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| − | {{Infobox sulfur}} | + | {{Elementbox |
| | + | |name=sulfur |
| | + | |number=16 |
| | + | |symbol=S |
| | + | |left=[[phosphorus]] |
| | + | |right=[[chlorine]] |
| | + | |above=[[oxygen|O]] |
| | + | |below=[[selenium|Se]] |
| | + | |series=nonmetal |
| | + | |series comment= |
| | + | |group=16 |
| | + | |period=3 |
| | + | |block=p |
| | + | |series color= |
| | + | |phase color= |
| | + | |appearance=Lemon yellow crystals. |
| | + | |image name=sulfur.jpg |
| | + | |image size= |
| | + | |image name comment= |
| | + | |image name 2= |
| | + | |image name 2 comment= |
| | + | |atomic mass=32.065 |
| | + | |atomic mass 2=5 |
| | + | |atomic mass comment= |
| | + | |electron configuration=[[[neon|Ne]]] 3s<sup>2</sup> 3p<sup>4</sup> |
| | + | |electrons per shell=2, 8, 6 |
| | + | |color= |
| | + | |phase=solid |
| | + | |phase comment= |
| | + | |density gplstp= |
| | + | |density gpcm3nrt=(alpha) 2.07 |
| | + | |density gpcm3nrt 2=(beta) 1.96 |
| | + | |density gpcm3nrt 3=(gamma) 1.92 |
| | + | |density gpcm3mp=1.819 |
| | + | |melting point K=388.36 |
| | + | |melting point C=115.21 |
| | + | |melting point F=239.38 |
| | + | |boiling point K=717.8 |
| | + | |boiling point C=444.6 |
| | + | |boiling point F=832.3 |
| | + | |triple point K= |
| | + | |triple point kPa= |
| | + | |critical point K=1314 |
| | + | |critical point MPa=20.7 |
| | + | |heat fusion=(mono) 1.727 |
| | + | |heat fusion 2= |
| | + | |heat vaporization=(mono) 45 |
| | + | |heat capacity=22.75 |
| | + | |vapor pressure 1=375 |
| | + | |vapor pressure 10=408 |
| | + | |vapor pressure 100=449 |
| | + | |vapor pressure 1 k=508 |
| | + | |vapor pressure 10 k=591 |
| | + | |vapor pressure 100 k=717 |
| | + | |vapor pressure comment= |
| | + | |crystal structure=orthorhombic |
| | + | |oxidation states='''6''', 5, '''4''', 3, ''' 2''', [[Disulfur dichloride|1]], -1, -2 |
| | + | |oxidation states comment=strongly [[acid]]ic oxide |
| | + | |electronegativity=2.58 |
| | + | |number of ionization energies=4 |
| | + | |1st ionization energy=999.6 |
| | + | |2nd ionization energy=2252 |
| | + | |3rd ionization energy=3357 |
| | + | |atomic radius= |
| | + | |atomic radius calculated= |
| | + | |covalent radius=[[1 E-10 m|105±3]] |
| | + | |Van der Waals radius=[[1 E-10 m|180]] |
| | + | |magnetic ordering=[[diamagnetic]]<ref name=magnet>{{cite book| url = http://www-d0.fnal.gov/hardware/cal/lvps_info/engineering/elementmagn.pdf | title = Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics| publisher = CRC press| isbn = 0849304814| year = 2000}}</ref> |
| | + | |electrical resistivity= |
| | + | |electrical resistivity at 0= |
| | + | |electrical resistivity at 20=(amorphous)<br />2×10<sup>15</sup> |
| | + | |thermal conductivity=(amorphous)<br />0.205 |
| | + | |thermal conductivity 2= |
| | + | |thermal diffusivity= |
| | + | |thermal expansion= |
| | + | |thermal expansion at 25= |
| | + | |speed of sound= |
| | + | |speed of sound rod at 20= |
| | + | |speed of sound rod at r.t.= |
| | + | |Young's modulus= |
| | + | |Shear modulus= |
| | + | |Bulk modulus=7.7 |
| | + | |Poisson ratio= |
| | + | |Mohs hardness=2.0 |
| | + | |Vickers hardness= |
| | + | |Brinell hardness= |
| | + | |CAS number=7704-34-9 |
| | + | |EC number=231-722-6 |
| | + | |isotopes= |
| | + | {{Elementbox_isotopes_stable | mn=32 | sym=S | na=95.02% | n=16 }} |
| | + | {{Elementbox_isotopes_stable | mn=33 | sym=S | na=0.75% | n=17 }} |
| | + | {{Elementbox_isotopes_stable | mn=34 | sym=S | na=4.21% | n=18 }} |
| | + | {{Elementbox_isotopes_decay | mn=35 | sym=S | na=[[synthetic radioisotope|syn]] | hl=[[1 E6 s|87.32 d]] | dm=[[beta emission|β<sup>−</sup>]] | de=0.167 | pn=35 | ps=[[chlorine|Cl]] }} |
| | + | {{Elementbox_isotopes_stable | mn=36 | sym=S | na=0.02% | n=20 }} |
| | + | |isotopes comment= |
| | + | }} |
