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| {{About|the chemical element}}
| | 46 yr old Corporate Standard Manager Cristopher Koehler from Lachute, has pastimes which includes curling, [http://xty.me/ganhardinheiro79158 como ganhar dinheiro] na internet and cave diving. Finds travel a terrific experience after building a journey to Citadel of the Ho Dynasty. |
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| {{Infobox sulfur}}
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| '''Sulfur''' [[American and British English spelling differences|or]] '''sulphur''' is a [[chemical element]] with the symbol '''S''' and [[atomic number]] 16. It is an [[Abundance of the chemical elements|abundant]], [[Valence (chemistry)|multivalent]] [[non-metal]]. Under [[Standard conditions for temperature and pressure|normal conditions]], sulfur atoms form cyclic octatomic molecules with chemical formula S<sub>8</sub>. Elemental sulfur is a bright yellow [[crystal]]line solid when at room temperature. Chemically, sulfur can react as either an [[oxidant]] or [[reducing agent]]. It [[oxidizes]] most [[metal]]s and several [[nonmetal]]s, including [[carbon]], which leads to its negative charge in most [[organosulfur compound]]s, but it [[Redox|reduces]] several strong [[oxidants]], such as [[oxygen]] and [[fluorine]].
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| Sulfur occurs [[nature|naturally]] as the pure [[Chemical element|element]] (native sulfur) and as [[Sulfide minerals|sulfide]] and [[sulfate minerals]]. Elemental sulfur crystals are commonly sought after by mineral collectors for their distinct, brightly colored [[polyhedron]] shapes. Being abundant in native form, sulfur was known in ancient times, mentioned for its uses in [[ancient India]], [[ancient Greece]], [[History of China#Ancient China|China]] and [[ancient Egypt|Egypt]]. Fumes from burning sulfur were used as fumigants, and sulfur-containing medicinal mixtures were used as balms and antiparasitics. Sulfur is referred to in the [[Bible]] as ''brimstone'' (burn stone) in [[English language|English]], with this name still used in several nonscientific tomes.<ref name=Greenwd/> It was needed to make the best quality of [[gunpowder|black gunpowder]]. In 1777, [[Antoine Lavoisier]] helped convince the scientific community that sulfur was a basic element rather than a compound.
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| Elemental sulfur was once extracted from [[salt dome]]s where it sometimes occurs in nearly pure form, but this method has been obsolete since the late 20th century. Today, almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants from [[natural gas]] and [[petroleum]]. The element's commercial uses are primarily in [[fertilizer]]s, because of the relatively high requirement of plants for it, and in the manufacture of [[sulfuric acid]], a primary industrial chemical. Other well-known uses for the element are in [[match]]es, [[insecticide]]s and [[fungicide]]s. Many sulfur compounds are odoriferous, and the smell of odorized natural gas, skunk scent, grapefruit, and garlic is due to sulfur compounds. [[Hydrogen sulfide]] produced by living organisms imparts the characteristic odor to rotting eggs and other biological processes.
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| Sulfur is an [[essential element]] for all life, and is widely used in biochemical processes. In metabolic reactions, sulfur compounds serve as both fuels ([[electron donor]]s) and respiratory (oxygen-alternative) materials ([[electron acceptors]]). Sulfur in organic form is present in the vitamins [[biotin]] and [[thiamine]], the latter being named for the Greek word for sulfur. Sulfur is an important part of many enzymes and in antioxidant molecules like [[glutathione]] and [[thioredoxin]]. Organically bonded sulfur is a component of all proteins, as the [[amino acid]]s [[cysteine]] and [[methionine]]. [[Disulfide]] bonds are largely responsible for the mechanical strength and insolubility of the protein [[keratin]], found in outer skin, hair, and feathers, and the element contributes to their pungent odor when burned.
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| ==Spelling and etymology==
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| ''Sulfur'' or ''sulphur'' comes via Old French from Latin ''sulphur'', which in turn is apparently formed on a root meaning "to burn".<ref>[[Online Etymology Dictionary]] entry for [http://www.etymonline.com/index.php?term=sulfur sulfur]. Retrieved 2011-08-18.</ref>
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| The element was traditionally spelled ''sulphur'' in the United Kingdom (since the 14th century),<ref name="PHorF">{{cite journal|last1 = Michie|first1 = C. A.|last2 = Langslow|first2 = D. R.|title = Sulphur or sulfur? A tale of two spellings|journal = Britisch Medical Journal|volume = 297|issue = 6664|pages = 1697–1699|year = 1988|doi = 10.1136/bmj.297.6664.1697}}</ref> and most of the [[Commonwealth of Nations|Commonwealth]] (including [[Australia]], [[India]], [[Malaysia]], [[South Africa]]), [[Hong Kong]], the [[Caribbean]] and [[Republic of Ireland|Ireland]]. ''Sulfur'' is used in the United States, while both spellings are used in Canada and the [[Philippines]].<ref name="PHorF"/>
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| However, the [[International Union of Pure and Applied Chemistry|IUPAC]] adopted the spelling ''sulfur'' in 1990, as did the [[Royal Society of Chemistry]] Nomenclature Committee in 1992.<ref>{{cite journal|last1 = McNaught|first1 = Alan|title = Journal style update|journal = The Analyst|volume = 116|issue = 11|page = 1094|year = 1991|doi = 10.1039/AN9911601094|bibcode = 1991Ana...116.1094M }}</ref> The [[Qualifications and Curriculum Authority]] for England and Wales recommended its use in 2000,<ref>[http://www.worldwidewords.org/topicalwords/tw-sul1.htm Sulphur], Worldwidewords</ref><!--http://books.google.com/books?id=1ZwoAQAAIAAJ&q=Qualifications+and+Curriculum+Authority++sulfur&dq=Qualifications+and+Curriculum+Authority++sulfur&hl=de&ei=369qTfGaN8T3sgbcnPzeDA&sa=X&oi=book_result&ct=result&resnum=1&ved=0CCoQ6AEwAA--> and it now appears in GCSE exams.<ref>{{cite web|url = http://store.aqa.org.uk/qual/gcse/qp-ms/AQA-CHY1AP-W-QP-MAR10.PDF|title = General Certificate of Secondary Education (Science A Chemistry| publisher = Foundation Tier and Higher Tier|date = March 2010|accessdate = 2011-02-27}}</ref> The Oxford Dictionaries note that ''"In chemistry... the -f- spelling is now the standard form in all related words in the field in both British and US contexts."''<ref>[http://oxforddictionaries.com/view/entry/m_en_gb0828080#m_en_gb0828080 Ask Oxford], accessed 12 November 2010.</ref>
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| In Latin, the word is variously written ''sulpur'', ''sulphur'', and ''sulfur'' (the Oxford Latin Dictionary lists the spellings in this order). It is an original Latin name and not a [[Ancient Greek language|Classical Greek]] loan, so the ''ph'' variant does not denote the Greek letter [[φ]] (phi). Sulfur in Greek is ''thion'' (θείον), whence comes the prefix [[thio-]]. The simplification of the Latin word's p or ph to an f appears to have taken place towards the end of the classical period.<ref>{{cite web|url = http://elements.vanderkrogt.net/element.php?sym=S|publisher = Vanderkrogt.net|title = Sulphuricum Sulphur|accessdate = 2011-02-27}}</ref><ref>{{cite journal |year=1995 |last1=Kelly |first1=Donovan P.|title= Sulfur and its Doppelgänger|journal=Archives of Microbiology|volume=163 |issue=3|pages=157–158|doi=10.1007/BF00305347}}</ref>
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| ==Characteristics==
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| [[File:Burning-sulfur.png|thumb|left|When burned, sulfur melts to a blood-red liquid and emits a blue flame that is best observed in the dark.]]
