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This article is about the chemical element. For other uses, seeSulfur (disambiguation).

Sulfur,16S
Sulfur
Alternative nameSulphur (British spelling)
AllotropesseeAllotropes of sulfur
AppearanceLemon yellowsinteredmicrocrystals
Standard atomic weightAr°(S)
Sulfur in theperiodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
O
↑
S
↓
Se
phosphorus←sulfur→chlorine
Atomic number(Z)16
Groupgroup 16 (chalcogens)
Periodperiod 3
Blockp-block
Electron configuration[Ne] 3s23p4
Electrons per shell2, 8, 6
Physical properties
PhaseatSTPsolid
Melting pointalpha (α-S8): 388.36K​(115.21 °C, ​239.38 °F)
Boiling point717.8 K ​(444.6 °C, ​832.3 °F)
Density(nearr.t.)alpha (α-S8): 2.07 g/cm3
beta (β-S8): 1.96 g/cm3
gamma (γ-S8): 1.92 g/cm3
when liquid (atm.p.)1.819 g/cm3
Critical point1314 K, 20.7 MPa
Heat of fusionbeta (β-S8): 1.727kJ/mol
Heat of vaporizationbeta (β-S8): 45 kJ/mol
Molar heat capacity22.75 J/(mol·K)
Vapor pressure
P(Pa) 1 10 100 1 k 10 k 100 k
atT(K) 375 408 449 508 591 717
Atomic properties
Oxidation states−2,−1,0,+1,+2,+3,+4,+5,+6(a stronglyacidicoxide)
ElectronegativityPauling scale: 2.58
Ionization energies
  • 1st: 999.6 kJ/mol
  • 2nd: 2252 kJ/mol
  • 3rd: 3357 kJ/mol
  • (more)
Covalent radius105±3pm
Van der Waals radius180 pm
Color lines in a spectral range
Spectral linesof sulfur
Other properties
Natural occurrenceprimordial
Crystal structurealpha (α-S8): ​orthorhombic(oF128)
Lattice constants
Orthorhombic crystal structure for alpha (α-S8): sulfur
a= 1.0460 nm
b= 1.2861 nm
c= 2.4481 nm (at 20 °C)[3]
Crystal structurebeta (β-S8): ​monoclinic(mP48)
Lattice constants
Monoclinic crystal structure for beta (β-S8): sulfur
a= 1.0923 nm
b= 1.0851 nm
c= 1.0787 nm
β = 95.905° (at 20 °C)[3]
Thermal conductivity0.205 W/(m⋅K) (amorphous)
Electrical resistivity2×1015Ω⋅m (at 20 °C) (amorphous)
Magnetic orderingdiamagnetic[4]
Molar magnetic susceptibilityalpha (α-S8):−15.5×10−6cm3/mol (298 K)[5]
Bulk modulus7.7 GPa
Mohs hardness2.0
CAS Number7704-34-9
History
Discoverybefore 2000 BCE[6]
Recognized as anelementbyAntoine Lavoisier(1777)
Isotopes of sulfur
Main isotopes Decay
abun­dance half-life(t1/2) mode pro­duct
32S 94.8% stable
33S 0.760% stable
34S 4.37% stable
35S trace 87.37 d β− 35Cl
36S 0.02% stable
34S abundances vary greatly (between 3.96 and 4.77 percent) in natural samples.
Category: Sulfur
|references

Sulfur(also spelledsulphurinBritish English) is achemical element;it hassymbolSandatomic number16. It isabundant,multivalentandnonmetallic.Undernormal conditions,sulfur atoms formcyclic octatomic moleculeswith the chemical formulaS8.Elemental sulfur is a bright yellow,crystallinesolid atroom 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,nativeform, sulfur on Earth usually occurs assulfideandsulfate minerals.Being abundant in native form, sulfur was known in ancient times, being mentioned for its uses inancient India,ancient Greece,China,andancient Egypt.Historically and in literature sulfur is also calledbrimstone,[7]which means "burning stone".[8]Today, almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants fromnatural gasandpetroleum.[9][10]The greatest commercial use of the element is the production ofsulfuric acidfor sulfate and phosphatefertilizers,and other chemical processes. Sulfur is used inmatches,insecticides,andfungicides.Many sulfur compounds are odoriferous, and the smells of odorized natural gas,skunkscent,bad breath,grapefruit,andgarlicare due toorganosulfurcompounds.Hydrogen sulfidegives the characteristic odor to rotting eggs and other biological processes.

