Calcium sulfate

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Calcium sulfate
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Template:Longitem 7487
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Template:Longitem CaSO4
Molar mass 136.141 g/mol (anhydrous)
145.149 g/mol (hemihydrate)
172.171 g/mol (dihydrate)
Appearance white solid
Odor odorless
Density 2.96 g/cm3 (anhydrous)
2.32 g/cm3 (dihydrate)
Melting point Template:Chembox CalcTemperatures
Template:Longitem 4.93 × 10−5 mol2L−2 (anhydrous)
3.14 × 10−5 (dihydrate)
[2]
Solubility in glycerol slightly soluble (dihydrate)
Acidity (pKa) 10.4 (anhydrous)
7.3 (dihydrate)
Template:Longitem −49.7·10−6 cm3/mol
Template:Longitem orthorhombic
Template:Longitem 107 J·mol−1·K−1 [3]
Template:Longitem −1433 kJ/mol[3]
NFPA 704 (fire diamond) Template:NFPA 704 diamond
Flash point Template:Chembox CalcTemperatures
Template:Longitem Magnesium sulfate
Strontium sulfate
Barium sulfate
Template:Longitem Calcium chloride
Magnesium sulfate
Template:Longitem Plaster of Paris
Gypsum

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Calcium sulfate (or calcium sulphate) is an inorganic salt with the chemical formula Template:Chem/link. It occurs in several hydrated forms; the anhydrous state (known as anhydrite) is a white crystalline solid often found in evaporite deposits. Its dihydrate form is the mineral gypsum, which may be dehydrated to produce bassanite, the hemihydrate state. Gypsum occurs in nature as crystals (selenite) or fibrous masses (satin spar), typically colorless to white, though impurities can impart other hues. All forms of calcium sulfate are sparingly soluble in water[5] and cause permanent hardness when dissolved therein.

Hydration states

Calcium sulfate occurs at three levels of hydration with different crystallographic structures: anhydrous, dihydrate, and hemihydrate.

The anhydrous Template:Chem/link (anhydrite)[6] crystallizes as an tightly-bound orthohombic lattice with space group Pnma, in which each Template:Chem/link is 8-coordinated, or surrounded, by 8 oxygen atoms from tetrahedral Template:Chem/link. It is similar in topology to zircon.

The dihydrate Template:Chem/link (gypsum)[7] forms a monoclinic crystal with space group C2/c. Its structure consists of alternating layers: one with Template:Chem/link coordinated with tetrahedral Template:Chem/link and another with interstitial water molecules.

The hemihydrate Template:Chem/link (bassanite) is also known as plaster of Paris. Specific hemihydrates are sometimes distinguished between α-hemihydrate and β-hemihydrate.[8]

File:CaSO4.tif
The structure of hemihydrate calcium sulfate is a network of Ca-O-S bonds. Color code: red (O), green (Ca), yellow (S).

Uses

Script error: No such module "Labelled list hatnote". The main use of calcium sulfate is to produce plaster of Paris and stucco. These applications exploit the fact that calcium sulfate which has been powdered and calcined forms a moldable paste upon hydration and hardens as crystalline calcium sulfate dihydrate. It is also convenient that calcium sulfate is poorly soluble in water and does not readily dissolve in contact with water after its solidification.

Hydration and dehydration reactions

With judicious heating, gypsum converts to the partially dehydrated mineral called bassanite or plaster of Paris. This material has the formula CaSO4·(nH2O), where 0.5 ≤ n ≤ 0.8.[8] Temperatures between Template:Cvt 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 Template:Cvt 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:

CaSO4 · 2 H2O   →   CaSO4 · Template:Sfrac H2O + Template:Sfrac H2O↑

The endothermic property of this reaction is relevant to the performance of drywall, conferring fire resistance to residential and other structures. In a fire, the structure behind a sheet of drywall will remain relatively cool as water is lost from the gypsum, thus preventing (or substantially retarding) damage to the framing (through combustion of wood members or loss of strength of steel at high temperatures) and consequent structural collapse. But at higher temperatures, calcium sulfate will release oxygen and act as an oxidizing agent. This property is used in aluminothermy. In contrast to most minerals, which when rehydrated simply form liquid or semi-liquid pastes, or remain powdery, calcined gypsum has an unusual property: when mixed with water at normal (ambient) temperatures, it quickly reverts chemically to the preferred dihydrate form, while physically "setting" to form a rigid and relatively strong gypsum crystal lattice:

CaSO4 · Template:Sfrac H2O + Template:Sfrac H2O   →   CaSO4 · 2 H2O

This reaction is exothermic and is responsible for the ease with which gypsum can be cast into various shapes including sheets (for drywall), sticks (for blackboard chalk), and molds (to immobilize broken bones, or for metal casting). Mixed with polymers, it has been used as a bone repair cement. Small amounts of calcined gypsum are added to earth to create strong structures directly from cast earth, an alternative to adobe (which loses its strength when wet). The conditions of dehydration can be changed to adjust the porosity of the hemihydrate, resulting in the so-called α- and β-hemihydrates (which are more or less chemically identical).