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| ===Physical properties===
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| Sulfur forms polyatomic molecules with different chemical formulas, with the best-known allotrope being [[octasulfur]], cyclo-S<sub>8</sub>. Octasulfur is a soft, bright-yellow solid with only a faint odor, similar to that of [[match]]es.<ref>A strong odor called "smell of sulfur" actually is given off by several sulfur compounds, such as [[hydrogen sulfide]] and [[organosulfur]] compounds.</ref> It melts at 115.21 °C, boils at 444.6 °C and sublimes easily.<ref name=Greenwd/> At 95.2 °C, below its melting temperature, cyclo-octasulfur changes from α-octasulfur to the β-[[polymorphism (materials science)|polymorph]].<ref name = "Greenwood">{{Greenwood&Earnshaw|pages = 645–662}}</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 [[viscosity]] due to the formation of [[polymer]]s.<ref name = "Greenwood"/> At even higher temperatures, however, the viscosity decreases as depolymerization occurs. Molten sulfur assumes a dark red color above 200 °C. The density of sulfur is about 2 g·cm<sup>−3</sup>, depending on the allotrope; all of its stable allotropes are excellent electrical insulators.
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| ===Chemical properties===
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| Sulfur burns with a blue flame concomitant with formation of [[sulfur dioxide]], notable for its peculiar suffocating odor. Sulfur is insoluble in water but soluble in [[carbon disulfide]] and, to a lesser extent, in other nonpolar organic solvents, such as [[benzene]] and [[toluene]]. The first and the second ionization energies of sulfur are 999.6 and 2252 kJ·mol<sup>−1</sup>, respectively. Despite such figures, the +2 oxidation state is rare, with +4 and +6 being more common. The fourth and sixth ionization energies are 4556 and 8495.8 kJ·mol<sup>−1</sup>, the magnitude of the figures caused by electron transfer between orbitals; these states are only stable with strong oxidants as [[fluorine]], [[oxygen]], and [[chlorine]].
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| ===Allotropes===
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| [[File:Cyclooctasulfur-above-3D-balls.png|thumb|left|The structure of the cyclooctasulfur molecule, S<sub>8</sub>.]]
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| {{Main|Allotropes of sulfur}}
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| Sulfur forms over 30 solid [[allotropy|allotropes]], more than any other element.<ref>{{cite journal |title = Solid Sulfur Allotropes Sulfur Allotropes| first1 = Ralf |last1 = Steudel|first2 = Bodo|last2 = Eckert|journal = Topics in Current Chemistry |year = 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 journal| doi=10.1007/3-540-11345-2_10 |last = Steudel|first = R. |title = Homocyclic Sulfur Molecules |journal = Topics in Current Chemistry |year = 1982 |volume = 102 |pages = 149–176| series=Topics in Current Chemistry| isbn=978-3-540-11345-4}}</ref> Removing one atom from the crown gives S<sub>7</sub>, which is more deeply 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|year = 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 |year = 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|year = 1976|volume = 76|issue = 3|pages = 367–388 |doi = 10.1021/cr60301a003}}</ref>
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| [[Amorphous]] or "plastic" sulfur is produced by rapid cooling of molten sulfur—for example, by pouring it into cold water. [[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 [[Elasticity (physics)|elastic]], and in bulk this form has the feel of crude rubber. This form is [[Metastability in molecules|metastable]] at room temperature and gradually reverts to crystalline molecular allotrope, which is no longer elastic. This process happens within a matter of hours to days, but can be rapidly catalyzed.
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| {{clear left}}
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| ===Isotopes===
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| {{Main|Isotopes of sulfur}}
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| Sulfur has 25 known [[isotope]]s, four of which are stable: <sup>32</sup>S (95.02%), <sup>33</sup>S (0.75%), <sup>34</sup>S (4.21%), and <sup>36</sup>S (0.02%). Other than <sup>35</sup>S, with a [[half-life]] of 87 days and formed in [[cosmic ray spallation]] of <sup>40</sup>[[Argon|Ar]], the [[radioactivity|radioactive]] isotopes of sulfur have half-lives less than {{abbr|170 minutes|Half-life of sulfur-38}}.
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| When [[sulfide mineral]]s are precipitated, isotopic equilibration among solids and liquid may cause small differences in the δS-34 values of co-genetic minerals. The differences between minerals can be used to estimate the temperature of equilibration. The δ[[carbon|C]]-13 and δS-34 of coexisting [[carbonate minerals]] and sulfides can be used to determine the [[pH]] and oxygen [[fugacity]] of the ore-bearing fluid during ore formation.
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| In most [[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 [[hydrology|hydrologic]] studies. Differences in the [[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 different δ<sup>34</sup>S values from lakes believed to be dominated by watershed sources of sulfate.
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| ===Natural occurrence===
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| [[File:Io highest resolution true color.jpg|thumb|Most of the yellow and orange hues of [[Io (moon)|Io]] are due to elemental sulfur and sulfur compounds, produced by active volcanoes.]]
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| [[File:Soufresicile3.jpg|thumb|Native sulfur crystals]]
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| [[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]]
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| <sup>32</sup>S is created inside massive stars, at a depth where the temperature exceeds 2.5×10<sup>9</sup> K, by the [[silicon burning|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=http://www.fas.org/sgp/eprint/CRL-41.pdf|year=1957}}</ref> As this is part of the [[alpha process]] that produces elements in abundance, sulfur is the 10th most common element in the universe.
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| Sulfur, usually as sulfide, is present in many types of [[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 [[troilite]] (FeS), but there are exceptions, with carbonaceous chondrites containing free sulfur, sulfates and other sulfur compounds.<ref>{{cite book|first=B.|last = Mason|title=Meteorites |location=New York |publisher=John Wiley & Sons|year=1962|page=160 |isbn=0-908678-84-3}}</ref> The distinctive colors of [[Jupiter]]'s [[volcano|volcanic]] moon [[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|year = 2005|doi = 10.1088/0034-4885/68/2/R02|bibcode=2005RPPh...68..303L}}</ref>
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| On Earth, elemental sulfur can be found near [[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. Such deposits are polycrystalline, with the largest documented single crystal measuring 22×16×11 cm.<ref>{{cite journal| url = http://www.minsocam.org/ammin/AM66/AM66_885.pdf| journal = American Mineralogist| volume = 66| pages = 885–907| year= 1981| title= The largest crystals| last = Rickwood|first = P. C.}}</ref> Historically, [[Sicily]] was a large source of sulfur in the [[Industrial Revolution]].<ref>{{cite book|last=Kutney|first=Gerald|title=Sulfur: history, technology, applications & industry|year=2007|publisher=ChemTec Publications|location=Toronto|isbn=978-1-895198-37-9|oclc=79256100|page=43}}</ref>
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| Significant deposits of elemental sulfur, believed to have been (and are still being) synthesised by [[anaerobic bacteria]] on [[Mineral#Sulfate class|sulfate minerals]] like [[gypsum]], exist in [[salt domes]] along the coast of the [[Gulf of Mexico]], and in [[evaporite]]s in eastern Europe and western Asia. Native sulfur may be produced by geological processes alone. Fossil-based sulfur deposits from salt domes have until recently been 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.