| | '''Sulfur''' or '''sulphur''' ({{pron-en|ˈsʌlfər}}, [[#Spelling and etymology|see spelling below]]) is the [[chemical element]] that has the [[atomic number]] 16. It is denoted with the symbol '''S'''. It is an abundant, [[Valence (chemistry)|multivalent]] [[non-metal]]. Sulfur in its native form is a yellow [[crystal]]line solid. In [[nature]], it can be found as the pure element and as [[sulfide]] and [[sulfate]] minerals. It is an essential element for life and is found in two [[amino acid]]s, [[cysteine]] and [[methionine]]. Its commercial uses are primarily in [[fertilizer]]s, but it is also widely used in black [[gunpowder]], [[match]]es, [[insecticide]]s and [[fungicide]]s. Elemental sulfur crystals are commonly sought after by mineral collectors for their brightly colored [[polyhedron]] shapes. In nonscientific context it can also be referred to as ''brimstone''. | | '''Sulfur''' or '''sulphur''' ({{pron-en|ˈsʌlfər}}, [[#Spelling and etymology|see spelling below]]) is the [[chemical element]] that has the [[atomic number]] 16. It is denoted with the symbol '''S'''. It is an abundant, [[Valence (chemistry)|multivalent]] [[non-metal]]. Sulfur in its native form is a yellow [[crystal]]line solid. In [[nature]], it can be found as the pure element and as [[sulfide]] and [[sulfate]] minerals. It is an essential element for life and is found in two [[amino acid]]s, [[cysteine]] and [[methionine]]. Its commercial uses are primarily in [[fertilizer]]s, but it is also widely used in black [[gunpowder]], [[match]]es, [[insecticide]]s and [[fungicide]]s. Elemental sulfur crystals are commonly sought after by mineral collectors for their brightly colored [[polyhedron]] shapes. In nonscientific context it can also be referred to as ''brimstone''. |
| | | | |
| − | ==History==
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| − | {{Expand-section|date=January 2008}}
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| − | [[File:SulphurCrystal.jpg|thumb|left|Rough sulfur crystal]]
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| − | [[File:Large Sulfur Crystal.jpg|thumb|left|Sulfur crystal from Agrigento, [[Sicily]].]]
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| − | Sulfur ([[Sanskrit]], ''sulvari''; [[Latin]] ''sulfur'' or ''sulpur'') was known in ancient times and is referred to in the [[Torah]] ([[Book of Genesis|Genesis]]).
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| − |
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| − | English translations of the [[Bible]] commonly referred to burning sulfur as "brimstone", giving rise to the name of '[[fire and brimstone]]' [[sermon]]s, in which listeners are reminded of the fate of eternal damnation that await the unbelieving and unrepentant. It is from this part of the Bible that [[Hell]] is implied to "smell of sulfur", although sulfur, in itself, is in fact odorless. The "smell of sulfur" usually refers to either the odor of [[hydrogen sulfide]], e.g. from rotten egg, or of burning sulfur, which produces [[sulfur dioxide]], the smell associated with burnt matches. The smell emanating from raw sulfur originates from a slow oxidation in the presence of air. Hydrogen sulfide is the principal odor of untreated [[sewage]] and is one of several unpleasant smelling sulfur-containing components of [[flatulence]] (along with sulfur-containing [[Thiol|mercaptans]]).
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| − |
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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 | doi = 10.1086/354207 | pages = 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. Early [[alchemy|alchemists]] gave sulfur its own [[alchemical symbol]] which was a triangle at the top of a cross.
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| − |
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| − | In 1777 [[Antoine Lavoisier]] helped convince the scientific community that sulfur was an element and not a compound. In 1867, sulfur was discovered in underground deposits in [[Louisiana]] and [[Texas]]. The overlying layer of earth was [[quicksand]], prohibiting ordinary mining operations, therefore the [[Frasch process]] was developed.