On heating to Template:Cvt, the nearly water-free form, called γ-anhydrite (CaSO4·nH2O where n = 0 to 0.05) is produced. γ-Anhydrite reacts slowly with water to return to the dihydrate state, a property exploited in some commercial desiccants. On heating above 250 °C, the completely anhydrous form called β-anhydrite or "natural" anhydrite is formed. Natural anhydrite does not react with water, even over geological timescales, unless very finely ground.

The variable composition of the hemihydrate and γ-anhydrite, and their easy inter-conversion, is due to their nearly identical crystal structures containing "channels" that can accommodate variable amounts of water, or other small molecules such as methanol.

Food industry

The calcium sulfate hydrates are used as a coagulant in products such as tofu.[9]

For the FDA, it is permitted in cheese and related cheese products; cereal flours, bakery products, frozen desserts, artificial sweeteners for jelly & preserves, condiment vegetables, and condiment tomatoes, and some candies.[10]

It is known in the E number series as E516, and the UN's FAO knows it as a firming agent, a flour treatment agent, a sequestrant, and a leavening agent.[10]

Dentistry

Calcium sulfate has a long history of use in dentistry.[11] 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.[12] It does not evoke a significant host response and creates a calcium-rich milieu in the area of implantation.[13]

Desiccant

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The desiccant Drierite

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 from anhydrous calcium sulfate.[14] 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 production.[15][16]

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Some component reactions pertaining to calcium sulfate:

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Production and occurrence

The main sources of calcium sulfate are naturally occurring gypsum and anhydrite, which occur at many locations worldwide as evaporites. These may be extracted by open-cast quarrying or by deep mining. World production of natural gypsum is around 127 million tonnes per annum.[17]

In addition to natural sources, calcium sulfate is produced as a by-product in a number of processes:

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Related sulfur-trapping methods use lime and some produces an impure calcium sulfite, which oxidizes on storage to calcium sulfate.

  • In the production of phosphoric acid from phosphate rock, calcium phosphate is treated with sulfuric acid and calcium sulfate precipitates. The product, called phosphogypsum is often contaminated with impurities making its use uneconomic.
  • In the production of hydrogen fluoride, calcium fluoride is treated with sulfuric acid, precipitating calcium sulfate.
  • In the refining of zinc, solutions of zinc sulfate are treated with hydrated lime to co-precipitate heavy metals such as barium.
  • 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 products 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.[19][20][21]

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

File:Temperature dependence calcium sulfate solubility.svg
Temperature dependence of the solubility of calcium sulfate (3 phases) in pure water.

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 scale in boilers along with the precipitation of calcium carbonate whose solubility also decreases when CO2 degasses from hot water or can escape out of the system.

See also

References

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  1. ↑ Script error: No such module "Citation/CS1".
  2. ↑ D.R. Linde (ed.) "CRC Handbook of Chemistry and Physics", 83rd Edition, CRC Press, 2002
  3. ↑ a b Script error: No such module "citation/CS1".
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  5. ↑ Franz Wirsching "Calcium Sulfate" in Ullmann's Encyclopedia of Industrial Chemistry, 2012 Wiley-VCH, Weinheim. Script error: No such module "CS1 identifiers".
  6. ↑ Script error: No such module "Citation/CS1".
  7. ↑ Script error: No such module "Citation/CS1".
  8. ↑ a b Taylor H.F.W. (1990) Cement Chemistry. Academic Press, Template:ISBN, pp. 186–187.
  9. ↑ Script error: No such module "citation/CS1".
  10. ↑ a b Script error: No such module "citation/CS1".
  11. ↑ Script error: No such module "Citation/CS1".
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  13. ↑ Script error: No such module "citation/CS1".
  14. ↑ Whitehaven Cement Plant
  15. ↑ Anhydrite Process
  16. ↑ 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
  17. ↑ Gypsum Script error: No such module "webarchive"., USGS, 2008
  18. ↑ Script error: No such module "citation/CS1".
  19. ↑ Script error: No such module "citation/CS1".
  20. ↑ Script error: No such module "citation/CS1".
  21. ↑ Script error: No such module "citation/CS1".

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External links

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