View source for Chemistry:Sulfur

{{Infobox sulfur}}
'''Sulfur''' (also spelled '''sulphur''' in British English) is a [[Chemistry:Chemical element|chemical element]]; it has [[Chemistry:Symbol|symbol]] '''S''' and [[Physics:Atomic number|atomic number]] 16. It is [[Chemistry:Abundance of the chemical elements|abundant]], [[Chemistry:Polyvalency|multivalent]] and [[Chemistry:Nonmetal|nonmetal]]lic. Under [[Standard conditions for temperature and pressure|normal conditions]], sulfur atoms form [[Chemistry:Octasulfur|cyclic octatomic molecules]] with the chemical formula [[Chemistry:Octasulfur|S<sub>8</sub>]]. Elemental sulfur is a bright yellow, [[Earth:Crystal|crystal]]line solid at [[Physics:Room temperature|room temperature]].

Sulfur is the tenth most abundant element by mass in the universe and the fifth most abundant on Earth. Though sometimes found in pure, [[Earth:Native element minerals|native]] form, sulfur on Earth usually occurs as [[Chemistry:Sulfide minerals|sulfide]] and [[Chemistry:Sulfate minerals|sulfate minerals]]. Being abundant in native form, sulfur was known in ancient times, being mentioned for its uses in [[Earth:Outline of ancient India|ancient India]], {{wipe|Ancient Greece}}, China, and ancient Egypt. Historically and in literature sulfur is also called '''brimstone''',<ref name="Greenwd" /> which means "burning stone".<ref>{{cite EB1911 |wstitle=Brimstone |volume=4 |page=571}}</ref> Today, almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants from [[Chemistry:Natural gas|natural gas]] and [[Chemistry:Petroleum|petroleum]].<ref name=BBC>{{cite web |last1=Laurence Knight |title=Sulphur surplus: Up to our necks in a diabolical element |url=https://www.bbc.com/news/magazine-28369829 |publisher=BBC |date=Jul 19, 2014}}</ref><ref name=podcast>{{cite web |title=Sulfur |url=https://www.bbc.co.uk/programmes/p027xnp2|website=Elements |publisher=BBC |date=Oct 11, 2014}}. Download [https://open.live.bbc.co.uk/mediaselector/6/redir/version/2.0/mediaset/audio-nondrm-download/proto/https/vpid/p02rc2q5.mp3 here].</ref> The greatest commercial use of the element is the production of [[Chemistry:Sulfuric acid|sulfuric acid]] for sulfate and phosphate [[Chemistry:Fertilizer|fertilizer]]s, and other chemical processes. Sulfur is used in [[Engineering:Match|match]]es, [[Biology:Insecticide|insecticide]]s, and [[Biology:Fungicide|fungicide]]s. Many sulfur compounds are odoriferous, and the smells of odorized [[Chemistry:Natural gas|natural gas]], [[Biology:Skunk|skunk]] scent, [[Medicine:Bad breath|bad breath]], [[Biology:Grapefruit|grapefruit]], and [[Biology:Garlic|garlic]] are due to organosulfur compounds. [[Chemistry:Hydrogen sulfide|Hydrogen sulfide]] gives the characteristic odor to rotting eggs and other biological processes.

Sulfur is an [[Chemistry:Mineral (nutrient)|essential element]] for all life, almost always in the form of [[Chemistry:Organosulfur compounds|organosulfur compounds]] or metal sulfides. [[Chemistry:Amino acid|Amino acid]]s (two [[Chemistry:Proteinogenic amino acid|proteinogenic]]: [[Chemistry:Cysteine|cysteine]] and [[Chemistry:Methionine|methionine]], and many other [[Chemistry:Non-proteinogenic amino acids|non-coded]]: [[Chemistry:Cystine|cystine]], [[Chemistry:Taurine|taurine]], etc.) and two vitamins ([[Chemistry:Biotin|biotin]] and [[Physics:Thiamine|thiamine]]) are organosulfur compounds crucial for life. Many [[Biology:Cofactor (biochemistry)|cofactors]] also contain sulfur, including [[Chemistry:Glutathione|glutathione]], and [[Biology:Iron–sulfur protein|iron–sulfur protein]]s. [[Biology:Disulfide|Disulfide]]s, S–S bonds, confer mechanical strength and insolubility of the (among others) protein [[Biology:Keratin|keratin]], found in outer skin, hair, and feathers. Sulfur is one of the core chemical elements needed for biochemical functioning and is an elemental macronutrient for all living organisms.

==Characteristics==

===Physical properties===
[[File:Burning-sulfur.png|thumb|upright|left|As a solid, sulfur is a characteristic lemon yellow; when burned, sulfur melts into a blood-red liquid and emits a blue flame.]]
Sulfur forms several polyatomic molecules. The best-known allotrope is [[Chemistry:Octasulfur|octasulfur]], cyclo-S<sub>8</sub>. The [[Point group|point group]] of cyclo-S<sub>8</sub> is D<sub>4d</sub> and its dipole moment is 0 D.<ref>{{cite journal|last1=Rettig|first1=S. J. |last2=Trotter |first2=J. |title=Refinement of the structure of orthorhombic sulfur, α-S8 |journal=Acta Crystallographica Section C|date=15 December 1987 |volume=43 |issue=12 |pages=2260–2262 |doi=10.1107/S0108270187088152 |bibcode=1987AcCrC..43.2260R |url=http://journals.iucr.org/c/issues/1987/12/00/a27614/a27614.pdf}}</ref> Octasulfur is a soft, bright-yellow solid that is odorless, but impure samples have an odor similar to that of [[Engineering:Match|match]]es.<ref>A strong odor called "smell of sulfur" actually is given off by several sulfur compounds, such as [[Chemistry:Hydrogen sulfide|hydrogen sulfide]] and organosulfur compounds.</ref> It melts at {{convert|115.21|C}}, boils at {{convert|444.6|C}}<ref name="Greenwd" /> and [[Sublimation (phase transition)|sublimes]] more or less between {{convert|20|C}} and {{convert|50|C}}.<ref>{{Cite journal |last=Tucker |first=Roy P. |date=1929-01-01 |title=Notes on the Sublimation of Sulfur between 25° and 50°C |url=https://pubs.acs.org/doi/abs/10.1021/ie50229a014 |journal=Industrial & Engineering Chemistry |language=en |volume=21 |issue=1 |pages=44–47 |doi=10.1021/ie50229a014 |issn=0019-7866}}</ref> At {{convert|95.2|C}}, below its melting temperature, cyclo-octasulfur changes from α-octasulfur to the β-[[Physics:Polymorphism (materials science)|polymorph]].<ref name="Greenwood">{{Greenwood&Earnshaw2nd|pages= 645–665}}</ref> The structure of the S<sub>8</sub> ring is virtually unchanged by this phase change, which affects the intermolecular interactions. Between its melting and boiling temperatures, octasulfur changes its allotrope again, turning from β-octasulfur to γ-sulfur, again accompanied by a lower density but increased [[Physics:Viscosity|viscosity]] due to the formation of [[Chemistry:Polymer|polymer]]s.<ref name="Greenwood" /> At higher temperatures, the viscosity decreases as depolymerization occurs. Molten sulfur assumes a dark red color above {{convert|200|C}}. The density of sulfur is about 2&nbsp;g/cm<sup>3</sup>, depending on the allotrope; all of the stable allotropes are excellent electrical insulators.

Sulfur is insoluble in water but soluble in [[Chemistry:Carbon disulfide|carbon disulfide]] and, to a lesser extent, in other [[Chemistry:Chemical polarity|nonpolar]] organic solvents, such as [[Chemistry:Benzene|benzene]] and [[Chemistry:Toluene|toluene]].

===Chemical properties===
Under normal conditions, sulfur [[Chemistry:Hydrolysis|hydrolyzes]] very slowly to mainly form [[Chemistry:Hydrogen sulfide|hydrogen sulfide]] and [[Chemistry:Sulfuric acid|sulfuric acid]]:

{{block indent|{{fraction|2}} {{chem|S|8}} + 4 {{chem|H|2|O}} → 3 {{chem|H|2|S}} + {{chem|H|2|SO|4|}}}}

The reaction involves adsorption of protons onto {{chem|S|8}} clusters, followed by [[Biology:Disproportionation|disproportionation]] into the reaction products.<ref>{{cite journal |last1=Maldonado-Zagal |first1=S. B. |last2=Boden |first2=P. J. |title=Hydrolysis of Elemental Sulphur in Water and its Effect on the Corrosion of Mild Steel |journal=British Corrosion Journal |date=1 January 1982 |volume=17 |issue=3 |pages=116–120 |doi=10.1179/000705982798274336 |url=https://www.tandfonline.com/doi/abs/10.1179/000705982798274336 |access-date=23 June 2022 |issn=0007-0599}}</ref>

The second, fourth and sixth [[Chemistry:Molar ionization energies of the elements|ionization energies]] of sulfur are 2252 kJ/mol, 4556 kJ/mol and 8495.8 kJ/mol, respectively. A composition of products of sulfur's reactions with oxidants (and its oxidation state) depends on that whether releasing out of a reaction energy overcomes these thresholds. Applying [[Chemistry:Catalysis|catalysts]] and / or [[Physics:Endothermic process|supply of outer energy]] may vary sulfur's oxidation state and a composition of reaction products. While reaction between sulfur and oxygen at normal conditions gives sulfur dioxide (oxidation state +4), formation of [[Chemistry:Sulfur trioxide|sulfur trioxide]] (oxidation state +6) requires temperature {{cvt|400|–|600|°C|sigfig=2}} and presence of a catalyst.

In reactions with elements of lesser [[Chemistry:Electronegativity|electronegativity]], it reacts as an oxidant and forms sulfides, where it has oxidation state −2.

Sulfur reacts with nearly all other elements with the exception of the noble gases, even with the notoriously unreactive metal [[Chemistry:Iridium|iridium]] (yielding [[Chemistry:Iridium disulfide|iridium disulfide]]).<ref name="Munson-1968">
{{cite journal
| url= https://htracyhall.org/ocr/HTH-Archives/Cabinet%208/Drawer%203%20(MATI%20-%20MOZ)/(Munson,%20R.A.)%20(Muntoni,%20C.)%20(Murase,%20K.)%20(linked)/(Munson,%20R.A.)%20(Muntoni,%20C.)%20(Murase,%20K.)-237_OCR.pdf
| author-last= Munson
| author-first= Ronald A.
| date= February 1968
| title= The synthesis of iridium disulfide and nickel diarsenide having the pyrite structure
| journal= Inorganic Chemistry
| volume= 7
| number= 2
| pages= 389–390
| doi= 10.1021/ic50060a047
}}</ref> Some of those reactions need elevated temperatures.<ref name="WibergWiberg2001">{{cite book|author1=Egon Wiberg|author2=Nils Wiberg|title=Inorganic Chemistry|url=https://books.google.com/books?id=Mtth5g59dEIC&pg=PA513|year=2001|publisher=Academic Press|isbn=978-0-12-352651-9|pages=513–}}</ref>

===Allotropes===
{{Main|Chemistry:Allotropes of sulfur}}
[[File:Cyclooctasulfur-above-3D-balls.png|thumb|left|The structure of the cyclooctasulfur molecule, S<sub>8</sub>]]
Sulfur forms over 30 solid [[Physics:Allotropy|allotropes]], more than any other element.<ref>{{cite book |title= Solid Sulfur Allotropes Sulfur Allotropes |first1= Ralf |last1= Steudel|first2= Bodo|last2= Eckert|date= 2003 |volume= 230 |pages= 1–80 |doi= 10.1007/b12110 |series= Topics in Current Chemistry |isbn= 978-3-540-40191-9}}</ref> Besides S<sub>8</sub>, several other rings are known.<ref>{{cite book |doi=10.1007/3-540-11345-2_10 |last= Steudel|first= R. |date= 1982 |volume= 102 |pages= 149–176 |series=Topics in Current Chemistry |isbn=978-3-540-11345-4|title= Inorganic Ring Systems|chapter= Homocyclic sulfur molecules}}</ref> Removing one atom from the crown gives S<sub>7</sub>, which is of a deeper yellow than S<sub>8</sub>. [[High-performance liquid chromatography|HPLC]] analysis of "elemental sulfur" reveals an equilibrium mixture of mainly S<sub>8</sub>, but with S<sub>7</sub> and small amounts of S<sub>6</sub>.<ref>{{cite journal |last1= Tebbe |first1= Fred N. |last2= Wasserman |first2= E. |last3= Peet |first3= William G. |last4= Vatvars |first4= Arturs |last5= Hayman |first5= Alan C. |title= Composition of Elemental Sulfur in Solution: Equilibrium of {{chem|S|6}}, S<sub>7</sub>, and S<sub>8</sub> at Ambient Temperatures |journal= Journal of the American Chemical Society|date= 1982 |volume= 104 |issue= 18 |pages= 4971–4972 |doi= 10.1021/ja00382a050}}</ref> Larger rings have been prepared, including S<sub>12</sub> and S<sub>18</sub>.<ref>{{cite journal|last1= Meyer|first1= Beat|title= Solid Allotropes of Sulfur|journal= Chemical Reviews |date= 1964|volume= 64|issue= 4|pages= 429–451|doi= 10.1021/cr60230a004}}</ref><ref>{{cite journal|last1= Meyer|first1= Beat|title= Elemental sulfur|journal= Chemical Reviews|date= 1976|volume= 76|issue= 3|pages= 367–388 |doi= 10.1021/cr60301a003}}</ref>

[[Physics:Amorphous|Amorphous]] or "plastic" sulfur is produced by rapid cooling of molten sulfur—for example, by pouring it into cold water. [[Physics:X-ray crystallography|X-ray crystallography]] studies show that the amorphous form may have a [[Helix|helical]] structure with eight atoms per turn. The long coiled polymeric molecules make the brownish substance [[Physics:Elasticity|elastic]], and in bulk this form has the feel of crude rubber. This form is metastable at room temperature and gradually reverts to the crystalline molecular allotrope, which is no longer elastic. This process happens within a matter of hours to days, but can be rapidly catalyzed.
{{clear left}}

===Isotopes===
{{Main|Physics:Isotopes of sulfur}}
Sulfur has 23 known [[Physics:Isotope|isotope]]s, four of which are stable: <sup>32</sup>S ({{val|94.99|0.26|u=%}}), <sup>33</sup>S ({{val|0.75|0.02|u=%}}), <sup>34</sup>S ({{val|4.25|0.24|u=%}}), and <sup>36</sup>S ({{val|0.01|0.01|u=%}}).<ref>[http://www.ciaaw.org/sulfur.htm Sulfur]. [[Organization:Commission on Isotopic Abundances and Atomic Weights|Commission on Isotopic Abundances and Atomic Weights]]</ref><ref name="b92">{{RubberBible92nd|page=1.14}}</ref> Other than <sup>35</sup>S, with a [[Physics:Half-life|half-life]] of 87 days, the radioactive isotopes of sulfur have half-lives less than 3&nbsp;hours.

