Template:Pp-pc1 Template:About Template:Pp-move-indef Template:Use American English Template:Use dmy dates Template:Infobox sulfur Sulfur (American spelling and the preferred IUPAC name) or sulphur (Commonwealth spelling)<ref>Nature Chemistry 1, 333 (2009). doi:10.1038/nchem.301</ref> is a chemical element; it has symbol S and atomic number 16. It is abundant, multivalent and nonmetallic. Under normal conditions, sulfur atoms form cyclic octatomic molecules with the chemical formula S8. Elemental sulfur is a bright yellow, crystalline solid at room temperature.
Sulfur is the tenth most abundant element by mass in the universe and the fifth most common on Earth. Though sometimes found in pure, native form, sulfur on Earth usually occurs as 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, China, and ancient Egypt. Historically and in literature sulfur is also called brimstone,<ref name="Greenwood-1997a" /> which means "burning stone".<ref>Template:Cite EB1911</ref> Almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants from natural gas and petroleum.<ref name="Knight-2014">{{#invoke:citation/CS1|citation |CitationClass=web }}</ref><ref name="BBC_Elements-2014">{{#invoke:citation/CS1|citation |CitationClass=web }}. Download here.</ref> The greatest commercial use of the element is the production of sulfuric acid for sulfate and phosphate fertilizers, and other chemical processes. Sulfur is used in matches, insecticides, and fungicides. Many sulfur compounds are odoriferous, and the smells of odorized natural gas, skunk scent, bad breath, 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, almost always in the form of organosulfur compounds or metal sulfides. Amino acids (two proteinogenic: cysteine and methionine, and many other non-coded: cystine, taurine, etc.) and two vitamins (biotin and thiamine) are organosulfur compounds crucial for life. Many cofactors also contain sulfur, including glutathione, and iron–sulfur proteins. Disulfides, S–S bonds, confer mechanical strength and insolubility of the (among others) 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.
CharacteristicsEdit
Physical propertiesEdit
Sulfur forms several polyatomic molecules. The best-known allotrope is octasulfur, cyclo-S8. The point group of cyclo-S8 is D4d and its dipole moment is 0 D.<ref>Template:Cite journal</ref> Octasulfur is a soft, bright-yellow solid that is odorless.Template:Efn It melts at Template:Convert,Template:Efn and boils at Template:Convert.<ref name="Greenwood-1997a" /> At Template:Convert, below its melting temperature, cyclo-octasulfur begins slowly changing from α-octasulfur to the β-polymorph.<ref name="Greenwood-1997b">Template:Greenwood&Earnshaw2nd</ref> The structure of the S8 ring is virtually unchanged by this phase transition, which affects the intermolecular interactions. Cooling molten sulfur freezes at Template:Convert,<ref>Template:Cite book</ref> as it predominantly consists of the β-S8 molecules.Template:Efn 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 polymers.<ref name="Greenwood-1997b" /> At higher temperatures, the viscosity decreases as depolymerization occurs. Molten sulfur assumes a dark red color above Template:Convert. The density of sulfur is about 2 g/cm3, depending on the allotrope; all of the stable allotropes are excellent electrical insulators.
The sublimation of sulfur becomes noticeable more or less between Template:Convert and Template:Convert, and occurs readily in boiling water at Template:Convert.<ref>Template:Cite journal</ref>
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.
Chemical propertiesEdit
Under normal conditions, sulfur hydrolyzes very slowly to mainly form hydrogen sulfide and sulfuric acid:
The reaction involves adsorption of protons onto Template:Chem clusters, followed by disproportionation into the reaction products.<ref>Template:Cite journal</ref>
The second, fourth and sixth ionization energies of sulfur are 2252 kJ/mol, 4556 kJ/mol and 8495.8 kJ/mol, respectively. The composition of reaction products of sulfur with oxidants (and its oxidation state) depends on whether releasing of reaction energy overcomes these thresholds. Applying catalysts and/or supply of external energy may vary sulfur's oxidation state and the composition of reaction products. While reaction between sulfur and oxygen under normal conditions gives sulfur dioxide (oxidation state +4), formation of sulfur trioxide (oxidation state +6) requires a temperature of Template:Cvt and presence of a catalyst.
In reactions with elements of lesser electronegativity, it reacts as an oxidant and forms sulfides, where it has oxidation state −2.