Sulfur is anessential elementfor all life, almost always in the form oforganosulfur compoundsor metal sulfides.Amino acids(twoproteinogenic:cysteineandmethionine,and many othernon-coded:cystine,taurine,etc.) and two vitamins (biotinandthiamine) are organosulfur compounds crucial for life. Manycofactorsalso contain sulfur, includingglutathione,andiron–sulfur proteins.Disulfides,S–S bonds, confer mechanical strength and insolubility of the (among others) proteinkeratin,found in outer skin, hair, and feathers. Sulfur is one of the core chemical elements needed forbiochemicalfunctioning and is an elementalmacronutrientfor all living organisms.

Characteristics

[edit]

Physical properties

[edit]
As a solid, sulfur is a characteristic lemon yellow; when burned, sulfur melts into a blood-red liquid and emits a blue flame.

Sulfur forms several polyatomic molecules. The best-known allotrope isoctasulfur,cyclo-S8.Thepoint groupof cyclo-S8is D4dand its dipole moment is 0 D.[11]Octasulfur is a soft, bright-yellow solid that is odorless.[a]It melts at 115.21 °C (239.38 °F),[b]boils at 444.6 °C (832.3 °F).[7]At 95.2 °C (203.4 °F), below its melting temperature, cyclo-octasulfur begins slow changing from α-octasulfur to the β-polymorph.[13]The structure of the S8ring is virtually unchanged by this phase change, which affects the intermolecular interactions. Cooling of molten sulfur gives freezing point in 119.6 °C (247.3 °F),[14]as it predominantly consists of the β-S8molecules.[c]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 increasedviscositydue to the formation ofpolymers.[13]At higher temperatures, the viscosity decreases as depolymerization occurs. Molten sulfur assumes a dark red color above 200 °C (392 °F). The density of sulfur is about 2 g/cm3,depending on the allotrope; all of the stable allotropes are excellent electrical insulators.

Sulfursublimesmore or less between 20 °C (68 °F) and 50 °C (122 °F).[18]

Sulfur is insoluble in water but soluble incarbon disulfideand, to a lesser extent, in othernonpolarorganic solvents, such asbenzeneandtoluene.

Chemical properties

[edit]

Under normal conditions, sulfurhydrolyzesvery slowly to mainly formhydrogen sulfideandsulfuric acid:

1⁄2S
8+ 4H
2O→ 3H
2S+H
2SO
4
Left: Liquid hydrogen sulfide inside a test tube. Right: A bottle of sulfuric acid.


The reaction involves adsorption of protons ontoS
8clusters, followed bydisproportionationinto the reaction products.[19]

The second, fourth and sixthionization energiesof sulfur are 2252 kJ/mol, 4556 kJ/mol and 8495.8 kJ/mol, respectively. A composition of products of sulfur's reactions with oxidants (and its oxidation state) depends on that whether releasing out of a reaction energy overcomes these thresholds. Applyingcatalystsand / orsupply of outer energymay vary sulfur's oxidation state and a composition of reaction products. While reaction between sulfur and oxygen at normal conditions gives sulfur dioxide (oxidation state +4), formation ofsulfur trioxide(oxidation state +6) requires temperature 400–600 °C (750–1,100 °F) and presence of a catalyst.

In reactions with elements of lesserelectronegativity,it reacts as an oxidant and forms sulfides, where it has oxidation state −2.

Sulfur reacts with nearly all other elements with the exception of the noble gases, even with the notoriously unreactive metaliridium(yieldingiridium disulfide).[20]Some of those reactions need elevated temperatures.[21]

Allotropes

[edit]
Main article:Allotropes of sulfur
The structure of the cyclooctasulfur molecule, S8

Sulfur forms over 30 solidallotropes,more than any other element.[22]Besides S8,several other rings are known.[23]Removing one atom from the crown gives S7,which is of a deeper yellow than S8.HPLCanalysis of "elemental sulfur" reveals an equilibrium mixture of mainly S8,but with S7and small amounts of S6.[24]Larger rings have been prepared, including S12and S18.[25][26]

Amorphousor "plastic" sulfur is produced by rapid cooling of molten sulfur—for example, by pouring it into cold water.X-ray crystallographystudies show that the amorphous form may have ahelicalstructure with eight atoms per turn. The long coiled polymeric molecules make the brownish substanceelastic,and in bulk this form has the feel of crude rubber. This form ismetastableat room temperature and gradually reverts to the crystalline molecular allotrope, which is no longer elastic. This process happens within a matter of hours to days, but can be rapidly catalyzed.

Isotopes

[edit]
Main article:Isotopes of sulfur

Sulfur has 23 knownisotopes,four of which are stable:32S (94.99%±0.26%),33S (0.75%±0.02%),34S (4.25%±0.24%), and36S (0.01%±0.01%).[27][28]Other than35S, with ahalf-lifeof 87 days, theradioactiveisotopes of sulfur have half-lives less than 3 hours.

