Disulfur dichloride

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Disulfur dichloride
Wireframe model of disulfur dichloride
Ball and stick model of disulfur dichloride
Spacefill model of disulfur dichloride
IUPAC name
Disulfur dichloride
Systematic IUPAC name
Other names

Dimeric sulfenic chloride

Sulfur monochloride
3D model (JSmol)
ECHA InfoCard 100.030.021
EC Number
  • 233-036-2
MeSH Sulfur+monochloride
RTECS number
  • WS4300000
UN number 3390
Molar mass 135.04 g/mol
Appearance Light-amber to yellow-red, oily liquid[1]
Odor pungent, nauseating, irritating[1]
Density 1.688 g/cm3
Melting point −80 °C (−112 °F; 193 K)
Boiling point 137.1 °C (278.8 °F; 410.2 K)
decomposes, with loss of HCl
Solubility soluble in ethanol, benzene, ether, chloroform, CCl4[2]
Vapor pressure 7 mmHg (20 °C)[1]
−62.2·10−6 cm3/mol
C2, gauche
1.60 D [2]
Safety data sheet ICSC 0958
Toxic (T)
Harmful (Xn)
Corrosive (C)
Dangerous for the environment (N)
R-phrases (outdated) R14, R20, R25, R29, R35, R50
S-phrases (outdated) (S1/2), S26, S36/37/39, S45, S61
NFPA 704 (fire diamond)
Flammability code 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g. canola oilHealth code 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g. chloroformReactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g. calciumSpecial hazards (white): no codeNFPA 704 four-colored diamond
Flash point 118.5 °C (245.3 °F; 391.6 K)
234 °C (453 °F; 507 K)
Lethal dose or concentration (LD, LC):
150 ppm (mouse, 1 min)[3]
NIOSH (US health exposure limits):
PEL (Permissible)
TWA 1 ppm (6 mg/m3)[1]
REL (Recommended)
C 1 ppm (6 mg/m3)[1]
IDLH (Immediate danger)
5 ppm[1]
Related compounds
Related sulfur chlorides
Sulfur dichloride
Thionyl chloride
Sulfuryl chloride
Related compounds
Disulfur difluoride
Disulfur dibromide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

Disulfur dichloride is the chemical compound of sulfur and chlorine with the formula S2Cl2.[4][5][6][7]

Some alternative names for this compound are sulfur monochloride (the name implied by its empirical formula, SCl), disulphur dichloride (British English Spelling) and sulphur monochloride (British English Spelling). S2Cl2 has the structure implied by the formula Cl−S−S−Cl, wherein the angle between the Cla−S−S and S−S−Clb planes is 90°. This structure is referred to as gauche, and is akin to that for H2O2. A different isomer of S2Cl2 is S=SCl2; this isomer forms transiently when S2Cl2 is exposed to UV-radiation (see thiosulfoxides).

This substance is listed in Schedule 3 Part B – Precursor Chemicals of the Chemical Weapons Convention (CWC). Facilities that produce and/or process and/or consume Scheduled chemicals may be subject to control, reporting mechanisms and inspection by the OPCW (Organisation for the Prohibition of Chemical Weapons).

Synthesis, basic properties, reactions[edit]

Pure disulfur dichloride is a yellow liquid that smokes in air due to reaction with water:

16 S2Cl2 + 16 H2O → 8 SO2 + 32 HCl + 3 S8

It is synthesized by partial chlorination of elemental sulfur. The reaction takes place at usable rates at room temperature. In the laboratory, chlorine gas is led into a flask containing elemental sulfur. As disulfur dichloride is formed, the contents become a golden yellow liquid:[8]

S8 + 4 Cl2 → 4 S2Cl2 ΔH = −58.2 kJ/mol

Excess chlorine produces sulfur dichloride which causes the liquid to become less yellow and more orange-red:

S2Cl2 + Cl2 ↔ 2 SCl2 ΔH = −40.6 kJ/mol

The reaction is reversible, and upon standing, SCl2 releases chlorine to revert to the disulfur dichloride. Disulfur dichloride has the ability to dissolve large quantities of sulfur, which reflects in part the formation of polysulfanes:

S2Cl2 + n S → S2+nCl2

Pure disulfur dichloride is obtained by distilling the yellow-orange liquid over excess elemental sulfur.

S2Cl2 also arises from the chlorination of CS2 as in the synthesis of thiophosgene.


S2Cl2 hydrolyzes to sulfur dioxide and elemental sulfur. When treated with hydrogen sulfide, polysulfanes are formed as indicated in the following idealized formula:

2 H2S + S2Cl2 → H2S4 + 2 HCl

It reacts with ammonia to give heptasulfur imide (S7NH) and related S−N rings S8−x(NH)x (x = 2, 3).


S2Cl2 has been used to introduce C−S bonds. In the presence of aluminium chloride (AlCl3), S2Cl2 reacts with benzene to give diphenyl sulfide:

S2Cl2 + 2 C6H6 → (C6H5)2S + 2 HCl + 1/8 S8

Anilines react with S2Cl2 in the presence of NaOH via the so-called Herz reaction to give ortho-aminothiophenolates. These species are precursors to thioindigo dyes. It is also used to prepare the sulfur mustard "gas" by reaction with ethylene at 60 °C (the Levinstein process):

S2Cl2 + 2 C2H4 → (ClC2H4)2S + 1/8 S8

Other uses of S2Cl2 include the manufacture of sulfur dyes, insecticides, and synthetic rubbers. It is also used in cold vulcanization of rubbers, as polymerization catalyst for vegetable oils and for hardening soft woods.


  1. ^ a b c d e f NIOSH Pocket Guide to Chemical Hazards. "#0578". National Institute for Occupational Safety and Health (NIOSH).
  2. ^ a b Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
  3. ^ "Sulfur monochloride". Immediately Dangerous to Life and Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  4. ^ Holleman, A. F.; Wiberg, E. Inorganic Chemistry Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  5. ^ Hartman, W. W.; Smith, L. A.; Dickey, J. B. (1934). "Diphenylsulfide". Organic Syntheses. 14: 36.; Collective Volume, 2, p. 242
  6. ^ R. J. Cremlyn An Introduction to Organosulfur Chemistry John Wiley and Sons: Chichester (1996). ISBN 0-471-95512-4
  7. ^ Garcia-Valverde M., Torroba T. (2006). "Heterocyclic chemistry of sulfur chlorides – Fast ways to complex heterocycles". European Journal of Organic Chemistry. 4 (4): 849–861. doi:10.1002/ejoc.200500786.
  8. ^ F. Fehér "Dichlorodisulfane" in Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 371.