Examine individual changes

'{{About|the chemical element}} {{pp-move-indef}} {{Use American English|date=March 2017}} {{Use dmy dates|date=March 2016}} {{Infobox sulfur}} '''Sulfur''' is a [[chemical element]] with symbol&nbsp;'''S''' and [[atomic number]]&nbsp;16. It is [[Abundance of the chemical elements|abundant]], [[Valence (chemistry)|multivalent]], and [[nonmetal]]lic. Under [[Standard conditions for temperature and pressure|normal conditions]], sulfur atoms form cyclic octatomic molecules with a chemical formula [[Octasulfur|S₈]]. Elemental sulfur is a bright yellow [[crystal]]line solid at room temperature. Chemically, sulfur reacts with all elements except for [[gold]], [[platinum]], [[iridium]], [[tellurium]], and the [[noble gas]]es. Sulfur is the tenth most common element by mass in the universe, and the fifth most common on Earth. Though sometimes found in pure, [[Native element minerals|native]] form, sulfur on Earth usually occurs as [[Sulfide minerals|sulfide]] and [[sulfate minerals]]. Being abundant in native form, sulfur was known in ancient times, being mentioned for its uses in [[ancient India]], [[ancient Greece]], [[History of China#Ancient China|China]], and [[ancient Egypt|Egypt]]. In the [[Bible]], sulfur is called '''brimstone'''.<ref name=Greenwd/> Today, almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants from [[natural gas]] and [[petroleum]]. The greatest commercial use of the element is the production of [[sulfuric acid]] for sulfate and phosphate [[fertilizer]]s, and other chemical processes. The element sulfur is used in [[match]]es, [[insecticide]]s, and [[fungicide]]s. Many sulfur compounds are odoriferous, and the smells of odorized natural gas, skunk scent, grapefruit, and garlic are due to [[organosulfur]] compounds. [[Hydrogen sulfide]] gives the characteristic odor to rotting eggs and other biological processes. Sulfur is an [[essential element]] for all life, but almost always in the form of [[Organosulfur compounds|organosulfur]] compounds or metal sulfides. Three [[amino acid]]s ([[cysteine]], [[cystine]], and [[methionine]]) and two vitamins ([[biotin]] and [[thiamine]]) are organosulfur compounds. Many [[Cofactor (biochemistry)|cofactors]] also contain sulfur including [[glutathione]] and [[thioredoxin]] and [[iron–sulfur protein]]s. [[Disulfide]]s, S–S bonds, confer mechanical strength and insolubility of the protein [[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== [[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.]] ===Physical properties=== Sulfur forms polyatomic molecules with different chemical formulas, the best-known allotrope being [[octasulfur]], cyclo-S₈. The [[point group]] of cyclo-S₈ is D_4d 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}}</ref> Octasulfur is a soft, bright-yellow solid that is odorless, but impure samples have an 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 {{convert|115.21|C}}, boils at {{convert|444.6|C}} and sublimes easily.<ref name=Greenwd/> At {{convert|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 = 645–665}}</ref> The structure of the S₈ 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 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 g·cm⁻³, depending on the allotrope; all of the stable allotropes are excellent electrical insulators. ===Chemical properties=== Sulfur burns with a blue flame with formation of [[sulfur dioxide]], which has a suffocating and irritating 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 second ionization energies of sulfur are 999.6 and 2252 kJ·mol⁻¹, 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⁻¹, the magnitude of the figures caused by electron transfer between orbitals; these states are only stable with strong oxidants such as [[fluorine]], [[oxygen]], and [[chlorine]]. {{citation needed|date=March 2016}} Sulfur reacts with nearly all other elements with the exception of gold, platinum, iridium, nitrogen, tellurium, iodine and the noble gases. 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=== [[File:Cyclooctasulfur-above-3D-balls.png|thumb|left|The structure of the cyclooctasulfur molecule, S₈]] {{Main article|Allotropes of sulfur}} 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 |date = 2003 |volume = 230 |pages = 1–80 |doi = 10.1007/b12110 |series = Topics in Current Chemistry |isbn = 978-3-540-40191-9}}</ref> Besides S₈, 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 |date = 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 S₇, which is more deeply yellow than S₈. [[High-performance liquid chromatography|HPLC]] analysis of "elemental sulfur" reveals an equilibrium mixture of mainly S₈, but with S₇ and small amounts of S₆.