The preponderance of32S is explained by its production in the so-called alpha-process (one of the main classes of nuclear fusion reactions) in exploding stars. Other stable sulfur isotopes are produced in the bypass processes related with34Ar, and their composition depends on a type of a stellar explosion. For example, proportionally more33S comes fromnovaethan fromsupernovae.[29]

On the planet Earth the sulfur isotopic composition was determined by the Sun. Though it is assumed that the distribution of different sulfur isotopes should be more or less equal, it has been found that proportions of two most abundant sulfur isotopes32S and34S varies in different samples. Assaying of these isotopes ratio (δ34S) in the samples allows to make suggestions about their chemical history, and with support of other methods, it allows to age-date the samples, estimate temperature of equilibrium between ore and water, determine pH and oxygen fugacity, identify the activity of sulfate-reducing bacteria in the time of formation of the sample, or suggest the main sources of sulfur in ecosystems.[30]However, there are ongoing discussions about what is the real reason of the δ34S shifts, biological activity or postdeposital alteration.[31]

For example, whensulfide mineralsare precipitated, isotopic equilibration among solids and liquid may cause small differences in theδ34Svalues of co-genetic minerals. The differences between minerals can be used to estimate the temperature of equilibration. Theδ13Cand δ34S of coexistingcarbonate mineralsand sulfides can be used to determine thepHand oxygenfugacityof the ore-bearing fluid during ore formation.

Scientists measure thesulfur isotopesofmineralsin rocks andsedimentsto study theredoxconditions in the oceans in the past.Sulfate-reducing bacteriain marine sediment fractionatesulfur isotopesas they take insulfateand producesulfide.Prior to the 2010s, it was thought that sulfate reduction could fractionatesulfur isotopesup to 46permil[32]and fractionation larger than 46 permil recorded in sediments must be due todisproportionationof sulfur compounds in the sediment. This view has changed since the 2010s as experiments show thatsulfate-reducing bacteriacan fractionate to 66 permil.[33]As substrates for disproportionation are limited by the product ofsulfate reduction,the isotopic effect of disproportionation should be less than 16 permil in most sedimentary settings.[34]

In mostforestecosystems, 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 inhydrologicstudies. Differences in thenatural abundancescan be used in systems where there is sufficient variation in the34S of ecosystem components.Rocky Mountainlakes thought to be dominated by atmospheric sources of sulfate have been found to have measurably different34S values than lakes believed to be dominated by watershed sources of sulfate.

The radioactive35S is formed incosmic ray spallationof the atmospheric40Ar.This fact may be used for proving the presence of recent (not more than 1 year) atmospheric sediments in various things. This isotope may be obtained artificially by different ways. In practice, the reaction35Cl+n→35S +pis used by irradiatingpotassium chloridewith neutrons.[35]The isotope35S is used in various sulfur-containing compounds as aradioactive tracerfor many biological studies, for example, theHershey-Chase experiment.

Because of the weakbeta activityof35S, its compounds are relatively safe as long as they are not ingested or absorbed by the body.[36]

Natural occurrence

[edit]
Sulfur vat from which railroad cars are loaded, Freeport Sulphur Co., Hoskins Mound, Texas (1943)
Most of the yellow and orange hues ofIoare due to elemental sulfur and sulfur compounds deposited by activevolcanoes.
Sulfur extraction, East Java
A man carrying sulfur blocks fromKawah Ijen,a volcano in East Java, Indonesia, 2009

32S is created inside massive stars, at a depth where the temperature exceeds 2.5×109K, by thefusionof one nucleus of silicon plus one nucleus of helium.[37]As this nuclear reaction is part of thealpha processthat produces elements in abundance, sulfur is the 10thmost common element in the universe.

Sulfur, usually as sulfide, is present in many types ofmeteorites.Ordinary chondritescontain on average 2.1% sulfur, andcarbonaceous chondritesmay contain as much as 6.6%. It is normally present astroilite(FeS), but there are exceptions, with carbonaceous chondrites containing free sulfur, sulfates and other sulfur compounds.[38]The distinctive colors ofJupiter'svolcanicmoonIoare attributed to various forms of molten, solid, and gaseous sulfur.[39]In July 2024, elemental sulfur was confirmed to exist onMarsby surprise, after theCuriosity roverran over and crushed a rock revealing sulfur crystals inside it.[40]

Sulfur is the fifth most common element by mass in the Earth. Elemental sulfur can be found nearhot springsandvolcanicregions in many parts of the world, especially along thePacific Ring of Fire;such volcanic deposits are currently mined in Indonesia, Chile, and Japan. These deposits are polycrystalline, with the largest documented single crystal measuring 22 cm × 16 cm × 11 cm (8.7 in × 6.3 in × 4.3 in).[41]Historically,Sicilywas a major source of sulfur in theIndustrial Revolution.[42]Lakes of molten sulfur up to about 200 m (660 ft) in diameter have been found on the sea floor, associated withsubmarine volcanoes,at depths where the boiling point of water is higher than the melting point of sulfur.[43]

Native sulfur is synthesized byanaerobic bacteriaacting onsulfate mineralssuch asgypsuminsalt domes.[44][45]Significant deposits in salt domes occur along the coast of theGulf of Mexico,and inevaporitesin 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.[46]Currently, commercial production is still carried out in theOsiekmine in Poland. Such sources are now of secondary commercial importance, and most are no longer worked.