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| − |
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| − | ===Spelling and etymology===
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| − | The element has traditionally been spelled ''sulphur'' in the [[United Kingdom]] (since the 14th Century)<ref>http://www.rod.beavon.clara.net/sulphur.htm, retrieved 2nd April 2009 18:29 GMT</ref>, most of [[the Commonwealth]] including [[India]], [[Malaysia]], [[South Africa]], and [[Hong Kong]], along with the rest of the [[Caribbean]] and [[Republic of Ireland|Ireland]], but ''sulfur'' in the [[United States]], while both spellings are used in [[Australia]], [[New Zealand]], [[Canada]], and the [[Philippines]]. [[IUPAC]] adopted the spelling “sulfur” in 1990, as did the [[Royal Society of Chemistry]] Nomenclature Committee in 1992<ref>[http://www.rsc.org/delivery/_ArticleLinking/DisplayArticleForFree.cfm?doi=JM99101FP055&JournalCode=JM Spelling of Sulfur (PDF)]</ref> and the [[Qualifications and Curriculum Authority]] for England and Wales recommended its use in 2000.<ref>[http://www.worldwidewords.org/topicalwords/tw-sul1.htm Worldwidewords], [[9 December]] [[2000]]</ref>
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| − |
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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 [[Classical Greek]] loan, so the ''ph'' variant does not denote the Greek letter φ. 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, with the f spelling becoming dominant in the medieval period.<ref>[http://elements.vanderkrogt.net/elem/s.html Vanderkrogt.net]</ref><ref>Kelly DP (1995) Sulfur and its Doppelgänger. ''Arch. Microbiol.'' '''163''': 157-158.</ref>
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| − |
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| − | ==Characteristics==
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| − | [[File:Burning-sulfur.png|thumb|right|Sulfur melts to a blood-red liquid. When burned, it emits a blue flame.]]
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| − | At room temperature, sulfur is a soft, bright-yellow solid. Elemental sulfur has only a faint odor, similar to that of [[matches]].
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| − | The odor associated with rotten eggs is due to [[hydrogen sulfide]] ({{chem|H|2|S}}) and organic sulfur compounds rather than elemental sulfur.
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| − | Sulfur burns with a blue flame that emits [[sulfur dioxide]], notable for its peculiar suffocating odor due to dissolving in the mucosa to form dilute [[sulfurous acid]]. Sulfur itself is insoluble in water, but [[solubility|soluble]] in [[carbon disulfide]] — and to a lesser extent in other non-polar organic solvents such as [[benzene]] and [[toluene]]. Common [[oxidation state]]s of sulfur include −2, +2, +4 and +6. Sulfur forms stable compounds with all elements except the [[noble gas]]es.
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| − | Sulfur in the solid state ordinarily exists as cyclic crown-shaped S<sub>8</sub> molecules.
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| − |
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| − | The [[crystallography]] of sulfur is complex. Depending on the specific conditions, the sulfur [[allotrope]]s form several distinct [[crystal structure]]s, with [[rhombic]] and [[monoclinic]] S<sub>8</sub> best known.
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| − |
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| − | A noteworthy property of sulfur is that its [[viscosity]] in its molten state, unlike most other liquids, increases above temperatures of 200 °C due to the formation of [[polymer]]s. The molten sulfur assumes a dark red color above this temperature. At higher temperatures, however, the viscosity is decreased as depolymerization occurs.
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| − |
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| − | [[Amorphous]] or "plastic" sulfur can be produced through the rapid cooling of molten sulfur. [[X-ray crystallography]] studies show that the amorphous form may have a [[helix|helical]] structure with eight atoms per turn. This form is [[Metastability in molecules|metastable]] at room temperature and gradually reverts back to crystalline form. This process happens within a matter of hours to days but can be rapidly catalyzed.
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| − |
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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 more than 30 solid [[allotropy|allotrope]]s, more than any other element.<ref>{{cite journal
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| − | | author = Ralf Steudel, Bodo Eckert
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| − | | title = Solid Sulfur Allotropes Sulfur Allotropes
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| − | | journal = Topics in Current Chemistry
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| − | | year = 2003
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| − | | volume = 230
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| − | | pages = 1–80
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| − | | doi = 10.1007/b12110}}</ref> Besides S<sub>8</sub>, several other rings are known.<ref>{{cite journal
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| − | | author = Steudel, R.
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| − | | title = Homocyclic Sulfur Molecules
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| − | | journal = Topics Curr. Chem.
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| − | | year = 1982
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| − | | volume = 102
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| − | | pages = 149}}</ref> Removing one atom from the crown gives S<sub>7</sub>, which is more deeply yellow than S<sub>8</sub>. [[HPLC]] analysis of "elemental sulfur" reveals an equilibrium mixture of mainly S<sub>8</sub>, but also S<sub>7</sub> and small amounts of S<sub>6</sub>.<ref>{{cite journal
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| − | | author = Tebbe, F. N.; Wasserman, E.; Peet, W. G.; Vatvars, A. and Hayman, A. C.
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| − | | title = Composition of Elemental Sulfur in Solution: Equilibrium of {{chem|S|6}}, S<sub>7</sub>, and S<sub>8</sub> at Ambient Temperatures
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| − | | journal = J. Am. Chem. Soc.