<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}}, S₇, and S₈ 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 S₁₂ and S₁₈.<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> [[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. {{clear left}} ===Isotopes=== {{Main article|Isotopes of sulfur}} Sulfur has 25 known [[isotope]]s, four of which are stable: ³²S (94.99 ± 0.26%), ³³S (0.75 ± 0.02%), ³⁴S (4.25 ± 0.24%), and ³⁶S (0.01 ± 0.01%).<ref>[http://www.ciaaw.org/sulfur.htm Sulfur]. [[Commission on Isotopic Abundances and Atomic Weights]]</ref><ref name=b92>{{RubberBible92nd|page=1.14}}</ref> Other than ³⁵S, with a [[half-life]] of 87 days and formed in [[cosmic ray spallation]] of ⁴⁰[[Argon|Ar]], the [[radioactivity|radioactive]] isotopes of sulfur have half-lives less than 3&nbsp;hours. 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. 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 ³⁴S of ecosystem components. [[Rocky Mountain]] lakes thought to be dominated by atmospheric sources of sulfate have been found to have different [[Δ34S|δ³⁴S]] values from lakes believed to be dominated by watershed sources of sulfate. ===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 (moon)|Io]] are due to elemental sulfur and sulfur compounds deposited by active [[volcano]]es.]] [[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]] ³²S is created inside massive stars, at a depth where the temperature exceeds 2.5×10⁹&nbsp;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=https://fas.org/sgp/eprint/CRL-41.pdf|date=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. 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|date=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|date = 2005|doi = 10.1088/0034-4885/68/2/R02|bibcode=2005RPPh...68..303L}}</ref> [[File:Fumarola Vulcano.jpg|thumb|left|Sulfur occurs in [[fumaroles]] such as this one in [[Vulcano]], Italy]] It is the fifth most common element by mass in the 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. 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, [[Sicily]] was a major source of sulfur in the [[Industrial Revolution]].<ref>{{cite book|last=Kutney|first=Gerald|title=Sulfur: history, technology, applications & industry|date=2007|publisher=ChemTec Publications|location=Toronto|isbn=978-1-895198-37-9|oclc=79256100|page=43}}</ref> Native sulfur is synthesised by [[anaerobic bacteria]] acting on [[sulfate minerals]] such as [[gypsum]] in [[salt domes]].<ref>Klein, Cornelis and Cornelius S. Hurlbut, Jr., ''Manual of Mineralogy,'' Wiley, 1985, 20th ed., p. 265-6 {{ISBN|0-471-80580-7}}</ref><ref>[http://www.mindat.org/min-3826.html Sulphur on Mindat.org]</ref> Significant deposits in salt domes occur 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 were until recently 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 [[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. ==Compounds== {{Category see also|Sulfur compounds}} Common [[oxidation state]]s of sulfur range from −2 to +6. Sulfur forms stable compounds with all elements except the [[noble gas]]es. ===Sulfur polycations=== Sulfur [[polycation]]s, S₈²⁺, S₄²⁺ and S₁₆²⁺ are produced when sulfur is reacted with mild 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 [[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. S₈²⁺ is deep blue, S₄²⁺ is yellow and S₁₆²⁺ is red.<ref name = "Greenwood"/> ===Sulfides=== 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 ofPORN S THE BEST WHEN YOU STICK THE DICK IN THE PUZZY AND CUMMING INSIDE OF HER OOHOHOHOHOHHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOOHOHOHOHOHOHOHOHOHOHOHOOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOOHOHOHOHOHOHOHOHOOHOHOHOHOHOHOHOHOHOHOHOHOOHHOHOHOOHOOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOHOH 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''). Reduction of elemental sulfur gives [[polysulfide]]s, which consist of chains of sulfur atoms terminated with S⁻ centers: :2 Na + S₈ → Na₂S₈ This reaction highlights a distinctive property of sulfur: its ability to [[Catenation|catenate]] (bind to itself by formation of chains). [[Protonation]] of these polysulfide anions produces the [[polysulfane]]s, H₂S_x 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> Ultimately, reduction of sulfur produces sulfide salts: :16 Na + S₈ → 8 Na₂S The interconversion of these species is exploited in the [[sodium-sulfur battery]]. The [[Trisulfur|radical anion]] S₃⁻ gives the blue color of the mineral [[lapis lazuli]]. [[File:Lapis lazuli block.jpg|thumb|upright|[[Lapis lazuli]] owes its blue color to a [[trisulfur]] radical anion ({{chem|S|3|-}})]] [[File:S@CNT.jpg|thumb|Two