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| Common naturally occurring sulfur compounds include the [[Mineral#Sulfide class|sulfide minerals]], such as [[pyrite]] (iron sulfide), [[cinnabar]] (mercury sulfide), [[galena]] (lead sulfide), [[sphalerite]] (zinc sulfide) and [[stibnite]] (antimony sulfide); and the sulfates, such as gypsum (calcium sulfate), [[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 [[hydrothermal vent]]s.
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| ==Production==
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| Sulfur may be found by itself and historically was usually obtained in this way, while pyrite has been a source of sulfur via sulfuric acid.<ref>{{cite book|last1=Riegel|first1=Emil|last2=Kent|first2=James|title=Kent and Riegel's handbook of industrial chemistry and biotechnology, Volume 1|year=2007|publisher=Springer|location=New York|isbn=978-0-387-27842-1|oclc=74650396|page=1171}}</ref> In volcanic regions in [[Sicily]], in ancient times, it was found on the surface of the Earth, and the ''"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 ''[[carusu|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 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 sulphur 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}}</ref>
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| Today's sulfur production is as a side product of other industrial processes such as oil refining; in these processes, sulfur often occurs as undesired or detrimental compounds that are extracted and converted to elemental sulfur. As a mineral, native sulfur under salt domes is thought to be a fossil mineral resource, produced by the action of ancient bacteria on sulfate deposits. It was removed from such salt-dome mines mainly by the [[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. However, 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|year = 2002|doi = 10.1002/ep.670210312}}</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|year = 2008|doi = 10.1002/ciuz.200800461}}</ref>
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| [[File:AlbertaSulfurAtVancouverBC.jpg|thumb|Sulfur recovered from hydrocarbons in [[Alberta]], stockpiled for shipment in [[North Vancouver]], [[British Columbia|B.C.]]]]
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| Today, sulfur is produced from petroleum, [[natural gas]], and related fossil resources, from which it is obtained mainly as [[hydrogen sulfide]]. [[Organosulfur compound]]s, undesirable impurities in petroleum, may be upgraded by subjecting them to [[hydrodesulfurization]], which cleaves the C–S bonds:<ref name="desulf1"/><ref name="desulf2"/>
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| :R-S-R + 2 H<sub>2</sub> → 2 RH + H<sub>2</sub>S
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| The resulting hydrogen sulfide from this process, and also as it occurs in natural gas, is converted into elemental sulfur by the [[Claus process]]. This process entails oxidation of some hydrogen sulfide to sulfur dioxide and then the [[comproportionation]] of the two:<ref name="desulf1"/><ref name="desulf2"/>
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| :3 O<sub>2</sub> + 2 H<sub>2</sub>S → 2 SO<sub>2</sub> + 2 H<sub>2</sub>O
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| :SO<sub>2</sub> + 2 H<sub>2</sub>S → 3 S + 2 H<sub>2</sub>O
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| [[File:SulfurPrice.png|thumb|Production and price (US market) of elemental sulfur]]
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| 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 journal|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|year = 1982|doi = 10.1021/bk-1982-0183.ch005|chapter = Sulfur Recovery from Oil Sands|series = ACS Symposium Series|isbn = 0-8412-0713-5}}</ref> Another way of storing sulfur is as a [[binder (material)|binder]] for concrete, the resulting product having many desirable properties (see [[sulfur concrete]]).<ref>{{cite book|last1=Mohamed|first1=Abdel-Mohsen|last2=El-Gamal|first2=Maisa|title=Sulfur concrete for the construction industry: a sustainable development approach|year=2010|publisher=J. Ross Publishing|location=[[Fort Lauderdale]]|isbn=978-1-60427-005-1|oclc=531718953|page=109}}</ref>
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| The world production of sulfur in 2011 amounted to 69 million tonnes (Mt), with more than 15 countries contributing more than 1 Mt each. Countries producing more than 5 Mt are China (9.6), US (8.8), Canada (7.1) and Russia (7.1).<ref>Apodaca, Lori E. (2012) [http://minerals.usgs.gov/minerals/pubs/commodity/sulfur/mcs-2012-sulfu.pdf Sulfur]. Mineral Commodity Summaries. USGS</ref> While the production has been slowly increasing from 1900 to 2010, the price was much less stable, especially in the 1980s and around 2010.<ref name=USGS/>
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| ==Compounds==
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| {{Category see also|Sulfur compounds}}
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| Common [[oxidation state]]s of sulfur range from −2 to +6. Sulfur forms stable compounds with all elements except the [[noble gas]]es.
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| ===Sulfides===
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| Treatment of sulfur with hydrogen gives [[hydrogen sulfide]]. When dissolved in water, hydrogen sulfide is mildly acidic:<ref name=Greenwd>Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.), Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.</ref>
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| :H<sub>2</sub>S <math>\overrightarrow{\leftarrow}</math> HS<sup>–</sup> + H<sup>+</sup>
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| Hydrogen sulfide gas and the hydrosulfide anion are extremely toxic to mammals, due to their inhibition of the oxygen-carrying capacity of hemoglobin and certain [[cytochrome]]s in a manner analogous to [[cyanide]] and [[azide]] (see below, under ''precautions'').
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| Reduction of elemental sulfur gives [[polysulfide]]s, which consist of chains of sulfur atoms terminated with S<sup>–</sup> centers:
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| :2 Na + S<sub>8</sub> → Na<sub>2</sub>S<sub>8</sub>
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| This reaction highlights arguably the single most distinctive property of sulfur: its ability to catenate (bind to itself by formation of chains). Protonation of these polysulfide anions gives the polysulfanes, H<sub>2</sub>S<sub>x</sub> where x = 2, 3, and 4.<ref>Handbook of Preparative Inorganic Chemistry, 2nd ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 421.</ref>
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| Ultimately reduction of sulfur gives sulfide salts:
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| :16 Na + S<sub>8</sub> → 8 Na<sub>2</sub>S
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| The interconversion of these species is exploited in the [[sodium-sulfur battery]]. The [[Trisulfur|radical anion]] S<sub>3</sub><sup>–</sup> gives the blue color of the mineral [[lapis lazuli]].
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| [[File:Lapis lazuli block.jpg|thumb|[[Lapis lazuli]] owes its blue color to a [[Trisulfur|sulfur radical]].]]
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| With very strong oxidants, S<sub>8</sub> can be oxidized, for example, to give bicyclic S<sub>8</sub><sup>2+</sup>.