The preponderance of <sup>32</sup>S is explained by its production in the so-called alpha-process (one of the main classes of nuclear fusion reactions) in exploding stars. Other stable sulfur isotopes are produced in the bypass processes related with <sup>34</sup>Ar, and their composition depends on a type of a stellar explosion. For example, proportionally more <sup>33</sup>S comes from novae than from supernovae.<ref>{{Cite web |title=Searching for the Origins of Presolar Grains |url=https://www.energy.gov/science/np/articles/searching-origins-presolar-grains |access-date=2023-02-04 |website=Energy.gov |language=en}}</ref>

On the planet Earth the sulfur isotopic composition was determined by the Sun. Though it is assumed that the distribution of different sulfur isotopes should be more or less equal, it has been found that proportions of two most abundant sulfur isotopes <sup>32</sup>S and <sup>34</sup>S varies in different samples. Assaying of these isotopes ratio (δ<sup>34</sup>S) in the samples allows to make suggestions about their chemical history, and with support of other methods, it allows to age-date the samples, estimate temperature of equilibrium between ore and water, determine pH and oxygen fugacity, identify the activity of sulfate-reducing bacteria in the time of formation of the sample, or suggest the main sources of sulfur in ecosystems.<ref>{{Cite book |last1=Paytan |first1=Adina |title=Geologic Time Scale |last2=Yao |first2=Weiqi |last3=Faul |first3=Kristina |last4=Gray |first4=E.T. |year=2020 |pages=259–278 |language=en |chapter=Sulfur Isotope Stratigraphy |doi=10.1016/B978-0-12-824360-2.00009-7 |isbn=9780128243602 |chapter-url=https://www.researchgate.net/publication/347656764}}</ref> However, there are ongoing discussions about what is the real reason of the δ<sup>34</sup>S shifts, biological activity or postdeposital alteration.<ref>{{Cite web |title=NASA Astrobiology |url=https://astrobiology.nasa.gov/news/the-first-sulfur-eaters/ |access-date=2023-02-04 |website=astrobiology.nasa.gov |language=en-EN}}</ref>

For example, when [[Chemistry:Sulfide mineral|sulfide mineral]]s are precipitated, isotopic equilibration among solids and liquid may cause small differences in the [[Biology:Δ34S|δ<sup>34</sup>S]] values of co-genetic minerals. The differences between minerals can be used to estimate the temperature of equilibration. The [[Biology:Δ13C|δ<sup>13</sup>C]] and δ<sup>34</sup>S of coexisting [[Chemistry:Carbonate minerals|carbonate minerals]] and sulfides can be used to determine the [[Chemistry:PH|pH]] and oxygen [[Physics:Fugacity|fugacity]] of the ore-bearing fluid during ore formation.

Scientists measure the [[Chemistry:Sulfur isotope biogeochemistry|sulfur isotopes]] of [[Earth:Mineral|minerals]] in rocks and [[Earth:Sediment|sediments]] to study the [[Chemistry:Redox|redox]] conditions in the oceans in the past. [[Earth:Sulfate-reducing microorganism|Sulfate-reducing bacteria]] in marine sediment fractionate [[Chemistry:Sulfur isotope biogeochemistry|sulfur isotopes]] as they take in [[Chemistry:Sulfate|sulfate]] and produce [[Chemistry:Sulfide|sulfide]]. Prior to 2010s, it was thought that sulfate reduction could fractionate [[Chemistry:Sulfur isotope biogeochemistry|sulfur isotopes]] up to 46 permil<ref>{{Cite journal |last1=Goldhaber |first1=M.B. |last2=Kaplan |first2=I.R. |date=April 1980 |title=Mechanisms of sulfur incorporation and isotope fractionation during early diagenesis in sediments of the gulf of California |url=https://linkinghub.elsevier.com/retrieve/pii/0304420380900638 |journal=Marine Chemistry |language=en |volume=9 |issue=2 |pages=95–143 |doi=10.1016/0304-4203(80)90063-8|bibcode=1980MarCh...9...95G }}</ref> and fractionation larger than 46 permil recorded in sediments must be due to [[Biology:Disproportionation|disproportionation]] of sulfur compounds in the sediment. This view has changed since the 2010s as experiments show that [[Earth:Sulfate-reducing microorganism|sulfate-reducing bacteria]] can fractionate to 66 permil.<ref>{{Cite journal |last1=Sim |first1=Min Sub |last2=Bosak |first2=Tanja |last3=Ono |first3=Shuhei |date=July 2011 |title=Large Sulfur Isotope Fractionation Does Not Require Disproportionation |url=https://www.science.org/doi/10.1126/science.1205103 |journal=Science |language=en |volume=333 |issue=6038 |pages=74–77 |doi=10.1126/science.1205103 |pmid=21719675 |bibcode=2011Sci...333...74S |s2cid=1248182 |issn=0036-8075}}</ref> As substrates for disproportionation are limited by the product of [[Earth:Sulfate-reducing microorganism|sulfate reduction]], the isotopic effect of disproportionation should be less than 16 permil in most sedimentary settings.<ref>{{Cite journal |last1=Tsang |first1=Man-Yin |last2=Böttcher |first2=Michael Ernst |last3=Wortmann |first3=Ulrich Georg |date=August 2023 |title=Estimating the effect of elemental sulfur disproportionation on the sulfur-isotope signatures in sediments |url=https://linkinghub.elsevier.com/retrieve/pii/S0009254123002334 |journal=Chemical Geology |language=en |volume=632 |pages=121533 |doi=10.1016/j.chemgeo.2023.121533|s2cid=258600480 }}</ref>

In most [[Earth:Forest|forest]] ecosystems, sulfate is derived mostly from the atmosphere; weathering of ore minerals and evaporites contribute some sulfur. Sulfur with a distinctive isotopic composition has been used to identify pollution sources, and enriched sulfur has been added as a tracer in [[Earth:Hydrology|hydrologic]] studies. Differences in the [[Chemistry:Natural abundance|natural abundance]]s can be used in systems where there is sufficient variation in the <sup>34</sup>S of ecosystem components. Rocky Mountain lakes thought to be dominated by atmospheric sources of sulfate have been found to have measurably different <sup>34</sup>S values than lakes believed to be dominated by watershed sources of sulfate.

The radioactive <sup>35</sup>S is formed in [[Physics:Cosmic ray spallation|cosmic ray spallation]] of the atmospheric <sup>40</sup>[[Chemistry:Argon|Ar]]. This fact may be used for proving the presence of recent (not more than 1 year) atmospheric sediments in various things. This isotope may be obtained artificially by different ways. In practice, the reaction <sup>35</sup>Cl + n → <sup>35</sup>S + p is used by irradiating potassium chloride with neutrons.<ref>{{Cite conference |last1=Kim |first1=Ik Soo |last2=Kwak |first2=Seung Im |last3=Park |first3=Ul Jae |last4=Bang |first4=Hong Sik |last5=Han |first5=Hyun Soo |date=2005-07-01 |title=Production of Sulfur-35 by the Cation Exchange Process |conference=2005 autumn meeting of the KNS, Busan (Korea, Republic of), 27–28 Oct 2005 |url=https://www.osti.gov/etdeweb/biblio/20765660 |language=en}}</ref> The isotope <sup>35</sup>S is used in various sulfur-containing compounds as a [[Biology:Radioactive tracer|radioactive tracer]] for many biological studies, for example, the Hershey-Chase experiment.

Because of the weak beta activity of <sup>35</sup>S, its compounds are relatively safe as long as they are not ingested or absorbed by the body.<ref>{{Cite web |title=Sulfur-35 (35 S) safety information and specific handling precautions |url=https://ehs.yale.edu/sites/default/files/files/radioisotope-s35.pdf |website=Yale Environmental Health & Safety}}</ref>

===Natural occurrence===
[[File:Nearly exhausted sulphur vat from which railroad cars are loaded, Freeport Sulphur Co., Hoskins Mound, Texas, 1a35438v.jpg|thumb|Sulfur vat from which railroad cars are loaded, Freeport Sulphur Co., Hoskins Mound, Texas (1943)]]
[[File:Io highest resolution true color.jpg|thumb|Most of the yellow and orange hues of Io are due to elemental sulfur and sulfur compounds deposited by active volcanoes.]]
[[File:Menambang_Belerang.jpg|thumb|Sulfur extraction, East Java]]
[[File:Kawah Ijen -East Java -Indonesia -sulphur-31July2009.jpg|thumb|A man carrying sulfur blocks from Kawah Ijen, a volcano in East Java, Indonesia, 2009]]
<sup>32</sup>S is created inside massive stars, at a depth where the temperature exceeds 2.5×10<sup>9</sup>&nbsp;K, by the fusion of one nucleus of silicon plus one nucleus of helium.<ref>{{cite journal|first= A. G. W.|last= Cameron |title=Stellar Evolution, Nuclear Astrophysics, and Nucleogenesis|journal=CRL-41|url=https://fas.org/sgp/eprint/CRL-41.pdf|date=1957}}</ref> As this nuclear reaction is part of the [[Astronomy:Alpha process|alpha process]] that produces elements in abundance, sulfur is the 10th most common element in the universe.

Sulfur, usually as sulfide, is present in many types of [[Astronomy:Meteorite|meteorite]]s. Ordinary chondrites contain on average 2.1% sulfur, and carbonaceous chondrites may contain as much as 6.6%. It is normally present as [[Chemistry:Troilite|troilite]] (FeS), but there are exceptions, with carbonaceous chondrites containing free sulfur, sulfates and other sulfur compounds.<ref>{{cite book |first=B. |last=Mason |date=1962 |title=Meteorites |location=New York |publisher=John Wiley & Sons |url-access=registration |url=https://archive.org/details/meteorites0000maso|page=[https://archive.org/details/meteorites0000maso/page/160 160] |isbn=978-0-908678-84-6}}</ref> The distinctive colors of [[Astronomy:Jupiter|Jupiter]]'s [[Earth:Volcano|volcanic]] moon [[Astronomy:Io (moon)|Io]] are attributed to various forms of molten, solid, and gaseous sulfur.<ref>{{cite journal|last1= Lopes|first1= Rosaly M. C.|last2= Williams|first2= David A.|title= Io after Galileo|journal= Reports on Progress in Physics|volume= 68|issue= 2|pages= 303–340 |date= 2005|doi= 10.1088/0034-4885/68/2/R02 |bibcode=2005RPPh...68..303L|s2cid= 44208045}}</ref>

It is the fifth most common element by mass in the Earth. Elemental sulfur can be found near [[Earth:Hot spring|hot spring]]s and volcanic regions in many parts of the world, especially along the Pacific Ring of Fire; such volcanic deposits are currently mined in Indonesia, Chile, and Japan. These deposits are polycrystalline, with the largest documented single crystal measuring 22×16×11&nbsp;cm.<ref>{{cite journal |url= http://www.minsocam.org/ammin/AM66/AM66_885.pdf |journal=American Mineralogist |volume=66 |pages=885–907 |date=1981 |title=The largest crystals |last=Rickwood |first=P. C.}}</ref> Historically, [[Place:Sicily|Sicily]] was a major source of sulfur in the [[History:Industrial Revolution|Industrial Revolution]].<ref>{{cite book |last=Kutney |first=Gerald |title=Sulfur: history, technology, applications & industry |date=2007 |publisher=ChemTec |location=Toronto |isbn=978-1-895198-37-9 |oclc=79256100 |page=43}}</ref> Lakes of molten sulfur up to ~200 m in diameter have been found on the sea floor, associated with [[Earth:Submarine volcano|submarine volcano]]es, at depths where the boiling point of water is higher than the melting point of sulfur.<ref name="ronde2015">{{cite book |first1=C. E. J. |last1=de Ronde |first2=W. W. Jr. |last2=Chadwick |first3=R. G. |last3=Ditchburn |first4=R. W. |last4=Embley |first5=V. |last5=Tunnicliffe |first6=E. T. |last6=Baker |first7=S. L. |last7=Walker |first8=V. L. |last8=Ferrini |first9=S. M. |last9=Merle |year=2015 |chapter=Molten Sulfur Lakes of Intraoceanic Arc Volcanoes |title=Volcanic Lakes |publisher=Springer |pages=261–288 |doi=10.1007/978-3-642-36833-2 |bibcode=2015vola.book.....R |isbn=978-3-642-36832-5|s2cid=199492543 }}</ref>

Native sulfur is synthesised by anaerobic bacteria acting on [[Chemistry:Sulfate minerals|sulfate minerals]] such as [[Chemistry:Gypsum|gypsum]] in [[Earth:Salt dome|salt dome]]s.<ref>{{cite book |last1=Klein |first1=Cornelis |first2=Cornelius S. Jr. |last2=Hurlbut |title=Manual of Mineralogy |publisher=Wiley |year=1985 |edition=20th |pages=265–66 |isbn=0-471-80580-7}}</ref><ref>{{Cite web|url=https://www.mindat.org/min-3826.html|title=Sulphur: Mineral information, data and localities|website=www.mindat.org}}</ref> Significant deposits in salt domes occur along the coast of the Gulf of Mexico, and in [[Earth:Evaporite|evaporite]]s in eastern Europe and western Asia. Native sulfur may be produced by geological processes alone. Fossil-based sulfur deposits from salt domes were once the basis for commercial production in the United States, Russia, Turkmenistan, and Ukraine.<ref name="Nehb" /> Currently, commercial production is still carried out in the Osiek mine in Poland. Such sources are now of secondary commercial importance, and most are no longer worked.

