Editing Sulfur (section)

==Compounds==
{{Main article|Sulfur compounds}}
Common [[oxidation state]]s of sulfur range from −2 to +6. Sulfur forms stable compounds with all elements except the [[noble gas]]es.

===Electron transfer reactions===
[[Image:Lapis lazuli block.jpg|thumb|upright|[[Lapis lazuli]] owes its blue color to a [[trisulfur]] radical anion ({{chem|S|3|-}})]]

Sulfur polycations, {{chem2|S8(2+)}}, {{chem2|S4(2+)}} and {{chem2|S19(2+)}} are produced when sulfur is reacted with oxidizing agents in a strongly acidic solution.<ref>Shriver, Atkins. Inorganic Chemistry, Fifth Edition. W. H. Freeman and Company, New York, 2010; pp 416</ref> The colored solutions produced by dissolving sulfur in [[oleum]] were first reported as early as 1804 by C.&nbsp;F. Bucholz, but the cause of the color and the structure of the polycations involved was only determined in the late 1960s. {{chem2|S8(2+)}} is deep blue, {{chem2|S4(2+)}} is yellow and {{chem2|S19(2+)}} is red.<ref name="Greenwood-1997b" />

Reduction of sulfur gives various [[polysulfide]]s with the formula {{chem|S|''x''|2-}}, many of which have been obtained in crystalline form. Illustrative is the production of [[sodium tetrasulfide]]:

{{block indent|{{chem2|4 Na + S8 -> 2 Na2S4}}}}

Some of these dianions dissociate to give [[radical anion]]s. For instance, {{chem2|S3-|link=Trisulfur}} gives the blue color of the rock [[lapis lazuli]].

[[File:S@CNT.jpg|thumb|Two parallel sulfur chains grown inside a single-wall [[carbon nanotube]] (CNT, a). Zig-zag (b) and straight (c) S chains inside double-wall CNTs<ref>{{cite journal |doi=10.1038/ncomms3162 |pmid=23851903 |pmc=3717502 |title=Conducting linear chains of sulphur inside carbon nanotubes |journal=Nature Communications |volume=4 |pages=2162 |year=2013 |last1=Fujimori |first1=Toshihiko |last2=Morelos-Gómez |first2=Aarón |last3=Zhu |first3=Zhen |last4=Muramatsu |first4=Hiroyuki |last5=Futamura |first5=Ryusuke |last6=Urita |first6=Koki |last7=Terrones |first7=Mauricio |last8=Hayashi |first8=Takuya |last9=Endo |first9=Morinobu |last10=Young Hong |first10=Sang |last11=Chul Choi |first11=Young |last12=Tománek |first12=David |last13=Kaneko |first13=Katsumi |bibcode=2013NatCo...4.2162F}}</ref>]]
This reaction highlights a distinctive property of sulfur: its ability to [[catenation|catenate]] (bind to itself by formation of chains). [[Protonation]] of these polysulfide anions produces the [[polysulfane]]s, H<sub>2</sub>S<sub>''x''</sub>, where ''x'' = 2, 3, and 4.<ref>{{cite book |title=Handbook of Preparative Inorganic Chemistry |edition=2nd |editor-first=G. |editor-last=Brauer |publisher=Academic Press |year=1963 |location=New York |volume=1 |page=421}}</ref> Ultimately, reduction of sulfur produces sulfide salts:

{{block indent|16 Na + S<sub>8</sub> → 8 Na<sub>2</sub>S}}

The interconversion of these species is exploited in the [[sodium–sulfur battery]].

===Hydrogenation===
Treatment of sulfur with hydrogen gives [[hydrogen sulfide]]. When dissolved in water, hydrogen sulfide is mildly acidic:<ref name="Greenwood-1997a">{{cite book |last1=Greenwood |first1=N. N. |last2=Earnshaw |first2=A. |year=1997 |title=Chemistry of the Elements |edition=2nd |location=Oxford |publisher=Butterworth-Heinemann |isbn=0-7506-3365-4}}</ref>

{{block indent|H<sub>2</sub>S {{eqm}} HS<sup>−</sup> + H<sup>+</sup>}}

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 [[cytochrome]]s in a manner analogous to [[cyanide]] and [[azide]] (see below, under ''precautions'').

===Combustion===
The two principal sulfur oxides are obtained by burning sulfur:

{{block indent|S + O<sub>2</sub> → SO<sub>2</sub> ([[sulfur dioxide]])}}
{{block indent|2 SO<sub>2</sub> + O<sub>2</sub> → 2 SO<sub>3</sub> ([[sulfur trioxide]])}}

Many other sulfur oxides are observed including the [[lower sulfur oxides|sulfur-rich oxides]] include [[sulfur monoxide]], [[disulfur monoxide]], disulfur dioxides, and [[higher sulfur oxides|higher oxides]] containing peroxo groups.