parallel sulfur chains grown inside a single-wall [[carbon nanotube]] (CNT, a). Zig-zag (b) and straight (c) S chains inside double-wall CNTs<ref>{{cite journal|doi=10.1038/ncomms3162|pmid=23851903|pmc=3717502|title=Conducting linear chains of sulphur inside carbon nanotubes|journal=Nature Communications|volume=4|pages=2162|year=2013|last1=Fujimori|first1=Toshihiko|last2=Morelos-Gómez|first2=Aarón|last3=Zhu|first3=Zhen|last4=Muramatsu|first4=Hiroyuki|last5=Futamura|first5=Ryusuke|last6=Urita|first6=Koki|last7=Terrones|first7=Mauricio|last8=Hayashi|first8=Takuya|last9=Endo|first9=Morinobu|last10=Young Hong|first10=Sang|last11=Chul Choi|first11=Young|last12=Tománek|first12=David|last13=Kaneko|first13=Katsumi|bibcode=2013NatCo...4E2162F}}</ref>]] ===Oxides, oxoacids and oxoanions=== The principal sulfur oxides are obtained by burning sulfur: :S + O₂ → SO₂ ([[sulfur dioxide]]) :2 SO₂ + O₂ → 2 SO₃ ([[sulfur trioxide]]) Multiple sulfur oxides are known; the [[lower sulfur oxides|sulfur-rich oxides]] include sulfur monoxide, disulfur monoxide, disulfur dioxides, and [[higher sulfur oxides|higher oxides]] containing peroxo groups. Sulfur forms [[oxoacids|sulfur oxoacids]], some of which cannot be isolated and are only known through the salts. [[Sulfur dioxide]] and [[sulfite]]s ({{chem|SO|3|2−}}) are related to the unstable [[sulfurous acid]] (H₂SO₃). [[Sulfur trioxide]] and [[sulfate]]s ({{chem|SO|4|2−}}) are related to [[sulfuric acid]] (H₂SO₄). Sulfuric acid and SO₃ combine to give oleum, a solution of [[pyrosulfuric acid]] (H₂S₂O₇) in sulfuric acid. <!--[[Peroxymonosulfuric acid]] (H₂SO₅) and [[peroxydisulfuric acid]]s (H₂S₂O₈), made from the action of SO₃ on concentrated [[hydrogen peroxide|H₂O₂]], and [[sulfuric acid|H₂SO₄]] on concentrated H₂O₂ respectively.--> [[Thiosulfate]] salts ({{chem|S|2|O|3|2−}}), sometimes referred as "hyposulfites", used in [[Photographic fixer|photographic fixing]] (hypo) and as reducing agents, feature sulfur in two oxidation states. [[Sodium dithionite]] ({{chem|Na|2|S|2|O|4}}), contains the more highly reducing [[dithionite]] anion ({{chem|S|2|O|4|2−}}). <!--[[Sodium dithionate]] (Na₂S₂O₆) contains the [[dithionate]] anion (S₂O₆²⁻) and is the first member of the [[polythionic acid]]s (H₂S_''n''O₆), 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.--> ===Halides and oxyhalides=== Several sulfur halides are important to modern industry. [[Sulfur hexafluoride]] is a dense gas used as an [[dielectric gas|insulator gas]] in high voltage [[transformer]]s; it is also a nonreactive and nontoxic propellant for pressurized containers. [[Sulfur tetrafluoride]] is 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|doi=10.1002/0471264180.os041.28}}</ref> [[Sulfur dichloride]] and [[disulfur dichloride]] are important industrial chemicals. [[Sulfuryl chloride]] and [[chlorosulfuric acid]] are derivatives of sulfuric acid; [[thionyl chloride]] (SOCl₂) 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|doi=10.15227/orgsyn.030.0062}}</ref> ===Pnictides=== An important S–N compound is the cage [[tetrasulfur tetranitride]] (S₄N₄). Heating this compound gives [[Polythiazyl|polymeric sulfur nitride]] ((SN)_x), which has metallic properties even though it does not contain any [[metal]] atoms. [[Thiocyanate]]s contain the SCN⁻ group. Oxidation of thiocyanate gives [[thiocyanogen]], (SCN)₂ with the connectivity NCS-SCN. [[Phosphorus sulfide]]s are numerous, the most important commercially being the cages P₄S₁₀ and P₄S₃.<ref>{{Cite book |last=Heal |first=H. G. |title=The Inorganic Heterocyclic Chemistry of Sulfur, Nitrogen, and Phosphorus |publisher=Academic Press |location=London |date=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 |date=2004 |isbn=981-256-095-5}}</ref> ===Metal sulfides=== {{Main article|Sulfide mineral}} 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 [[Geochemical cycle|geochemically]] and in the laboratory, by the reaction of hydrogen sulfide with metal salts. The mineral [[galena]] (PbS) was the first demonstrated semiconductor and was used 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 FeS₂.