Sulfur reacts with nearly all other elements except noble gases, even with the notoriously unreactive metal iridium (yielding iridium disulfide).<ref name="Munson-1968">Template:Cite journal</ref> Some of those reactions require elevated temperatures.<ref name="Wiberg-2001">Template:Cite book</ref>
AllotropesEdit
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Sulfur forms over 30 solid allotropes, more than any other element.<ref>Template:Cite book</ref> Besides S8, several other rings are known.<ref>Template:Cite book</ref> Removing one atom from the crown gives S7, which is of a deeper yellow than S8. HPLC analysis of "elemental sulfur" reveals an equilibrium mixture of mainly S8, but with S7 and small amounts of S6.<ref>Template:Cite journal</ref> Larger rings have been prepared, including S12 and S18.<ref>Template:Cite journal</ref><ref>Template:Cite journal</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 helical structure with eight atoms per turn. The long coiled polymeric molecules make the brownish substance elastic, and in bulk it has the feel of crude rubber. This form is metastable at room temperature and gradually reverts to the crystalline molecular allotrope, which is no longer elastic. This process happens over a matter of hours to days, but can be rapidly catalyzed. Template:Clear left
IsotopesEdit
{{#invoke:Labelled list hatnote|labelledList|Main article|Main articles|Main page|Main pages}} Sulfur has 23 known isotopes, four of which are stable: 32S (Template:Val), 33S (Template:Val), 34S (Template:Val), and 36S (Template:Val).<ref name="CIAAWsulfur"/><ref name="CRC-2011">Template:RubberBible92nd</ref> Other than 35S, with a half-life of 87 days, the radioactive isotopes of sulfur have half-lives less than 3 hours.
The preponderance of 32S is explained by its production in the so-called alpha process (one of the main classes of nuclear fusion reactions) in exploding stars. The other stable sulfur isotopes are produced in the bypass processes related with 34Ar, and their composition depends on a type of a stellar explosion. For example, proportionally more 33S comes from novae than from supernovae.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref>
On the planet Earth the sulfur isotopic composition was determined by the Sun. Though it was assumed that the distribution of different sulfur isotopes would be more or less equal, it has been found that proportions of the two most abundant sulfur isotopes 32S and 34S varies in different samples. Assaying of the isotope ratio (δ34S) in the samples suggests their chemical history, and with support of other methods, it allows to age-date the samples, estimate temperature of equilibrium between ore and water, determine pH and oxygen fugacity, identify the activity of sulfate-reducing bacteria in the time of formation of the sample, or suggest the main sources of sulfur in ecosystems.<ref>Template:Cite book</ref> However, there are ongoing discussions over the real reason for the δ34S shifts, biological activity or postdeposit alteration.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref>
For example, when sulfide minerals are precipitated, isotopic equilibration among solids and liquid may cause small differences in the δ34S values of co-genetic minerals. The differences between minerals can be used to estimate the temperature of equilibration. The δ13C and δ34S of coexisting carbonate minerals and sulfides can be used to determine the pH and oxygen fugacity of the ore-bearing fluid during ore formation.
Scientists measure the sulfur isotopes of minerals in rocks and sediments to study the redox conditions in past oceans. Sulfate-reducing bacteria in marine sediment fractionate sulfur isotopes as they take in sulfate and produce sulfide. Prior to the 2010s, it was thought that sulfate reduction could fractionate sulfur isotopes up to 46 permil<ref>Template:Cite journal</ref> and fractionation larger than 46 permil recorded in sediments must be due to disproportionation of sulfur compounds in the sediment. This view has changed since the 2010s as experiments showed that sulfate-reducing bacteria can fractionate to 66 permil.<ref>Template:Cite journal</ref> As substrates for disproportionation are limited by the product of sulfate reduction, the isotopic effect of disproportionation should be less than 16 permil in most sedimentary settings.<ref>Template:Cite journal</ref>
In 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 hydrologic studies. Differences in the natural abundances 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 measurably different 34S values than lakes believed to be dominated by watershed sources of sulfate.
The radioactive 35S is formed in cosmic ray spallation of the atmospheric 40Ar. This fact may be used to verify the presence of recent (up to 1 year) atmospheric sediments in various materials. This isotope may be obtained artificially by different ways. In practice, the reaction 35Cl + n → 35S + p is used by irradiating potassium chloride with neutrons.<ref>Template:Cite conference</ref> The isotope 35S is used in various sulfur-containing compounds as a radioactive tracer for many biological studies, for example, the Hershey-Chase experiment.
