User:Praseodymium-141/Sulfur compounds

Sulfur compounds are chemical compounds formed the element sulfur (S). Common oxidation states of sulfur range from −2 to +6. Sulfur forms stable compounds with all elements except the noble gases.

Hydrogen sulfide

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Treatment of sulfur with hydrogen gives hydrogen sulfide. When dissolved in water, hydrogen sulfide is mildly acidic:[1]

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 cytochromes in a manner analogous to cyanide and azide.

Oxides

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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

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Sulfur reacts with fluorine to give the highly reactive sulfur tetrafluoride and the highly inert Sulfur hexafluoride.[2] 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 (SOCl2) is a common reagent in organic synthesis.[3]

Pseudohalides

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Sulfur oxidizes cyanide and sulfite to give thiocyanate and thiosulfate, respectively.

Metal sulfides

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Sulfur reacts with many metals. Electropositive metals give polysulfide salts. Copper, zinc and silver are tarnished by slfur. Although many metal sulfides are known, most are prepared by high temperature reactions of the elements.[4] Sulfide minerals contain the sulfide (S2-) or disulfide (S22-) anions. Typical examples are:

  • Acanthite Ag2S
  • Chalcocite Cu2S
  • Galena PbS
  • Sphalerite ZnS
  • Chalcopyrite CuFeS2
  • Millerite NiS
  • Cinnabar HgS
  • Stibnite Sb2S3
  • Pyrite FeS2
  • Molybdenite MoS2

Organic compounds

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Some of the main classes of sulfur-containing organic compounds include the following:[5]

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.[6]

Organosulfur compounds are responsible for some of the unpleasant odors of decaying organic matter. They are widely known as the odorant in domestic natural gas, garlic odor, and skunk spray. Not all organic sulfur compounds smell unpleasant at all concentrations: the sulfur-containing monoterpenoid (grapefruit mercaptan) in small concentrations is the characteristic scent of grapefruit, but has a generic thiol odor at larger concentrations. Sulfur mustard, a potent vesicant, was used in World War I as a disabling agent.[7]

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

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References

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  1. ^ Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.), Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.
  2. ^ Hasek, W. R. (1961). "1,1,1-Trifluoroheptane". Organic Syntheses. 41: 104. doi:10.1002/0471264180.os041.28.
  3. ^ Rutenberg, M. W.; Horning, E. C. (1950). "1-Methyl-3-ethyloxindole". Organic Syntheses. 30: 62. doi:10.15227/orgsyn.030.0062.
  4. ^ Vaughan, David J.; Craig, James R. (1978). Mineral chemistry of metal sulfides. Cambridge earth science series. Cambridge London New york [etc.]: Cambridge university press. ISBN 978-0-521-21489-6.
  5. ^ Cremlyn R. J. (1996). An Introduction to Organosulfur Chemistry. Chichester: John Wiley and Sons. ISBN 0-471-95512-4.
  6. ^ Wilson, R. W.; Penzias, A. A.; Wannier, P. G.; Linke, R. A. (15 March 1976). "Isotopic abundances in interstellar carbon monosulfide". Astrophysical Journal. 204: L135–L137. Bibcode:1976ApJ...204L.135W. doi:10.1086/182072.
  7. ^ Banoub, Joseph (2011). Detection of Biological Agents for the Prevention of Bioterrorism. NATO Science for Peace and Security Series A: Chemistry and Biology. p. 183. Bibcode:2011dbap.book.....B. doi:10.1007/978-90-481-9815-3. ISBN 978-90-481-9815-3. OCLC 697506461. {{cite book}}: |journal= ignored (help)