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| ===Oxides and oxyanions===
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| The principal sulfur oxides are obtained by burning sulfur:
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| :S + O<sub>2</sub> → SO<sub>2</sub>
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| :2 SO<sub>2</sub> + O<sub>2</sub> → 2 SO<sub>3</sub>
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| Other oxides are known, e.g. sulfur monoxide and disulfur mono- and dioxides, but they are unstable.
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| The sulfur oxides form numerous oxyanions with the formula SO<sub>n</sub><sup>2–</sup>. Sulfur dioxide and [[sulfite]]s ({{chem|SO|3|2−}}) are related to the unstable sulfurous acid (H<sub>2</sub>SO<sub>3</sub>). [[Sulfur trioxide]] and [[sulfate]]s ({{chem|SO|4|2−}}) are related to [[sulfuric acid]]. Sulfuric acid and SO<sub>3</sub> combine to give oleum, a solution of [[pyrosulfuric acid]] (H<sub>2</sub>S<sub>2</sub>O<sub>7</sub>) in sulfuric acid.
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| ::[[File:Peroxydisulfuric-acid-3D-balls.png|right|thumb|[[Peroxydisulfuric acid]]]]
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| Peroxides convert sulfur into unstable oxides such as S<sub>8</sub>O, a sulfoxide. [[Peroxymonosulfuric acid]] (H<sub>2</sub>SO<sub>5</sub>) and [[peroxydisulfuric acid]]s (H<sub>2</sub>S<sub>2</sub>O<sub>8</sub>), made from the action of SO<sub>3</sub> on concentrated [[hydrogen peroxide|H<sub>2</sub>O<sub>2</sub>]], and [[sulfuric acid|H<sub>2</sub>SO<sub>4</sub>]] on concentrated H<sub>2</sub>O<sub>2</sub> respectively.
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| [[File:Sulfate-3D-vdW.png|thumb|The sulfate anion, {{chem|SO|4|2−}}]]
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| [[Thiosulfate]] salts ({{chem|S|2|O|3|2−}}), sometimes referred as "hyposulfites", used in photographic fixing (HYPO) and as reducing agents, feature sulfur in two oxidation states. [[Sodium dithionite]], ({{chem|S|2|O|4|2−}}), contains the more highly reducing dithionite anion. [[Sodium dithionate]] (Na<sub>2</sub>S<sub>2</sub>O<sub>6</sub>) is the first member of the [[polythionic acid]]s (H<sub>2</sub>S<sub>''n''</sub>O<sub>6</sub>), where ''n'' can range from 3 to many. '''Thiosulfurous acid''' HS-S(=O)-OH is formed in trace amounts when hydrogen sulfide and sulfur dioxide gases are mixed at room temperature, but its salts (thiosulfites) are unknown.
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| ===Halides and oxyhalides===
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| The two main sulfur fluorides are [[sulfur hexafluoride]], a dense gas used as nonreactive and nontoxic propellant, and [[sulfur tetrafluoride]], a rarely used organic reagent that is highly toxic.<ref>{{OrgSynth|last=Hasek|first=W. R.|title=1,1,1-Trifluoroheptane|volume=41|page=104|year=1961|url=http://www.orgsyn.org/orgsyn/pdfs/CV5P1082.pdf}}</ref> Their chlorinated analogs are [[sulfur dichloride]] and [[sulfur monochloride]]. [[Sulfuryl chloride]] and [[chlorosulfuric acid]] are derivatives of sulfuric acid; [[thionyl chloride]] (SOCl<sub>2</sub>) is a common reagent in [[organic synthesis]].<ref>{{OrgSynth|last1=Rutenberg|first1=M. W|last2=Horning|first2=E. C.|title = 1-Methyl-3-ethyloxindole|volume=30|page=62|year=1950|url=http://www.orgsyn.org/orgsyn/pdfs/CV4P0620.pdf}}</ref>
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| ===Pnictides===
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| An important S–N compound is the cage [[tetrasulfur tetranitride]] (S<sub>4</sub>N<sub>4</sub>). Heating this compound gives [[Polythiazyl|polymeric sulfur nitride]] ((SN)<sub>x</sub>), which has metallic properties even though it does not contain any [[metal]] atoms. [[Thiocyanate]]s contain the SCN<sup>−</sup> group. Oxidation of thiocyanate gives [[thiocyanogen]], (SCN)<sub>2</sub> with the connectivity NCS-SCN. [[Phosphorus sulfide]]s are numerous, the most important commercially being the cages P<sub>4</sub>S<sub>10</sub> and P<sub>4</sub>S<sub>3</sub>.<ref>{{Cite book |last=Heal |first=H. G. |title=The Inorganic Heterocyclic Chemistry of Sulfur, Nitrogen, and Phosphorus |publisher=Academic Press |location=London |year=1980 |isbn=0-12-335680-6}}</ref><ref name=Chivers>{{Cite book |last=Chivers |first=T. |title=A Guide To Chalcogen-Nitrogen Chemistry |publisher=World Scientific |location=Singapore |year=2004 |isbn=981-256-095-5}}</ref>
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| ===Metal sulfides===
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| {{Main|Sulfide mineral}}
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| The principal ores of copper, zinc, nickel, cobalt, molybdenum, and other metals are sulfides. These materials tend to be dark-colored [[semiconductor]]s that are not readily attacked by water or even many acids. They are formed, both geochemically and in the laboratory, by the reaction of hydrogen sulfide with metal salts. The mineral [[galena]] (PbS) was the first demonstrated semiconductor and found a use as a signal [[rectifier]] in the [[Cat's-whisker detector|cat's whiskers]] of early [[crystal radio]]s. The iron sulfide called [[pyrite]], the so-called "fool's gold," has the formula FeS<sub>2</sub>.<ref>Vaughan, D. J.; Craig, J. R. "Mineral Chemistry of Metal Sulfides" Cambridge University Press, Cambridge (1978) ISBN 0-521-21489-0</ref> The upgrading of these ores, usually by [[smelting|roasting]], is costly and environmentally hazardous. Sulfur corrodes many metals via the process called [[tarnishing]].
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| ===Organic compounds===
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| {{Main|Organosulfur compounds}}
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| <gallery caption="Illustrative organosulfur compounds">
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| File:R-allicin-2D-skeletal.png|[[Allicin]], the active ingredient in garlic
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| File:Cysteine.svg| R-[[cysteine]], an [[amino acid]] containing a thiol group
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| File:Methionin - Methionine.svg|[[Methionine]], an [[amino acid]] containing a thioether
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| File:Diphenyl disulfide.png|[[Diphenyl disulfide]], a representative disulfide
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| File:Perfluorooctanesulfonic acid.png|[[Perfluorooctanesulfonic acid]], a controversial surfactant
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| File:Dibenzothiophen - Dibenzothiophene.svg|[[Dibenzothiophene]], a component of crude oil
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| File:Penicillin core.svg|[[Penicillin]]
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| </gallery>
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| Some of the main classes of sulfur-containing organic compounds include the following:<ref name=Cremlyn>Cremlyn R. J.; "An Introduction to Organosulfur Chemistry" John Wiley and Sons: Chichester (1996). ISBN 0-471-95512-4.</ref>
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| * [[Thiol]]s or mercaptans (as they are '''mer'''cury '''capt'''urers as chelators) are the sulfur analogs of [[alcohol]]s; treatment of thiols with base gives [[thiolate]] ions.