Because of the weak beta activity of 35S, its compounds are relatively safe as long as they are not ingested or absorbed by the body.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref>
Natural occurrenceEdit
32S is created inside massive stars, at a depth where the temperature exceeds 2.5×109 K, by the fusion of one nucleus of silicon plus one nucleus of helium.<ref>Template:Cite journal</ref> As this nuclear reaction 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 meteorites. 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>Template:Cite book</ref> The distinctive colors of Jupiter's volcanic moon Io are attributed to various forms of molten, solid, and gaseous sulfur.<ref>Template:Cite journal</ref> In July 2024, elemental sulfur was accidentally discovered to exist on Mars after the Curiosity rover drove over and crushed a rock, revealing sulfur crystals inside it.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref>
Sulfur is the fifth most common element by mass in the Earth. Elemental sulfur can be found near hot springs and volcanic regions in many parts of the world, especially along the Pacific Ring of Fire; such volcanic deposits are mined in Indonesia, Chile, and Japan. These deposits are polycrystalline, with the largest documented single crystal measuring Template:Cvt.<ref>Template:Cite journal</ref> Historically, Sicily was a major source of sulfur in the Industrial Revolution.<ref>Template:Cite book</ref> Lakes of molten sulfur up to about Template:Convert in diameter have been found on the sea floor, associated with submarine volcanoes, at depths where the boiling point of water is higher than the melting point of sulfur.<ref name="de Ronde-2015">Template:Cite book</ref>
Native sulfur is synthesized by anaerobic bacteria acting on sulfate minerals such as gypsum in salt domes.<ref>Template:Cite book</ref><ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref> Significant deposits in salt domes occur along the coast of the Gulf of Mexico, and in evaporites in eastern Europe and western Asia. Native sulfur may be produced by geological processes alone. Fossil-based sulfur deposits from salt domes were once the basis for commercial production in the United States, Russia, Turkmenistan, and Ukraine.<ref name="Nehb-2006" /> Such sources have become of secondary commercial importance, and most are no longer worked but commercial production is still carried out in the Osiek mine in Poland.
Common naturally occurring sulfur compounds include the sulfide minerals, such as pyrite (iron sulfide), cinnabar (mercury sulfide), galena (lead sulfide), sphalerite (zinc sulfide), and stibnite (antimony sulfide); and the sulfate minerals, 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 vents.
The main industrial source of sulfur has become petroleum and natural gas.<ref name="Knight-2014"/>
CompoundsEdit
Template:Main article Common oxidation states of sulfur range from −2 to +6. Sulfur forms stable compounds with all elements except the noble gases.
Electron transfer reactionsEdit
Sulfur polycations, Template:Chem2, Template:Chem2 and Template:Chem2 are produced when sulfur is reacted with oxidizing agents in a strongly acidic solution.<ref>Shriver, Atkins. Inorganic Chemistry, Fifth Edition. W. H. Freeman and Company, New York, 2010; pp 416</ref> The colored solutions produced by dissolving sulfur in oleum were first reported as early as 1804 by C. F. Bucholz, but the cause of the color and the structure of the polycations involved was only determined in the late 1960s. Template:Chem2 is deep blue, Template:Chem2 is yellow and Template:Chem2 is red.<ref name="Greenwood-1997b" />
Reduction of sulfur gives various polysulfides with the formula Template:Chem, many of which have been obtained in crystalline form. Illustrative is the production of sodium tetrasulfide:
Some of these dianions dissociate to give radical anions. For instance, Template:Chem2 gives the blue color of the rock lapis lazuli.
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This reaction highlights a distinctive property of sulfur: its ability to catenate (bind to itself by formation of chains). Protonation of these polysulfide anions produces the polysulfanes, H2Sx, where x = 2, 3, and 4.<ref>Template:Cite book</ref> Ultimately, reduction of sulfur produces sulfide salts:
The interconversion of these species is exploited in the sodium–sulfur battery.
HydrogenationEdit
Treatment of sulfur with hydrogen gives hydrogen sulfide. When dissolved in water, hydrogen sulfide is mildly acidic:<ref name="Greenwood-1997a">Template:Cite book</ref>
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 cytochromes in a manner analogous to cyanide and azide (see below, under precautions).
CombustionEdit
The two principal sulfur oxides are obtained by burning sulfur:
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Many other sulfur oxides are observed including the sulfur-rich oxides include sulfur monoxide, disulfur monoxide, disulfur dioxides, and higher oxides containing peroxo groups.
HalogenationEdit
Sulfur reacts with fluorine to give the highly reactive sulfur tetrafluoride and the highly inert sulfur hexafluoride.<ref>Template:OrgSynth</ref> Whereas fluorine gives S(IV) and S(VI) compounds, chlorine gives S(II) and S(I) derivatives. Thus, sulfur dichloride, disulfur dichloride, and higher chlorosulfanes arise from the chlorination of sulfur. Sulfuryl chloride and chlorosulfuric acid are derivatives of sulfuric acid; thionyl chloride (SOCl2) is a common reagent in organic synthesis.<ref name="Lauss">Template:Ullmann</ref> Bromine also oxidizes sulfur to form sulfur dibromide and disulfur dibromide.<ref name="Lauss" />
PseudohalidesEdit
Sulfur oxidizes cyanide and sulfite to give thiocyanate and thiosulfate, respectively.