Common naturally occurring sulfur compounds include thesulfide minerals,such aspyrite(iron sulfide),cinnabar(mercury sulfide),galena(lead sulfide),sphalerite(zinc sulfide), andstibnite(antimony sulfide); and thesulfate minerals,such asgypsum(calcium sulfate),alunite(potassium aluminium sulfate), andbarite(barium sulfate). On Earth, just as upon Jupiter's moon Io, elemental sulfur occurs naturally in volcanic emissions, including emissions fromhydrothermal vents.

The main industrial source of sulfur is nowpetroleumandnatural gas.[9]

Compounds

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Main article:Sulfur compounds

Commonoxidation statesof sulfur range from −2 to +6. Sulfur forms stable compounds with all elements except thenoble gases.

Electron transfer reactions

[edit]
Lapis lazuliowes its blue color to atrisulfurradical anion (S−
3)

Sulfur polycations,S2+8,S2+4andS2+16are produced when sulfur is reacted with oxidizing agents in a strongly acidic solution.[47]The colored solutions produced by dissolving sulfur inoleumwere 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.S2+8is deep blue,S2+4is yellow andS2+16is red.[13]

Reduction of sulfur gives variouspolysulfideswith the formulaS2−
x,many of which have been obtained in crystalline form. Illustrative is the production ofsodium tetrasulfide:

4 Na + S8→ 2 Na2S4

Some of these dianions dissociate to giveradical anions,such asS−3gives the blue color of the rocklapis lazuli.

Two parallel sulfur chains grown inside a single-wallcarbon nanotube(CNT, a). Zig-zag (b) and straight (c) S chains inside double-wall CNTs[48]

This reaction highlights a distinctive property of sulfur: its ability tocatenate(bind to itself by formation of chains).Protonationof these polysulfide anions produces thepolysulfanes,H2Sx,wherex= 2, 3, and 4.[49]Ultimately, reduction of sulfur produces sulfide salts:

16 Na + S8→ 8 Na2S

The interconversion of these species is exploited in thesodium–sulfur battery.

Hydrogenation

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Treatment of sulfur with hydrogen giveshydrogen sulfide.When dissolved in water, hydrogen sulfide is mildly acidic:[7]

H2S ⇌ HS−+ H+

Hydrogen sulfide gas and the hydrosulfide anion are extremely toxic to mammals, due to their inhibition of the oxygen-carrying capacity ofhemoglobinand certaincytochromesin a manner analogous tocyanideandazide(see below, underprecautions).

Combustion

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The two principal sulfur oxides are obtained by burning sulfur:

S + O2→ SO2(sulfur dioxide)
2 SO2+ O2→ 2 SO3(sulfur trioxide)

Many other sulfur oxides are observed including thesulfur-rich oxidesincludesulfur monoxide,disulfur monoxide,disulfur dioxides, andhigher oxidescontaining peroxo groups.

Halogenation

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Sulfur reacts withfluorineto give the highly reactivesulfur tetrafluorideand the highly inertsulfur hexafluoride.[50]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 chlorideandchlorosulfuric acidare derivatives of sulfuric acid;thionyl chloride(SOCl2) is a common reagent inorganic synthesis.[51]Bromine also oxidizes sulfur to formsulfur dibromideanddisulfur dibromide.[51]

Pseudohalides

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Sulfur oxidizescyanideandsulfiteto givethiocyanateandthiosulfate,respectively.

Metal sulfides

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Sulfur reacts with many metals. Electropositive metals give polysulfide salts. Copper, zinc, and silver are attacked by sulfur; seetarnishing.Although manymetal sulfidesare known, most are prepared by high temperature reactions of the elements.[52]Geoscientists also study the isotopes of metal sulfides in rocks and sediment to study environmental conditions in the Earth's past.[53]

Organic compounds

[edit]
Main article:Organosulfur compounds

Some of the main classes of sulfur-containing organic compounds include the following:[54]

Compounds with carbon–sulfur multiple bonds are uncommon, an exception beingcarbon disulfide,a volatile colorless liquid that is structurally similar to carbon dioxide. It is used as a reagent to make the polymerrayonand many organosulfur compounds. Unlikecarbon monoxide,carbon monosulfideis stable only as an extremely dilute gas, found between solar systems.[55]