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| * [[Thioether]]s are the sulfur analogs of [[ether]]s.
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| * [[Sulfonium]] ions have three groups attached to a cationic sulfur center. [[Dimethylsulfoniopropionate]] (DMSP) is one such compound, important in the marine organic [[sulfur cycle]].
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| * [[Sulfoxide]]s and [[sulfone]]s are thioethers with one and two oxygen atoms attached to the sulfur atom, respectively. The simplest sulfoxide, [[dimethyl sulfoxide]], is a common solvent; a common sulfone is [[sulfolane]].
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| * [[Sulfonic acid]]s are used in many detergents.
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| Compounds with carbon–sulfur bonds are uncommon with the notable exception of [[carbon disulfide]], a volatile colorless liquid that is structurally similar to carbon dioxide. It is used as a reagent to make the polymer [[rayon]] and many organosulfur compounds. Unlike [[carbon monoxide]], [[carbon monosulfide]] is only stable as a dilute gas, as in the interstellar medium.<ref>{{cite journal|last=Wilson|first=R. W.|last2=Penzias|first2=A. A.|last3=Wannier|first3=P. G.|last4=Linke|first4=R. A.|authorlink=Robert Woodrow Wilson|authorlink2=Arno Allan Penzias|title=Isotopic abundances in interstellar carbon monosulfide|journal=Astrophysical Journal|date=March 15, 1976|volume=204|pages=L135–L137|doi=10.1086/182072|bibcode=1976ApJ...204L.135W}}</ref>
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| Organosulfur compounds are responsible for the some of the unpleasant odors of decaying organic matter. They are used in the odoration of natural gas and cause the odor of garlic and skunk spray. Not all organic sulfur compounds smell unpleasant at all concentrations: the sulfur-containing [[terpene|monoterpenoid]] [[grapefruit mercaptan]] in small concentrations is responsible for the characteristic scent of grapefruit, but has a generic thiol odor at larger concentrations. [[Sulfur mustard]], a potent [[blister agent|vesicant]], was [[Chemical weapons in World War I|used]] in [[World War I]] as a disabling agent.<ref>{{cite book|last=Banoub|first=Joseph|title=Detection of Biological Agents for the Prevention of Bioterrorism|year=2011|publisher=Springer|location=Dordrecht|isbn=978-90-481-9815-3|oclc=697506461|page=183}}</ref>
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| Sulfur-sulfur bonds are a structural component to stiffen rubber, in a way similar to the biological role of disulfide bridges to rigidify proteins (see biological below). In the most common type of industrial "curing" or hardening and strengthening of natural [[rubber]], elemental sulfur is heated with the rubber to the point that chemical reactions form [[disulfide]] bridges between [[isoprene]] units of the polymer. This process, patented in 1843, allowed rubber to become a major industrial product, especially automobile tires. Because of the heat and sulfur, the process was named [[vulcanization]], after the Roman god of the forge and volcanism.
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| ==History==
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| ===Antiquity===
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| [[File:MODOAzufre.jpg|thumb|Pharmaceutical container for sulfur from the first half of the 20th century. From the [[Museo del Objeto del Objeto]] collection]]
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| Being abundantly available in native form, sulfur ([[Latin]] ''sulphur'') was known in ancient times and is referred to in the [[Torah]] ([[Book of Genesis|Genesis]]). [[English translations of the Bible]] commonly referred to burning sulfur as "brimstone", giving rise to the name of '[[fire and brimstone|fire-and-brimstone]]' [[sermon]]s, in which listeners are reminded of the fate of [[Damnation|eternal damnation]] that await the unbelieving and unrepentant. It is from this part of the Bible that [[Hell]] is implied to "smell of sulfur" (likely due to its association with volcanic activity). According to the [[Ebers Papyrus]], a sulfur ointment was used in ancient [[Egypt]] to treat granular eyelids. Sulfur was used for fumigation in preclassical [[Greece]];<ref>{{cite book | url = http://books.google.com/books?id=ed0yC98aAKYC&pg=PA242 | title = Archaeomineralogy | isbn = 978-3-540-78593-4 | page = 242 | author1 = Rapp | first1 = George Robert | date = 2009-02-04}}</ref> this is mentioned in the ''[[Odyssey]]''.<ref>[http://www.perseus.tufts.edu/hopper/text?doc=Hom.+Od.+22.480 ''Odyssey'', book 22, lines 480–495] • www.perseus.tufts.edu. Retrieved on 2012-08-16.</ref> [[Pliny the Elder]] discusses sulfur in book 35 of his ''[[Natural History (Pliny)|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>
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| A natural form of sulfur known as ''shiliuhuang'' was known in China since the 6th century BC and found in [[Hanzhong]].<ref name="yunming 487">{{cite journal|author = Zhang, Yunming|year = 1986|title = The History of Science Society: Ancient Chinese Sulfur Manufacturing Processes|journal = [[Isis (journal)|Isis]]|volume = 77|issue = 3|doi = 10.1086/354207|page=487}}</ref> By the 3rd century, the Chinese discovered that sulfur could be extracted from [[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 [[traditional Chinese medicine]].<ref name="yunming 487"/> A [[Song Dynasty]] military treatise of 1044 AD described different formulas for Chinese [[black powder]], which is a mixture of [[potassium nitrate]] ({{chem|K||N||O|3}}), [[charcoal]], and sulfur.
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| Indian alchemists, practitioners of "the science of mercury" ([[sanskrit]] rasaśāstra, रसशास्त्र), 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|year=1996|publisher=University of Chicago Press|location=Chicago|isbn=978-0-226-89499-7|pages=passim}}</ref> In the [[rasa shastra|rasaśāstra]] tradition, sulfur is called "the smelly" (sanskrit gandhaka, गन्धक).
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| Early [[Europe]]an [[alchemy|alchemists]] gave sulfur its own [[alchemical symbol]], a triangle at the top of a cross.
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| In traditional skin treatment before the modern era of scientific medicine, elemental sulfur was used, mainly in creams, to alleviate conditions such as [[scabies]], [[ringworm]], [[psoriasis]], [[eczema]], and [[acne]]. The mechanism of action is unknown—though elemental sulfur does oxidize slowly to sulfurous acid, which in turn (through the action of [[sulfite]]) acts as 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|year = 1988|pmid = 2450900}}</ref><ref>{{cite journal|doi = 10.1016/S0190-9622(08)81225-X|last1 = Maibach|first1 = HI|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|year = 1990| pmid = 2365870}}</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 : JDD|volume = 3|issue = 4|pages = 427–31|year = 2004| pmid = 15303787}}</ref>
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| ===Modern times===
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| [[File:Soufre extraction 1.jpg|thumb|Sicilian kiln used to obtain sulfur from volcanic rock.]]
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| In 1777, [[Antoine Lavoisier]] helped convince the scientific community that sulfur was an element, not a compound.