Metal sulfidesEdit
Sulfur reacts with many metals. Electropositive metals give polysulfide salts. Copper, zinc, and silver are attacked by sulfur; see tarnishing. Although many metal sulfides are known, most are prepared by high temperature reactions of the elements.<ref>Template:Cite book</ref> Geoscientists also study the isotopes of metal sulfides in rocks and sediment to study environmental conditions in the Earth's past.<ref>Template:Cite journal</ref>
Organic compoundsEdit
{{#invoke:Labelled list hatnote|labelledList|Main article|Main articles|Main page|Main pages}}
- Illustrative organosulfur compounds
- L-Cystein - L-Cysteine.svg
(L)-cysteine, an amino acid containing a thiol group
- Methionin - Methionine.svg
Methionine, an amino acid containing a thioether
- Thiamin.svg
Thiamine or vitamin B1
- Biotin structure.svg
Biotin or vitamin B7
- Penicillin core.svg
Penicillin, an antibiotic ("R" is the variable group)
- Allicin skeletal.svg
Allicin, a chemical compound in garlic
- Diphenyl disulfide.svg
Diphenyl disulfide, a representative disulfide
- Dibenzothiophen - Dibenzothiophene.svg
Dibenzothiophene, a component of crude oil
- Perfluorooctanesulfonic acid structure.svg
Perfluorooctanesulfonic acid (PFOS), a surfactant
Some of the main classes of sulfur-containing organic compounds include the following:<ref name="Cremlyn-1996">Template:Cite book</ref>
- Thiols or mercaptans (so called because they capture mercury as chelators) are the sulfur analogs of alcohols; treatment of thiols with base gives thiolate ions.
- Thioethers are the sulfur analogs of ethers.
- Sulfonium ions have three groups attached to a cationic sulfur center. Dimethylsulfoniopropionate (DMSP) is one such compound, important in the marine organic sulfur cycle.
- Sulfoxides and sulfones 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 acids are used in many detergents.
Compounds with carbon–sulfur multiple bonds are uncommon, an exception being carbon disulfide, a volatile colorless liquid that is structurally similar to carbon dioxide. It is used as a reagent to make the polymer rayon and many organosulfur compounds.<ref>Template:Cite journal</ref> Unlike carbon monoxide, carbon monosulfide is stable only as an extremely dilute gas, found between solar systems.<ref>Template:Cite journal</ref>
Organosulfur compounds are responsible for some of the unpleasant odors of decaying organic matter. They are widely known as the odorant in domestic natural gas, garlic odor, and skunk spray, as well as a component of bad breath odor. Not all organic sulfur compounds smell unpleasant at all concentrations: the sulfur-containing 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 vesicant, was used in World War I as a disabling agent.<ref>Template:Cite book</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,Template:Citation needed 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.
HistoryEdit
AntiquityEdit
Being abundantly available in native form, sulfur was known in ancient times and is referred to in the Torah.<ref>Template:Bibleverse, Template:Bibleverse</ref> English translations of the Christian Bible commonly referred to burning sulfur as "brimstone", giving rise to the term "fire-and-brimstone" sermons, 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<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref> 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>Template:Cite book</ref> this is mentioned in the Odyssey.<ref>Odyssey, book 22, lines 480–495. www.perseus.tufts.edu. Retrieved on 16 August 2012.</ref> Pliny the Elder discusses sulfur in book 35 of his Natural History, saying that its best-known source is the island of Melos. He mentions its use for fumigation, medicine, and bleaching cloth.<ref>Pliny the Elder on science and technology, John F. Healy, Oxford University Press, 1999, Template:ISBN, pp. 247–249.</ref>
A natural form of sulfur known as Template:Transliteration ({{#invoke:Lang|lang}}) was known in China since the 6th century BC and found in Hanzhong.<ref name="Zhang-1986">Template:Cite journal</ref> By the 3rd century, the Chinese had discovered that sulfur could be extracted from pyrite.<ref name="Zhang-1986" /> 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="Zhang-1986" /> The Wujing Zongyao of 1044 AD described various formulas for Chinese black powder, which is a mixture of potassium nitrate (Template:Chem), charcoal, and sulfur.<ref>Template:Cite book</ref>
Template:Multiple image Indian alchemists, practitioners of the "science of chemicals" (Template:Langx), wrote extensively about the use of sulfur in alchemical operations with mercury, from the eighth century AD onwards.<ref name="White-1996">Template:Cite book</ref> In the [[rasa shastra|Template:Transliteration]] tradition, sulfur is called "the smelly" ({{#invoke:Lang|lang}}, Template:Transliteration).