Common naturally occurring sulfur compounds include the [[Earth:Mineral#Sulfides|sulfide minerals]], such as [[Chemistry:Pyrite|pyrite]] (iron sulfide), [[Chemistry:Cinnabar|cinnabar]] (mercury sulfide), [[Chemistry:Galena|galena]] (lead sulfide), [[Chemistry:Sphalerite|sphalerite]] (zinc sulfide), and [[Chemistry:Stibnite|stibnite]] (antimony sulfide); and the [[Earth:Mineral#Sulfates|sulfate minerals]], such as [[Chemistry:Gypsum|gypsum]] (calcium sulfate), [[Chemistry:Alunite|alunite]] (potassium aluminium sulfate), and barite (barium sulfate). On Earth, just as upon Jupiter's moon Io, elemental sulfur occurs naturally in volcanic emissions, including emissions from [[Earth:Hydrothermal vent|hydrothermal vent]]s.

The main industrial source of sulfur is now [[Chemistry:Petroleum|petroleum]] and [[Chemistry:Natural gas|natural gas]].<ref name=BBC/>

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

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

Sulfur polycations, S<sub>8</sub><sup>2+</sup>, S<sub>4</sub><sup>2+</sup> and S<sub>16</sub><sup>2+</sup> are produced when sulfur is reacted with oxidising 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 [[Chemistry:Oleum|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<sub>8</sub><sup>2+</sup> is deep blue, S<sub>4</sub><sup>2+</sup> is yellow and S<sub>16</sub><sup>2+</sup> is red.<ref name="Greenwood" />

Reduction of sulfur gives various [[Chemistry:Polysulfide|polysulfide]]s with the formula S<sub>''x''</sub><sup>2−</sup>, 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 [[Physics:Radical anion|radical anion]]s, such as [[Chemistry:Trisulfur|S<sub>3</sub><sup>−</sup>]] gives the blue color of the rock [[Chemistry:Lapis lazuli|lapis lazuli]].

[[File:S@CNT.jpg|thumb|Two parallel sulfur chains grown inside a single-wall [[Physics:Carbon nanotube|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 [[Chemistry:Catenation|catenate]] (bind to itself by formation of chains). [[Chemistry:Protonation|Protonation]] of these polysulfide anions produces the [[Chemistry:Polysulfane|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 [[Engineering:Sodium–sulfur battery|sodium–sulfur battery]].

===Hydrogenation===
Treatment of sulfur with hydrogen gives [[Chemistry:Hydrogen sulfide|hydrogen sulfide]]. When dissolved in water, hydrogen sulfide is mildly acidic:<ref name="Greenwd">{{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 [[Biology:Cytochrome|cytochrome]]s in a manner analogous to [[Chemistry:Cyanide|cyanide]] and [[Chemistry:Azide|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> ([[Chemistry:Sulfur dioxide|sulfur dioxide]])}}
{{block indent|2 SO<sub>2</sub> + O<sub>2</sub> → 2 SO<sub>3</sub> ([[Chemistry:Sulfur trioxide|sulfur trioxide]])}}

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

===Halogenation===
Sulfur reacts with fluorine to give the highly reactive [[Chemistry:Sulfur tetrafluoride|sulfur tetrafluoride]] and the highly inert [[Chemistry:Sulfur hexafluoride|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, [[Chemistry:Sulfur dichloride|sulfur dichloride]], [[Chemistry:Disulfur dichloride|disulfur dichloride]], and higher chlorosulfanes arise from the chlorination of sulfur. [[Chemistry:Sulfuryl chloride|Sulfuryl chloride]] and [[Chemistry:Chlorosulfuric acid|chlorosulfuric acid]] are derivatives of sulfuric acid; [[Chemistry:Thionyl chloride|thionyl chloride]] (SOCl<sub>2</sub>) is a common reagent in [[Chemistry:Organic synthesis|organic synthesis]].<ref name="Ullmann Sulfur Halides">{{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 [[Chemistry:Sulfur dibromide|sulfur dibromide]] and [[Chemistry:Disulfur dibromide|disulfur dibromide]].<ref name="Ullmann Sulfur Halides" />

===Pseudohalides===
Sulfur oxidizes [[Chemistry:Cyanide|cyanide]] and [[Chemistry:Sulfite|sulfite]] to give [[Chemistry:Thiocyanate|thiocyanate]] and [[Chemistry:Thiosulfate|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 [[Chemistry: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|Chemistry:Organosulfur compounds}}
<gallery caption="Illustrative organosulfur compounds">
File:Allicin skeletal.svg|[[Chemistry:Allicin|Allicin]], a chemical compound in garlic
File:L-Cystein - L-Cysteine.svg |(''L'')-[[Chemistry:Cysteine|cysteine]], an [[Chemistry:Amino acid|amino acid]] containing a thiol group
File:Methionin - Methionine.svg|[[Chemistry:Methionine|Methionine]], an [[Chemistry:Amino acid|amino acid]] containing a thioether
File:Diphenyl disulfide.svg|[[Chemistry:Diphenyl disulfide|Diphenyl disulfide]], a representative disulfide
File:Perfluorooctanesulfonic acid structure.svg|[[Chemistry:Perfluorooctanesulfonic acid|Perfluorooctanesulfonic acid]], a surfactant
File:Dibenzothiophen - Dibenzothiophene.svg|[[Chemistry:Dibenzothiophene|Dibenzothiophene]], a component of crude oil
File:Penicillin core.svg|[[Chemistry:Penicillin|Penicillin]], an antibiotic where "R" is the variable group
</gallery>

Some of the main classes of sulfur-containing organic compounds include the following:<ref name="Cremlyn">{{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>
* [[Chemistry:Thiol|Thiol]]s or mercaptans (so called because they capture mercury as [[Chemistry:Chelation|chelators]]) are the sulfur analogs of [[Chemistry:Alcohol|alcohol]]s; treatment of thiols with base gives thiolate ions.
* [[Chemistry:Thioether|Thioether]]s are the sulfur analogs of [[Chemistry:Ether|ether]]s.
* [[Physics:Sulfonium|Sulfonium]] ions have three groups attached to a cationic sulfur center. [[Chemistry:Dimethylsulfoniopropionate|Dimethylsulfoniopropionate]] (DMSP) is one such compound, important in the marine organic [[Biology:Sulfur cycle|sulfur cycle]].
* [[Chemistry:Sulfoxide|Sulfoxide]]s and [[Chemistry:Sulfone|sulfone]]s are thioethers with one and two oxygen atoms attached to the sulfur atom, respectively. The simplest sulfoxide, [[Chemistry:Dimethyl sulfoxide|dimethyl sulfoxide]], is a common solvent; a common sulfone is [[Chemistry:Sulfolane|sulfolane]].
* [[Chemistry:Sulfonic acid|Sulfonic acid]]s are used in many detergents.

Compounds with carbon–sulfur multiple bonds are uncommon, an exception being [[Chemistry:Carbon disulfide|carbon disulfide]], a volatile colorless liquid that is structurally similar to carbon dioxide. It is used as a reagent to make the polymer [[Chemistry:Rayon|rayon]] and many organosulfur compounds. Unlike [[Chemistry:Carbon monoxide|carbon monoxide]], [[Chemistry:Carbon monosulfide|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-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 [[Physics:Odorizer|odorant]] in domestic natural gas, garlic odor, and skunk spray, as well as a component of [[Medicine:Bad breath|bad breath]] odor. Not all organic sulfur compounds smell unpleasant at all concentrations: the sulfur-containing [[Biology:Terpene|monoterpenoid]] [[Chemistry:Grapefruit mercaptan|grapefruit mercaptan]] in small concentrations is the characteristic scent of grapefruit, but has a generic thiol odor at larger concentrations. [[Chemistry:Sulfur mustard|Sulfur mustard]], a potent [[Chemistry:Blister agent|vesicant]], was used in World War 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 [[Chemistry:Rubber|rubber]], elemental sulfur is heated with the rubber to the point that chemical reactions form [[Biology:Disulfide|disulfide]] bridges between [[Chemistry:Isoprene|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 [[Chemistry:Sulfur vulcanization|vulcanization]], after the Roman god of the forge and volcanism.

==History==

===Antiquity===
[[File:MODOAzufre.jpg|thumb|Pharmaceutical container for sulfur from the first half of the 20th century. From the Museo del Objeto del Objeto collection]]

Being abundantly available in native form, sulfur was known in ancient times and is referred to in the Torah ([[Religion:Book of Genesis|Genesis]]). English translations of the Christian Bible commonly referred to burning sulfur as "brimstone", giving rise to the term "[[Chemistry:Fire and brimstone|fire-and-brimstone]]" [[Religion:Sermon|sermon]]s, in which listeners are reminded of the fate of [[Religion:Damnation|eternal damnation]] that await the unbelieving and unrepentant. It is from this part of the Bible<ref>{{Cite web |title=Sulfur in the Bible (14 instances) |url=https://bible.knowing-jesus.com/words/Sulfur |access-date=2022-05-19 |website=bible.knowing-jesus.com}}</ref> that [[Religion:Hell|Hell]] is implied to "smell of sulfur" (likely due to its association with volcanic activity). According to the [[Chemistry:Ebers Papyrus|Ebers Papyrus]], a sulfur ointment was used in ancient [[Place:Egypt|Egypt]] to treat granular eyelids. Sulfur was used for fumigation in preclassical Greece;<ref>{{cite book |url= https://books.google.com/books?id=ed0yC98aAKYC&pg=PA242 |title= Archaeomineralogy |isbn= 978-3-540-78593-4 |page= 242 |last1= Rapp |first1= George Robert |date= 4 February 2009|publisher= Springer }}</ref> this is mentioned in the ''[[Religion:Odyssey|Odyssey]]''.<ref>[https://www.perseus.tufts.edu/hopper/text?doc=Hom.+Od.+22.480 ''Odyssey'', book 22, lines 480–495]. www.perseus.tufts.edu. Retrieved on 16 August 2012.</ref> Pliny the Elder discusses sulfur in book 35 of his ''Natural History'', saying that its best-known source is the island of Melos. He mentions its use for fumigation, medicine, and bleaching cloth.<ref>''Pliny the Elder on science and technology'', John F. Healy, Oxford University Press, 1999, {{ISBN|0-19-814687-6}}, pp. 247–249.</ref>

A natural form of sulfur known as {{transliteration|zh|shiliuhuang}} ({{lang|zh|石硫黄}}) was known in China since the 6th century BC and found in Hanzhong.<ref name="yunming 487">{{cite journal|author= Zhang, Yunming|date= 1986|title= The History of Science Society: Ancient Chinese Sulfur Manufacturing Processes|journal= Isis|volume= 77|issue= 3|doi= 10.1086/354207|page=487|s2cid= 144187385}}</ref> By the 3rd century, the Chinese had discovered that sulfur could be extracted from [[Chemistry:Pyrite|pyrite]].<ref name="yunming 487" /> Chinese Daoists were interested in sulfur's flammability and its reactivity with certain metals, yet its earliest practical uses were found in [[Medicine:Traditional Chinese medicine|traditional Chinese medicine]].<ref name="yunming 487" /> The ''Wujing Zongyao'' of 1044 AD described various formulas for Chinese black powder, which is a mixture of [[Chemistry:Potassium nitrate|potassium nitrate]] ({{chem|K|N|O|3}}), [[Chemistry:Charcoal|charcoal]], and sulfur.<ref>{{cite book |last1=Needham |first1=Joseph |first2=Robin |last2=Yates |year=1994 |title=Science and Civilisation in China, Volume 5: Chemistry and Chemical Technology, Part 6, Military Technology: Missiles and Sieges |location=Cambridge |publisher=Cambridge University Press |isbn=9780521327275 |oclc=489677531 |pages=120}}</ref>

{{multiple image
| align=
| direction=
| footer=Alchemical signs for sulfur, or the [[Chemistry:Combustibility and flammability|combustible]] elements, and brimstone, an older/archaic name for sulfur.<ref>{{cite book |last1=Koch |first1=Rudolf |title=The book of signs : which contains all manner of symbols used from the earliest times to the Middle Ages by primitive peoples and early Christians |date=1955 |location=New York |isbn=0-486-20162-7 |publisher=Dover Publications }}</ref>
| image1= Sulphur symbol (fixed width).svg
| caption1= Sulfur
| image2= Black sulfur symbol (fixed width).svg
| caption2=Brimstone
| total_width=280
}}
Indian alchemists, practitioners of the "science of chemicals" ({{lang-sa|रसशास्त्र|rasaśāstra}}), wrote extensively about the use of sulfur in alchemical operations with mercury, from the eighth century AD onwards.<ref name="white-alchemical">{{cite book |last=White |first=David Gordon |title=The Alchemical Body — Siddha Traditions in Medieval India |date=1996 |publisher=University of Chicago Press |location=Chicago |isbn=978-0-226-89499-7 |pages=passim}}</ref> In the [[Chemistry:Rasa shastra|{{transliteration|sa|rasaśāstra}}]] tradition, sulfur is called "the smelly" ({{lang|sa|गन्धक}}, {{transliteration|sa|gandhaka}}).