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| Sulfur deposits in [[Sicily]] were the dominant supply source for over half a century. Approximately 2000 tons per year of sulfur were imported into [[Marseilles]], [[France]] for the production of [[sulphuric acid]] via the [[Leblanc process]] by the late 18th century. In [[Industrial Revolution|industrializing]] Britain, with the repeal of [[tarrif]]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 [[Ferdinand II of the Two Sicilies|King Ferdinand II]] gave a monopoly of the sulfur industry to a French firm, violating an earlier 1816 trade agreement with Britain. A peaceful negotiated solution was eventually mediated by France.<ref>{{cite book|url=http://books.google.com/?id=wZg4ecXXmNYC|title=Sicily and the Unification of Italy: Liberal Policy and Local Power, 1859–1866|author=Riall, Lucy|year=1998|publisher=Oxford University Press|accessdate=2013-02-07|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|pages=163–180|doi=10.1177/026569149502500201|last1=Thomson|first1=D. W.|issue=2}}</ref>
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| In 1867, sulfur was discovered in underground deposits in [[Louisiana]] and [[Texas]]. The highly successful [[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|year = 2001|volume = 35|issue = 5|language = German|pages = 324–331|doi = 10.1002/1521-3781(200110)35:5<324::AID-CIUZ324>3.0.CO;2-9}}</ref>
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| In the late 18th century, [[furniture]] makers used molten sulfur to produce decorative [[inlay]]s in their craft. Because of the [[sulfur dioxide]] produced during the process of melting sulfur, the craft of sulfur inlays was soon abandoned. 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 [[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|year=2006|publisher=Littleton|location=Colorado|isbn=978-0-87335-233-8|edition=7th|page=935|oclc=62805047}}</ref>
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| ==Applications==
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| ===Sulfuric acid===
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| Elemental sulfur is mainly used as a precursor to other chemicals. Approximately 85% (1989) is converted to [[sulfuric acid]] ([[hydrogen|H]]<sub>2</sub>SO<sub>4</sub>):
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| :2 S + 3 O<sub>2</sub> + 2 H<sub>2</sub>O → 2 H<sub>2</sub>SO<sub>4</sub>
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| With sulfuric acid being of central importance to the [[world economy|world's economies]], its production and consumption is an indicator of a nation's industrial development.<ref>[http://www.pafko.com/history/h_s_acid.html Sulfuric Acid Growth]. Pafko.com. Retrieved on 2012-08-16.</ref> For example with 32.5 million tonnes in 2010, the United States produces more sulfuric acid every year 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 = http://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/>
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| [[File:2000sulphuric acid.PNG|thumb|760 px|center|Sulfuric acid production in 2000]]
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| ===Other large-scale sulfur chemicals===
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| Sulfur reacts directly with methane to give [[carbon disulfide]], which is used to manufacture [[cellophane]] and [[rayon]].<ref name=Nehb>{{cite encyclopedia |last=Nehb |first=Wolfgang|coauthors=Vydra, Karel |encyclopedia=Ullmann's Encyclopedia of Industrial Chemistry |title=Ullmann's Encyclopedia of Industrial Chemistry|year=2006|publisher=Wiley-VCH Verlag|doi=10.1002/14356007.a25_507.pub2 |chapter=Sulfur |isbn=3-527-30673-0}}</ref> One of the direct uses of sulfur is in [[vulcanization]] of rubber, where [[polysulfide]]s crosslink organic polymers.
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| <!--need something on pulping-->
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| [[Sulfite]]s are heavily used to [[Bleach (chemical)|bleach]] [[paper]] and as preservatives in dried [[fruit]]. Many [[surfactant]]s and [[detergents]], e.g. [[sodium lauryl sulfate]], are produced are sulfate derivatives. [[Calcium sulfate]], gypsum, (CaSO<sub>4</sub><sup>·</sup>2H<sub>2</sub>O) is mined on the scale of 100 million tons each year for use in [[Portland cement]] and fertilizers.
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| When silver-based [[photography]] was widespread, sodium and ammonium [[sodium thiosulfate|thiosulfate]] were widely used as "fixing agents."
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| Sulfur is a component of [[gunpowder]].
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| ===Fertilizer===
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| Sulfur is increasingly used as a component of [[fertilizer]]s. The most important form of sulfur for fertilizer is the mineral [[calcium sulfate]]. Elemental sulfur is [[hydrophobic]] (that is, it is not soluble in water) and, therefore, cannot be directly utilized by plants. Over time, soil bacteria can convert it to soluble derivatives, which can then be utilized by plants. Sulfur improves the use efficiency of other essential plant nutrients, particularly nitrogen and phosphorus.<ref>[http://www.sulphurinstitute.org/learnmore/faq.cfm#plants Sulfur as a fertilizer]. Sulphurinstitute.org. Retrieved on 2012-08-16.</ref> Biologically produced sulfur particles are naturally hydrophilic due to a biopolymer coating. This sulfur is, therefore, easier to disperse over the land (via spraying as a diluted slurry), and results in a faster release.
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| Plant requirements for sulfur are equal to or exceed those for phosphorus. It is one of the major nutrients essential for plant growth, root nodule formation of legumes and plants protection mechanisms. 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|year = 1999|last1 = Zhao|first1 = F.|journal = Journal of Cereal Science|volume = 30|issue = 1|pages = 1–17|last2 = Hawkesford|first2 = MJ|last3 = McGrath|first3 = SP}}</ref><ref>{{cite journal|doi =10.1023/A:1026503812267|year =2000|last1 =Blake-Kalff|first1 =M. M. A.|journal =Plant and Soil|volume =225|issue =1/2|pages =95–107}}</ref><ref>{{cite journal|doi =10.1007/BF00747690|title =Plant nutrient sulphur-a review of nutrient balance, environmental impact and fertilizers|year =1996|last1 =Ceccotti|first1 =S. P.|journal =Fertilizer Research|volume =43|issue =1–3|pages =117–125}}</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.
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| ===Fine chemicals===
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| [[File:Malathion-3D-vdW.png|thumb|left|A molecular model of the pesticide [[malathion]].]]
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| Organosulfur compounds are used in [[pharmaceutical]]s, [[dyestuff]]s, and agrochemicals. Many drugs contain sulfur, early examples being antibacterial [[sulfonamide (medicine)|sulfonamides]], known as ''sulfa drugs''. Sulfur is a part of many bacterial defense molecules. Most [[β-lactam]] antibiotics, including the [[penicillin]]s, [[cephalosporins]] and [[monolactam]]s contain sulfur.<ref name=Cremlyn/>
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| [[Magnesium sulfate]], known as Epsom salts when in hydrated crystal form, can be used as a [[laxative]], a bath additive, an [[exfoliant]], [[magnesium]] supplement for plants, or (when in dehydrated form) as a [[desiccant]].
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| ===Fungicide and pesticide===<!--[[Wettable Sulfur]] redirs here-->
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| [[File:Sulphur Candle.jpg|thumb|Sulfur candle originally sold for home fumigation]]
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| 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 [[Organic agriculture|organically]] farmed apple production against the main disease [[apple scab]] under colder conditions. Biosulfur (biologically produced elemental sulfur with hydrophilic characteristics) can be used well for these applications.