Early European alchemists gave sulfur a unique alchemical symbol, a triangle atop a cross (🜍). (This is sometimes confused with the astronomical crossed-spear symbol ⚴ for 2 Pallas.) The variation known as brimstone has a symbol combining a two-barred cross atop a lemniscate (🜏). 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>Template:Cite journal</ref><ref>Template:Cite journal</ref><ref>Template:Cite journal</ref>
Modern timesEdit
Sulfur appears in a column of fixed (non-acidic) alkali in a chemical table of 1718.<ref>Template:Cite book</ref> Antoine Lavoisier used sulfur in combustion experiments, writing of some of these in 1777.<ref>Template:Cite book</ref>
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 Marseille, France, for the production of sulfuric acid for use in the Leblanc process. In industrializing Britain, with the repeal of tariffs on salt in 1824, demand for sulfur from Sicily surged. The increasing British control and exploitation of the mining, refining, and transportation of sulfur, coupled with the failure of this lucrative export to transform Sicily's backward and impoverished economy, led to the Sulfur Crisis of 1840, when King Ferdinand II gave a monopoly of the sulfur industry to a French firm, violating an earlier 1816 trade agreement with Britain. A peaceful solution was eventually negotiated by France.<ref>Template:Cite book</ref><ref>Template:Cite journal</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="Botsch-2001">Template:Cite journal</ref>
In the late 18th century, furniture makers used molten sulfur to produce decorative inlays.<ref name="Mass-2003">Template:Cite journal</ref> Molten sulfur is sometimes still used for setting steel bolts into drilled concrete holes where high shock resistance is desired for floor-mounted equipment attachment points. Pure powdered sulfur was used as a medicinal tonic and laxative.<ref name="Nehb-2006" />
Since the advent of the contact process, the majority of sulfur is used to make sulfuric acid for a wide range of uses, particularly fertilizer.<ref>Template:Cite book</ref>
In recent times, the main source of sulfur has become petroleum and natural gas. This is due to the requirement to remove sulfur from fuels in order to prevent acid rain, and has resulted in a surplus of sulfur.<ref name="Knight-2014"/>
Spelling and etymologyEdit
Sulfur is derived from the Latin word {{#invoke:Lang|lang}}, which was Hellenized to {{#invoke:Lang|lang}} in the erroneous belief that the Latin word came from Greek. This spelling was later reinterpreted as representing an /f/ sound and resulted in the spelling {{#invoke:Lang|lang}}, which appears in Latin toward the end of the Classical period. The true Ancient Greek word for sulfur, {{#invoke:Lang|lang}}, theîon (from earlier {{#invoke:Lang|lang}}, théeion), is the source of the international chemical prefix thio-. The Modern Standard Greek word for sulfur is θείο, theío.
In 12th-century Anglo-French, it was {{#invoke:Lang|lang}}. In the 14th century, the erroneously Hellenized Latin {{#invoke:Lang|lang}} was restored in Middle English {{#invoke:Lang|lang}}. By the 15th century, both full Latin spelling variants sulfur and sulphur became common in English. The parallel f~ph spellings continued in Britain until the 19th century, when the word was standardized as sulphur.<ref name="OED">Template:Cite OED</ref> On the other hand, sulfur was the form eventually chosen in the United States, though multiple place names (such as White Sulphur Springs) use -ph-. Canada uses both spellings.
IUPAC adopted the spelling sulfur in 1990<ref>Template:Cite journal</ref><ref>Template:Cite book</ref> as did the Nomenclature Committee of the Royal Society of Chemistry in 1992, restoring the spelling sulfur to Britain.<ref>Template:Cite journal</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>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref>
ProductionEdit
Sulfur may be found by itself and historically was usually obtained in this form; pyrite has also been a source of sulfur.<ref>Template:Cite book</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 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 sulfur mine in Sicily is about the nearest thing to hell that I expect to see in this life."<ref>Template:Cite book</ref> Sulfur is still mined from surface deposits in poorer nations with volcanoes, such as Indonesia, and problems with working conditions still exist.<ref>Template:Cite journal</ref>
Elemental sulfur was extracted from salt domes (where it sometimes occurs in nearly pure form) until the late 20th century, when it became a side product of other industrial processes such as in oil refining, in which sulfur is undesirable. As a mineral, native sulfur under salt domes is thought to be a fossil mineral resource, produced by the action of anaerobic bacteria on sulfate deposits. It was removed from such salt-dome mines mainly by the Frasch process.<ref name="Nehb-2006" /> 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 had significant use anywhere in the world since 2002.<ref name="Eow-2002">Template:Cite journal</ref><ref name="Schreiner-2008">Template:Cite journal</ref>
Since then, sulfur has typically been produced from petroleum, natural gas, and related fossil resources, from which it is obtained mainly as hydrogen sulfide.<ref name="Knight-2014"/> Organosulfur compounds, undesirable impurities in petroleum, may be upgraded by subjecting them to hydrodesulfurization, which cleaves the C–S bonds:<ref name="Eow-2002" /><ref name="Schreiner-2008" />
The resulting hydrogen sulfide from this process, and also as it occurs in natural gas, is converted into elemental sulfur by the Claus process, which entails oxidation of some hydrogen sulfide to sulfur dioxide and then the comproportionation of the two:<ref name="Eow-2002" /><ref name="Schreiner-2008" />
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Due to the high sulfur content of the Athabasca Oil Sands, stockpiles of elemental sulfur from this process exist throughout Alberta, Canada.<ref name="Hyndman-1982">Template:Cite book</ref> Another way of storing sulfur is as a binder for concrete, the resulting product having some desirable properties (see sulfur concrete).<ref name="Mohamed-2010">Template:Cite book</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), the United States (8.8), Canada (7.1) and Russia (7.1).<ref>Apodaca, Lori E. (2012) "Sulfur". Mineral Commodity Summaries. United States Geological Survey.</ref> Production has been slowly increasing from 1900 to 2010; the price was unstable in the 1980s and around 2010.<ref name="Apodaca-2010" />
ApplicationsEdit
Sulfuric acidEdit
Elemental sulfur is used mainly as a precursor to other chemicals. Approximately 85% (1989) is converted to sulfuric acid (H2SO4):
In 2010, the United States produced more sulfuric acid than any other inorganic industrial chemical.<ref name="Apodaca-2010">{{#invoke:citation/CS1|citation |CitationClass=web }}</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-2006" />
Other important sulfur chemistryEdit
Sulfur reacts directly with methane to give carbon disulfide, which is used to manufacture cellophane and rayon.<ref name="Nehb-2006">Template:Cite book</ref> One of the uses of elemental sulfur is in vulcanization of rubber, where polysulfide chains crosslink organic polymers. Large quantities of sulfites are used to bleach paper and to preserve dried fruit. Many surfactants and detergents (e.g. sodium lauryl sulfate) are sulfate derivatives. Calcium sulfate, gypsum (CaSO4·2H2O) is mined on the scale of 100 million tonnes each year for use in Portland cement and fertilizers.
