Sulfur compounds are chemical compounds formed with the element sulfur (S). Common of sulfur range from −2 to +6. Sulfur forms stable compounds with all elements except the .
Electron transfer reactions
Sulfur polycations, S
82+, S
42+ and S
162+ are produced when sulfur is reacted with oxidising agents in a strongly acidic solution.
[Shriver, Atkins. Inorganic Chemistry, Fifth Edition. W. H. Freeman and Company, New York, 2010; pp 416] The colored solutions produced by dissolving sulfur in
oleum were first reported as early as 1804 by C.F. Bucholz, but the cause of the color and the structure of the polycations involved was only determined in the late 1960s. S
82+ is deep blue, S
42+ is yellow and S
162+ is red.
Reduction of sulfur gives various with the formula Sx2-, many of which have been obtained in crystalline form. Illustrative is the production of sodium tetrasulfide:
Some of these dianions dissociate to give
, such as
Trisulfur gives the blue color of the rock
lapis lazuli.
This reaction highlights a distinctive property of sulfur: its ability to
catenation (bind to itself by formation of chains).
Protonation of these polysulfide anions produces the
, H
2S
x where x= 2, 3, and 4.
[Handbook of Preparative Inorganic Chemistry, 2nd ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 421.] Ultimately, reduction of sulfur produces sulfide salts:
- 16 Na + S8 → 8 Na2S
The interconversion of these species is exploited in the sodium–sulfur battery.
Hydrogen sulfide
Treatment of sulfur with hydrogen gives
hydrogen sulfide. When dissolved in water, hydrogen sulfide is mildly acidic:
[Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.), Oxford:Butterworth-Heinemann. .]
- H2S HS− + H+
Hydrogen sulfide gas and the hydrosulfide anion are extremely toxic to mammals, due to their inhibition of the oxygen-carrying capacity of hemoglobin and certain in a manner analogous to cyanide and azide.
Oxides
The two principal sulfur oxides are obtained by burning sulfur:
- S + O2 → SO2 (sulfur dioxide)
- 2 SO2 + O2 → 2 SO3 (sulfur trioxide)
Many other sulfur oxides are observed including the sulfur-rich oxides include sulfur monoxide, disulfur monoxide, disulfur dioxides, and higher oxides containing peroxo groups.
Halides
Sulfur reacts with fluorine to give the highly reactive sulfur tetrafluoride and the highly inert Sulfur hexafluoride.
Whereas fluorine gives S(IV) and S(VI) compounds, chlorine gives S(II) and S(I) derivatives. Thus, sulfur dichloride, disulfur dichloride, and higher chlorosulfanes arise from the chlorination of sulfur. Sulfuryl chloride and chlorosulfuric acid are derivatives of sulfuric acid;
thionyl chloride (SOCl
2) is a common reagent in organic synthesis.
Sulfur halides are precursors to a variety of metal complexes.
Pseudohalides
Sulfur oxidizes
cyanide and
sulfite to give
thiocyanate and
thiosulfate, respectively.
Metal sulfides
Sulfur reacts with many metals. Electropositive metals give polysulfide salts.
Copper,
zinc and
silver are
tarnished by sulfur. Although many
sulfide mineral are known, most are prepared by high temperature reactions of the elements.
Sulfide minerals contain the
sulfide (S
2-) or
disulfide (S
22-) anions. Typical examples are:
Organic compounds
File:Allicin skeletal.svg|Allicin, a chemical compound in garlic
File:L-Cystein - L-Cysteine.svg |( R)-cysteine, an amino acid containing a thiol group
File:L-Methionin - L-Methionine.svg|( S)-Methionine, an amino acid containing a thioether
File:Diphenyl disulfide.svg|Diphenyl disulfide, a representative disulfide
File:Perfluorooctanesulfonic acid structure.svg|Perfluorooctanesulfonic acid, a surfactant
File:Dibenzothiophen - Dibenzothiophene.svg|Dibenzothiophene, a component of crude oil
File:Penicillin core.svg|Penicillin, an antibiotic where "R" is the variable group
Some of the main classes of sulfur-containing organic compounds include the following:
-
or mercaptans (so called because they capture mercury as chelation) are the sulfur analogs of alcohols; treatment of thiols with base gives thiolate ions.
-
are the sulfur analogs of .
-
Sulfonium ions have three groups attached to a cationic sulfur center. Dimethylsulfoniopropionate (DMSP) is one such compound, important in the marine organic sulfur cycle.
-
and are thioethers with one and two oxygen atoms attached to the sulfur atom, respectively. The simplest sulfoxide, dimethyl sulfoxide, is a common solvent; a common sulfone is sulfolane.
-
are used in many detergents.
Compounds with carbon–sulfur multiple bonds are uncommon, an exception being carbon disulfide, a volatile colorless liquid that is structurally similar to carbon dioxide. It is used as a reagent to make the polymer rayon and many organosulfur compounds. Unlike carbon monoxide, carbon monosulfide is stable only as an extremely dilute gas, found between solar systems.
Organosulfur compounds are responsible for some of the unpleasant odors of decaying organic matter. They are widely known as the odorizer in domestic natural gas, garlic odor, and skunk spray. Not all organic sulfur compounds smell unpleasant at all concentrations: the sulfur-containing terpene (grapefruit mercaptan) in small concentrations is the characteristic scent of grapefruit, but has a generic thiol odor at larger concentrations. Sulfur mustard, a potent blister agent, was used in World War I as a disabling agent.
Sulfur–sulfur bonds are a structural component used to stiffen rubber, similar to the disulfide bridges that rigidify proteins (see biological below). In the most common type of industrial "curing" or hardening and strengthening of natural rubber, elemental sulfur is heated with the rubber to the point that chemical reactions form disulfide bridges between isoprene units of the polymer. This process, patented in 1843, made rubber a major industrial product, especially in automobile tires. Because of the heat and sulfur, the process was named vulcanization, after the Roman god of the forge and volcanism.
See also