Organosulfur compounds are responsible for some of the unpleasant odors of decaying organic matter. They are widely known as theodorantin domestic natural gas, garlic odor, and skunk spray, as well as a component ofbad breathodor. Not all organic sulfur compounds smell unpleasant at all concentrations: the sulfur-containingmonoterpenoidgrapefruit mercaptanin small concentrations is the characteristic scent of grapefruit, but has a generic thiol odor at larger concentrations.Sulfur mustard,a potentvesicant,wasused in World War Ias a disabling agent.[56]

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 naturalrubber,elemental sulfur is heated with the rubber to the point that chemical reactions formdisulfidebridges betweenisopreneunits of the polymer. This process, patented in 1843,[citation needed]made rubber a major industrial product, especially in automobile tires. Because of the heat and sulfur, the process was namedvulcanization,after the Roman god of the forge andvolcanism.

History

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Antiquity

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Pharmaceutical container for sulfur from the first half of the 20th century. From theMuseo del Objeto del Objetocollection

Being abundantly available in native form, sulfur was known in ancient times and is referred to in theTorah(Genesis).English translations of the Christian Biblecommonly referred to burning sulfur as "brimstone", giving rise to the term "fire-and-brimstone"sermons,in which listeners are reminded of the fate ofeternal damnationthat await the unbelieving and unrepentant. It is from this part of the Bible[57]thatHellis implied to "smell of sulfur" (likely due to its association with volcanic activity). According to theEbers Papyrus,a sulfur ointment was used in ancientEgyptto treat granular eyelids. Sulfur was used forfumigationin preclassicalGreece;[58]this is mentioned in theOdyssey.[59]Pliny the Elderdiscusses sulfur in book 35 of hisNatural History,saying that its best-known source is the island ofMelos.He mentions its use for fumigation, medicine, and bleaching cloth.[60]

A natural form of sulfur known asshiliuhuang(Thạch lưu hoàng) was known in China since the 6th century BC and found inHanzhong.[61]By the 3rd century, the Chinese had discovered that sulfur could be extracted frompyrite.[61]ChineseDaoistswere interested in sulfur's flammability and its reactivity with certain metals, yet its earliest practical uses were found intraditional Chinese medicine.[61]TheWujing Zongyaoof 1044 AD described various formulas for Chineseblack powder,which is a mixture ofpotassium nitrate(KNO
3),charcoal,and sulfur.[62]

Sulfur
Brimstone
Alchemical signs for sulfur, or thecombustibleelements, and brimstone, an older/archaic name for sulfur[63]

Indian alchemists, practitioners of the "science of chemicals" (Sanskrit:रसशास्त्र,romanized:rasaśāstra), wrote extensively about the use of sulfur in alchemical operations with mercury, from the eighth century AD onwards.[64]In therasaśāstratradition, sulfur is called "the smelly" (गन्धक,gandhaka).

EarlyEuropeanalchemistsgave sulfur a uniquealchemical symbol,a triangle atop a cross (🜍). (This is sometimes confused with the astronomical crossed-spear symbol ⚴ for2 Pallas.) The variation known as brimstone has a symbol combining atwo-barred crossatop alemniscate(🜏). In traditional skin treatment, elemental sulfur was used (mainly in creams) to alleviate such conditions asscabies,ringworm,psoriasis,eczema,andacne.The mechanism of action is unknown—though elemental sulfur does oxidize slowly to sulfurous acid, which is (through the action ofsulfite) a mild reducing and antibacterial agent.[65][66][67]

Modern times

[edit]
Today sulfur is known to have antifungal, antibacterial, andkeratolyticactivity; in the past it was used against acne vulgaris, rosacea, seborrheic dermatitis, dandruff, pityriasis versicolor, scabies, and warts.[68]This 1881 advertisement baselessly claims efficacy against rheumatism, gout, baldness, and graying of hair.

Sulfur appears in a column of fixed (non-acidic)alkaliin a chemical table of 1718.[69]Antoine Lavoisierused sulfur in combustion experiments, writing of some of these in 1777.[70]

Sulfur deposits inSicilywere the dominant source for more than a century. By the late 18th century, about 2,000 tonnes per year of sulfur were imported intoMarseille,France, for the production ofsulfuric acidfor use in theLeblanc process.InindustrializingBritain, with the repeal oftariffson salt in 1824, demand for sulfur from Sicily surged upward. The increasing British control and exploitation of the mining, refining, and transportation of the sulfur, coupled with the failure of this lucrative export to transform Sicily's backward and impoverished economy, led to theSulfur Crisis of 1840,whenKing Ferdinand IIgave 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.[71][72]

In 1867, elemental sulfur was discovered in underground deposits inLouisianaandTexas.The highly successfulFrasch processwas developed to extract this resource.[73]

In the late 18th century,furnituremakers used molten sulfur to producedecorative inlays.[74]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.[46]

