 2022}}</ref>
 
Calcium sulfate is also a common component of [[fouling]] deposits in industrial heat exchangers, because its solubility decreases with increasing temperature (see the specific section on the retrograde solubility).
 
==Solubility==
==Retrograde solubility==
[[File:Temperature dependence calcium sulfate solubility.svg|thumb|400px|left|Temperature dependence of the solubility of calcium sulfate (3 phases) in pure water.]]{{clear-left}}▼
The dissolution of the different crystalline phases of calcium sulfate in water is [[exothermic]] and releases [[heat]] (decrease in [[Enthalpy]]: ΔH < 0). As an immediate consequence, to proceed, the dissolution reaction needs to evacuate this heat that can be considered as a product of reaction. If the system is cooled, the dissolution equilibrium will evolve towards the right according to the [[Le Chatelier principle]] and calcium sulfate will dissolve more easily. Thus the solubility of calcium sulfate increases as the temperature decreases and vice versa. If the temperature of the system is raised, the reaction heat cannot dissipate and the equilibrium will regress towards the left according to Le Chatelier principle. The solubility of calcium sulfate decreases as temperature increases. This counter-intuitive solubility behaviour is called retrograde solubility. It is less common than for most of the salts whose dissolution reaction is [[endothermic]] (i.e., the reaction consumes heat: increase in [[Enthalpy]]: ΔH > 0) and whose solubility increases with temperature. Another calcium compound, [[calcium hydroxide]] (Ca(OH)<sub>2</sub>, [[portlandite]]) also exhibits a retrograde solubility for the same thermodynamic reason: because its dissolution reaction is also exothermic and releases heat. So, to dissolve the maximum amount of calcium sulfate or calcium hydroxide in water, it is necessary to cool the solution down close to its freezing point instead of increasing its temperature.
The solubility of calcium sulfate decreases as temperature increases. This behaviour ("retrograde solubility") is uncommon: dissolution of most of the salts is [[endothermic]] and their solubility increases with temperature.The retrograde solubility of calcium sulfate is also responsible for its precipitation in the hottest zone of heating systems and for its contribution to the formation of [[Fouling#Precipitation fouling|scale]] in [[boiler]]s along with the precipitation of [[calcium carbonate]] whose [[solubility]] also decreases when [[Carbon dioxide|CO<sub>2</sub>]] degasses from hot water or can escape out of the system.▼
 
▲[[File:Temperature dependence calcium sulfate solubility.svg|thumb|400px|left|Temperature dependence of the solubility of calcium sulfate (3 phases) in pure water.]]{{clear-left}}
 
▲The retrograde solubility of calcium sulfate is also responsible for its precipitation in the hottest zone of heating systems and for its contribution to the formation of [[Fouling#Precipitation fouling|scale]] in [[boiler]]s along with the precipitation of [[calcium carbonate]] whose [[solubility]] also decreases when [[Carbon dioxide|CO<sub>2</sub>]] degasses from hot water or can escape out of the system.
 
==On planet Mars==
2011 findings by the [[Opportunity (rover)|''Opportunity'']] rover on the planet [[Mars]] show a form of calcium sulfate in a vein on the surface. Images suggest the mineral is [[gypsum]].<ref>{{cite web|title=NASA Mars Opportunity rover finds mineral vein deposited by water|url=http://www.jpl.nasa.gov/news/news.php?release=2011-377&cid=release_2011-377&msource=11377&tr=y&auid=9976954|publisher=NASA Jet Propulsion Laboratory|access-date=April 23, 2013|date=December 7, 2011}}</ref>
 
==See also==
Retrieved from "https://en.wikipedia.org/wiki/Calcium_sulfate"