When silver-based photography was widespread, sodium and ammonium thiosulfate were widely used as "fixing agents". Sulfur is a component of gunpowder ("black powder").
FertilizerEdit
Amino acids synthesized by living organisms such as methionine and cysteine contain organosulfur groups (thioester and thiol respectively). The antioxidant glutathione protecting many living organisms against free radicals and oxidative stress also contains organic sulfur. Some crops such as onion and garlic also produce different organosulfur compounds such as syn-propanethial-S-oxide responsible of lacrymal irritation (onions), or diallyl disulfide and allicin (garlic). Sulfates, commonly found in soils and groundwaters are often a sufficient natural source of sulfur for plants and bacteria. Atmospheric deposition of sulfur dioxide (SO2) is also a common artificial source (coal combustion) of sulfur for the soils. Under normal circumstances, in most agricultural soils, sulfur is not a limiting nutrient for plants and microorganisms (see Liebig's barrel). However, in some circumstances, soils can be depleted in sulfate, e.g. if this later is leached by meteoric water (rain) or if the requirements in sulfur for some types of crops are high. This explains that sulfur is increasingly recognized and used as a component of fertilizers. The most important form of sulfur for fertilizer is calcium sulfate, commonly found in nature as the mineral gypsum (CaSO4·2H2O). Elemental sulfur is hydrophobic (not soluble in water) and cannot be used directly by plants. Elemental sulfur (ES) is sometimes mixed with bentonite to amend depleted soils for crops with high requirement in organo-sulfur. Over time, oxidation abiotic processes with atmospheric oxygen and soil bacteria can oxidize and convert elemental sulfur to soluble derivatives, which can then be used by microorganisms and plants. Sulfur improves the efficiency of other essential plant nutrients, particularly nitrogen and phosphorus.<ref name="sulphurinstitute-2020">{{#invoke:citation/CS1|citation |CitationClass=web }}</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 by plants.
The plants requirement for sulfur equals or exceeds the requirement for phosphorus. 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>Template:Cite journal</ref><ref>Template:Cite journal</ref><ref>Template:Cite journal</ref> Because atmospheric inputs of sulfur continue to decrease, the deficit in the sulfur input/output is likely to increase unless sulfur fertilizers are used. Atmospheric inputs of sulfur decrease because of actions taken to limit acid rains.<ref name="NASA-2011">Template:Citation</ref><ref name="sulphurinstitute-2020" />
Fungicide and pesticideEdit
Elemental sulfur is one of the oldest fungicides and pesticides. "Dusting sulfur", elemental sulfur in powdered form, is a common fungicide for grapes, strawberry, many vegetables and several other crops. It has a good efficacy against a wide range of powdery mildew diseases as well as black spot. In organic production, sulfur is the most important fungicide. It is the only fungicide used in organically farmed apple production against the main disease 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 miscible.<ref name="Mohamed-2010" /><ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</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" (i.e., "green") insecticide (actually an acaricide) against ticks 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 (fungus), mange, and other dermatoses and parasites.
Sulfur candles of almost pure sulfur were burned to fumigate structures and wine barrels, but are now considered too toxic for residences.
