Calcium sulfide is the chemical compound with the formula CaS. This white material crystallizes in cubes like rock salt. CaS has been studied as a component in a process that would recycle gypsum, a product of flue-gas desulfurization. Like many salts containing sulfide ions, CaS typically has an odour of H2S, which results from small amount of this gas formed by hydrolysis of the salt.
3D model (JSmol)
|Molar mass||72.143 g/mol|
|Appearance||white crystals |
|Melting point||2,525 °C (4,577 °F; 2,798 K)|
|Solubility||insoluble in alcohol |
reacts with acid
Refractive index (nD)
|Halite (cubic), cF8|
|Fm3m, No. 225|
|Octahedral (Ca2+); octahedral (S2−)|
|Main hazards||H2S source|
Dangerous for the environment (N)
|R-phrases (outdated)||R31, R36/37/38, R50|
|S-phrases (outdated)||(S2), S28, S61|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
In terms of its atomic structure, CaS crystallizes in the same motif as sodium chloride indicating that the bonding in this material is highly ionic. The high melting point is also consistent with its description as an ionic solid. In the crystal, each S2− ion is surrounded by an octahedron of six Ca2+ ions, and complementarily, each Ca2+ ion surrounded by six S2− ions.
- CaSO4 + 2 C → CaS + 2 CO2
and can react further:
- 3 CaSO4 + CaS → 4 CaO + 4 SO2
In the second reaction the sulfate (+6 oxidation state) oxidizes the sulfide (-2 oxidation state) to sulfur dioxide (+4 oxidation state), while it is being reduced to sulfur dioxide itself (+4 oxidation state).
- Na2S + CaCO3 → CaS + Na2CO3
Millions of tons of this calcium sulfide byproduct was discarded, causing extensive pollution and controversy.
Reactivity and usesEdit
Calcium sulfide decomposes upon contact with water, including moist air, giving a mixture of Ca(SH)2, Ca(OH)2, and Ca(SH)(OH).
- CaS + H2O → Ca(SH)(OH)
- Ca(SH)(OH) + H2O → Ca(OH)2 + H2S
- CaS + 2 HCl → CaCl2 + H2S
Oldhamite is the name for mineralogical form of CaS. It is a rare component of some meteorites and has scientific importance in solar nebula research. Burning of coal dumps can also produce the compound.
- Christian Thieme (2000). "Sodium Carbonates". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a24_299. ISBN 978-3527306732.
- Kiefer, David M. (January 2002). "It was all about alkali". Today's Chemist at Work. 11 (1): 45–6.
- Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.