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| − | | year = 1982
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| − | | volume = 104
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| − | | pages = 4971
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| − | | doi = 10.1021/ja00382a050}}</ref> Larger rings have been prepared, including S<sub>12</sub> and S<sub>18</sub>.<ref>{{cite journal
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| − | | author = Beat Meyer
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| − | | title = Solid Allotropes of Sulfur
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| − | | journal = Chem. Rev.
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| − | | year = 1964
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| − | | volume = 64
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| − | | issue = 4
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| − | | pages = 429–451
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| − | | doi = 10.1021/cr60230a004}}</ref><ref>{{cite journal
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| − | | author = Beat Meyer
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| − | | title = Elemental sulfur
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| − | | journal = Chem. Rev.
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| − | | year = 1976
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| − | | volume = 76
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| − | | pages = 367–388
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| − | | doi = 10.1021/cr60301a003}}</ref> By contrast, sulfur's lighter neighbor [[oxygen]] only exists in two states of allotropic significance: O<sub>2</sub> and O<sub>3</sub>. [[Selenium]], the heavier analogue of sulfur, can form rings but is more often found as a polymer chain.
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| − |
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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, the [[radioactive isotopes]] of sulfur are all short lived. <sup>35</sup>S is formed from [[cosmic ray]] [[spallation]] of <sup>40</sup>[[argon]] in the [[Earth's atmosphere|atmosphere]]. It has a [[half-life]] of 87 days.
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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]]s 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 also 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 also 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 δS-34 values from lakes believed to be dominated by watershed sources of sulfate.
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| − |
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| − | ===Occurrence===
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| − | [[File:NZ sulfur NI.jpg|thumb|left|Sulfur crystallites at [[Waiotapu]] [[hot springs]], [[New Zealand]]]]
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| − | 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]]. [[Sicily]] is also famous for its sulfur mines. Sulfur deposits are polycrystalline, and the largest documented single crystal measured 22x16x11 cm<sup>3</sup>.<ref>{{cite journal| url = http://www.minsocam.org/ammin/AM66/AM66_885.pdf| journal = American Mineralogist| volume = 66| page = 885-907| year= 1981| title= The largest crystals| author = P. C. Rickwood}}</ref><ref>{{cite web|url = http://giantcrystals.strahlen.org/europe/perticara.htm| title =The giant crystal project site| accessdate = 2009-06-06}}</ref>
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| − |
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| − | Significant deposits of elemental sulfur also exist in [[salt domes]] along the coast of the [[Gulf of Mexico]], and in [[evaporite]]s in eastern Europe and western Asia. The sulfur in these deposits is believed to come from the action of [[anaerobic bacteria]] on [[Mineral#Sulfate class|sulfate minerals]], especially [[gypsum]], although apparently native sulfur may be produced by geological processes alone, without the aid of living organisms (see below). However, fossil-based sulfur deposits from salt domes are the basis for commercial production in the [[United States]], [[Poland]], [[Russia]], [[Turkmenistan]], and [[Ukraine]].
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| − | [[File:AlbertaSulfurAtVancouverBC.jpg|thumb|right|Sulfur recovered from hydrocarbons in [[Alberta]], stockpiled for shipment at [[Vancouver]], [[British Columbia|B.C.]]]]
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| − | Sulfur production through [[hydrodesulfurization]] of oil, gas, and the [[Athabasca Oil Sands]] has produced a surplus — huge stockpiles of sulfur now exist throughout Alberta, Canada.
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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). It occurs naturally in volcanic emissions, such as from [[hydrothermal vent]]s, and from bacterial action on decaying sulfur-containing organic matter.
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| − | The distinctive colors of [[Jupiter]]'s [[volcano|volcanic]] moon, [[Io (moon)|Io]], are from various forms of molten, solid and gaseous sulfur. There is also a dark area near the [[Moon|Lunar]] [[Impact crater|crater]] [[Aristarchus (crater)|Aristarchus]] that may be a sulfur deposit.
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| − | Sulfur 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%. Sulfur in meteorites is normally present entirely as [[troilite]] (FeS), but other sulfides are found in some meteorites, and carbonaceous chondrites contain free sulfur, sulfates, and possibly other sulfur compounds.<ref>B. Mason, ''Meteorites'', (New York: John Wiley & Sons, 1962), p. 160.</ref>
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| − | ==Extraction and production==
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| − | ===Extraction from natural resources===
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| − | Sulfur is extracted by mainly two processes: the Sicilian process and the [[Frasch process]]. The Sicilian process, which was first used in [[Sicily]], was used in ancient times to get sulfur from rocks present in volcanic regions. In this process, the sulfur deposits are piled and stacked in brick kilns built on sloping hillsides, and with airspaces between them. Then powdered sulfur is put on top of the sulfur deposit and ignited. As the sulfur burns, the heat melts the sulfur deposits, causing the molten sulfur to flow down the sloping hillside. The molten sulfur can then be collected in wooden buckets.