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| 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 [[miscibility|miscible]].<ref>{{cite book |url = http://books.google.com/books?id=OYecyRmnTEkC&pg=PA104|pages = 104–105 |title = Sulfur Concrete for the Construction Industry: A Sustainable Development Approach |isbn = 978-1-60427-005-1 |author1 = Mohamed, Abdel-Mohsen Onsy |author2 = El Gamal, M. M |date = 2010-07-13}}</ref><ref>{{cite web
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| |url=http://www.freepatentsonline.com/3398227.pdf
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| |title=Method for Preparation of Wettable Sulfur
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| |accessdate= 2010-05-20
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| |author= Every, Richard L., ''et al.''
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| |date= 1968-08-20
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| }}</ref> It has similar applications and is used as a [[fungicide]] against [[mildew]] and other mold-related problems with plants and soil.
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| Elemental sulfur powder is used as an "[[organic farming|organic]]" (i.e. "green") [[insecticide]] (actually an [[acaricide]]) against [[tick]]s and [[mites]]. A common method of use is to dust clothing or limbs with sulfur powder.
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| Diluted solutions of [[lime sulfur]] (made by combinding [[calcium hydroxide]] with elemental sulfur in water), are used as a dip for pets to destroy [[ringworm|ringworm (fungus)]], [[mange]] and other [[cutaneous conditions|dermatoses]] and [[parasitism|parasites]]. Sulfur candles consist of almost pure sulfur in blocks or pellets that are burned to [[fumigant|fumigate]] structures. It is no longer used in the home due to the toxicity of the products of combustion.
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| ===Bactericide in winemaking and food preservation===
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| Small amounts of [[sulfur dioxide]] gas addition (or equivalent [[potassium metabisulfite]] addition) to fermented wine to produce traces of [[sulfurous acid]] (produced when SO<sub>2</sub> reacts with water) and its [[sulfite]] salts in the mixture, has been called "the most powerful tool in winemaking."<ref>Spencer, Benjamin [http://www.intowine.com/sulfur-wine-demystified Sulfur in wine demystified]. intowine.com. Retrieved Oct 26, 2011.</ref> After the yeast-fermentation stage in [[winemaking]], sulfites absorb oxygen and inhibit [[aerobic organism|aerobic]] bacterial growth that otherwise would turn ethanol into acetic acid, souring the wine. Without this preservative step, indefinite refrigeration of the product before consumption is usually required. Similar methods go back into antiquity but modern historical mentions of the practice go to the fifteenth century. The practice is used by large industrial wine producers and small organic wine producers alike.
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| Sulfur dioxide and various sulfites have been used for their antioxidant antibacterial preservative properties in many other parts of the food industry also. The practice has declined since reports of an allergy-like reaction of some persons to sulfites in foods.
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| ==Biological role==
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| ===Protein and organic cofactors===
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| Sulfur is an essential component of all living [[cell (biology)|cells]]. It is the seventh or eighth most abundant element in the human body by weight, being about as common as [[potassium]], and a little more common than sodium or chlorine. A 70 kg human body contains about 140 grams of sulfur.
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| In [[plant]]s and [[animal]]s, the [[amino acid]]s [[cysteine]] and [[methionine]] contain most of the sulfur. The element is thus present in all [[polypeptide]]s, [[protein]]s, and [[enzyme]]s that contain these amino acids. In humans, [[methionine]] is an essential amino acid that must be ingested. However, save for the vitamins [[biotin]] and [[thiamine]], cysteine and all sulfur-containing compounds in the human body can be synthesized from methionine. The enzyme [[sulfite oxidase]] is needed for the metabolism of methionine and cysteine in humans and animals.
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| [[Disulfide bond]]s (S-S bonds) formed between cysteine residues in peptide chains are very important in protein assembly and structure. These covalent bonds between peptide chains confer extra toughness and rigidity.<ref name=Lehn/> For example, the high strength of feathers and hair is in part due to their high content of S-S bonds and their high content of cysteine and sulfur. Eggs are high in sulfur because large amounts of the element are necessary for feather formation, and the characteristic odor of rotting eggs is due to [[hydrogen sulfide]]. The high disulfide bond content of hair and feathers contributes to their indigestibility and to their characteristic disagreeable odor when burned.
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| [[Homocysteine]] and [[taurine]] are other sulfur-containing acids that are similar in structure, but not coded by [[DNA]], and are not part of the [[primary structure]] of proteins. 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 [[coenzyme A]] and [[alpha-lipoic acid]].<ref name=Lehn>{{cite book|isbn = 1-57259-153-6|last1 = Nelson|first1 = D. L.|last2 = Cox|first2 = M. M.|title = Lehninger, Principles of Biochemistry|edition= 3rd |publisher = Worth Publishing|place = New York|year = 2000}}</ref> Two of the 13 classical vitamins, [[biotin]] and [[thiamine]] contain sulfur, with the latter being named for its sulfur content. Sulfur plays an important part, as a carrier of reducing hydrogen and its electrons, for cellular repair of oxidation. Reduced [[glutathione]], a sulfur-containing tripeptide, is a reducing agent through its sulfhydryl (-SH) moiety derived from [[cysteine]]. The [[thioredoxin]]s, a class of small protein essential to all known life, using neighboring pairs of reduced cysteines to act as general protein reducing agents, to similar effect.
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| [[Methanogenesis]], the route to most of the world's methane, is a multistep biochemical transformation of [[carbon dioxide]]. This conversion requires several organosulfur cofactors. These include [[coenzyme M]], CH<sub>3</sub>SCH<sub>2</sub>CH<sub>2</sub>SO<sub>3</sub><sup>–</sup>, the immediate precursor to [[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|year = 1998|pmid = 9782487|doi = 10.1099/00221287-144-9-2377}}</ref>
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| ===Metalloproteins and inorganic cofactors===
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| Inorganic sulfur forms a part of [[iron-sulfur cluster]]s as well as many copper, nickel, and iron proteins. Most pervasive are the ferrodoxins, which serve as electron shuttles in cells. In bacteria, the important [[nitrogenase]] enzymes contains an Fe-Mo-S cluster, is a [[catalyst]] that performs the important function of [[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 =0-935702-73-3| first1 = S. J.|last1 = Lippard|first2 = J. M.|last2 = Berg|title = Principles of Bioinorganic Chemistry|publisher = University Science Books|year =1994}}</ref>
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| :[[File:FdRedox.png|center|500px]]
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| ===Sulfur metabolism and the sulfur cycle===
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| {{Main|Sulfur metabolism|Sulfur cycle}}
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| The sulfur cycle was the first of the [[biogeochemical cycle]]s to be discovered. In the 1880s, while studying [[Beggiatoa]] (a bacterium living in a sulfur rich environment), [[Sergei Winogradsky]] found that it oxidized [[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). He continued to study it together with [[Selman Waksman]] until the 1950s.