<ref>Vaughan, D. J.; Craig, J. R. "Mineral Chemistry of Metal Sulfides" Cambridge University Press, Cambridge (1978) {{ISBN|0-521-21489-0}}</ref> Processing these ores, usually by [[smelting|roasting]], is costly and environmentally hazardous. Sulfur corrodes many metals through [[tarnishing]]. ===Organic compounds=== {{Main article|Organosulfur compounds}} <gallery caption="Illustrative organosulfur compounds"> File:R-allicin-2D-skeletal.png|[[Allicin]], the active ingredient in garlic File:Cysteine.svg| (''R'')-[[cysteine]], an [[amino acid]] containing a thiol group File:Methionin - Methionine.svg|[[Methionine]], an [[amino acid]] containing a thioether File:Diphenyl disulfide.png|[[Diphenyl disulfide]], a representative disulfide File:Perfluorooctanesulfonic acid.png|[[Perfluorooctanesulfonic acid]], a controversial surfactant File:Dibenzothiophen - Dibenzothiophene.svg|[[Dibenzothiophene]], a component of crude oil File:Penicillin core.svg|[[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>Cremlyn R. J.; "An Introduction to Organosulfur Chemistry" John Wiley and Sons: Chichester (1996). {{ISBN|0-471-95512-4}}.</ref> * [[Thiol]]s or mercaptans (so called because they capture mercury as [[Chelation|chelators]]) are the sulfur analogs of [[alcohol]]s; treatment of thiols with base gives [[thiolate]] ions. * [[Thioether]]s are the sulfur analogs of [[ether]]s. * [[Sulfonium]] ions have three groups attached to a cationic sulfur center. [[Dimethylsulfoniopropionate]] (DMSP) is one such compound, important in the marine organic [[sulfur cycle]]. * [[Sulfoxide]]s and [[sulfone]]s are thioethers with one and two oxygen atoms attached to the sulfur atom, respectively. The simplest sulfoxide, [[dimethyl sulfoxide]], is a common solvent; a common sulfone is [[sulfolane]]. * [[Sulfonic acid]]s are used in many detergents. Compounds with carbon-sulfur multiple bonds are uncommon, an exception being [[carbon disulfide]], a volatile colorless liquid that is structurally similar to carbon dioxide. It is used as a reagent to make the polymer [[rayon]] and many organosulfur compounds. Unlike [[carbon monoxide]], [[carbon monosulfide]] is stable only as an extremely dilute gas, found between solar systems.<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=15 March 1976|volume=204|pages=L135–L137|doi=10.1086/182072|bibcode=1976ApJ...204L.135W}}</ref> Organosulfur compounds are responsible for some of the unpleasant odors of decaying organic matter. They are widely known as the [[Odorizer|odorant]] in domestic natural gas, garlic odor, and skunk spray. Not all organic sulfur compounds smell unpleasant at all concentrations: the sulfur-containing [[terpene|monoterpenoid]] ([[grapefruit mercaptan]]) in small concentrations is the characteristic scent of grapefruit, but has a generic thiol odor at larger concentrations. [[Sulfur mustard]], a potent [[blister agent|vesicant]], was [[Chemical weapons in World War I|used in World War I]] as a disabling agent.<ref>{{cite book|last=Banoub|first=Joseph|title=Detection of Biological Agents for the Prevention of Bioterrorism|date=2011|publisher=Springer|location=Dordrecht|isbn=978-90-481-9815-3|oclc=697506461|page=183}}</ref> Sulfur-sulfur bonds are a structural component used to stiffen rubber, similar to the disulfide bridges that rigidify proteins (see biological below). In the most common type of industrial "curing" or hardening and strengthening of natural [[rubber]], elemental sulfur is heated with the rubber to the point that chemical reactions form [[disulfide]] bridges between [[isoprene]] units of the polymer. This process, patented in 1843, made rubber a major industrial product, especially in automobile tires. Because of the heat and sulfur, the process was named [[vulcanization]], after the Roman god of the forge and [[volcanism]]. ==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]] ([[Book of Genesis|Genesis]]). [[English translations of the Bible]] commonly referred to burning sulfur as "brimstone", giving rise to the term "[[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 = https://books.google.com/books?id=ed0yC98aAKYC&pg=PA242 | title = Archaeomineralogy | isbn = 978-3-540-78593-4 | page = 242 | author1 = Rapp | first1 = George Robert | date = 4 February 2009}}</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 16 August 2012.</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> A natural form of sulfur known as ''shiliuhuang'' (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 (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. It remains an ingredient of [[gunpowder|black gunpowder]]. 