Early [[Earth:Europe|Europe]]an [[Philosophy:Alchemy|alchemists]] gave sulfur a unique [[Unsolved:Alchemical symbol|alchemical symbol]], a triangle atop a cross (🜍). (This is sometimes confused with the astronomical crossed-spear symbol ⚴ for [[Astronomy:2 Pallas|2 Pallas]].) The variation known as brimstone has a symbol combining a two-barred cross atop a [[Infinity symbol|lemniscate]] (🜏). In traditional skin treatment, elemental sulfur was used (mainly in creams) to alleviate such conditions as [[Medicine:Scabies|scabies]], ringworm, [[Medicine:Psoriasis|psoriasis]], eczema, and [[Engineering:Acne|acne]]. The mechanism of action is unknown—though elemental sulfur does oxidize slowly to sulfurous acid, which is (through the action of [[Chemistry:Sulfite|sulfite]]) a mild reducing and antibacterial agent.<ref>{{cite journal|doi= 10.1016/S0190-9622(88)70079-1|last1= Lin|first1= A. N.|last2= Reimer|first2= R. J.|last3= Carter|first3= D. M.|title= Sulfur revisited|journal= Journal of the American Academy of Dermatology|volume= 18|issue= 3|pages= 553–558|date= 1988|pmid= 2450900}}</ref><ref>{{cite journal|doi= 10.1016/S0190-9622(08)81225-X|last1= Maibach|first1= H. I.|last2= Surber|first2= C.|last3= Orkin|first3= M.|title= Sulfur revisited|journal= Journal of the American Academy of Dermatology|volume= 23|issue= 1|pages= 154–156|date= 1990 |pmid= 2365870|doi-access= free}}</ref><ref>{{cite journal|last1= Gupta|first1= A. K.|last2= Nicol|first2= K.|title= The use of sulfur in dermatology|journal= Journal of Drugs in Dermatology|volume= 3|issue= 4|pages= 427–31|date= 2004 |pmid= 15303787}}</ref>

===Modern times===
[[File:18810703 Sulphur Soap - advertisement in The Helena Independent.png|thumb|upright=0.8|Today sulfur is known to have antifungal, antibacterial, and keratolytic activity; in the past it was used against acne vulgaris, rosacea, seborrheic dermatitis, dandruff, pityriasis versicolor, scabies, and warts.<ref name="SulfurInDerm_2004">{{cite journal |last1=Gupta |first1=Aditya K |last2=Nicol |first2=Karyn |title=The Use of Sulfur in Dermatology |journal=Journal of Drugs in Dermatology |date=Jul–Aug 2004 |volume=3 |issue=4 |pages=427–431 |pmid=15303787 |url=https://pubmed.ncbi.nlm.nih.gov/15303787/}}</ref> This 1881 advertisement baselessly claims efficacy against rheumatism, gout, baldness, and graying of hair.]]
Sulfur appears in a column of fixed (non-acidic) [[Chemistry:Alkali|alkali]] in a chemical table of 1718.<ref>{{Cite book |last=Donovan |first=Arthur |url=https://books.google.com/books?id=0-fOasz2GUMC&pg=PA66|title=Antoine Lavoisier: Science, Administration and Revolution |publisher=Cambridge University Press|year=1996|isbn=978-0-521-56672-8|pages=66|language=en}}</ref> [[Biography:Antoine Lavoisier|Antoine Lavoisier]] used sulfur in combustion experiments, writing of some of these in 1777.<ref>{{Cite book |last=Poirier|first=Jean-Pierre |url=https://books.google.com/books?id=u5h0BAAAQBAJ&pg=PA107 |title=Lavoisier: Chemist, Biologist, Economist |publisher=University of Pennsylvania Press|year=1998 |isbn=978-0-8122-1649-3 |pages=107–8 |language=en}}</ref>

Sulfur deposits in [[Place:Sicily|Sicily]] were the dominant source for more than a century. By the late 18th century, about 2,000 tonnes per year of sulfur were imported into Marseille, France, for the production of [[Chemistry:Sulfuric acid|sulfuric acid]] for use in the [[Chemistry:Leblanc process|Leblanc process]]. In [[History:Industrial Revolution|industrializing]] Britain, with the repeal of [[Finance:Tariff|tariff]]s on salt in 1824, demand for sulfur from Sicily surged upward. The increasing British control and exploitation of the mining, refining, and transportation of the sulfur, coupled with the failure of this lucrative export to transform Sicily's backward and impoverished economy, led to the Sulfur Crisis of 1840, when King Ferdinand II gave a monopoly of the sulfur industry to a French firm, violating an earlier 1816 trade agreement with Britain. A peaceful solution was eventually negotiated by France.<ref>{{cite book |url=https://books.google.com/books?id=wZg4ecXXmNYC|title=Sicily and the Unification of Italy: Liberal Policy and Local Power, 1859–1866|author=Riall, Lucy|date=1998|publisher=Oxford University Press|access-date=7 February 2013 |isbn=9780191542619}}</ref><ref>{{cite journal|title=Prelude to the Sulphur War of 1840: The Neapolitan Perspective |journal=European History Quarterly|date=April 1995|volume=25|issue=2|pages=163–180 |doi=10.1177/026569149502500201|last1=Thomson|first1=D. W. |s2cid=145807900}}</ref>

In 1867, elemental sulfur was discovered in underground deposits in Louisiana and {{wipe|Texas}}. The highly successful [[Engineering:Frasch process|Frasch process]] was developed to extract this resource.<ref name="Frasch">{{cite journal|first= Walter|last= Botsch|title= Chemiker, Techniker, Unternehmer: Zum 150. Geburtstag von Hermann Frasch|journal= Chemie in unserer Zeit|date= 2001|volume= 35|issue= 5|language= de|pages= 324–331|doi= 10.1002/1521-3781(200110)35:5<324::AID-CIUZ324>3.0.CO;2-9}}</ref>

In the late 18th century, [[Engineering:Furniture|furniture]] makers used molten sulfur to produce decorative inlays.<ref name="Mass 2003">{{cite journal|doi=10.1088/0957-0233/14/9/311|issn=0957-0233 |title= Pennsylvania German sulfur-inlaid furniture: characterization, reproduction, and ageing phenomena of the inlays|journal= Measurement Science and Technology|volume= 14|issue= 9|pages= 1598|year= 2003|last1= Mass|first1= Jennifer L|last2= Anderson|first2= Mark J|s2cid=250882259 }}</ref> Molten sulfur is sometimes still used for setting steel bolts into drilled concrete holes where high shock resistance is desired for floor-mounted equipment attachment points. Pure powdered sulfur was used as a medicinal tonic and laxative.<ref name="Nehb" />

With the advent of the [[Chemistry:Contact process|contact process]], the majority of sulfur today is used to make sulfuric acid for a wide range of uses, particularly fertilizer.<ref>{{cite book|last=Kogel|first=Jessica|title=Industrial minerals & rocks: commodities, markets, and uses|date=2006|publisher=Littleton|location=Colorado|isbn=978-0-87335-233-8|edition=7th|page=935|oclc=62805047}}</ref>

In recent times, the main source of sulfur has become [[Chemistry:Petroleum|petroleum]] and [[Chemistry:Natural gas|natural gas]]. This is due to the requirement to remove sulfur from fuels in order to prevent [[Earth:Acid rain|acid rain]], and has resulted in a surplus of sulfur.<ref name=BBC/>

===Spelling and etymology===
''Sulfur'' is derived from the Latin word ''{{lang|la|sulpur}}'', which was [[Social:Hellenization|Hellenized]] to ''{{lang|la|sulphur}}'' in the erroneous belief that the Latin word came from Greek. This spelling was later reinterpreted as representing an /f/ sound and resulted in the spelling ''{{lang|la|sulfur}}'', which appears in Latin toward the end of the [[History:Classical antiquity|Classical period]]. The true Ancient Greek word for sulfur, {{lang|grc|θεῖον}}, ''theîon'' (from earlier {{lang|grc|θέειον}}, ''théeion''), is the source of the international chemical prefix ''thio-''. The Modern Standard Greek word for sulfur is θείο, ''theío''.

In 12th-century [[Social:Anglo-Norman language|Anglo-French]], it was ''{{lang|xno|sulfre}}''. In the 14th century, the erroneously Hellenized Latin ''{{lang|la|-ph-}}'' was restored in Middle English ''{{lang|enm|sulphre}}''. By the 15th century, both full Latin spelling variants ''sulfur'' and ''sulphur'' became common in English. The parallel ''f~ph'' spellings continued in Britain until the 19th century, when the word was standardized as ''sulphur''.<ref name="OED">{{Cite OED|sulphur}}</ref> On the other hand, ''sulfur'' was the form chosen in the United States, whereas Canada uses both.

The [[Organization:International Union of Pure and Applied Chemistry|IUPAC]] adopted the spelling ''sulfur'' in 1990 or 1971, depending on the source cited,<ref>{{Cite journal|title=So long sulphur|date=4 August 2009|journal=Nature Chemistry |volume=1 |issue=5 |pages=333 |doi=10.1038/nchem.301 |pmid=21378874 |bibcode=2009NatCh...1Q.333. |doi-access=free}}</ref> as did the Nomenclature Committee of the [[Organization:Royal Society of Chemistry|Royal Society of Chemistry]] in 1992, restoring the spelling ''sulfur'' to Britain.<ref>{{cite journal |last1=McNaught |first1=Alan |title=Journal style update |journal=The Analyst |volume=116 |issue=11 |page=1094 |date=1991 |doi=10.1039/AN9911601094 |bibcode=1991Ana...116.1094M}}</ref> Oxford Dictionaries note that "in chemistry and other technical uses ... the ''-f-'' spelling is now the standard form for this and related words in British as well as US contexts, and is increasingly used in general contexts as well."<ref>{{cite web |url=https://en.oxforddictionaries.com/definition/sulphur |archive-url=https://web.archive.org/web/20161120085132/https://en.oxforddictionaries.com/definition/sulphur |url-status=dead |archive-date=20 November 2016 |title=sulphur – definition of sulphur in English |website=Oxford Dictionaries |access-date=2016-11-19}}</ref>

==Production==
[[File:Soufre extraction 1.jpg|thumb|left|Sicilian kiln used to obtain sulfur from volcanic rock (diagram from a 1906 chemistry book)]]
[[File:Bergelut dengan asap nan beracun.jpg|thumb|upright |Traditional sulfur mining at [[Earth:Ijen|Ijen Volcano]], East Java, Indonesia. This image shows the dangerous and rugged conditions the miners face, including toxic smoke and high drops, as well as their lack of protective equipment. The pipes over which they are standing are for condensing sulfur vapors.]]
Sulfur may be found by itself and historically was usually obtained in this form; [[Chemistry:Pyrite|pyrite]] has also been a source of sulfur.<ref>{{cite book |last1=Riegel |first1=Emil |last2=Kent |first2=James |title=Kent and Riegel's Handbook of Industrial Chemistry and Biotechnology |volume=1 |date=2007 |isbn=978-0-387-27842-1 |oclc=74650396 |page=1171 |publisher=Springer |bibcode=2007karh.book......}}</ref> In volcanic regions in [[Place:Sicily|Sicily]], in ancient times, it was found on the surface of the Earth, and the "[[Engineering:Sicilian method|Sicilian process]]" was used: sulfur deposits were piled and stacked in brick kilns built on sloping hillsides, with airspaces between them. Then, some sulfur was pulverized, spread over the stacked ore and ignited, causing the free sulfur to melt down the hills. Eventually the surface-borne deposits played out, and miners excavated veins that ultimately dotted the Sicilian landscape with labyrinthine mines. Mining was unmechanized and labor-intensive, with pickmen freeing the ore from the rock, and mine-boys or ''carusi'' carrying baskets of ore to the surface, often through a mile or more of tunnels. Once the ore was at the surface, it was reduced and extracted in smelting ovens. The conditions in [[Chemistry:Sulfur mining in Sicily|Sicilian sulfur mines]] were horrific, prompting Booker T. Washington to write "I am not prepared just now to say to what extent I believe in a physical hell in the next world, but a sulfur mine in Sicily is about the nearest thing to hell that I expect to see in this life."<ref>{{cite book |last=Washington|first=Booker T.|title=The Man Farthest Down: A Record of Observation and Study in Europe |year=1912 |url=https://archive.org/stream/manfarthestdownr00wash#page/214 |publisher=Doubleday, Page.|page=214}}</ref> Sulfur is still mined from surface deposits in poorer nations with volcanoes, such as Indonesia, and worker conditions have not improved much since Booker T. Washington's days.<ref>{{cite journal |last=McElvaney|first=Kevin|date=25 February 2015 |title=The Men Who Mine Volcanos |journal=The Atlantic |url=https://www.theatlantic.com/features/archive/2015/02/the-men-who-mine-volcanoes-indonesia/385913/|access-date=26 February 2015}}</ref>

