Open main menu

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.

Calcium sulfide
Calcium sulfide
Names
IUPAC name
Calcium sulfide
Other names
Calcium monosulfide,
Hepar calcies,
Sulfurated lime
Oldhamite
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.039.869
EC Number 243-873-5
KEGG
UNII
Properties
CaS
Molar mass 72.143 g/mol
Appearance white crystals
hygroscopic
Density 2.59 g/cm3
Melting point 2,525 °C (4,577 °F; 2,798 K)
hydrolyses
Solubility insoluble in alcohol
reacts with acid
2.137
Structure
Halite (cubic), cF8
Fm3m, No. 225
Octahedral (Ca2+); octahedral (S2−)
Hazards
Main hazards H2S source
Irritant (Xi)
Dangerous for the environment (N)
R-phrases (outdated) R31, R36/37/38, R50
S-phrases (outdated) (S2), S28, S61
NFPA 704
Flammability code 0: Will not burn. E.g., waterHealth code 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g., chloroformReactivity code 3: Capable of detonation or explosive decomposition but requires a strong initiating source, must be heated under confinement before initiation, reacts explosively with water, or will detonate if severely shocked. E.g., fluorineSpecial hazards (white): no codeNFPA 704 four-colored diamond
0
2
3
Related compounds
Other anions
Calcium oxide
Other cations
Magnesium sulfide
Strontium sulfide
Barium sulfide
Related sulfides
Sodium sulfide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is ☑Y☒N ?)
Infobox references

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.

Contents

ProductionEdit

CaS is produced by "carbothermic reduction" of calcium sulfate, which entails the conversion of carbon, usually as charcoal, to carbon dioxide:

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).

CaS is also a byproduct in the Leblanc process, a once major industrial process for producing sodium carbonate. In that process sodium sulfide reacts with calcium carbonate:[1]

Na2S + CaCO3 → CaS + Na2CO3

Millions of tons of this calcium sulfide byproduct was discarded, causing extensive pollution and controversy.[2]

Milk of lime, Ca(OH)2, reacts with elemental sulfur to give a "lime-sulfur", which has been used as an insecticide. The active ingredient is probably a calcium polysulfide, not CaS.[3]

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

It reacts with acids such as hydrochloric acid to release toxic hydrogen sulfide gas.

CaS + 2 HCl → CaCl2 + H2S

Natural occurrenceEdit

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.

See alsoEdit

ReferencesEdit

  1. ^ Christian Thieme (2000). "Sodium Carbonates". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a24_299. ISBN 978-3527306732.
  2. ^ Kiefer, David M. (January 2002). "It was all about alkali". Today's Chemist at Work. 11 (1): 45–6.
  3. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.