PharmaceuticalsEdit
{{#invoke:Labelled list hatnote|labelledList|Main article|Main articles|Main page|Main pages}} 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="Springer-1978">Template:Cite book</ref> Precipitated sulfur and colloidal sulfur are used, in form of lotions, creams, powders, soaps, and bath additives, for the treatment of acne vulgaris, acne rosacea, and seborrhoeic dermatitis.<ref name="Arzneibuch-Kommentar">Template:Cite book</ref>
Many drugs contain sulfur.<ref>Template:Citation</ref> Early examples include antibacterial sulfonamides, known as sulfa drugs. A more recent example is mucolytic acetylcysteine. Sulfur is a part of many bacterial defense molecules. Most β-lactam antibiotics, including the penicillins, cephalosporins and monobactams contain sulfur.<ref name="Cremlyn-1996" />
BatteriesEdit
Due to their high energy density and the availability of sulfur, there is ongoing research in creating rechargeable lithium–sulfur batteries. Until now, carbonate electrolytes have caused failures in such batteries after a single cycle. In February 2022, researchers at Drexel University have not only created a prototypical battery that lasted 4000 recharge cycles, but also found the first monoclinic gamma sulfur that remained stable below 95 degrees Celsius.<ref>Template:Cite journal</ref>
Biological roleEdit
Sulfur is an essential component of all living cells. It is the eighth most abundant element in the human body by weight,<ref name="sulphurinstitute-2021">Template:Cite news</ref> about equal in abundance to potassium, and slightly greater than sodium and chlorine.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref> A Template:Convert human body contains about Template:Convert of sulfur.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref> The main dietary source of sulfur for humans is sulfur-containing amino acids,<ref name="Parcell-2002">Template:Cite journal</ref> which can be found in plant and animal proteins.<ref name="Ingenbleek-2013">Template:Cite journal</ref>
Transferring sulfur between inorganic and biomoleculesEdit
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).<ref>Template:Cite journal</ref> Another contributor, who continued to study it was Selman Waksman.<ref>Template:Cite journal</ref> Primitive bacteria that live around deep ocean volcanic vents oxidize hydrogen sulfide for their nutrition, as discovered by Robert Ballard.<ref name="BBC_Elements-2014" />
Sulfur oxidizers can use as energy sources reduced sulfur compounds, including hydrogen sulfide, elemental sulfur, sulfite, thiosulfate, and various polythionates (e.g., tetrathionate).<ref>Template:Cite journal</ref> They depend on enzymes such as sulfur oxygenase and sulfite oxidase to oxidize sulfur to sulfate. Some lithotrophs 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 uses oxygen as the electron acceptor. Sulfur-based chemosynthesis may be simplifiedly compared with photosynthesis:
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There are bacteria combining these two ways of nutrition: green sulfur bacteria and purple sulfur bacteria.<ref>Template:Citation</ref> Also sulfur-oxidizing bacteria can go into symbiosis with larger organisms, enabling the later to use hydrogen sulfide as food to be oxidized. Example: the giant tube worm.<ref>Template:Cite journal</ref>
There are sulfate-reducing bacteria, that, by contrast, "breathe sulfate" instead of oxygen. They use organic compounds or molecular hydrogen as the energy source. They use sulfur as the electron acceptor, and reduce various oxidized sulfur compounds back into sulfide, often into hydrogen sulfide. They can grow on other partially oxidized sulfur compounds (e.g. thiosulfates, thionates, polysulfides, sulfites).
There are studies pointing that many deposits of native sulfur in places that were the bottom of the ancient oceans have biological origin.<ref>Template:Cite journal</ref><ref>Template:Cite journal</ref><ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref> These studies indicate that this native sulfur have been obtained through biological activity, but what is responsible for that (sulfur-oxidizing bacteria or sulfate-reducing bacteria) is still unknown for sure.
Sulfur is absorbed by plants roots 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-1996">Template:Cite book</ref>
While the plants' role in transferring sulfur to animals by food chains is more or less understood, the role of sulfur bacteria is just getting investigated.<ref>Template:Cite journal</ref><ref>Template:Cite journal</ref>
Protein and organic metabolitesEdit
In all forms of life, most of the sulfur is contained in two proteinogenic amino acids (cysteine and methionine), thus the element is present in all proteins that contain these amino acids.<ref>Template:Cite journal</ref> Some of the sulfur is present in certain metabolites—many of which are cofactors—and sulfated polysaccharides of connective tissue (chondroitin sulfates, heparin).
The functionality of a given protein is heavily dependent on its structure. Proteins reach this structure through the process of protein folding, which is facilitated by a variety of intra- and inter-molecular bonds. While much of the folding is driven by the formation of hydrogen bonds, covalent bonding of cysteine residues into disulfide bridges imposes constraints that stabilize particular conformations while preventing others from forming. As the bond energy of a covalent disulfide bridge is higher than the energy of a coordinate bond or hydrophobic interaction, greater numbers of disulfide bridges lead to higher energies required for protein denaturation. Disulfide bonds often serve to stabilize protein structures in the more oxidizing conditions of the extracellular environment.<ref>Template:Cite book</ref> Within the cytoplasm, disulfide bonds may instead be reduced (i.e. in -SH form) to their constituent cysteine residues by thioredoxins.<ref>Template:Cite journal</ref>
Many important cellular enzymes use prosthetic groups ending with sulfhydryl (-SH) moieties to handle reactions involving acyl-containing biochemicals: two common examples from basic metabolism are coenzyme A and alpha-lipoic acid.<ref name="Nelson-2000">Template:Cite book</ref> Cysteine-related metabolites homocysteine and taurine are other sulfur-containing amino acids that are similar in structure, but not coded by DNA, and are not part of the primary structure of proteins, take part in various locations of mammalian physiology.<ref>Template:Cite journal</ref><ref>Template:Cite journal</ref> Two of the 13 classical vitamins, biotin and thiamine, contain sulfur, and serve as cofactors to several enzymes.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref><ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref> 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.