Elemental sulfur was extracted from [[Earth:Salt dome|salt dome]]s (in which it sometimes occurs in nearly pure form) until the late 20th century. Sulfur is now produced as a side product of other industrial processes such as in oil refining, in which sulfur is undesired. As a mineral, native sulfur under salt domes is thought to be a fossil mineral resource, produced by the action of anaerobic bacteria on sulfate deposits. It was removed from such salt-dome mines mainly by the [[Engineering:Frasch process|Frasch process]].<ref name="Nehb" /> In this method, superheated water was pumped into a native sulfur deposit to melt the sulfur, and then compressed air returned the 99.5% pure melted product to the surface. Throughout the 20th century this procedure produced elemental sulfur that required no further purification. Due to a limited number of such sulfur deposits and the high cost of working them, this process for mining sulfur has not been employed in a major way anywhere in the world since 2002.<ref name="desulf1">{{cite journal|last1= Eow|first1= John S.|title= Recovery of sulfur from sour acid gas: A review of the technology|journal= Environmental Progress|volume= 21|issue= 3|pages= 143–162|date= 2002|doi= 10.1002/ep.670210312|bibcode= 2002EnvPr..21..143E}}</ref><ref name="desulf2">{{cite journal |last1= Schreiner|first1= Bernhard|title= Der Claus-Prozess. Reich an Jahren und bedeutender denn je|journal= Chemie in unserer Zeit |volume= 42|issue= 6|pages= 378–392|date= 2008|doi= 10.1002/ciuz.200800461}}</ref>

[[File:AlbertaSulfurAtVancouverBC.jpg|thumb|left|Sulfur recovered from hydrocarbons in Alberta, stockpiled for shipment in North Vancouver, British Columbia]]
Today, sulfur is produced from petroleum, [[Chemistry:Natural gas|natural gas]], and related fossil resources, from which it is obtained mainly as [[Chemistry:Hydrogen sulfide|hydrogen sulfide]].<ref name=BBC/> Organosulfur compounds, undesirable impurities in petroleum, may be upgraded by subjecting them to [[Chemistry:Hydrodesulfurization|hydrodesulfurization]], which cleaves the C–S bonds:<ref name="desulf1" /><ref name="desulf2" />

{{block indent|R-S-R + 2 H<sub>2</sub> → 2 RH + H<sub>2</sub>S}}

The resulting hydrogen sulfide from this process, and also as it occurs in natural gas, is converted into elemental sulfur by the [[Chemistry:Claus process|Claus process]]. This process entails oxidation of some hydrogen sulfide to sulfur dioxide and then the [[Chemistry:Comproportionation|comproportionation]] of the two:<ref name="desulf1" /><ref name="desulf2" />

{{block indent|3 O<sub>2</sub> + 2 H<sub>2</sub>S → 2 SO<sub>2</sub> + 2 H<sub>2</sub>O}}
{{block indent|SO<sub>2</sub> + 2 H<sub>2</sub>S → 3 S + 2 H<sub>2</sub>O}}

[[File:Sulfur price world production.svg|thumb|upright=1.35 |Production and price (US market) of elemental sulfur]]
Owing to the high sulfur content of the Athabasca Oil Sands, stockpiles of elemental sulfur from this process now exist throughout Alberta, Canada.<ref name="Atha">{{cite book|last1= Hyndman|first1= A. W.|last2= Liu|first2= J. K.|last3= Denney|first3= D. W.|title= Sulfur: New Sources and Uses|volume= 183|pages= 69–82|date= 1982|doi= 10.1021/bk-1982-0183.ch005|chapter= Sulfur Recovery from Oil Sands|series= ACS Symposium Series|isbn= 978-0-8412-0713-4}}</ref> Another way of storing sulfur is as a [[Chemistry:Binder (material)|binder]] for concrete, the resulting product having some desirable properties (see [[Biology:Sulfur concrete|sulfur concrete]]).<ref name="Mohamed_2010">{{cite book|last1=Mohamed|first1=Abdel-Mohsen Onsy|last2=El-Gamal|first2=Maisa M.|title=Sulfur concrete for the construction industry: a sustainable development approach|date=2010|publisher=J. Ross |location=Fort Lauderdale|isbn=978-1-60427-005-1|oclc=531718953 |url= https://books.google.com/books?id=OYecyRmnTEkC&pg=PA104|pages=104–105, 109}}</ref>

The world production of sulfur in 2011 amounted to 69&nbsp;million tonnes (Mt), with more than 15 countries contributing more than 1 Mt each. Countries producing more than 5 Mt are {{wipe|China}} (9.6), the {{wipe|United States}} (8.8), {{wipe|Canada}} (7.1) and {{wipe|Russia}} (7.1).<ref>Apodaca, Lori E. (2012) [https://minerals.usgs.gov/minerals/pubs/commodity/sulfur/mcs-2012-sulfu.pdf "Sulfur"]. Mineral Commodity Summaries. United States Geological Survey.</ref> Production has been slowly increasing from 1900 to 2010; the price was unstable in the 1980s and around 2010.<ref name="USGS" />

==Applications==

===Sulfuric acid===
Elemental sulfur is used mainly as a precursor to other chemicals. Approximately 85% (1989) is converted to [[Chemistry:Sulfuric acid|sulfuric acid]] (H<sub>2</sub>SO<sub>4</sub>):

{{block indent|{{frac|1|8}} {{chem2|S8}} + {{frac|3|2}} {{chem2|O2}} + {{chem2|H2O}} → {{chem2|H2SO4}}}}

[[File:Sulfuric acid 2000.png|thumb|upright=1.35|Sulfuric acid production in 2000]]
In 2010, the United States produced more sulfuric acid than any other inorganic industrial chemical.<ref name="USGS">{{cite web|title= Mineral Yearbook 2010: Sulfur|author=Apodaca, Lori E. |publisher= United States Geological Survey|url= https://minerals.usgs.gov/minerals/pubs/commodity/sulfur/myb1-2010-sulfu.pdf}}</ref> The principal use for the acid is the extraction of phosphate ores for the production of fertilizer manufacturing. Other applications of sulfuric acid include oil refining, wastewater processing, and mineral extraction.<ref name="Nehb" />

===Other important sulfur chemistry===
Sulfur reacts directly with methane to give [[Chemistry:Carbon disulfide|carbon disulfide]], which is used to manufacture [[Chemistry:Cellophane|cellophane]] and [[Chemistry:Rayon|rayon]].<ref name="Nehb">{{cite book |last1=Nehb |first1=Wolfgang|last2=Vydra|first2=Karel |title=Ullmann's Encyclopedia of Industrial Chemistry |date=2006|publisher=Wiley-VCH Verlag|doi=10.1002/14356007.a25_507.pub2 |chapter=Sulfur |isbn=978-3-527-30673-2}}</ref> One of the uses of elemental sulfur is in [[Chemistry:Vulcanization|vulcanization]] of rubber, where [[Chemistry:Polysulfide|polysulfide]] chains crosslink organic polymers.<!--need something on pulping--> Large quantities of [[Chemistry:Sulfite|sulfite]]s are used to bleach [[Chemistry:Paper|paper]] and to preserve dried fruit. Many [[Biology:Surfactant|surfactant]]s and detergents (e.g. sodium lauryl sulfate) are sulfate derivatives. [[Chemistry:Calcium sulfate|Calcium sulfate]], gypsum (CaSO<sub>4</sub>·2H<sub>2</sub>O) is mined on the scale of 100&nbsp;million [[Tonne|tonne]]s each year for use in Portland cement and fertilizers.

When silver-based [[Photography|photography]] was widespread, sodium and ammonium [[Chemistry:Sodium thiosulfate|thiosulfate]] were widely used as "fixing agents". Sulfur is a component of [[Chemistry:Gunpowder|gunpowder]] ("black powder").

===Fertilizer===
[[Chemistry:Amino acid|Amino acid]]s synthesized by living organisms such as [[Chemistry:Methionine|methionine]] and [[Chemistry:Cysteine|cysteine]] contain [[Chemistry:Organosulfur compounds|organosulfur]] groups ([[Biology:Thioester|thioester]] and [[Chemistry:Thiol|thiol]] respectively). The [[Biology:Antioxidant|antioxidant]] [[Chemistry:Glutathione|glutathione]] protecting many living organisms against [[Biology:Free radical|free radical]]s and [[Biology:Oxidative stress|oxidative stress]] also contains organic sulfur. Some crops such as [[Biology:Onion|onion]] and [[Biology:Garlic|garlic]] also produce different [[Chemistry:Organosulfur compounds|organosulfur compounds]] such as [[Chemistry:Syn-Propanethial-S-oxide|syn-propanethial-S-oxide]] responsible of lacrymal irritation (onions), or [[Chemistry:Diallyl disulfide|diallyl disulfide]] and [[Chemistry:Allicin|allicin]] (garlic). [[Chemistry:Sulfate|Sulfate]]s, commonly found in [[Chemistry:Soil|soil]]s and [[Earth:Groundwater|groundwater]]s are often a sufficient natural source of sulfur for plants and bacteria. Atmospheric deposition of [[Chemistry:Sulfur dioxide|sulfur dioxide]] (SO<sub>2</sub>) is also a common artificial source (coal combustion) of sulfur for the soils. Under normal circumstances, in most agricultural soils, sulfur is not a [[Earth:Limiting factor|limiting nutrient]] for plants and [[Biology:Microorganism|microorganism]]s (see [[Earth:Liebig's law of the minimum#Liebig's barrel|Liebig's barrel]]). However, in some circumstances, soils can be depleted in [[Chemistry:Sulfate|sulfate]], e.g. if this later is leached by meteoric water (rain) or if the requirements in sulfur for some types of crops are high. This explains that sulfur is increasingly recognized and used as a component of [[Chemistry:Fertilizer|fertilizer]]s. The most important form of sulfur for fertilizer is [[Chemistry:Calcium sulfate|calcium sulfate]], commonly found in nature as the mineral [[Chemistry:Gypsum|gypsum]] (CaSO<sub>4</sub>·2H<sub>2</sub>O). Elemental sulfur is [[Chemistry:Hydrophobic|hydrophobic]] (not soluble in water) and cannot be used directly by plants. Elemental sulfur (ES) is sometimes mixed with [[Chemistry:Bentonite|bentonite]] to amend depleted soils for crops with high requirement in organo-sulfur. Over time, [[Chemistry:Redox|oxidation]] [[Earth:Abiotic component|abiotic]] processes with atmospheric [[Chemistry:Oxygen|oxygen]] and [[Biology:Soil microbiology|soil bacteria]] can [[Chemistry:Redox|oxidize]] and convert elemental sulfur to soluble derivatives, which can then be used by microorganisms and plants. Sulfur improves the efficiency of other essential plant nutrients, particularly [[Chemistry:Nitrate|nitrogen]] and phosphorus.<ref name="Sulfur as a fertilizer">{{Cite web |url=https://www.sulphurinstitute.org/about-sulphur/faq/#plants |title=FAQ – The Sulphur Institute |date=2020 |website=sulphurinstitute.org |publisher=[[Organization:The Sulphur Institute|The Sulphur Institute]] |access-date=27 February 2020}}</ref> Biologically produced sulfur particles are naturally [[Chemistry:Hydrophile|hydrophilic]] due to a [[Biology:Biopolymer|biopolymer]] coating and are easier to disperse over the land in a spray of diluted slurry, resulting in a faster uptake by plants.