With the advent of thecontact process,the majority of sulfur today is used to make sulfuric acid for a wide range of uses, particularly fertilizer.[75]

In recent times, the main source of sulfur has becomepetroleumandnatural gas.This is due to the requirement to remove sulfur from fuels in order to preventacid rain,and has resulted in a surplus of sulfur.[9]

Spelling and etymology

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Sulfuris derived from the Latin wordsulpur,which wasHellenizedtosulphurin the erroneous belief that the Latin word came from Greek. This spelling was later reinterpreted as representing an /f/ sound and resulted in the spellingsulfur,which appears in Latin toward the end of theClassical period.The true Ancient Greek word for sulfur,θεῖον,theîon(from earlierθέειον,théeion), is the source of the international chemical prefixthio-.The Modern Standard Greek word for sulfur is θείο,theío.

In 12th-centuryAnglo-French,it wassulfre.In the 14th century, the erroneously Hellenized Latin-ph-was restored in Middle Englishsulphre.By the 15th century, both full Latin spelling variantssulfurandsulphurbecame common in English. The parallelf~phspellings continued in Britain until the 19th century, when the word was standardized assulphur.[76]On the other hand,sulfurwas the form chosen in the United States, whereas Canada uses both.

TheIUPACadopted the spellingsulfurin 1990[77][78]as did the Nomenclature Committee of theRoyal Society of Chemistryin 1992, restoring the spellingsulfurto Britain.[79]Oxford Dictionariesnote 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. "[80]

Production

[edit]
Sicilian kiln used to obtain sulfur from volcanic rock (diagram from a 1906 chemistry book)
Traditional sulfur mining atIjen Volcano,East Java, Indonesia. This image shows the dangerous and rugged conditions the miners face, including toxic smoke and high drops, as well as their lack of protective equipment. The pipes over which they are standing are for condensing sulfur vapors.

Sulfur may be found by itself and historically was usually obtained in this form;pyritehas also been a source of sulfur.[81]In volcanic regions inSicily,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 orcarusicarrying 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 inSicilian sulfur mineswere horrific, promptingBooker T. Washingtonto 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."[82]Sulfur is still mined from surface deposits in poorer nations with volcanoes, such asIndonesia,and worker conditions have not improved much since Booker T. Washington's days.[83]

Elemental sulfur was extracted fromsalt domes(in which it sometimes occurs in nearly pure form) until the late 20th century. Sulfur is now produced as a side product of other industrial processes such as in oil refining, in which sulfur is undesired. As a mineral, native sulfur under salt domes is thought to be a fossil mineral resource, produced by the action of anaerobic bacteria on sulfate deposits. It was removed from such salt-dome mines mainly by theFrasch process.[46]In this method, superheated water was pumped into a native sulfur deposit to melt the sulfur, and then compressed air returned the 99.5% pure melted product to the surface. Throughout the 20th century this procedure produced elemental sulfur that required no further purification. Due to a limited number of such sulfur deposits and the high cost of working them, this process for mining sulfur has not been employed in a major way anywhere in the world since 2002.[84][85]

Sulfur recovered from hydrocarbons inAlberta,stockpiled for shipment inNorth Vancouver,British Columbia

Today, sulfur is produced from petroleum,natural gas,and related fossil resources, from which it is obtained mainly ashydrogen sulfide.[9]Organosulfur compounds,undesirable impurities in petroleum, may be upgraded by subjecting them tohydrodesulfurization,which cleaves the C–S bonds:[84][85]

R-S-R + 2 H2→ 2 RH + H2S

The resulting hydrogen sulfide from this process, and also as it occurs in natural gas, is converted into elemental sulfur by theClaus process.This process entails oxidation of some hydrogen sulfide to sulfur dioxide and then thecomproportionationof the two:[84][85]

3 O2+ 2 H2S → 2 SO2+ 2 H2O
SO2+ 2 H2S → 3 S + 2 H2O
Production and price (US market) of elemental sulfur

Owing to the high sulfur content of theAthabasca Oil Sands,stockpiles of elemental sulfur from this process now exist throughoutAlberta,Canada.[86]Another way of storing sulfur is as abinderfor concrete, the resulting product having some desirable properties (seesulfur concrete).[87]

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 areChina(9.6), theUnited States(8.8),Canada(7.1) andRussia(7.1).[88]Production has been slowly increasing from 1900 to 2010; the price was unstable in the 1980s and around 2010.[89]

Applications

[edit]

Sulfuric acid

[edit]

Elemental sulfur is used mainly as a precursor to other chemicals. Approximately 85% (1989) is converted tosulfuric acid(H2SO4):

1⁄8S8+3⁄2O2+H2O→H2SO4
Sulfuric acid production in 2000

In 2010, the United States produced more sulfuric acid than any other inorganic industrial chemical.[89]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.[46]

Other important sulfur chemistry

[edit]

Sulfur reacts directly with methane to givecarbon disulfide,which is used to manufacturecellophaneandrayon.[46]One of the uses of elemental sulfur is invulcanizationof rubber, wherepolysulfidechains crosslink organic polymers. Large quantities ofsulfitesare used tobleachpaperand to preservedried fruit.Manysurfactantsanddetergents(e.g.sodium lauryl sulfate) are sulfate derivatives.Calcium sulfate,gypsum (CaSO4·2H2O) is mined on the scale of 100 milliontonneseach year for use inPortland cementand fertilizers.