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| − |
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| − | The second process used to obtain sulfur is the Frasch process. In this method, three concentric pipes are used: the outermost pipe contains superheated water, which melts the sulfur, and the innermost pipe is filled with hot compressed air, which serves to create foam and pressure. The resulting sulfur foam is then expelled through the middle pipe.<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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| − |
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| − | The Frasch process produces sulfur with a 99.5% purity content, and which needs no further purification. The sulfur produced by the Sicilian process must be purified by distillation.
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| − |
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| − | ===Production from hydrogen sulfide===
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| − | ====Chemically====
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| − | The [[Claus process]] is used to extract elemental sulfur from [[hydrogen sulfide]] produced in [[hydrodesulfurization]] of petroleum or from [[natural gas]].
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| − |
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| − | ====Biologically====
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| − | In the biological route, hydrogen sulfide (H<sub>2</sub>S) from natural gas or refinery gas is absorbed with a slight alkaline solution in a wet scrubber. Or the sulfide is produced by biological sulfate reduction. In the subsequent process step, the dissolved sulfide is biologically converted to elemental sulfur. This solid sulfur is removed from the reactor. This process has been built on commercial scale. The main advantages of this process are:
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| − | # no use of expensive chemicals,
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| − | # the process is safe as the H<sub>2</sub>S is directly absorbed in an alkaline solution,
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| − | # no production of a polluted waste stream,
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| − | # re-usable sulfur is produced, and
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| − | # the process occurs under ambient conditions.
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| − |
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| − | The biosulfur product is different from other processes in which sulfur is produced because the sulfur is hydrophillic. Next to straightforward reuses as source for sulfuric acid production, it can also be applied as sulfur fertilizer.<ref>{{cite journal
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| − | | author = Zessen, E. van, et al.
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| − | | title = Application of THIOPAQ(TM) biosulphur in agriculture
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| − | | journal = Proceedings of Sulphur 2004, Barcelona (Spain), 24 - 27 Oct.
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| − | | year = 2004
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| − | | volume = 57 - 68}}</ref>
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| − |
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| − | ==Chemistry==
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| − | {{Expand-section|date=January 2008}}
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| − | ===Inorganic compounds===
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| − | [[File:Sulfur powder.jpg|right|thumb|Sulfur powder.]]
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| − | When dissolved in water, hydrogen sulfide is acidic and will react with metals to form a series of metal sulfides. Natural metal sulfides are common, especially those of iron. Iron sulfide is called [[pyrite]], the so-called ''fool's gold''. Pyrite can show semiconductor properties.<ref>{{cite web | author = Nyle Steiner | date = 22 February 1 | title = Iron Pyrites Negative Resistance Oscillator | url = http://home.earthlink.net/~lenyr/iposc.htm | accessdate = 2007-08-15}}</ref> [[Galena]], a naturally occurring lead sulfide, was the first [[semiconductor]] discovered, and found a use as a signal [[rectifier]] in the "cat's whiskers" of early [[crystal radio]]s.
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| − |
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| − | [[Polythiazyl|Polymeric sulfur nitride]] has metallic properties even though it does not contain any [[metal]] atoms. This compound also has unusual electrical and optical properties. This polymer can be made from [[tetrasulfur tetranitride]] S<sub>4</sub>N<sub>4</sub>.
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| − |
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| − | Phosphorus sulfides are useful in synthesis. For example, P<sub>4</sub>S<sub>10</sub> and its derivatives [[Lawesson's reagent]] and [[naphthalen-1,8-diyl 1,3,2,4-dithiadiphosphetane 2,4-disulfide]] are used to replace oxygen from some organic molecules with sulfur.
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| − |
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| − | [[File:Sulfate-3D-vdW.png|thumb|right|The '''sulfate''' anion, SO<sub>4</sub><sup>2−</sup>]]
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| − | * [[Sulfide]]s (S<sup>2−</sup>), a complex family of compounds usually derived from S<sup>2−</sup>. [[Cadmium sulfide]] (CdS) is an example.
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| − | * [[Sulfites]] (SO<sub>3</sub><sup>2−</sup>), the salts of [[sulfurous acid]] (H<sub>2</sub>SO<sub>3</sub>) which is generated by dissolving SO<sub>2</sub> in water. Sulfurous acid and the corresponding sulfites are fairly strong reducing agents. Other compounds derived from SO<sub>2</sub> include the pyrosulfite or metabisulfite ion (S<sub>2</sub>O<sub>5</sub><sup>2−</sup>).
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| − | * [[Sulfate]]s (SO<sub>4</sub><sup>2−</sup>), the salts of [[sulfuric acid]]. Sulfuric acid also reacts with SO<sub>3</sub> in equimolar ratios to form [[pyrosulfuric acid]] (H<sub>2</sub>S<sub>2</sub>O<sub>7</sub>).