The plants requirement for sulfur equals or exceeds the requirement for [[Chemistry:Phosphorus|phosphorus]]. It is an [[Biology:Plant nutrition|essential nutrient for plant]] growth, root nodule formation of legumes, and immunity and defense systems. Sulfur deficiency has become widespread in many countries in Europe.<ref>{{cite journal|doi= 10.1006/jcrs.1998.0241|title= Sulphur Assimilation and Effects on Yield and Quality of Wheat|date= 1999|last1= Zhao|first1= F.|journal= Journal of Cereal Science|volume= 30|issue= 1|pages= 1–17|last2= Hawkesford|first2= M. J.|last3= McGrath|first3= S. P.}}</ref><ref>{{cite journal|title= Diagnosing sulfur deficiency in field-grown oilseed rape (Brassica napus L.) and wheat (Triticum aestivum L.)|doi= 10.1023/A:1026503812267|date= 2000|last1= Blake-Kalff|first1= M. M. A.|journal= Plant and Soil|volume= 225|issue= 1/2|pages= 95–107|s2cid= 44208638}}</ref><ref>{{cite journal|doi= 10.1007/BF00747690|title= Plant nutrient sulphur-a review of nutrient balance, environmental impact and fertilizers|date= 1996|last1= Ceccotti|first1= S. P.|journal= Fertilizer Research|volume= 43|issue= 1–3|pages= 117–125|s2cid= 42207099}}</ref> Because atmospheric inputs of sulfur continue to decrease, the deficit in the sulfur input/output is likely to increase unless sulfur fertilizers are used. Atmospheric inputs of sulfur decrease because of actions taken to limit [[Earth:Acid rain|acid rain]]s.<ref name="NASA Glossary">{{Citation|title=Glossary |at=acid rain |publisher=[[Astronomy:NASA Earth Observatory|NASA Earth Observatory]] |location=United States |url=http://earthobservatory.nasa.gov/Glossary/?mode=all |access-date=February 15, 2013 |archive-url=https://web.archive.org/web/20111213175357/http://earthobservatory.nasa.gov/Glossary/?mode=all |archive-date=December 13, 2011 |url-status=live }}</ref><ref name="Sulfur as a fertilizer" />

===Fungicide and pesticide===<!--Wettable Sulfur redirs here-->
[[File:Sulphur Candle.jpg|thumb|upright=0.8|Sulfur candle originally sold for home fumigation]]
Elemental sulfur is one of the oldest fungicides and pesticides. "Dusting sulfur", elemental sulfur in powdered form, is a common fungicide for grapes, strawberry, many vegetables and several other crops. It has a good efficacy against a wide range of powdery mildew diseases as well as black spot. In organic production, sulfur is the most important fungicide. It is the only fungicide used in organically farmed apple production against the main disease [[Biology:Apple scab|apple scab]] under colder conditions. Biosulfur (biologically produced elemental sulfur with hydrophilic characteristics) can also be used for these applications.

Standard-formulation dusting sulfur is applied to crops with a sulfur duster or from a dusting plane. Wettable sulfur is the commercial name for dusting sulfur formulated with additional ingredients to make it water [[Chemistry:Miscibility|miscible]].<ref name="Mohamed_2010" /><ref>{{cite web|url=http://www.freepatentsonline.com/3398227.pdf|title=Method for Preparation of Wettable Sulfur|access-date= 20 May 2010|author= Every, Richard L.|display-authors=0|date= 20 August 1968}}</ref> It has similar applications and is used as a [[Biology:Fungicide|fungicide]] against [[Biology:Mildew|mildew]] and other mold-related problems with plants and soil.

Elemental sulfur powder is used as an "organic" (i.e., "green") [[Biology:Insecticide|insecticide]] (actually an [[Chemistry:Acaricide|acaricide]]) against [[Biology:Tick|tick]]s and [[Biology:Mite|mite]]s. A common method of application is dusting the clothing or limbs with sulfur powder.

A diluted solution of [[Chemistry:Lime sulfur|lime sulfur]] (made by combining [[Chemistry:Calcium hydroxide|calcium hydroxide]] with elemental sulfur in water) is used as a dip for pets to destroy ringworm (fungus), mange, and other dermatoses and [[Earth:Parasitism|parasites]].

Sulfur candles of almost pure sulfur were burned to fumigate structures and wine barrels, but are now considered too toxic for residences.

===Pharmaceuticals===
{{main|Medicine:Sulfur (pharmacy)}}
Sulfur (specifically [[Chemistry:Octasulfur|octasulfur]], S<sub>8</sub>) is used in pharmaceutical skin preparations for the treatment of [[Engineering:Acne|acne]] and other conditions. It acts as a keratolytic agent and also kills bacteria, fungi, [[Medicine:Scabies|scabies]] mites, and other parasites.<ref name="Hager">{{cite book|title=Hagers Handbuch der Pharmazeutischen Praxis|edition=4th|publisher=Springer|location=Berlin–Heidelberg–New York|language=de|date=1978|volume=6B|pages=672–9|isbn=978-3-540-07738-1}}</ref> Precipitated sulfur and colloidal sulfur are used, in form of [[Medicine:Lotion|lotion]]s, creams, powders, soaps, and bath additives, for the treatment of acne vulgaris, acne rosacea, and [[Medicine:Seborrhoeic dermatitis|seborrhoeic dermatitis]].<ref name="Arzneibuch-Kommentar">{{cite book|title=Arzneibuch-Kommentar. Wissenschaftliche Erläuterungen zum Europäischen Arzneibuch und zum Deutschen Arzneibuch|trans-title=Pharmacopoeia Commentary. Scientific annotations to the European Pharmacopoeia and the German Pharmacopoeia|publisher=Wissenschaftliche Verlagsgesellschaft|location=Stuttgart|date=2004|language=de|isbn=978-3-8047-2575-1|at=Monographie ''Schwefel zum äußerlichen Gebrauch'' [Monograph ''Sulfur for external use'']|edition=23rd}}</ref>

Many drugs contain sulfur.<ref>{{Citation |last1=Scott |first1=Kevin A. |last2=Njardarson |first2=Jon T. |title=Analysis of US FDA-Approved Drugs Containing Sulfur Atoms |date=2019 |url=https://par.nsf.gov/servlets/purl/10218659 |format=PDF |work=Sulfur Chemistry |series=Topics in Current Chemistry Collections |pages=1–34 |editor-last=Jiang |editor-first=Xuefeng |access-date=2023-03-08 |publisher=Springer International Publishing |language=en |doi=10.1007/978-3-030-25598-5_1 |isbn=978-3-030-25598-5 }}</ref> Early examples include antibacterial [[Medicine:Sulfonamide|sulfonamides]], known as ''sulfa drugs''. A more recent example is mucolytic [[Chemistry:Acetylcysteine|acetylcysteine]]. Sulfur is a part of many bacterial defense molecules. Most β-lactam antibiotics, including the [[Chemistry:Penicillin|penicillin]]s, cephalosporins and [[Chemistry:Monobactam|monobactam]]s contain sulfur.<ref name="Cremlyn" />

=== Batteries ===
Due to their high energy density and the availability of sulfur, there is ongoing research in creating rechargeable [[Engineering:Lithium–sulfur battery|lithium–sulfur batteries]]. Until now, carbonate electrolytes have caused failures in such batteries after a single cycle. In February 2022, researchers at [[Organization:Drexel University|Drexel University]] have not only created a prototypical battery that lasted 4000 recharge cycles, but also found the first monoclinic gamma sulfur that remained stable below 95 degrees Celsius.<ref>{{Cite journal |title=Stabilization of gamma sulfur at room temperature to enable the use of carbonate electrolyte in Li–S batteries |journal=Communications Chemistry |last=Pai |first=Rahul |date=2022-02-10 |volume=5 |issue=1 |page=17 |doi=10.1038/s42004-022-00626-2|pmid=36697747 |pmc=9814344 |s2cid=246704531 |doi-access=free }}</ref>

==Biological role==
Sulfur is an essential component of all living [[Biology:Cell|cells]]. It is the eighth most abundant element in the human body by weight,<ref name="tsihb">{{cite news |title=Sulphur and the Human Body |url=https://www.sulphurinstitute.org/pub/?id=8c64bf34-bc30-5bd9-0719-f6de83f7e841 |access-date=3 April 2021 |publisher=The Sulfur Institute}}</ref> about equal in abundance to [[Chemistry:Potassium|potassium]], and slightly greater than [[Chemistry:Sodium|sodium]] and [[Chemistry:Chlorine|chlorine]].<ref>{{Cite web |title=What is the body made of? |url=https://www.newscientist.com/question/what-is-the-body-made-of/ |url-status=live |archive-url=https://web.archive.org/web/20211103173257/https://www.newscientist.com/question/what-is-the-body-made-of/ |archive-date=November 3, 2021 |access-date=November 9, 2021 |website=New Scientist}}</ref> A {{convert|70|kg|abbr=on}} human body contains about 140&nbsp;grams of sulfur.<ref>{{Cite web |last=Helmenstine |first=Anne |date=February 3, 2019 |title=Elemental Composition of the Human Body by Mass |url=https://www.thoughtco.com/elemental-composition-human-body-by-mass-608192 |url-status=live |archive-url=https://web.archive.org/web/20210413013702/https://www.thoughtco.com/elemental-composition-human-body-by-mass-608192 |archive-date=April 13, 2021 |access-date=November 21, 2021 |website=ThoughtCo.}}</ref> The main dietary source of sulfur for humans is sulfur-containing amino-acids,<ref name="Parcell-2002">{{Cite journal |last=Parcell |first=Stephen |date=February 2002 |title=Sulfur in human nutrition and applications in medicine |url=https://pubmed.ncbi.nlm.nih.gov/11896744/ |journal=Alternative Medicine Review |volume=7 |issue=1 |pages=22–44 |issn=1089-5159 |pmid=11896744}}</ref> which can be found in plant and animal proteins.<ref name="Ingenbleek-2013">{{Cite journal |last1=Ingenbleek |first1=Yves |last2=Kimura |first2=Hideo |date=July 2013 |title=Nutritional essentiality of sulfur in health and disease |journal=Nutrition Reviews |volume=71 |issue=7 |pages=413–432 |doi=10.1111/nure.12050 |issn=1753-4887 |pmid=23815141|doi-access=free }}</ref>

=== Transferring sulfur between inorganic and biomolecules ===

In the 1880s, while studying [[Biology:Beggiatoa|Beggiatoa]] (a bacterium living in a sulfur rich environment), Sergei Winogradsky found that it oxidized [[Chemistry:Hydrogen sulfide|hydrogen sulfide]] (H<sub>2</sub>S) as an energy source, forming intracellular sulfur droplets. Winogradsky referred to this form of metabolism as inorgoxidation (oxidation of inorganic compounds).<ref>{{Cite journal |last=Dworkin |first=Martin |date=March 2012 |title=Sergei Winogradsky: a founder of modern microbiology and the first microbial ecologist |journal=FEMS Microbiology Reviews |volume=36 |issue=2 |pages=364–379 |doi=10.1111/j.1574-6976.2011.00299.x |issn=1574-6976 |pmid=22092289|doi-access=free }}</ref> Another contributor, who continued to study it was [[Biography:Selman Waksman|Selman Waksman]].<ref>{{Cite journal |last1=Waksman |first1=S. A. |last2=Starkey |first2=R. L. |title=On the Growth and Respiration of Sulfur-Oxidizing Bacteria |date=1923-01-20 |journal=The Journal of General Physiology |volume=5 |issue=3 |pages=285–310 |doi=10.1085/jgp.5.3.285 |issn=0022-1295 |pmc=2140527 |pmid=19871997}}</ref> Primitive bacteria that live around deep ocean [[Earth:Hydrothermal vent|volcanic vents]] oxidize hydrogen sulfide for their nutrition, as discovered by Robert Ballard.<ref name="podcast" />

Sulfur oxidizers can use as energy sources reduced sulfur compounds, including hydrogen sulfide, elemental sulfur, [[Chemistry:Sulfite|sulfite]], [[Chemistry:Thiosulfate|thiosulfate]], and various polythionates (e.g., [[Chemistry:Tetrathionate|tetrathionate]]).<ref>{{cite journal |author= Pronk JT |author2= Meulenberg R |author3= Hazeu W |author4= Bos P |author5= Kuenen JG |date= 1990 |title= Oxidation of reduced inorganic sulphur compounds by acidophilic thiobacilli |journal= FEMS Microbiology Letters |volume= 75 |issue= 2–3 |pages= 293–306 |doi= 10.1111/j.1574-6968.1990.tb04103.x |df= dmy-all|doi-access= free }}</ref> They depend on enzymes such as [[Biology:Sulfur dioxygenase|sulfur oxygenase]] and [[Biology:Sulfite oxidase|sulfite oxidase]] to oxidize sulfur to sulfate. Some [[Biology:Lithotroph|lithotroph]]s can even use the energy contained in sulfur compounds to produce sugars, a process known as [[Earth:Chemosynthesis|chemosynthesis]]. Some [[Biology:Bacteria|bacteria]] and [[Biology:Archaea|archaea]] use hydrogen sulfide in place of water as the [[Chemistry:Electron donor|electron donor]] in chemosynthesis, a process similar to [[Earth:Photosynthesis|photosynthesis]] that produces sugars and uses oxygen as the [[Chemistry:Electron acceptor|electron acceptor]]. Sulfur-based chemosynthesis may be simplifiedly compared with photosynthesis:

{{block indent|H<sub>2</sub>S + CO<sub>2</sub> → sugars + S}}
{{block indent|H<sub>2</sub>O + CO<sub>2</sub> → sugars + O<sub>2</sub>}}

There are bacteria combining these two ways of nutrition: [[Biology:Green sulfur bacteria|green sulfur bacteria]] and [[Biology:Purple sulfur bacteria|purple sulfur bacteria]].<ref>{{Citation |last1=Frigaard |first1=Niels-Ulrik |title=Sulfur Metabolism in Phototrophic Sulfur Bacteria |date=2008-01-01 |url=https://www.sciencedirect.com/science/article/pii/S0065291108000027 |volume=54 |pages=103–200 |editor-last=Poole |editor-first=Robert K. |publisher=Academic Press |language=en |access-date=2022-05-17 |last2=Dahl |first2=Christiane|series=Advances in Microbial Physiology |doi=10.1016/S0065-2911(08)00002-7 |pmid=18929068 |isbn=9780123743237 }}</ref> Also sulfur-oxidizing bacteria can go into symbiosis with larger organisms, enabling the later to use hydrogen sulfide as food to be oxidized. Example: the giant tube worm.<ref>{{Cite journal |last=Cavanaugh |first=Colleen M. |date=1994 |title=Microbial Symbiosis: Patterns of Diversity in the Marine Environment |journal=American Zoologist |volume=34 |pages=79–89 |doi=10.1093/icb/34.1.79 |doi-access=free }}</ref>

There are sulfate-reducing bacteria, that, by contrast, "breathe sulfate" instead of oxygen. They use organic compounds or molecular hydrogen as the energy source. They use sulfur as the electron acceptor, and reduce various oxidized sulfur compounds back into sulfide, often into hydrogen sulfide. They can grow on other partially oxidized sulfur compounds (e.g. thiosulfates, thionates, polysulfides, sulfites).