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| Sulfur oxidizers can use as energy sources reduced sulfur compounds, including hydrogen sulfide, elemental sulfur, [[sulfite]], [[thiosulfate]], and various polythionates (e.g., [[tetrathionate]]).<ref>{{cite journal | author = Pronk JT, Meulenberg R, Hazeu W, Bos P, Kuenen JG | year = 1990 | title = Oxidation of reduced inorganic sulphur compounds by acidophilic thiobacilli | journal =FEMS Microbiology letters | volume = 75 | issue = 2–3 |pages = 293–306 |url=http://repository.tudelft.nl/assets/uuid:9592868a-b999-4712-a233-191b615da6c6/864579.pdf | doi = 10.1111/j.1574-6968.1990.tb04103.x}}</ref> They depend on enzymes such as [[sulfur dioxygenase|sulfur oxygenase]] and [[sulfite oxidase]] to oxidize sulfur to sulfate. Some [[lithotroph]]s can even use the energy contained in sulfur compounds to produce sugars, a process known as [[chemosynthesis]]. Some [[bacteria]] and [[archaea]] use hydrogen sulfide in place of water as the [[electron donor]] in chemosynthesis, a process similar to [[photosynthesis]] that produces sugars and utilizes oxygen as the [[electron acceptor]]. The [[photosynthesis|photosynthetic]] [[green sulfur bacteria]] and [[purple sulfur bacteria]] and some [[lithotroph]]s use elemental oxygen to carry out such oxidization of hydrogen sulfide to produce elemental sulfur (S<sup>0</sup>), oxidation state = 0. Primitive bacteria that live around deep ocean [[hydrothermal vent|volcanic vents]] oxidize hydrogen sulfide in this way with oxygen; see [[giant tube worm]] for an example of large organisms that use hydrogen sulfide (via bacteria) as food to be oxidized.
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| The so-called [[sulfate-reducing bacteria]], 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 a number of other partially oxidized sulfur compounds (e.g. thiosulfates, thionates, polysulfides, sulfites). The hydrogen sulfide produced by these bacteria is responsible for some of the smell of intestinal gases ([[flatus]]) and decomposition products.
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| Sulfur is absorbed by [[plant]]s via the [[root]]s from soil as the [[sulfate]] and transported as a phosphate ester. Sulfate is reduced to sulfide via sulfite before it is incorporated into [[cysteine]] and other organosulfur compounds.<ref name="Heldt">{{cite book|isbn = 3-8274-0103-8| pages = 321–333|first = Hans-Walter|last = Heldt|title = Pflanzenbiochemie|publisher = Spektrum Akademischer Verlag|place = Heidelberg|year =1996}}</ref>
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| :SO<sub>4</sub><sup>2–</sup> → SO<sub>3</sub><sup>2–</sup> → H<sub>2</sub>S → cysteine → methionine
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| ==Precautions==
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| {{refimprove section|date=May 2012}}
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| [[File:Acid rain woods1.JPG|Effect of acid rain on a forest, Jizera Mountains, Czech Republic|thumb]]
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| Elemental sulfur is non-toxic, as generally are the soluble [[sulfate]] salts, such as [[Epsom salt]]s. Soluble sulfate salts are poorly absorbed and laxative. However, when injected parenterally, they are freely filtered by the kidneys and eliminated with very little toxicity in multi-gram amounts.
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| When sulfur burns in air, it produces [[sulfur dioxide]]. In water, this gas produces sulfurous acid and sulfites, which are antioxidants that inhibit growth of aerobic bacteria and allow its use as a [[food additive]] in small amounts. However, at high concentrations these acids harm the [[Human lungs|lungs]], [[Human eyes|eyes]] or other [[Biological tissue|tissues]]. In organisms without lungs such as insects or plants, it otherwise prevents [[Respiration (physiology)|respiration]] in high concentrations. [[Sulfur trioxide]] (made by catalysis from sulfur dioxide) and [[sulfuric acid]] are similarly highly corrosive, due to the strong acids that form on contact with water.
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| The burning of [[coal]] and/or [[petroleum]] by industry and [[power plants]] generates sulfur dioxide (SO<sub>2</sub>), which reacts with atmospheric water and oxygen to produce sulfuric acid (H<sub>2</sub>SO<sub>4</sub>) and [[sulfurous acid]] (H<sub>2</sub>SO<sub>3</sub>). These acids are components of [[acid rain]], which lower the [[pH]] of [[soil]] and freshwater bodies, sometimes resulting in substantial damage to the [[environment (biophysical)|environment]] and [[chemical weathering]] of statues and structures. Fuel standards increasingly require that fuel producers extract sulfur from [[fossil fuel]]s to prevent acid rain formation. This extracted and refined sulfur represents a large portion of sulfur production. In coal-fired power plants, [[flue gases]] are sometimes purified. More modern power plants that use [[synthesis gas]] extract the sulfur before they burn the gas.
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| [[Hydrogen sulfide]] is as [[toxic]] as [[hydrogen cyanide]], and kills by the same mechanism, though hydrogen sulfide is less likely to cause surprise poisonings from small inhaled amounts, because of its disagreeable warning odor. Though pungent at first, however, hydrogen sulfide quickly deadens the sense of smell—so a victim may breathe increasing quantities and be unaware of its presence until severe symptoms occur, which can quickly lead to death. Dissolved [[sulfide]] and [[hydrosulfide]] salts are also toxic by the same mechanism.
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| ==See also==
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| {{colbegin|2}}
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| * [[Cysteine]]
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| * [[Disulfide bond]]
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| * [[Methionine]]
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| * [[Stratospheric sulfur aerosols]]
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| * [[Sulfur assimilation]]
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| * [[Sulfur cycle]]
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| * [[Sulfur metabolism]]
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| * [[Ultra-low sulfur diesel]]
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| {{colend}}
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| {{Subject bar
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| |book1=Sulfur
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| |book2=Period 3 elements
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| |book3=Chalcogens
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| |book4=Chemical elements (sorted alphabetically)
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| |book5=Chemical elements (sorted by number)
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| |portal=Chemistry
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| |commons=y
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| |wikt=y
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| |wikt-search=sulfur
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| |v=y
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| |v-search=Sulfur atom
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| |b=y
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| |b-search=Wikijunior:The Elements/Sulfur
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| }}
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| ==References==
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| {{Reflist|colwidth=30em}}
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| ==External links==
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| * [http://www.periodicvideos.com/videos/016.htm Sulfur] at ''[[The Periodic Table of Videos]]'' (University of Nottingham)
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| * [http://physics.nist.gov/PhysRefData/Handbook/Tables/sulfurtable1.htm Atomic Data for Sulfur], [[NIST]] Physical Measurement Laboratory
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| * [http://library.tedankara.k12.tr/chemistry/vol2/allotropy/z129.htm Sulfur phase diagram], Introduction to Chemistry For Ages 13–17
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| * [http://www.stromboli.net/perm/vulcano/sulphur-vulcano-en.html Crystalline, liquid and polymerization of sulfur on Vulcano Island, Italy]
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| * [http://extoxnet.orst.edu/pips/sulfur.htm Sulfur and its use as a pesticide]
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| * [http://www.sulphurinstitute.org/ The Sulphur Institute]
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| {{Clear}}
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| {{Compact periodic table}}
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| {{Acne_Agents}}
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| [[Category:Sulfur| ]]
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| {{Link FA|fa}}
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