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|date=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). {{multiple image | align = | direction = | width = |footer=Various alchemical symbols for sulfur | image1 = Sulphur.svg |width1=162 | caption1 = | image2 = Sulfur symbol 2.svg |width2=100 | caption2 = {{Citation needed|date=November 2016}} | image3 = Alchemy sulfur symbol.svg |width3=146 |total_width=250 | height1 = 252 | height2 = 100 | height3 = 150 }}Early [[Europe]]an [[alchemy|alchemists]] gave sulfur a unique [[alchemical symbol]], a triangle at the top of a cross. In traditional skin treatment, elemental sulfur was used (mainly in creams) to alleviate such conditions as [[scabies]], [[ringworm]], [[psoriasis]], [[eczema]], and [[acne]]. The mechanism of action is unknown—though elemental sulfur does oxidize slowly to sulfurous acid, which is (through the action of [[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}}</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|date = 2004| pmid = 15303787}}</ref> ===Modern times=== [[File:Soufre extraction 1.jpg|thumb|Sicilian kiln used to obtain sulfur from volcanic rock]] In 1777, [[Antoine Lavoisier]] helped convince the scientific community that sulfur was an element, not a compound. Sulfur deposits in [[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 [[Marseilles]], France, for the production of [[sulfuric acid]] for use in the [[Leblanc process]]. In [[Industrial Revolution|industrializing]] Britain, with the repeal of [[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 [[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 solution was eventually negotiated by France.<ref>{{cite book|url=https://books.google.com/?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|accessdate=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|pages=163–180|doi=10.1177/026569149502500201|last1=Thomson|first1=D. W.|issue=2}}</ref> In 1867, elemental 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|date = 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> In the late 18th century, [[furniture]] makers used molten sulfur to produce [[sulfur inlay|decorative inlays]] 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|date=2006|publisher=Littleton|location=Colorado|isbn=978-0-87335-233-8|edition=7th|page=935|oclc=62805047}}</ref> ===Spelling and etymology=== ''Sulfur'' is derived from the Latin word ''{{lang|la|sulpur}}'', which was [[Hellenized]] to ''{{lang|la|sulphur}}''. The spelling ''{{lang|la|sulfur}}'' appears toward the end of the [[Classical antiquity|Classical period]]. (The true Greek word for sulfur, {{lang|grc|θεῖον}}, is the source of the international chemical prefix ''[[thio-]]''.) In 12th-century [[Anglo-Norman language|Anglo-French]], it was ''{{lang|xno|sulfre}}''; in the 14th century the Latin ''{{lang|la|-ph-}}'' was restored, for ''{{lang|enm|sulphre}}''; and by the 15th century the full Latin spelling was restored, for ''sulfur, sulphur''. The parallel ''f~ph'' spellings continued in Britain until the 19th century, when the word was standardized as ''sulphur''.<ref name="OED">{{OED|sulphur}}</ref> ''Sulfur'' was the form chosen in the United States, whereas Canada uses both. The [[International Union of Pure and Applied Chemistry|IUPAC]] adopted the spelling ''sulfur'' in 1990, as did the Nomenclature Committee of the [[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 |title=sulphur – definition of sulphur in English |website=Oxford Dictionaries |access-date=2016-11-19 }}</ref> ==Production== [[File:Bergelut dengan asap nan beracun.jpg|thumb|Traditional sulfur mining at [[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; 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|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| url =https://archive.org/stream/manfarthestdownr00wash#page/214 |year=1912|publisher=Doubleday, Page.|page=214}}</ref> Elemental sulfur was extracted from [[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 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. 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}}</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|Sulfur recovered from hydrocarbons in [[Alberta]], stockpiled for shipment in North Vancouver, British Columbia]] 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"/> :R-S-R + 2 H₂ → 2 RH + H₂S 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"/> :3 O₂ + 2 H₂S → 2 SO₂ + 2 H₂O :SO₂ + 2 H₂S → 3 S + 2 H₂O [[File:SulfurPrice.png|thumb|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 = 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|date=2010|publisher=J. Ross Publishing|location=[[Fort Lauderdale]]|isbn=978-1-60427-005-1|oclc=531718953|page=109}}</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 web|last1=McElvaney|first1=Kevin|title=The Men Who Mine Volcanos|url=https://www.theatlantic.com/features/archive/2015/02/the-men-who-mine-volcanoes-indonesia/385913/|website=theatlantic.com|publisher=The Atlantic|accessdate=26 February 2015}}</ref> 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> 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 [[sulfuric acid]] ([[hydrogen|H]]₂SO₄): :2 S + 3 O₂ + 2 H₂O → 2 H₂SO₄[[File:2000sulphuric acid.PNG|thumb|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 = 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/> ===Other important sulfur chemistry=== Sulfur reacts directly with methane to give [[carbon disulfide]], used to manufacture [[cellophane]] and [[rayon]].