When silver-basedphotographywas widespread, sodium and ammoniumthiosulfatewere widely used as "fixing agents". Sulfur is a component ofgunpowder( "black powder" ).

Fertilizer

[edit]

Amino acidssynthesized byliving organismssuch asmethionineandcysteinecontainorganosulfurgroups (thioesterandthiolrespectively). Theantioxidantglutathioneprotecting many living organisms againstfree radicalsandoxidative stressalso contains organic sulfur. Somecropssuch asonionandgarlicalso produce differentorganosulfur compoundssuch assyn-propanethial-S-oxideresponsible of lacrymal irritation (onions), ordiallyl disulfideandallicin(garlic).Sulfates,commonly found insoilsandgroundwatersare often a sufficient natural source of sulfur for plants and bacteria.Atmospheric depositionofsulfur 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 alimiting nutrientfor plants andmicroorganisms(seeLiebig's barrel). However, in some circumstances, soils can be depleted insulfate,e.g. if this later is leached bymeteoric 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 offertilizers.The most important form of sulfur for fertilizer iscalcium sulfate,commonly found in nature as the mineralgypsum(CaSO4·2H2O). Elemental sulfur ishydrophobic(not soluble in water) and cannot be used directly by plants. Elemental sulfur (ES) is sometimes mixed withbentoniteto amend depleted soils for crops with high requirement in organo-sulfur. Over time,oxidationabioticprocesses withatmosphericoxygenandsoil bacteriacanoxidizeand convert elemental sulfur to soluble derivatives, which can then be used by microorganisms and plants. Sulfur improves the efficiency of other essential plant nutrients, particularlynitrogenand phosphorus.[90]Biologically produced sulfur particles are naturallyhydrophilicdue to abiopolymercoating 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 forphosphorus.It is anessential nutrient for plantgrowth,root noduleformation of legumes, and immunity and defense systems. Sulfur deficiency has become widespread in many countries in Europe.[91][92][93]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 limitacid rains.[94][90]

Fungicide and pesticide

[edit]
Sulfur candle originally sold for home fumigation

Elemental sulfur is one of the oldest fungicides andpesticides."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 ofpowdery mildewdiseases as well as black spot. In organic production, sulfur is the most important fungicide. It is the only fungicide used inorganicallyfarmed apple production against the main diseaseapple scabunder 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 orfrom a dusting plane.Wettable sulfur is the commercial name for dusting sulfur formulated with additional ingredients to make it watermiscible.[87][95]It has similar applications and is used as afungicideagainstmildewand other mold-related problems with plants and soil.

Elemental sulfur powder is used as an "organic"(i.e.," green ")insecticide(actually anacaricide) againstticksandmites.A common method of application is dusting the clothing or limbs with sulfur powder.

A diluted solution oflime sulfur(made by combiningcalcium hydroxidewith elemental sulfur in water) is used as a dip for pets to destroyringworm (fungus),mange,and otherdermatosesandparasites.

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

Pharmaceuticals

[edit]
Main article:Sulfur (pharmacy)

Sulfur (specificallyoctasulfur,S8) is used in pharmaceutical skin preparations for the treatment ofacneand other conditions. It acts as akeratolyticagent and also kills bacteria, fungi,scabiesmites, and other parasites.[96]Precipitated sulfur and colloidal sulfur are used, in form oflotions,creams, powders, soaps, and bath additives, for the treatment ofacne vulgaris,acne rosacea,andseborrhoeic dermatitis.[97]

Many drugs contain sulfur.[98]Early examples include antibacterialsulfonamides,known assulfa drugs.A more recent example is mucolyticacetylcysteine.Sulfur is a part of many bacterial defense molecules. Mostβ-lactamantibiotics, including thepenicillins,cephalosporinsandmonobactamscontain sulfur.[54]

Batteries

[edit]

Due to their high energy density and the availability of sulfur, there is ongoing research in creating rechargeablelithium–sulfur batteries.Until now, carbonate electrolytes have caused failures in such batteries after a single cycle. In February 2022, researchers atDrexel Universityhave 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.[99]

Biological role

[edit]