There are studies pointing that many deposits of native sulfur in places that were the bottom of [[Earth:Tethys Ocean|the ancient oceans]] have biological origin.<ref>{{Cite journal |last1=Jones |first1=Galen E. |last2=Starkey |first2=Robert L. |last3=Feely |first3=Herbert W. |last4=Kulp |first4=J. Laurence |date=1956-06-22 |title=Biological Origin of Native Sulfur in Salt Domes of Texas and Louisiana |url=https://www.science.org/doi/10.1126/science.123.3208.1124 |journal=Science |language=en |volume=123 |issue=3208 |pages=1124–1125 |doi=10.1126/science.123.3208.1124 |pmid=17793426 |bibcode=1956Sci...123.1124J |issn=0036-8075}}</ref><ref>{{Cite journal |last1=Philip |first1=G. |last2=Wali |first2=A. M. A. |last3=Aref |first3=M. A. M. |date=1994-09-01 |title=On the origin of native sulfur deposits in Gebel El Zeit, Gulf of Suez, Egypt |url=https://doi.org/10.1007/BF03175232 |journal=Carbonates and Evaporites |language=en |volume=9 |issue=2 |pages=223–232 |doi=10.1007/BF03175232 |bibcode=1994CarEv...9..223P |s2cid=128827551 |issn=1878-5212}}</ref><ref>{{Cite web |title=Petrography and mineralogy of the crystalline limestone of Fatha Formation from Mishraq area, Iraq |url=https://www.researchgate.net/publication/330038794 |access-date=2022-04-15 |website=ResearchGate |language=en}}</ref> These studies indicate that this native sulfur have been obtained through biological activity, but what is responsible for that (sulfur-oxidizing bacteria or sulfate-reducing bacteria) is still unknown for sure.

Sulfur is absorbed by [[Biology:Plant|plant]]s [[Biology:Root|root]]s from soil as [[Chemistry:Sulfate|sulfate]] and transported as a phosphate ester. Sulfate is reduced to sulfide via sulfite before it is incorporated into [[Chemistry:Cysteine|cysteine]] and other organosulfur compounds.<ref name="Heldt">{{cite book |last=Heldt |first=Hans-Walter |title=Pflanzenbiochemie |publisher=Spektrum Akademischer Verlag |year=1996 |isbn=978-3-8274-0103-8 |place=Heidelberg |pages=321–333 |language=de}}</ref>

{{block indent|SO<sub>4</sub><sup>2−</sup> → SO<sub>3</sub><sup>2−</sup> → H<sub>2</sub>S → cysteine (thiol) → methionine (thioether)}}

While the plants' role in transferring sulfur to animals by food chains is more or less understood, the role of sulfur bacteria is just getting investigated.<ref>{{Cite journal |last1=Kuenen |first1=J. G. |last2=Beudeker |first2=R. F. |date=1982-09-13 |title=Microbiology of thiobacilli and other sulphur-oxidizing autotrophs, mixotrophs and heterotrophs |url=https://pubmed.ncbi.nlm.nih.gov/6127737/ |journal=Philosophical Transactions of the Royal Society of London. Series B, Biological Sciences |volume=298 |issue=1093 |pages=473–497 |doi=10.1098/rstb.1982.0093 |issn=0962-8436 |pmid=6127737|bibcode=1982RSPTB.298..473K }}</ref><ref>{{Cite journal |last1=Wasmund |first1=Kenneth |last2=Mußmann |first2=Marc |last3=Loy |first3=Alexander |date=August 2017 |title=The life sulfuric: microbial ecology of sulfur cycling in marine sediments: Microbial sulfur cycling in marine sediments |journal=Environmental Microbiology Reports |language=en |volume=9 |issue=4 |pages=323–344 |doi=10.1111/1758-2229.12538 |pmc=5573963 |pmid=28419734}}</ref>

===Protein and organic metabolites===
In all forms of life, most of the sulfur is contained in two [[Chemistry:Proteinogenic amino acid|proteinogenic amino acid]]s ([[Chemistry:Cysteine|cysteine]] and [[Chemistry:Methionine|methionine]]), thus the element is present in all [[Biology:Protein|protein]]s that contain these amino acids, as well as in respective [[Chemistry:Peptide|peptide]]s.<ref>{{Cite journal |last1=Gutiérrez-Preciado |first1=A. |last2=Romero |first2=H. |last3=Peimbert |first3=M. |date=2010 |title=An Evolutionary Perspective on Amino Acids |url=https://www.nature.com/scitable/topicpage/an-evolutionary-perspective-on-amino-acids-14568445/ |journal=Nature Education |volume=3 |issue=9 |page=29}}</ref> Some of the sulfur is comprised in certain metabolites — many of which are [[Biology:Cofactor (biochemistry)|cofactors]] — and sulfated polysaccharides of connective tissue ([[Chemistry:Chondroitin sulfate|chondroitin sulfate]]s, [[Chemistry:Heparin|heparin]]).
[[File:Disulfide-bridges-made-in-Avogadro.png|alt=Disulfide bonds between two alpha-helix|left|thumb|Schematic representation of disulfide bridges (in yellow) between two protein helices]]
Proteins, to execute their [[Biology:Protein#Cellular functions|biological function]], need to have specific space geometry. Formation of this geometry is performed in a process called [[Biology:Protein folding|protein folding]], and is provided by intra- and inter-molecular bonds. The process has several stages. While at premier stages a polypeptide chain folds due to [[Chemistry:Hydrogen bond|hydrogen bond]]s, at later stages folding is provided (apart from hydrogen bonds) by [[Chemistry:Covalent bond|covalent bond]]s between two sulfur atoms of two cysteine residues (so called disulfide bridges) at different places of a chain (tertiary protein structure) as well as between two cysteine residues in two separated protein subunits (quaternary protein structure). Both structures easily may be seen in [[Biology:Insulin|insulin]]. As the [[Physics:Bond energy|bond energy]] of a covalent disulfide bridge is higher than the energy of a [[Chemistry:Coordinate covalent bond|coordinate bond]] or hydrophobic interaction, higher disulfide bridges content leads to higher energy needed for protein [[Biology:Denaturation (biochemistry)|denaturation]]. In general disulfide bonds are necessary in proteins functioning outside cellular space, and they do not change proteins' conformation (geometry), but serve as its stabilizers.<ref>{{Cite book |last1=Alberts |first1=Bruce |url=https://www.ncbi.nlm.nih.gov/books/NBK26830/#_A436_ |title=Molecular Biology of the Cell. 4th edition. |last2=Johnson |first2=Alexander |last3=Lewis |first3=Julian |last4=Raff |first4=Martin |last5=Roberts |first5=Keith |last6=Walter |first6=Peter |publisher=Garland Science |year=2002 |isbn=978-0-8153-3218-3 |location=New York |language=en}}</ref> Within [[Biology:Cytoplasm|cytoplasm]] cysteine residues of proteins are saved in reduced state (i.e. in -SH form) by [[Biology:Thioredoxin|thioredoxin]]s.<ref>{{Cite journal |last1=Arnér |first1=Elias S. J. |last2=Holmgren |first2=Arne |date=25 December 2001 |title=Physiological functions of thioredoxin and thioredoxin reductase: Thioredoxin and thioredoxin reductase |journal=European Journal of Biochemistry |language=en |volume=267 |issue=20 |pages=6102–6109 |doi=10.1046/j.1432-1327.2000.01701.x |pmid=11012661 |doi-access=free }}</ref>

This property manifests in following examples. [[Biology:Lysozyme|Lysozyme]] is stable enough to be applied as a drug.<ref>{{Cite web |title=Lysozyme |url=https://www.drugs.com/international/lysozyme.html |access-date=2022-05-19 |website=Drugs.com |language=en}}</ref> Feathers and hair have relative strength, and consisting in them [[Biology:Keratin|keratin]] is considered indigestible by most organisms. However, there are fungi and bacteria containing [[Biology:Keratinase|keratinase]], and are able to destruct keratin.

Many important cellular enzymes use prosthetic groups ending with -SH moieties to handle reactions involving acyl-containing biochemicals: two common examples from basic metabolism are [[Chemistry:Coenzyme A|coenzyme A]] and alpha-lipoic acid.<ref name="Lehn">{{cite book|isbn= 978-1-57259-153-0|last1= Nelson|first1= D. L.|last2= Cox|first2= M. M.|title= Lehninger, Principles of Biochemistry|edition= 3rd|publisher= Worth Publishing|place= New York|date= 2000|url-access= registration|url= https://archive.org/details/lehningerprincip01lehn}}</ref> Cysteine-related metabolites [[Chemistry:Homocysteine|homocysteine]] and [[Chemistry:Taurine|taurine]] are other sulfur-containing amino acids that are similar in structure, but not coded by [[Biology:DNA|DNA]], and are not part of the primary structure of proteins, take part in various locations of mammalian physiology.<ref>{{Cite journal |last=Selhub |first=J. |date=1999-07-01 |title=Homocysteine metabolism |url=https://www.annualreviews.org/doi/10.1146/annurev.nutr.19.1.217 |journal=Annual Review of Nutrition |volume=19 |issue=1 |pages=217–246 |doi=10.1146/annurev.nutr.19.1.217 |pmid=10448523 |issn=0199-9885}}</ref><ref>{{Cite journal |last=Huxtable |first=R. J. |date=1992-01-01 |title=Physiological actions of taurine |url=https://journals.physiology.org/doi/abs/10.1152/physrev.1992.72.1.101 |journal=Physiological Reviews |volume=72 |issue=1 |pages=101–163 |doi=10.1152/physrev.1992.72.1.101 |pmid=1731369 |issn=0031-9333}}</ref> Two of the 13 classical vitamins, [[Chemistry:Biotin|biotin]] and [[Physics:Thiamine|thiamine]], contain sulfur, and serve as cofactors to several enzymes.<ref>{{Cite web |title=The Function of Biotin |url=https://www.chem.uwec.edu/webpapers2001/barkacs/pages/function.html |access-date=2022-06-10 |website=www.chem.uwec.edu}}</ref><ref>{{Cite web |last=Edwards |first=Katie A. |title=Thiamine Biochemistry |url=http://thiamine.dnr.cornell.edu/Thiamine_biochemistry.html |access-date=2022-06-10 |website=thiamine.dnr.cornell.edu}}</ref>

In intracellular chemistry, sulfur operates as a carrier of reducing hydrogen and its electrons for cellular repair of oxidation. Reduced [[Chemistry:Glutathione|glutathione]], a sulfur-containing tripeptide, is a reducing agent through its sulfhydryl (-SH) moiety derived from [[Chemistry:Cysteine|cysteine]].

[[Engineering:Methanogenesis|Methanogenesis]], the route to most of the world's methane, is a multistep biochemical transformation of [[Chemistry:Carbon dioxide|carbon dioxide]]. This conversion requires several organosulfur cofactors. These include [[Chemistry:Coenzyme M|coenzyme M]], CH<sub>3</sub>SCH<sub>2</sub>CH<sub>2</sub>SO<sub>3</sub><sup>−</sup>, the immediate precursor to [[Chemistry:Methane|methane]].<ref>{{cite journal|last1= Thauer|first1= R. K.|title= Biochemistry of methanogenesis: a tribute to Marjory Stephenson:1998 Marjory Stephenson Prize Lecture|journal= Microbiology|volume= 144|issue= 9|pages= 2377–2406|date= 1998|pmid= 9782487|doi= 10.1099/00221287-144-9-2377|doi-access= free}}</ref>

===Metalloproteins and inorganic cofactors===
Metalloproteins — in which the active site is a transition metal ion (or metal-sulfide cluster) often coordinated by sulfur atoms of cysteine residues<ref>{{Cite journal |last1=Pace |first1=Nicholas J. |last2=Weerapana |first2=Eranthie |date=2014-04-17 |title=Zinc-binding cysteines: diverse functions and structural motifs |journal=Biomolecules |volume=4 |issue=2 |pages=419–434 |doi=10.3390/biom4020419 |issn=2218-273X |pmc=4101490 |pmid=24970223|doi-access=free }}</ref> — are essential components of enzymes involved in electron transfer processes. Examples include [[Biology:Plastocyanin|plastocyanin]] (Cu<sup>2+</sup>) and [[Biology:Nitrous-oxide reductase|nitrous oxide reductase]] (Cu–S). The function of these enzymes is dependent on the fact that the transition metal ion can undergo redox reactions. Other examples include many zinc proteins,<ref>{{Cite journal |last1=Giles |first1=Niroshini M |last2=Watts |first2=Aaron B |last3=Giles |first3=Gregory I |last4=Fry |first4=Fiona H |last5=Littlechild |first5=Jennifer A |last6=Jacob |first6=Claus |date=2003-08-01 |title=Metal and Redox Modulation of Cysteine Protein Function |journal=Chemistry & Biology |language=en |volume=10 |issue=8 |pages=677–693 |doi=10.1016/S1074-5521(03)00174-1 |pmid=12954327 |issn=1074-5521|doi-access=free }}</ref> as well as [[Chemistry:Iron–sulfur cluster|iron–sulfur cluster]]s. Most pervasive are the ferrodoxins, which serve as electron shuttles in cells. In bacteria, the important [[Biology:Nitrogenase|nitrogenase]] enzymes contain an Fe–Mo–S cluster and is a [[Chemistry:Catalyst|catalyst]] that performs the important function of [[Biology:Nitrogen fixation|nitrogen fixation]], converting atmospheric nitrogen to ammonia that can be used by microorganisms and plants to make proteins, DNA, RNA, alkaloids, and the other organic nitrogen compounds necessary for life.<ref>{{cite book|isbn= 978-0-935702-73-6 |first1= S. J.|last1= Lippard|first2= J. M.|last2= Berg|title= Principles of Bioinorganic Chemistry|publisher= University Science Books|date= 1994}}</ref>
:[[File:FdRedox.png|center|upright=3|Easiness of electron flow in a cluster provides catalytic effect of a respective enzyme.|thumb]]