<ref name=Nehb>{{cite book |last=Nehb |first=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=3-527-30673-0}}</ref> One of the direct uses of sulfur is in [[vulcanization]] of rubber, where [[polysulfide]] chains crosslink organic polymers.<!--need something on pulping--> Large quantities of [[sulfite]]s are used to [[Bleach (chemical)|bleach]] [[paper]] and to preserve dried [[fruit]]. Many [[surfactant]]s and [[detergents]] (e.g. [[sodium lauryl sulfate]]) are sulfate derivatives. [[Calcium sulfate]], gypsum, (CaSO₄·2H₂O) is mined on the scale of 100 million [[tonne]]s each year for use in [[Portland cement]] and fertilizers. When silver-based [[photography]] was widespread, sodium and ammonium [[sodium thiosulfate|thiosulfate]] were widely used as "fixing agents." Sulfur is a component of [[gunpowder]] ("black powder"). ===Fertilizer=== 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]] (not soluble in water) and cannot be used directly by plants. Over time, soil bacteria can convert it to soluble derivatives, which can then be used by plants. Sulfur improves the 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 16 August 2012.</ref> Biologically produced sulfur particles are naturally hydrophilic due to a biopolymer coating and are easier to disperse over the land in a spray of diluted slurry, resulting in a faster uptake. The botanical requirement for sulfur equals or exceeds the requirement for phosphorus. [[Plant nutrition|It is an 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}}</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}}</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. ===Fine chemicals=== [[File:Malathion-3D-vdW.png|thumb|left|A molecular model of the pesticide [[malathion]]]] 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/> [[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]]. ===Fungicide and pesticide===<!--[[Wettable Sulfur]] redirs here--> [[File:Sulphur Candle.jpg|thumb|upright|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 [[Organic agriculture|organically]] farmed apple production against the main disease [[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 [[miscibility|miscible]].<ref>{{cite book |url = https://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 = 13 July 2010}}</ref><ref>{{cite web|url=http://www.freepatentsonline.com/3398227.pdf|title=Method for Preparation of Wettable Sulfur|accessdate= 20 May 2010|author= Every, Richard L.|display-authors= etal|date= 20 August 1968}}</ref> It has similar applications and is used as a [[fungicide]] against [[mildew]] and other mold-related problems with plants and soil. 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 application is dusting the clothing or limbs with sulfur powder. A diluted solution of [[lime sulfur]] (made by combining [[calcium hydroxide]] with elemental sulfur in water) is used as a dip for pets to destroy [[ringworm|ringworm (fungus)]], [[mange]], and other [[cutaneous conditions|dermatoses]] and [[parasitism|parasites]]. Sulfur candles of almost pure sulfur were burned to [[fumigant|fumigate]] structures and wine barrels, but are now considered too toxic for residences. ===Bactericide in winemaking and food preservation=== Small amounts of [[sulfur dioxide]] gas addition (or equivalent [[potassium metabisulfite]] addition) to fermented wine to produce traces of [[sulfurous acid]] (produced when SO₂ 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 26 October 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. Sulfur dioxide and various sulfites have been used for their antioxidant antibacterial preservative properties in many other parts of the food industry. The practice has declined since reports of an allergy-like reaction of some persons to sulfites in foods. ===Pharmaceuticals=== {{Infobox drug | drug_name = Octasulfur | IUPAC_name = Octathiocane | image = Molecular Sulfur (S8) V.1.svg | alt = | caption = <!-- Clinical data --> | tradename = | Drugs.com = {{Drugs.com|MTM|sulfur-topical}} | MedlinePlus = | pregnancy_AU = <!-- A / B1 / B2 / B3 / C / D / X --> | pregnancy_US = <!-- A / B / C / D / X --> | pregnancy_category= | legal_AU = <!-- S2, S3, S4, S5, S6, S7, S8, S9 or Unscheduled--> | legal_CA = <!-- OTC, Rx-only, Schedule I, II, III, IV, V, VI, VII, VIII --> | legal_UK = <!-- GSL, P, POM, CD, CD Lic, CD POM, CD No Reg POM, CD (Benz) POM, CD (Anab) POM or CD Inv POM --> | legal_US = OTC | legal_status = OTC | routes_of_administration = Topical, rarely oral <!