Sulfur compounds arechemical compounds formed the elementsulfur (S). Commonoxidation states of sulfur range from −2 to +6. Sulfur forms stable compounds with all elements except thenoble gases.
Sulfur polycations, S82+, S42+ and S162+ are produced when sulfur is reacted with oxidising agents in a strongly acidic solution.[1] The colored solutions produced by dissolving sulfur inoleum 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. S82+ is deep blue, S42+ is yellow and S162+ is red.[2]
Reduction of sulfur gives variouspolysulfides 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 giveradical anions, such asS3− gives the blue color of the rocklapis lazuli.
This reaction highlights a distinctive property of sulfur: its ability tocatenate (bind to itself by formation of chains).Protonation of these polysulfide anions produces thepolysulfanes, H2Sx where x= 2, 3, and 4.[4] Ultimately, reduction of sulfur produces sulfide salts:
The interconversion of these species is exploited in thesodium–sulfur battery.
Treatment of sulfur with hydrogen giveshydrogen sulfide. When dissolved in water, hydrogen sulfide is mildly acidic:[5]
Hydrogen sulfide gas and the hydrosulfide anion are extremely toxic to mammals, due to their inhibition of the oxygen-carrying capacity of hemoglobin and certaincytochromes in a manner analogous tocyanide andazide.
The two principal sulfur oxides are obtained by burning sulfur:
Many other sulfur oxides are observed including thesulfur-rich oxides include sulfur monoxide, disulfur monoxide, disulfur dioxides, andhigher oxides containing peroxo groups.
Sulfur reacts with fluorine to give the highly reactivesulfur tetrafluoride and the highly inertSulfur hexafluoride.[6] 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 andchlorosulfuric acid are derivatives of sulfuric acid;thionyl chloride (SOCl2) is a common reagent inorganic synthesis.[7] Sulfur halides are precursors to a variety of metal complexes.[8]
Sulfur oxidizescyanide andsulfite to givethiocyanate andthiosulfate, respectively.
Sulfur reacts with many metals. Electropositive metals give polysulfide salts.Copper,zinc andsilver aretarnished by sulfur. Although manymetal sulfides are known, most are prepared by high temperature reactions of the elements.[9] Sulfide minerals contain thesulfide (S2-) ordisulfide (S22-) anions. Typical examples are:
Some of the main classes of sulfur-containing organic compounds include the following:[10]
Compounds with carbon–sulfur multiple bonds are uncommon, an exception beingcarbon disulfide, a volatile colorless liquid that is structurally similar to carbon dioxide. It is used as a reagent to make the polymerrayon and many organosulfur compounds. Unlikecarbon monoxide,carbon monosulfide is stable only as an extremely dilute gas, found between solar systems.[11]
Organosulfur compounds are responsible for some of the unpleasant odors of decaying organic matter. They are widely known as theodorant in domestic natural gas, garlic odor, and skunk spray. Not all organic sulfur compounds smell unpleasant at all concentrations: the sulfur-containingmonoterpenoid (grapefruit mercaptan) in small concentrations is the characteristic scent of grapefruit, but has a generic thiol odor at larger concentrations.Sulfur mustard, a potentvesicant, wasused in World War I as a disabling agent.[12]
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 naturalrubber, elemental sulfur is heated with the rubber to the point that chemical reactions formdisulfide bridges betweenisoprene 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 namedvulcanization, after the Roman god of the forge andvolcanism.
