Carbon disulfide

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Carbon disulfide
Carbon disulfide syringe sample and original bottle by Thoisoi2.png
Sample of carbon disulfide in a syringe
IUPAC name
Other names
Carbon bisulfide
Carbon disulphide
Dithiocarbonic anhydride
Molar mass 76.13 g/mol
Appearance Colorless liquid
Odor Chloroform-like (pure)
Foul, rotten eggs-like (technical grade)
Density 1.539 g/cm3 (-186°C)
1.2927 g/cm3 (0 °C)
1.266 g/cm3 (25 °C)
Melting point −111.61 °C (−168.90 °F; 161.54 K)
Boiling point 46.24 °C (115.23 °F; 319.39 K)
0.258 g/100 ml (0 °C)
0.239 g/100 ml (10 °C)
0.217 g/100 ml (20 °C)
Solubility Miscible with acetone, benzene, CCl4, chloroform, dichloromethane, diethyl ether, ethanol, ethyl acetate, mineral oil, toluene, xylene
Soluble in DMSO
Slightly soluble in acetic acid, formic acid
Solubility in formic acid 4.66 g/100 g
Solubility in dimethyl sulfoxide 45 g/100 g (20.3 °C)
Vapor pressure 48.1 kPa (25 °C)
82.4 kPa (40 °C)
151 J·mol-1·K-1
88.7 kJ/mol
Safety data sheet Sigma-Aldrich
Flash point −43 °C (−45 °F; 230 K)
Lethal dose or concentration (LD, LC):
3188 mg/kg (rat, oral)
>1670 ppm (rat, 1 hr)
15500 ppm (rat, 1 hr)
3000 ppm (rat, 4 hr)
3500 ppm (rat, 4 hr)
7911 ppm (rat, 2 hr)
3165 ppm (mouse, 2 hr)
Related compounds
Related compounds
Carbon dioxide
Hydrogen sulfide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Infobox references

Carbon disulfide is a colorless volatile liquid with the formula CS2, useful as a non-polar solvent.



Carbon disulfide is highly flammable and will burn in air to release carbon and sulfur dioxides.

CS2 + 3 O2 → CO2 + 2 SO2

Chlorination of CS2 yields carbon tetrachloride and sulfur dichloride:

CS2 + 3 Cl2 → CCl4 + S2Cl2

When sulfur is dissolved in carbon disulfide, it will react with copper metal at room temperature, forming copper sulfide.[1]

Carbon disulfide will react with aq. alkali bases to form dithiocarbonates:

CS2 + 2 NaOH(aq) → Na2COS2 + H2O


Carbon disulfide is a colorless liquid, with an chloroform like smell when pure. Impure CS2 has a yellowish color and has a putrid smell. It is insoluble in water, but soluble in many organic solvents, such as benzene, ethanol, diethyl ether, carbon tetrachloride, chloroform. It's poorly soluble in formic acid. CS2 boils at 46.24 °C and freezes at −111.61 °C. It has a low autoignition temperature for a solvent, of only 102 °C.


Carbon disulfide is sold by chemical suppliers, however there appears to be no sellers on eBay or Amazon. It was used as insecticide years ago, but its use has fallen out.

Carbon disulfide is very difficult to find in most parts of the world.


There are many ways to synthesize CS2, most if not all tend to produce plenty of side products, with most synthesis routes producing hydrogen sulfide as the main side product. Some processes have low yield.

The first method, involves the reaction of sulfur vapors with carbon (or coke, anthracite, or any other coal) at 900°C. The resulting vapors are condensed and sulfur disulfide is purified. BromicAcid managed to synthesize only a few ml using this method. A few years later, garagechemist in collaboration with Len1, tried to improve the said method, using a tube furnace. It yielded around 44 g of CS2. myst32YT mas also made a small amount of carbon disulfide.[2]

Another method, that works at lower temperatures (600 °C), utilizes methane as the carbon source in the presence of silica gel or alumina catalysts:

2 CH4 + S8 → 2 CS2 + 4 H2S[3]

Another method involves the reaction of carbon tetrachloride with sulfur, between 120-220 °C, in the presence of a catalyst such as copper chloride, iron chloride, aluminium chloride.[4]

The thermal decomposition of ammonium thiocyanate, preferably in an inert atmosphere yields carbon disulfide, ammonia, hydrogen sulfide, leaving a residue of guanidinium thiocyanate. This reaction produces a fairly pure compound, though the starting products are not particularly cheap. Other thiocyanates can also be used. An approximate reaction of the decomposition is shown below:

NH4SCN → CS2 + NH3 + H2S + CH6N3SCN

Reacting sulfur or sulfur dioxide vapors with carbon dioxide at 800 °C in the presence of alumina catalyst will also yield carbon disulfide. Carbonyl sulfide may also form as a side product.[5][6]

A mixture of acetylene and sulfur vapor at temperatures between 325-650 °C will also yield carbon disulfide. Pyrite can also be used instead of sulfur.[7]

A different way involves the reaction of carbon monoxide with sulfur, in the presence of a catalyst such as iron, iron(III) sulfide, at temperatures between 400 - 500 °C.[8] The same reaction can also be carried out at 500 °C in the presence of silica gel.[9]

The reaction of calcium carbide with sulfur at at 500 °C.[10] The same reaction can also occur at temperatures between 250 - 360 °C, with a different yield.[11]

Another method described in literature is the reaction of lead(II) sulfide with carbon monoxide, in a furnace, at high temperature.[12]

Carbon disulfide can also be obtained in traces by reacting benzene with sulfur dioxide, at 500°C. It also produces many byproducts,such as dithiobenzene, carbon dioxide, carbonyl sulfide, some free oxygen, and some unknown polymeric residue.[13] The reaction can be improved by using a vanadium pentoxide/alumina catalyst, the reaction taking place at 1000-1200 °C.[14]

Pyrolysis of scrap tires also gives carbon disulfide, although separating it from the mixture is not very practical.[15]




Carbon disulfide has moderate toxicity. It is an irritant to the eyes and skin. CS2 is however extremely volatile and flammable.[16]


Carbon disulfide should be stored in closed bottles away from any heat and light sources. Due to it's high volatility its best kept in a cold place. The bottle should be open periodically to prevent a pressure build-up.


Carbon disulfide can be burned, although this will produce toxic sulfur dioxide gas.


  4. Fomin, W. A., Zhurnal Obshchei Khimii, 1936, Vol. 6, p. 852 - 854
  7. Bulletin de la Societe Chimique de France, 1908, Vol. <4> 3, p. 151
  8. DE398322 C
  11. Moissan, Comptes Rendus Hebdomadaires des Seances de l'Academie des Sciences (1894), Vol. 118, p. 502

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