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| − | * [[sodium thiosulfate|Thiosulfates]] (S<sub>2</sub>O<sub>3</sub><sup>2−</sup>). Sometimes referred as thiosulfites or "hyposulfites", Thiosulfates are used in photographic fixing (HYPO) as reducing agents. Ammonium thiosulfate is being investigated as a [[cyanide]] replacement in leaching [[gold]].[http://doccopper.tripod.com/gold/AltLixiv.html]
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| − | * [[Sodium dithionite]], {{chem|Na|2|S|2|O|4}}, is the highly reducing dianion derived from hyposulfurous/dithionous acid.
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| − | * [[Sodium dithionate]] (Na<sub>2</sub>S<sub>2</sub>O<sub>6</sub>).
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| − | * [[Polythionic acid]]s (H<sub>2</sub>S<sub>''n''</sub>O<sub>6</sub>), where ''n'' can range from 3 to 80.
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| − | * [[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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| − | * [[Sodium polysulfide]]s (Na<sub>2</sub>S<sub>x</sub>)
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| − | * [[Sulfur hexafluoride]], SF<sub>6</sub>, a dense gas at ambient conditions, is used as nonreactive and nontoxic propellant
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| − | * Sulfur nitrides are chain and cyclic compounds containing only S and N. [[Tetrasulfur tetranitride]] S<sub>4</sub>N<sub>4</sub> is an example.
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| − | * [[Thiocyanate]]s contain the SCN<sup>−</sup> group. Oxidation of thiocyanoate gives [[thiocyanogen]], (SCN)<sub>2</sub> with the connectivity NCS-SCN.
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| − |
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| − | ===Organic compounds===
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| − | Many of the unpleasant odors of organic matter are based on sulfur-containing compounds such as [[methanethiol|methyl mercaptan]] and dimethyl sulfide. Thiols and sulfides are used in the odoriation of natural gas, notably, 2-methyl-2-propanethiol (t-butyl mercaptan). The odor of [[garlic]] and "[[skunk]] stink" are also caused by sulfur-containing organic compounds. Not all organic sulfur compounds smell unpleasant; for example, [[grapefruit mercaptan]], a sulfur-containing [[terpene|monoterpenoid]] is responsible for the characteristic scent of grapefruit. It should be noted that this thiol is present in very low concentrations. In larger concentrations, the odor of this compound is that typical of all thiols, unpleasant.
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| − |
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| − | Sulfur-containing organic compounds include the following (R, R', and R'' are organic groups such as CH<sub>3</sub>):
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| − | [[File:Dithiane33d.png|thumb|right|An organic sulfur compound, [[dithiane]].]]
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| − | * [[Thioether]]s have the form ''R''-S-''R′''. These compounds are the sulfur equivalents of [[ether]]s.
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| − | * [[Sulfonium]] ions have the formula RR'S-'R'", i.e. where three groups are attached to the cationic sulfur center. [[Dimethylsulfoniopropionate]] (DMSP; (CH<sub>3</sub>)<sub>2</sub>S<sup>+</sup>CH<sub>2</sub>CH<sub>2</sub>COO<sup>−</sup>) is a sulfonium ion, which is important in the marine organic [[sulfur cycle]].
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| − | * [[Thiol]]s (also known as mercaptans) have the form R-SH. These are the sulfur equivalents of [[alcohol]]s.
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| − | * [[Thiolate]]s ions have the form R-S<sup>-</sup>. Such anions arise upon treatment of [[thiol]]s with base.
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| − | * [[Sulfoxide]]s have the form ''R''-S(=O)-''R''′. The simplest sulfoxide, [[dimethyl sulfoxide|DMSO]], is a common solvent.
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| − | * [[Sulfone]]s have the form ''R''-S(=O)<sub>2</sub>-''R''′. A common sulfone is sulfolane C<sub>4</sub>H<sub>8</sub>SO<sub>2</sub>.
