Sulfuryl chloride

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Sulfuryl chloride
Structure and dimensions of sulfuryl chloride
Ball-and-stick model of sulfuryl chloride
Sulphuryl chloride 25ml.jpg
CAS number 7791-25-5 YesY
PubChem 24648
ChemSpider 23050 YesY
EC number 232-245-6
ChEBI CHEBI:29291 YesY
Jmol-3D images Image 1
Molecular formula SO2Cl2
Molar mass 134.9698 g mol−1
Appearance Colorless liquid with a pungent odor. Yellows upon standing.
Density 1.67 g cm−3 (20 °C)
Melting point −54.1 °C (−65.4 °F; 219.1 K)
Boiling point 69.4 °C (156.9 °F; 342.5 K)
Solubility in water hydrolyzes
Solubility miscible with benzene, toluene, chloroform, CCl4, glacial acetic acid
Refractive index (nD) 1.4437 (20 °C) [1]
EU Index 016-016-00-6
EU classification Corrosive (C)
R-phrases R14, R34, R37
S-phrases (S1/2), S26, S30, S45
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas Reactivity code 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g., phosphorus Special hazard W: Reacts with water in an unusual or dangerous manner. E.g., cesium, sodiumNFPA 704 four-colored diamond
Flash point Not flammable
Related compounds
Related sulfuryl halides Sulfuryl fluoride
Related compounds Thionyl chloride
Chlorosulfonic acid
Sulfuric acid
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
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Infobox references

Sulfuryl chloride is an inorganic compound with the formula SO2Cl2. At room temperature, it is a colorless liquid with a pungent odor. Sulfuryl chloride is not found in nature, as can be inferred from its rapid hydrolysis.

Sulfuryl chloride is commonly confused with thionyl chloride, SOCl2. The properties of these two sulfur oxychlorides are quite different: sulfuryl chloride is a source of chlorine whereas thionyl chloride is a source of chloride ions. An alternative IUPAC name is sulfuroyl dichloride.


Sulfur is tetrahedral in SO2Cl2, being bound to two oxygen atoms via bonds intermediate of a dative bond and a polarized double bond (which does not utilize d-orbitals[2]) and to two chlorine atoms via polarized single bonds. The oxidation state of the sulfur atom is +6, as in H2SO4.


SO2Cl2 is prepared by the reaction of sulfur dioxide and chlorine in the presence of a catalyst, such as activated carbon.

SO2 + Cl2 → SO2Cl2

The crude product can be purified by fractional distillation. It is uncommon to prepare SO2Cl2 in the laboratory because it is commercially available. Sulfuryl chloride can also be considered a derivative of sulfuric acid.[3]

Sulfuryl chloride was first prepared in 1838 by the French chemist Henri Victor Regnault.[4]


Sulfuryl chloride reacts with water, releasing hydrogen chloride gas and sulfuric acid:

2 H2O + SO2Cl2 → 2 HCl + H2SO4

SO2Cl2 will also decompose when heated to or above 100 °C, about 30 °C above its boiling point.

Upon standing, SO2Cl2 decomposes to sulfur dioxide and chlorine, which gives the older samples a slightly yellowish color.


Sulfuryl chloride is often used as a source of Cl2. Because it is a pourable liquid, it is considered more convenient than Cl2 to measure, store, and dispense. SO2Cl2 is widely used as a reagent in the conversion of C-H → C-Cl adjacent to activating substituents such as carbonyls and sulfoxides. It also chlorinates alkanes, alkenes, alkynes, aromatics, and epoxides. Such reactions occur under free radical conditions using an initiator such as AIBN. It can also be used to convert disulfides into their corresponding sulfenyl chlorides. SO2Cl2 can also convert alcohols to alkyl chlorides. In industry, sulfuryl chloride is most used in producing pesticides.

SO2Cl2 can also be used to treat wool to prevent shrinking.


SO2Cl2 is toxic, corrosive, and acts as a lachrymator. It can form fuming mixtures with water, as well as donor solvents such as DMSO and DMF.


  1. ^ Patnaik, P. (2002). Handbook of Inorganic Chemicals. McGraw-Hill. ISBN 0-07-049439-8. 
  2. ^ Cunningham, T. P.; Cooper, D. L.; Gerratt, J.; Karadakov, P. B.; Raimondi, M. (1997). "Chemical bonding in oxofluorides of hypercoordinatesulfur". Journal of the Chemical Society, Faraday Transactions 93 (13): 2247–2254. doi:10.1039/A700708F. 
  3. ^ Hogan, C. M. (2011). "Sulfur". In Jorgensen, A.; Cleveland, C. J. Encyclopedia of Earth. Washington DC: National Council for Science and the Environment. 
  4. ^ See:

Further reading[edit]