-- Pharmacokinetic data --> | bioavailability = | protein_bound = | metabolism = | elimination_half-life = | excretion = <!-- Identifiers --> | CAS_number = 10544-50-0 | ATCvet = | ATC_prefix = D10 | ATC_suffix = AB02 | PubChem = 66348 | DrugBank = | ChemSpiderID = 59726 | ChEBI = 29385 <!-- Chemical data --> | S = 8 | molecular_weight = 256.52 g/mol | smiles = S1SSSSSSS1 | StdInChI=1S/S8/c1-2-4-6-8-7-5-3-1 | StdInChIKey = JLQNHALFVCURHW-UHFFFAOYSA-N }} Sulfur (specifically [[octasulfur]], S₈) is used in pharmaceutical skin preparations for the treatment of [[acne]] and other conditions. It acts as a [[keratolytic]] agent and also kills bacteria, fungi, [[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=German|date=1978|volume=6B|pages=672–9|isbn=3-540-07738-3}}</ref> Precipitated sulfur and colloidal sulfur are used, in form of [[lotion]]s, creams, powders, soaps, and bath additives, for the treatment of [[acne vulgaris]], [[acne rosacea]], and [[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=German|isbn=978-3-8047-2575-1|at=Monographie ''Schwefel zum äußerlichen Gebrauch'' [Monograph ''Sulfur for external use'']|edition=23rd}}</ref> Common adverse effects include irritation of the skin at the application site, such as dryness, stinging, itching and peeling.<ref name="MTM">{{Drugs.com|MTM|sulfur-topical}}: Sulfur topical.</ref> ====Mechanism of action==== {{expand section|1=the keratolysis mechanism|date=March 2014}} Sulfur is converted to hydrogen sulfide (H₂S) through [[Redox|reduction]], partly by bacteria. H₂S kills bacteria (possibly including ''[[Propionibacterium acnes]]'' which plays a role in acne,<ref name="Drugs.com">{{Drugs.com|pro|perrigo-sodium-sulfacetamide-and-sulfur}}: Perrigo Sodium Sulfacetamide and Sulfur.</ref>) fungi, and parasites such as scabies mites.<ref name="Hager" /> ===Furniture=== {{main|Sulfur inlay}} Sulfur can be used to create decorative [[inlay]]s in wooden furniture. After a design has been cut into the wood, molten sulfur is poured in and then scraped away so it is flush. Sulfur inlays were particularly popular in the late 18th and early 19th centuries, notably amongst [[Pennsylvania German]] cabinetmakers. The practice soon died out, as less toxic and flammable substances were substituted. However, some modern craftsmen have occasionally revived the technique in the creation of replica pieces.<ref>[http://www.popsci.com/diy/article/2005-01/worst-way-inlay The Worst Way To Inlay], ''[[Popular Science]]'', January 1, 2005.</ref><ref>{{cite journal|doi=10.1088/0957-0233/14/9/311|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}}</ref> ==Biological role== ===Protein and organic cofactors=== 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, about equal in abundance to [[potassium]], and slightly greater than sodium and chlorine. A {{convert|70|kg|abbr=on}} human body contains about 140&nbsp;grams of sulfur. In [[plant]]s and [[animal]]s, the [[amino acid]]s [[cysteine]] and [[methionine]] contain most of the sulfur, and the element is 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. [[Disulfide bond]]s (S-S bonds) 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 due in part to the high content of S-S bonds with cysteine and sulfur. Eggs are high in sulfur to nourish feather formation in chicks, 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. [[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|date = 2000}}</ref> Two of the 13 classical vitamins, [[biotin]] and [[thiamine]], contain sulfur, with the latter being named for its sulfur content. In intracellular chemistry, sulfur operates 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 proteins essential to all known life, use neighboring pairs of reduced cysteines to work as general protein reducing agents, with similar effect. [[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]], CH₃SCH₂CH₂SO₃⁻, 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|date = 1998|pmid = 9782487|doi = 10.1099/00221287-144-9-2377}}</ref> ===Metalloproteins and inorganic cofactors=== Inorganic sulfur forms a part of [[iron–sulfur cluster]]s as well as many copper, nickel, and iron proteins. Most pervasive are the [[ferrodoxin]]s, which serve as electron shuttles in cells. In bacteria, the important [[nitrogenase]] enzymes contains an Fe–Mo–S cluster and 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|date =1994}}</ref> :[[File:FdRedox.png|center|500px]] ===Sulfur metabolism and the sulfur cycle=== {{Main article|Sulfur metabolism|Sulfur cycle}} 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]] (H₂S) 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. 