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| − |
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| − | ''See also [[:category:Sulfur compounds|Category: sulfur compounds]] and [[organosulfur chemistry]]''
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| − |
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| − | ==Applications==
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| − | {{Expand-section|date=January 2008}}
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| − | One of the direct uses of sulfur is in [[vulcanization]] of rubber, where [[polysulfide]]s crosslink organic polymers. Sulfur is a component of [[gunpowder]]. It 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 |authorlink= |coauthors=Vydra, Karel |editor= |encyclopedia=Ullmann's Encyclopedia of Industrial Chemistry |title=Sulfur |edition= |date= |year=2006 |month= |publisher=Wiley-VCH Verlag |volume= |location= |id= |isbn= |doi=10.1002/14356007.a25_507.pub2 |pages= |quote= }}</ref>
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| − |
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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>S[[oxygen|O]]<sub>4</sub>), which is of such prime importance to the [[world economy|world's economies]] that the production and consumption of sulfuric acid is an indicator of a nation's industrial development.<ref>[http://www.pafko.com/history/h_s_acid.html Sulfuric Acid Growth]</ref> For example, more sulfuric acid is produced in the [[United States]] every year than any other industrial chemical.{{Fact|date=August 2008}} 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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| − |
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| − | Sulfur compounds are also used in [[detergents]], [[fungicide]]s, [[dyestuff]]s, and agrichemicals. In silver-based [[photography]] sodium and ammonium [[sodium thiosulfate|thiosulfate]] are used as "fixing agents."
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| − |
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| − | Sulfur is an ingredient in some acne treatments.
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| − |
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| − | An increasing application is as fertilizer. Standard sulfur is hydrophobic and therefore has to be covered with a surfactant by bacteria in the ground before it can be oxidized to sulfate. This makes it a slow release fertilizer, which cannot be taken up by the plants instantly, but has to be oxidized to sulfate over the growth season. Sulfur also 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]</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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| − |
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| − | [[Sulfite]]s, derived from burning sulfur, are heavily used to [[Bleach (chemical)|bleach]] [[paper]]. They are also used as preservatives in dried [[fruit]].
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| − |
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| − | [[Magnesium sulfate]], better known as [[Epsom salts]], can be used as a [[laxative]], a bath additive, an [[exfoliant]], a [[magnesium]] supplement for plants, or a [[desiccant]].
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| − |
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| − | ===Specialized applications===
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| − | Sulfur is used as a light-generating medium in the rare lighting fixtures known as [[sulfur lamp]]s.
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| − | ===Historical applications===
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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 also used as a medicinal tonic and laxative. Sulfur was also used in baths for people who had fits.
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| − |
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| − | ===Fungicide and pesticide===<!--[[Wettable Sulfur]] redirs here-->
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| − | <!-- Image with unknown copyright status removed: [[File:Biosulfur as fungicide2.jpg|right|thumb|200px|Apple on which biosulfur has just been applied as a natural fungicide.]] -->
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| − |
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| − | 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|organic]]ally farmed apple production against the main disease [[apple scab]] under colder conditions. Biosulfur (biologically produced elemental sulfur with hydrophillic characteristics) can be used well for these applications.
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| − |
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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 soluble. 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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| − |
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| − | Sulfur is also 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. Some [[livestock]] owners set out a sulfur salt block as a [[salt lick]].
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| − |
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| − | ==Biological role==
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| − | {{main|Sulfur assimilation}}
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| − |
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| − | See [[sulfur cycle]] for more on the inorganic and organic natural transformations of sulfur.
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| − |
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| − | Sulfur is an essential component of all living [[cell (biology)|cells]].
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| − |
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| − | Inorganic sulfur forms a part of [[iron-sulfur cluster]]s, and sulfur is the bridging ligand in the [[copper|Cu]]<sub>A</sub> site of [[cytochrome c oxidase]], a basic substance involved in utilization of oxygen by all aerobic life.
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| − |
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| − | Sulfur may also serve as chemical food source for some primitive organisms: some forms of [[bacterium|bacteria]] use [[hydrogen sulfide]] (H<sub>2</sub>S) in the place of water as the [[electron]] donor in a primitive [[photosynthesis]]-like process in which oxygen is the electron receptor. The [[photosynthesis|photosynthetic]] green and purple sulfur [[bacteria]] and some [[chemolithotroph]]s use elemental oxygen to carry out such oxidization of hydrogen sulfide to produce elemental sulfur (S<sup>o</sup>), oxidation state = 0. Primitive bacteria which live around deep ocean volcanic vents oxidize hydrogen sulfide in this way with oxygen: see [[giant tube worm]] for an example of large organisms (via bacteria) making metabolic use of hydrogen sulfide as food to be oxidized.
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| − |
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| − | The so-called [[sulfur bacteria]], by contrast, "breathe sulfate" instead of oxygen. They use sulfur as the electron acceptor, and reduce various oxidized sulfur compounds back into sulfide, often into hydrogen sulfide. They also 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 the smell of some intestinal gases and decomposition products.
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| − |
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| − | Sulfur is a part of many bacterial defense molecules. For example, though sulfur is not a part of the [[lactam]] ring, it is a part of most [[beta lactam]] antibiotics, including the [[penicillin]]s, [[cephalosporins]], and [[monobactam]]s.
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| − |
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| − | Sulfur is absorbed by [[plant]]s via the [[root]]s from soil as the [[sulfate]] [[ion]] and reduced to sulfide before it is incorporated into [[cysteine]] and other organic sulfur compounds (see [[sulfur assimilation]] for details of this process).