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, Template:Chem2, the immediate precursor to methane.<ref>Template:Cite journal</ref>
Metalloproteins and inorganic cofactorsEdit
Metalloproteins—in which the active site is a transition metal ion (or metal-sulfide cluster) often coordinated by sulfur atoms of cysteine residues<ref>Template:Cite journal</ref>—are essential components of enzymes involved in electron transfer processes. Examples include plastocyanin (Cu2+) and nitrous oxide reductase (Cu–S). The function of these enzymes is dependent on the fact that the transition metal ion can undergo redox reactions. Other examples include many zinc proteins,<ref>Template:Cite journal</ref> as well as iron–sulfur clusters. Most pervasive are the ferrodoxins, which serve as electron shuttles in cells. In bacteria, the important nitrogenase enzymes contain 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>Template:Cite book</ref>
Sulfur is also present in molybdenum cofactor.<ref>Template:Cite journal</ref>
- File:FdRedox.pngEasiness of electron flow in a cluster provides catalytic effect of a respective enzyme.
SulfateEdit
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DeficiencyEdit
In humans methionine is an essential amino acid; cysteine is conditionally essential and may be synthesized from non-essential serine via sulfur salvaged from methionine. Sulfur deficiency is uncommon due to the ubiquity of cysteine and methionine in food.Template:Cn
Isolated sulfite oxidase deficiency is a rare, fatal genetic disease caused by mutations to sulfite oxidase, which is needed to metabolize sulfites to sulfates.<ref>Template:Cite journal</ref>
PrecautionsEdit
Though elemental sulfur is only minimally absorbed through the skin and is of low toxicity to humans, inhalation of sulfur dust or contact with eyes or skin may cause irritation. Excessive ingestion of sulfur can cause a burning sensation or diarrhea,<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref> and cases of life-threatening metabolic acidosis have been reported after patients deliberately consumed sulfur as a folk remedy.<ref>Template:Cite journal</ref><ref>Template:Cite journal</ref>
Toxicity of sulfur compoundsEdit
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 lungs, eyes, or other tissues.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref> In organisms without lungs such as insects, sulfite in high concentration prevents respiration.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref>
Sulfur trioxide (made by catalysis from sulfur dioxide) and sulfuric acid are similarly highly acidic and corrosive in the presence of water. Concentrated sulfuric acid is a strong dehydrating agent that can strip available water molecules and water components from sugar and organic tissue.<ref>Template:Cite magazine</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 sulfurous acid (H2SO3).<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref> These acids are components of acid rain, lowering the pH of soil and freshwater bodies, sometimes resulting in substantial damage to the environment and chemical weathering of statues and structures. Fuel standards increasingly require that fuel producers extract sulfur from fossil fuels to prevent acid rain formation. This extracted and refined sulfur represents a large portion of sulfur production. In coal-fired power plants, flue gases are sometimes purified. More modern power plants that use synthesis gas extract the sulfur before they burn the gas.
Hydrogen sulfide is about one-half as toxic as hydrogen cyanide, and intoxicates by the same mechanism (inhibition of the respiratory enzyme cytochrome oxidase),<ref>Template:Cite journal</ref> though hydrogen sulfide is less likely to cause sudden poisonings from small inhaled amounts (near its permissible exposure limit (PEL) of 20 ppm) because of its disagreeable odor.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref> However, its presence in ambient air at concentration over 100–150 ppm quickly deadens the sense of smell,<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref> and a victim may breathe increasing quantities without noticing until severe symptoms cause death. Dissolved sulfide and hydrosulfide salts are toxic by the same mechanism.
NotesEdit
See alsoEdit
- Blue lava
- Stratospheric sulfur aerosols
- Sulfur assimilation
- Sulfur isotope biogeochemistry
- Ultra-low sulfur diesel
ReferencesEdit
Further readingEdit
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External linksEdit
- Sulfur at The Periodic Table of Videos (University of Nottingham)
- Atomic Data for Sulfur, NIST Physical Measurement Laboratory
- Sulfur phase diagram Template:Webarchive, Introduction to Chemistry for Ages 13–17
- Crystalline, liquid and polymerization of sulfur on Vulcano Island, Italy
- Sulfur and its use as a pesticide
- The Sulphur Institute
- Nutrient Stewardship and The Sulphur Institute
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