|Molar mass||136.14 g/mol (anhydrous)
145.15 g/mol (hemihydrate)
172.172 g/mol (dihydrate)
|Density||2.96 g/cm3 (anhydrous)
2.32 g/cm3 (dihydrate)
|Melting point||1,460 °C (2,660 °F; 1,730 K) (anhydrous)|
|0.21g/100ml at 20 °C (anhydrous)
0.24 g/100ml at 20 °C (dihydrate)
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)
|107 J·mol−1·K−1 |
Std enthalpy of
|EU Index||Not listed|
|US health exposure limits (NIOSH):|
|TWA 15 mg/m3 (total) TWA 5 mg/m3 (resp) [for anhydrous form only]|
|TWA 10 mg/m3 (total) TWA 5 mg/m3 (resp) [anhydrous only]|
LDLH (Immediate danger)
|Plaster of Paris
|Supplementary data page|
|Refractive index (n),
Dielectric constant (εr), etc.
|UV, IR, NMR, MS|
Except where noted otherwise, data is given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
|what is: / ?)(|
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. 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. 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.
Commercial production and recovery
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.
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 stations and other processes (e.g. cement manufacture) are scrubbed to reduce their sulfur oxide content, by injecting finely ground limestone or lime. This 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.
- 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 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 is particularly the case with the phosphate by-product, since phosphate rocks naturally contain actinides.
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 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.
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
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 produces Calcium silicate, a precursor in cement production.
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
- Calcium sulfate (data page)
- Bathybius haeckelii
- Chalk (calcium carbonate)
- Gypsum plaster
- Flue-gas desulfurization
- S. Gangolli (1999). The Dictionary of Substances and Their Effects: C. Royal Society of Chemistry. p. 71. ISBN 0-85404-813-8.
- 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.
- D.R. Linde (ed.) "CRC Handbook of Chemistry and Physics", 83rd Edition, CRC Press, 2002
- Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A21. ISBN 0-618-94690-X.
- "NIOSH Pocket Guide to Chemical Hazards". National Institute for Occupational Safety and Health (NIOSH). id=0095.
- About Tofu Coagulant Retrieved 9 Jan. 2008.
- H F W Taylor, Cement Chemistry, Academic Press, 1990, ISBN 0-12-683900-X, pp. 186-187
- What the heck is plaster anyway?
- Gypsum, USGS, 2008
- "NASA Mars Rover Finds Mineral Vein Deposited by Water". NASA Jet Propulsion Laboratory. December 7, 2011. Retrieved April 23, 2013.
|Salts and the ester of the Sulfate ion|