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| − |
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| − | Sulfur is regarded as secondary nutrient although plant requirements for sulfur are equal to and sometimes exceed those for phosphorus. However sulfur is recognized as 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.{{Fact|date=November 2008}} Because atmospheric inputs of sulfur will continue to decrease, the deficit in the sulfur input/output is likely to increase, unless sulfur fertilizers are used.
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| − |
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| − | In [[plants]] and [[animals]] the [[amino acid]]s [[cysteine]] and [[methionine]] contain sulfur, as do all [[polypeptide]]s, [[protein]]s, and [[enzyme]]s which contain these amino acids. [[Homocysteine]] and [[taurine]] are other sulfur-containing acids which are similar in structure, but which are not coded for by [[DNA]], and are not part of the [[primary structure]] of proteins. [[Glutathione]] is an important sulfur-containing tripeptide which plays a role in cells as a source of chemical reduction potential in the cell, through its sulfhydryl (-SH) moiety. 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]].
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| − |
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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 strong covalent bonds between peptide chains give proteins a great deal of extra toughness and resiliency. 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). The high disulfide content of hair and feathers contributes to their indigestibility, and also their odor when burned.
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| − |
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| − | ===Traditional medical role for elemental sulfur===
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| − | In traditional medical skin treatment which predates modern era of scientific medicine, elemental sulfur has been used mainly as part of creams to alleviate various conditions such as psoriasis, eczema and acne. The mechanism of action is not known, although elemental sulfur does oxidize slowly to sulfurous acid, which in turn (though the action of [[sulfite]]) acts as a mild reducing and antibacterial agent.
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| − |
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| − | ==Precautions==
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| − | {{Expand-section|date=January 2008}}
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| − | Carbon disulfide, carbon oxysulfide, hydrogen sulfide, and sulfur dioxide should all be handled with care.
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| − |
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| − | Although [[sulfur dioxide]] is sufficiently safe to be used as a [[food additive]] in small amounts, at high concentrations it reacts with moisture to form [[sulfurous acid]] which in sufficient quantities may harm the [[lungs]], [[eyes]] or other [[Biological tissue|tissues]]. In organisms without lungs such as insects or plants, it otherwise prevents [[Respiration (physiology)|respiration]].
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| − |
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| − | [[Hydrogen sulfide]] is [[toxic]]. Although very pungent at first, it quickly deadens the sense of smell, so potential victims may be unaware of its presence until death or other symptoms occur.
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| − |
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| − | ===Environmental impact===
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| − | The burning of [[coal]] and/or [[petroleum]] by industry and [[power plants]] generates [[sulfur dioxide]] (S[[oxygen|O]]<sub>2</sub>), which reacts with atmospheric water and oxygen to produce [[sulfuric acid]] (H<sub>2</sub>SO<sub>4</sub>). This sulfuric acid is a component of [[acid rain]], which lowers 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 sulfur to be extracted from [[fossil fuel]]s to prevent the formation of acid rain. This extracted sulfur is then refined and represents a large portion of sulfur production. In coal fired power plants, the flue gases are sometimes purified. In more modern power plants that use [[synthesis gas|syngas]] the sulfur is extracted before the gas is burned.
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| − |
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| − | ==See also==
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| − | *[[Sulfur cycle]]
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| − | *[[Stratospheric sulfur aerosols]]
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| − | *[[Disulfide bond]]
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| − | *[[Sulfonium]] S<sup>+</sup>, S<sup>+</sup>R<sub>3</sub>
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| − | *[[Ultra-low sulfur diesel]]
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| − | *[[Claus process]]
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| − | *[[Shell-Paques sulfide removal/sulfur recovery process]]
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| | | | |
| | ==References== | | ==References== |
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| | ==External links== | | ==External links== |
| | {{Commons|Sulfur}} | | {{Commons|Sulfur}} |
| − | {{wiktionary|sulfur}}
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| | * [http://library.tedankara.k12.tr/chemistry/vol2/allotropy/z129.htm Sulfur phase diagram] | | * [http://library.tedankara.k12.tr/chemistry/vol2/allotropy/z129.htm Sulfur phase diagram] |
| | * [http://www.webelements.com/webelements/elements/text/S/index.html WebElements.com – Sulfur] | | * [http://www.webelements.com/webelements/elements/text/S/index.html WebElements.com – Sulfur] |
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| | * [http://extoxnet.orst.edu/pips/sulfur.htm Sulfur and its use as a pesticide] | | * [http://extoxnet.orst.edu/pips/sulfur.htm Sulfur and its use as a pesticide] |
| | * [http://www.sulphurinstitute.org/ The Sulphur Institute] | | * [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:Chalcogens]] | | [[Category:Chalcogens]] |