===Halogenation===
Sulfur reacts with [[fluorine]] to give the highly reactive [[sulfur tetrafluoride]] and the highly inert [[sulfur hexafluoride]].<ref>{{OrgSynth|last=Hasek|first=W. R.|title=1,1,1-Trifluoroheptane|volume=41|page=104|year=1961|doi=10.1002/0471264180.os041.28}}</ref> 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<sub>2</sub>) is a common reagent in [[organic synthesis]].<ref name="Lauss">{{Ullmann|first1=H.-D.|last1=Lauss|first2=W.|last2=Steffens |title=Sulfur Halides|doi=10.1002/14356007.a25_623}}</ref> Bromine also oxidizes sulfur to form [[sulfur dibromide]] and [[disulfur dibromide]].<ref name="Lauss" />

===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 attacked by sulfur; see [[tarnishing]]. Although many [[sulfide mineral|metal sulfides]] are known, most are prepared by high temperature reactions of the elements.<ref>{{cite book |last1=Vaughan |first1=D. J. |last2=Craig |first2=J. R. |title=Mineral Chemistry of Metal Sulfides |publisher=Cambridge University Press |location=Cambridge |year=1978 |isbn=0-521-21489-0}}</ref> Geoscientists also study the isotopes of metal sulfides in rocks and sediment to study environmental conditions in the Earth's past.<ref>{{Cite journal |last1=Tsang |first1=Man-Yin |last2=Inagaki |first2=Fumio |date=2020-05-29 |title=Microbial Life Deep Under the Seafloor—A Story of Not Giving Up |journal=Frontiers for Young Minds |volume=8 |pages=70 |doi=10.3389/frym.2020.00070 |issn=2296-6846 |doi-access=free }}</ref>

===Organic compounds===
{{Main|Organosulfur compounds}}
<gallery caption="Illustrative organosulfur compounds">
File:L-Cystein - L-Cysteine.svg |(''L'')-[[cysteine]], an [[amino acid]] containing a thiol group
File:Methionin - Methionine.svg|[[Methionine]], an amino acid containing a thioether
File:Thiamin.svg|[[Thiamine]] or vitamin B<sub>1</sub>
File:Biotin_structure.svg|[[Biotin]] or vitamin B<sub>7</sub>
File:Penicillin core.svg|[[Penicillin]], an antibiotic ("R" is the variable group)
File:Allicin skeletal.svg|[[Allicin]], a chemical compound in garlic
File:Diphenyl disulfide.svg|[[Diphenyl disulfide]], a representative disulfide
File:Dibenzothiophen - Dibenzothiophene.svg|[[Dibenzothiophene]], a component of crude oil
File:Perfluorooctanesulfonic acid structure.svg|[[Perfluorooctanesulfonic acid]] (PFOS), a surfactant
</gallery>

Some of the main classes of sulfur-containing organic compounds include the following:<ref name="Cremlyn-1996">{{cite book | last = Cremlyn |first=R. J. | title = An Introduction to Organosulfur Chemistry | publisher = John Wiley and Sons | location = Chichester | date = 1996 | isbn = 0-471-95512-4 }}</ref>
* [[Thiol]]s or mercaptans (so called because they capture mercury as [[chelation|chelators]]) are the sulfur analogs of [[Alcohol (chemistry)|alcohol]]s; treatment of thiols with base gives [[thiolate]] ions.
* [[Thioether]]s are the sulfur analogs of [[ether]]s.
* [[Sulfonium]] ions have three groups attached to a cationic sulfur center. [[Dimethylsulfoniopropionate]] (DMSP) is one such compound, important in the marine organic [[sulfur cycle]].
* [[Sulfoxide]]s and [[sulfone]]s 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]].
* [[Sulfonic acid]]s 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.<ref>{{Cite journal |last1=DeMartino |first1=Anthony W. |last2=Zigler |first2=David F. |last3=Fukuto |first3=Jon M. |last4=Ford |first4=Peter C. |date=2017 |title=Carbon disulfide. Just toxic or also bioregulatory and/or therapeutic? |url=https://xlink.rsc.org/?DOI=C6CS00585C |journal=Chemical Society Reviews |language=en |volume=46 |issue=1 |pages=21–39 |doi=10.1039/C6CS00585C |pmid=27722688 |issn=0306-0012}}</ref> Unlike [[carbon monoxide]], [[carbon monosulfide]] is stable only as an extremely dilute gas, found between solar systems.<ref>{{cite journal |last1=Wilson |first1=R. W. |last2=Penzias |first2=A. A. |last3=Wannier |first3=P. G. |last4=Linke |first4=R. A. |author-link=Robert Woodrow Wilson |author-link2=Arno Allan Penzias |title=Isotopic abundances in interstellar carbon monosulfide |journal=Astrophysical Journal |date=15 March 1976 |volume=204 |pages=L135–L137 |doi=10.1086/182072 |bibcode=1976ApJ...204L.135W |doi-access=free}}</ref>

Organosulfur compounds are responsible for some of the unpleasant odors of decaying organic matter. They are widely known as the [[odorizer|odorant]] in domestic natural gas, garlic odor, and skunk spray, as well as a component of [[bad breath]] odor. Not all organic sulfur compounds smell unpleasant at all concentrations: the sulfur-containing [[terpene|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 [[blister agent|vesicant]], was [[Chemical weapons in World War I|used in World War&nbsp;I]] as a disabling agent.<ref>{{cite book |last=Banoub |first=Joseph |title=Detection of Biological Agents for the Prevention of Bioterrorism |date=2011 |isbn=978-90-481-9815-3 |oclc=697506461 |page=183 |bibcode=2011dbap.book.....B |doi=10.1007/978-90-481-9815-3 |series=NATO Science for Peace and Security Series A: Chemistry and Biology}}</ref>

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,{{citation needed|date=September 2024}} 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]].