===Deficiency===
In humans [[Chemistry:Methionine|methionine]] is an [[Chemistry:Essential amino acid|essential amino acid]]; [[Chemistry:Cysteine|cysteine]] is conditionally essential and may be synthesized from non-essential [[Chemistry:Serine|serine]] (sulfur donor would be methionine in this case). Dietary deficiency rarely happens in common conditions. Artificial methionine deficiency is attempted to apply in cancer treatment,<ref>{{Cite journal |last1=Mladenović |first1=Dušan |last2=Radosavljević |first2=Tatjana |last3=Hrnčić |first3=Dragan |last4=Rasic-Markovic |first4=Aleksandra |last5=Stanojlović |first5=Olivera |date=2019-07-26 |title=The effects of dietary methionine restriction on the function and metabolic reprogramming in the liver and brain – implications for longevity |url=https://pubmed.ncbi.nlm.nih.gov/30817309/ |journal=Reviews in the Neurosciences |volume=30 |issue=6 |pages=581–593 |doi=10.1515/revneuro-2018-0073 |issn=2191-0200 |pmid=30817309|s2cid=73470156 }}</ref> but the method is still potentially dangerous.<ref>{{Cite journal |last1=Binz |first1=Regina L. |last2=Sadhukhan |first2=Ratan |last3=Miousse |first3=Isabelle R. |last4=Garg |first4=Sarita |last5=Koturbash |first5=Igor |last6=Zhou |first6=Daohong |last7=Hauer-Jensen |first7=Martin |last8=Pathak |first8=Rupak |date=2021-02-27 |title=Dietary Methionine Deficiency Enhances Genetic Instability in Murine Immune Cells |journal=International Journal of Molecular Sciences |volume=22 |issue=5 |pages=2378 |doi=10.3390/ijms22052378 |issn=1422-0067 |pmc=7956689 |pmid=33673497|doi-access=free }}</ref>

Isolated sulfite oxidase deficiency is a rare, fatal genetic disease preventing production of [[Biology:Sulfite oxidase|sulfite oxidase]], needed to metabolize sulfites to sulfates.<ref>{{Cite journal |last1=Karakas |first1=Erkan |last2=Kisker |first2=Caroline |date=2005-10-18 |title=Structural analysis of missense mutations causing isolated sulfite oxidase deficiency |url=https://pubs.rsc.org/en/content/articlelanding/2005/dt/b505789m |journal=Dalton Transactions |language=en |issue=21 |pages=3459–3463 |doi=10.1039/B505789M |pmid=16234925 |issn=1477-9234}}</ref>

==Precautions==
{{Chembox
| container_only = yes
| Section7= {{Chembox Hazards
| ExternalSDS=
| GHSPictograms= {{GHS07}} {{GHS02}}
| GHSSignalWord= Warning
| HPhrases= {{H-phrases|315}}<ref>{{Cite web|url=https://www.sigmaaldrich.com/catalog/product/sigald/84683|title=Sulfur 84683|website=S}}</ref>
| NFPA-H= 2
| NFPA-F= 1
| NFPA-R= 0
| NFPA-S=
| NFPA_ref= <ref>{{Cite web|url=https://cameochemicals.noaa.gov/chemical/4562|title=Chemical Datasheet. SULFUR, MOLTEN}}</ref>
}}
}}
[[File:Acid rain woods1.JPG|Effect of acid rain on a forest, Jizera Mountains, Czech Republic|thumb]]
Though elemental sulfur is only minimally absorbed through the skin and is of low toxicity to humans, inhalation of sulfur dust or contact with eyes or skin may cause irritation. Excessive ingestion of sulfur can cause a burning sensation or diarrhea,<ref>{{Cite web |title=Sulfur General Fact Sheet |url=http://npic.orst.edu/factsheets/sulfurgen.html |access-date=2022-09-02 |website=npic.orst.edu}}</ref> and cases of life-threatening metabolic acidosis have been reported after patients deliberately consumed sulfur as a folk remedy.<ref>{{Cite journal |last1=Schwartz |first1=Steven M. |last2=Carroll |first2=Hugh M. |last3=Scharschmidt |first3=Linda A. |date=1986-07-01 |title=Sublimed (Inorganic) Sulfur Ingestion: A Cause of Life-Threatening Metabolic Acidosis With a High Anion Gap |url=https://doi.org/10.1001/archinte.1986.00360190229034 |journal=Archives of Internal Medicine |volume=146 |issue=7 |pages=1437–1438 |doi=10.1001/archinte.1986.00360190229034 |pmid=3718141 |issn=0003-9926}}</ref><ref>{{Cite journal |last1=Blum |first1=J. Eric |last2=Coe |first2=Fredric L. |date=2010-01-13 |title=Metabolic Acidosis after Sulfur Ingestion |journal=New England Journal of Medicine |url=https://www.nejm.org/doi/10.1056/NEJM197710202971606 |access-date=2022-09-02 |volume=297 |issue=16 |pages=869–870 |language=EN |doi=10.1056/nejm197710202971606|pmid=904661 }}</ref>

===Toxicity of sulfur compounds===
When sulfur burns in air, it produces [[Chemistry:Sulfur dioxide|sulfur dioxide]]. In water, this gas produces sulfurous acid and sulfites; sulfites are antioxidants that inhibit growth of aerobic bacteria and a useful [[Chemistry:Food additive|food additive]] in small amounts. At high concentrations these acids harm the lungs, eyes, or other tissues.<ref>{{Cite web |title=Sulfur Dioxide {{!}} Toxicological Profile {{!}} ATSDR |url=https://wwwn.cdc.gov/TSP/ToxProfiles/ToxProfiles.aspx?id=253&tid=46 |date= March 26, 2014 |access-date=2023-10-24 |website=Centers for Disease Control and Prevention }}</ref> In organisms without lungs such as insects, sulfite in high concentration prevents [[Biology:Respiration (physiology)|respiration]].<ref>{{Cite web |title=Sulfur General Fact Sheet |url=http://npic.orst.edu/factsheets/sulfurgen.html#:~:text=However,%20the%20main%20theory%20is,cells%20and%20affects%20cell%20respiration.&text=Sulfur%20can%20kill%20insects%20if%20they%20touch%20it%20or%20eat%20it.&text=It%20disrupts%20their%20normal%20body,their%20ability%20to%20produce%20energy. |access-date=2024-01-20 |website=npic.orst.edu}}</ref>

[[Chemistry:Sulfur trioxide|Sulfur trioxide]] (made by catalysis from sulfur dioxide) and [[Chemistry:Sulfuric acid|sulfuric acid]] are similarly highly acidic and corrosive in the presence of water. Concentrated sulfuric acid is a strong dehydrating agent that can strip available water molecules and water components from sugar and organic tissue.<ref>{{cite magazine |last=Baker |first=Colin |date=1 March 2007 |title=The dehydration of sucrose |url=https://eic.rsc.org/exhibition-chemistry/the-dehydration-of-sucrose/2020073.article |magazine=[[Chemistry:Education in Chemistry|Education in Chemistry]] |publisher=[[Organization:Royal Society of Chemistry|Royal Society of Chemistry]] |access-date=14 June 2018}}</ref>

The burning of [[Earth:Coal|coal]] and/or [[Chemistry:Petroleum|petroleum]] by industry and power plants generates sulfur dioxide (SO<sub>2</sub>) that reacts with atmospheric water and oxygen to produce sulfuric acid (H<sub>2</sub>SO<sub>4</sub>) and [[Chemistry:Sulfurous acid|sulfurous acid]] (H<sub>2</sub>SO<sub>3</sub>).<ref>{{Cite web |date=2021-02-28 |title=4.1: Fossil Fuels and Products of Combustion |url=https://eng.libretexts.org/Bookshelves/Environmental_Engineering_(Sustainability_and_Conservation)/Energy_Conservation_for_Environmental_Protection_(Pisupati)/04%3A_Energy_and_the_Environment/4.01%3A_New_Page |access-date=2023-10-24 |website=Engineering LibreTexts |language=en}}</ref> These acids are components of [[Earth:Acid rain|acid rain]], lowering the [[Chemistry:PH|pH]] of [[Chemistry:Soil|soil]] and freshwater bodies, sometimes resulting in substantial damage to the environment and chemical weathering of statues and structures. Fuel standards increasingly require that fuel producers extract sulfur from fossil fuels to prevent acid rain formation. This extracted and refined sulfur represents a large portion of sulfur production. In coal-fired power plants, [[Physics:Flue gas|flue gas]]es are sometimes purified. More modern power plants that use synthesis gas extract the sulfur before they burn the gas.

[[Chemistry:Hydrogen sulfide|Hydrogen sulfide]] is about one-half as toxic as [[Chemistry:Hydrogen cyanide|hydrogen cyanide]], and intoxicates by the same mechanism (inhibition of the respiratory enzyme cytochrome oxidase),<ref>{{Cite journal|journal=Medscape|url=https://emedicine.medscape.com/article/815139-overview#a4|title=Hydrogen Sulfide Toxicity: Practice Essentials, Pathophysiology, Etiology|date=30 March 2017|via=eMedicine}}</ref> though hydrogen sulfide is less likely to cause sudden poisonings from small inhaled amounts (near its [[Earth:Permissible exposure limit|permissible exposure limit]] (PEL) of 20 ppm) because of its disagreeable odor.<ref>{{Cite web |last=Summers |first=Vincent |date=2017-04-08 |title=Hydrogen Sulfide or Hydrogen Cyanide: Which is More Dangerous? |url=https://www.quirkyscience.com/hydrogen-sulfide-or-hydrogen-cyanide/ |access-date=2022-08-23 |website=Quirky Science |language=en-US}}</ref> However, its presence in ambient air at concentration over 100–150 ppm quickly deadens the sense of smell,<ref>{{Cite web |title=Hydrogen Sulfide – Hazards {{!}} Occupational Safety and Health Administration |url=https://www.osha.gov/hydrogen-sulfide/hazards |access-date=2022-08-23 |website=www.osha.gov}}</ref> and a victim may breathe increasing quantities without noticing until severe symptoms cause death. Dissolved [[Chemistry:Sulfide|sulfide]] and hydrosulfide salts are toxic by the same mechanism.

==See also==
*[[Chemistry:Blue lava|Blue lava]]
*[[Earth:Stratospheric sulfur aerosols|Stratospheric sulfur aerosols]]
*[[Biology:Sulfur assimilation|Sulfur assimilation]]
*[[Chemistry:Sulfur isotope biogeochemistry|Sulfur isotope biogeochemistry]]
*Ultra-low sulfur diesel

==References==
{{Reflist}}

==Further reading==
{{cite book |editor1-last=Sigel |editor1-first=Astrid |editor2-last=Freisinger |editor2-first=Eva |editor3-last=Sigel |editor3-first=Roland K.O. |title=Transition Metals and Sulfur: A Strong Relationship for Life |date=2020 |publisher=de Gruyter |location=Berlin/Boston |isbn=978-3-11-058889-7 |pages=xlv+455|others=Guest Editors Martha E Sosa Torres and Peter M.H.Kroneck}}

==External links==
*[http://www.periodicvideos.com/videos/016.htm Sulfur] at ''The Periodic Table of Videos'' (University of Nottingham)
*[https://physics.nist.gov/PhysRefData/Handbook/Tables/sulfurtable1.htm Atomic Data for Sulfur], NIST Physical Measurement Laboratory
*[http://library.tedankara.k12.tr/chemistry/vol2/allotropy/z129.htm Sulfur phase diagram] {{Webarchive|url=https://web.archive.org/web/20100223201329/http://library.tedankara.k12.tr/chemistry/vol2/allotropy/z129.htm |date=23 February 2010 }}, Introduction to Chemistry for Ages 13–17
*[https://www.swisseduc.ch/stromboli/perm/vulcano/sulphur-vulcano-en.html Crystalline, liquid and polymerization of sulfur on Vulcano Island, Italy]
*[http://extoxnet.orst.edu/pips/sulfur.htm Sulfur and its use as a pesticide]
*[https://www.sulphurinstitute.org/ The Sulphur Institute]
*[https://web.archive.org/web/20140714221502/http://www.nutrientstewardship.com/partners/products-and-services/sulfur-institute Nutrient Stewardship and The Sulphur Institute]

{{Periodic table (navbox)}}
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[[Category:Sulfur| ]]
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{{Sourceattribution|Sulfur}}