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 | 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 |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 (S⁰), oxidation state = 0. Primitive bacteria that live around deep ocean [[hydrothermal vent|volcanic vents]] oxidize hydrogen sulfide in this way with oxygen; the [[giant tube worm]] is an example of a large organism that uses hydrogen sulfide (via bacteria) as food to be oxidized. 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 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. Sulfur is absorbed by [[plant]]s [[root]]s from soil as [[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|date =1996}}</ref> :SO₄²⁻ → SO₃²⁻ → H₂S → cysteine → methionine ==Precautions== {{refimprove section|date=May 2012}} {{NFPA 704|Health = 0|Flammability = 2|Reactivity = 0|S= |caption=Fire diamond hazard sign for elemental sulfur<ref>[http://periodictable.com/Elements/016/data.html Technical data for Sulfur]. periodictable.com</ref>}} [[File:Acid rain woods1.JPG|Effect of acid rain on a forest, Jizera Mountains, Czech Republic|thumb]] Elemental sulfur is non-toxic, as are most of the soluble [[sulfate]] salts, such as [[Epsom salt]]s. Soluble sulfate salts are poorly absorbed and laxative. When injected parenterally, they are freely filtered by the kidneys and eliminated with very little toxicity in multi-gram amounts. When sulfur burns in air, it produces [[sulfur dioxide]]. In water, this gas produces sulfurous acid and sulfites; sulfites are antioxidants that inhibit growth of aerobic bacteria and a useful [[food additive]] in small amounts. 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, sulfite in high concentration prevents [[Respiration (physiology)|respiration]]. [[Sulfur trioxide]] (made by catalysis from sulfur dioxide) and [[sulfuric acid]] are similarly highly acidic and corrosive in the presence of water. Sulfuric acid is a strong dehydrating agent that can strip available water molecules and water components from sugar and organic tissue.<ref>{{Cite web|url=http://www.rsc.org/education/eic/issues/2007March/ExhibitionChemistry.asp|title=EiC March 2007 – Feature – Exhibition chemistry: the dehydration of source|website=www.rsc.org|access-date=2016-04-30}}</ref> The burning of [[coal]] and/or [[petroleum]] by industry and [[power plants]] generates sulfur dioxide (SO₂) that reacts with atmospheric water and oxygen to produce sulfuric acid (H₂SO₄) and [[sulfurous acid]] (H₂SO₃). These acids are components of [[acid rain]], lowering 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. [[Hydrogen sulfide]] is as [[toxic]] as [[hydrogen cyanide]], and kills by the same mechanism (inhibition of the respiratory enzyme [[cytochrome oxidase]]),<ref>[http://emedicine.medscape.com/article/815139-overview#a4 Gresham et al, Medscape: Hydrogen sulfide toxicity]</ref> though hydrogen sulfide is less likely to cause surprise poisonings from small inhaled amounts because of its disagreeable odor. Hydrogen sulfide quickly deadens the sense of smell and a victim may breathe increasing quantities without noticing the increase until severe symptoms cause death. Dissolved [[sulfide]] and [[hydrosulfide]] salts are toxic by the same mechanism. ==See also== * [[Stratospheric sulfur aerosols]] * [[Sulfur assimilation]] * [[Ultra-low sulfur diesel]] {{Subject bar |book1=Sulfur |book2=Period 3 elements |book3=Chalcogens |book4=Chemical elements (sorted&nbsp;alphabetically) |book5=Chemical elements (sorted by number) |portal=Chemistry |commons=y |wikt=y |wikt-search=sulfur |v=y |v-search=Sulfur atom |b=y |b-search=Wikijunior:The Elements/Sulfur }} ==References== {{Reflist|30em}} ==External links== * [http://www.periodicvideos.com/videos/016.htm Sulfur] at ''[[The Periodic Table of Videos]]'' (University of Nottingham) * [http://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], Introduction to Chemistry For Ages 13–17 * [http://www.stromboli.net/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] * [http://www.sulphurinstitute.org/ The Sulphur Institute] * [http://www.nutrientstewardship.com/partners/products-and-services/sulfur-institute Nutrient Stewardship and The Sulphur Institute] {{Compact periodic table}} {{Sulfur compounds}} {{Authority control}} [[Category:Sulfur| ]] [[Category:Chemical elements]] [[Category:Chalcogens]] [[Category:Polyatomic nonmetals]] [[Category:Native element minerals]] [[Category:Dietary minerals]] [[Category:Inorganic polymers]] [[Category:Pyrotechnic fuels]] [[Category:Agricultural chemicals]] [[Category:Biology and pharmacology of chemical elements]] [[Category:Anti-acne preparations]] [[Category:Orthorhombic minerals]] [[Category:Industrial minerals]]'