Calcium sulfate

From Wikipedia, the free encyclopedia
Jump to: navigation, search
Calcium sulfate
Calcium sulfate hemihydrate
CAS number 7778-18-9 YesY
10034-76-1 (hemihydrate),
10101-41-4 (dihydrate)
PubChem 24928
ChemSpider 22905 YesY
UNII E934B3V59H YesY
KEGG D09201 N
ChEBI CHEBI:31346 YesY
RTECS number WS6920000
Jmol-3D images Image 1
Molecular formula CaSO4
Molar mass 136.14 g/mol (anhydrous)
145.15 g/mol (hemihydrate)
172.172 g/mol (dihydrate)
Appearance white solid
Odor odorless
Density 2.96 g/cm3 (anhydrous)
2.32 g/cm3 (dihydrate)
Melting point 1460 °C (anhydrous)
Solubility in water 0.21g/100ml at 20 °C (anhydrous)[1]
0.24 g/100ml at 20 °C (dihydrate)[2]
Solubility product, Ksp 4.93 × 10−5 mol2L−2 (anhydrous)
3.14 × 10−5 (dihydrate)
Solubility in glycerol slightly soluble (dihydrate)
Acidity (pKa) 10.4 (anhydrous)
7.3 (dihydrate)
Crystal structure orthorhombic
Std molar
107 J·mol−1·K−1 [4]
Std enthalpy of
-1433 kJ/mol[4]
MSDS External MSDS
EU Index Not listed
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 1: Exposure would cause irritation but only minor residual injury. E.g., turpentine Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
Flash point Non-flammable
Related compounds
Other cations Magnesium sulfate
Strontium sulfate
Barium sulfate
Related desiccants Calcium chloride
Magnesium sulfate
Related compounds Plaster of Paris
Supplementary data page
Structure and
n, εr, etc.
Phase behaviour
Solid, liquid, gas
Spectral data UV, IR, NMR, MS
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
 N (verify) (what is: YesY/N?)
Infobox references

Calcium sulfate (or calcium sulphate) is a common laboratory and industrial chemical. In the form of γ-anhydrite (the nearly anhydrous form), it is used as a desiccant. It is also used as a coagulant in products like tofu.[5] In the natural state, unrefined calcium sulfate is a translucent, crystalline white rock. When sold as a color-indicating variant under the name Drierite, it appears blue or pink due to impregnation with Cobalt(II) chloride, which functions as a moisture indicator. The hemihydrate (CaSO4·~0.5H2O) is better known as plaster of Paris, while the dihydrate (CaSO4·2H2O) occurs naturally as gypsum. The anhydrous form occurs naturally as β-anhydrite. Depending on the method of calcination of calcium sulfate dihydrate, specific hemihydrates are sometimes distinguished: alpha-hemihydrate and beta-hemihydrate.[6] They appear to differ only in crystal shape. Alpha-hemihydrate crystals are more prismatic than beta-hemihydrate crystals and, when mixed with water, form a much stronger and harder superstructure.[7]

Commercial production and recovery[edit]

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.[8]

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

These precipitation processes tend to concentrate radioactive elements in the calcium sulfate product. This is particularly the case with the phosphate by-product, since phosphate rocks naturally contain actinides.

Dehydration reactions[edit]

Heating gypsum to between 100 °C and 150 °C (302 °F) partially dehydrates the mineral by driving off approximately 75% of the water contained in its chemical structure. The temperature and time needed depend on ambient partial pressure of H2O. Temperatures as high as 170 °C are used in industrial calcination, but at these temperatures γ-anhydrite begins to form. The reaction for the partial dehydration is:

CaSO4·2H2O + heat → CaSO4·½H2O + 1½H2O (steam)

The partially dehydrated mineral is called calcium sulfate hemihydrate or calcined gypsum (commonly known as plaster of Paris) (CaSO4·nH2O), where n is in the range 0.5 to 0.8.[6]

The dehydration (specifically known as calcination) begins at approximately 80 °C (176 °F), although in dry air, some dehydration will take place already at 50 °C. 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 endothermic property of this reaction is exploited by drywall to confer 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:

CaSO4·½H2O + 1½ H2O → CaSO4·2H2O

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 alpha and beta hemihydrates (which are more or less chemically identical).

On heating to 180 °C, 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.

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

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

Commercial use in the synthesis of sulfuric acid[edit]

Up to the 1970s, commercial quantities of sulfuric acid were produced from the anhydrite of 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 produced Calcium oxide, a precursor in cement production.[9]

Fouling deposits[edit]

Calcium sulfate is a common component of fouling deposits in industrial heat exchangers, because its solubility decreases with increasing temperature (see the figure).

Discovery on Mars[edit]

2011 findings by the Mars Opportunity rover show a form of calcium sulfate in a vein on the surface. Images suggest the mineral is gypsum .[10]

See also[edit]


  1. ^ S. Gangolli (1999). The Dictionary of Substances and Their Effects: C. Royal Society of Chemistry. p. 71. ISBN 0-85404-813-8. 
  2. ^ American Chemical Society (2006). Reagent chemicals: specifications and procedures : American Chemical Society specifications, official from January 1, 2006. Oxford University Press. p. 242. ISBN 0-8412-3945-2. 
  3. ^ D.R. Linde (ed.) "CRC Handbook of Chemistry and Physics", 83rd Edition, CRC Press, 2002
  4. ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A21. ISBN 0-618-94690-X. 
  5. ^ About Tofu Coagulant Retrieved 9 Jan. 2008.
  6. ^ a b H F W Taylor, Cement Chemistry, Academic Press, 1990, ISBN 0-12-683900-X, pp. 186-187
  7. ^ What the heck is plaster anyway?
  8. ^ Gypsum, USGS, 2008
  9. ^
  10. ^ "NASA Mars Rover Finds Mineral Vein Deposited by Water". NASA Jet Propulsion Laboratory. December 7, 2011. Retrieved April 23, 2013. 

External links[edit]