Sulfur is an essential component of all livingcells.It is the eighth most abundant element in the human body by weight,[100]about equal in abundance topotassium,and slightly greater thansodiumandchlorine.[101]A 70 kg (150 lb) human body contains about 140 grams (4.9 oz) of sulfur.[102]The main dietary source of sulfur for humans is sulfur-containing amino-acids,[103]which can be found in plant and animal proteins.[104]

Transferring sulfur between inorganic and biomolecules

[edit]
See also:Sulfur cycleandSulfur metabolism

In the 1880s, while studyingBeggiatoa(a bacterium living in a sulfur rich environment),Sergei Winogradskyfound that it oxidizedhydrogen sulfide(H2S) as an energy source, forming intracellular sulfur droplets. Winogradsky referred to this form of metabolism as inorgoxidation (oxidation of inorganic compounds).[105]Another contributor, who continued to study it wasSelman Waksman.[106]Primitive bacteria that live around deep oceanvolcanic ventsoxidize hydrogen sulfide for their nutrition, as discovered byRobert Ballard.[10]

Sulfur oxidizers can use as energy sources reduced sulfur compounds, including hydrogen sulfide, elemental sulfur,sulfite,thiosulfate,and variouspolythionates(e.g.,tetrathionate).[107]They depend on enzymes such assulfur oxygenaseandsulfite oxidaseto oxidize sulfur to sulfate. Somelithotrophscan even use the energy contained in sulfur compounds to produce sugars, a process known aschemosynthesis.Somebacteriaandarchaeause hydrogen sulfide in place of water as theelectron donorin chemosynthesis, a process similar tophotosynthesisthat produces sugars and uses oxygen as theelectron acceptor.Sulfur-based chemosynthesis may be simplifiedly compared with photosynthesis:

H2S + CO2→ sugars + S
H2O + CO2→ sugars + O2

There are bacteria combining these two ways of nutrition:green sulfur bacteriaandpurple sulfur bacteria.[108]Also sulfur-oxidizing bacteria can go into symbiosis with larger organisms, enabling the later to use hydrogen sulfide as food to be oxidized. Example: thegiant tube worm.[109]

There aresulfate-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 ofthe ancient oceanshave biological origin.[110][111][112]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 byplantsrootsfrom soil assulfateand transported as a phosphate ester. Sulfate is reduced to sulfide via sulfite before it is incorporated intocysteineand other organosulfur compounds.[113]

SO2−4→SO2−3→H2S→ cysteine (thiol) → methionine (thioether)

While the plants' role in transferring sulfur to animals byfood chainsis more or less understood, the role of sulfur bacteria is just getting investigated.[114][115]

Protein and organic metabolites

[edit]

In all forms of life, most of the sulfur is contained in twoproteinogenic amino acids(cysteineandmethionine), thus the element is present in allproteinsthat contain these amino acids, as well as in respectivepeptides.[116]Some of the sulfur is comprised in certain metabolites—many of which arecofactors—and sulfated polysaccharides ofconnective tissue(chondroitin sulfates,heparin).

Disulfide bonds between two alpha-helix
Schematic representation of disulfide bridges (in yellow) between two protein helices

Proteins, to execute theirbiological function,need to have specific space geometry. Formation of this geometry is performed in a process calledprotein folding,and is provided by intra- and inter-molecular bonds. The process has several stages. While at premier stages a polypeptide chain folds due tohydrogen bonds,at later stages folding is provided (apart from hydrogen bonds) bycovalent bondsbetween two sulfur atoms of two cysteine residues (so called disulfide bridges) at different places of a chain (tertiary protein structure) as well as between two cysteine residues in two separated protein subunits (quaternary protein structure). Both structures easily may be seen ininsulin.As thebond energyof a covalent disulfide bridge is higher than the energy of acoordinate bondor hydrophobic interaction, higher disulfide bridges content leads to higher energy needed for proteindenaturation.In general disulfide bonds are necessary in proteins functioning outside cellular space, and they do not change proteins' conformation (geometry), but serve as its stabilizers.[117]Withincytoplasmcysteine residues of proteins are saved in reduced state (i.e. in -SH form) bythioredoxins.[118]

This property manifests in following examples.Lysozymeis stable enough to be applied as a drug.[119]Feathers and hair have relative strength, and consisting in themkeratinis considered indigestible by most organisms. However, there are fungi and bacteria containingkeratinase,and are able to destruct keratin.

Many important cellular enzymes use prosthetic groups ending with -SH moieties to handle reactions involving acyl-containing biochemicals: two common examples from basic metabolism arecoenzyme Aandalpha-lipoic acid.[120]Cysteine-related metaboliteshomocysteineandtaurineare other sulfur-containing amino acids that are similar in structure, but not coded byDNA,and are not part of theprimary structureof proteins, take part in various locations of mammalian physiology.[121][122]Two of the 13 classical vitamins,biotinandthiamine,contain sulfur, and serve as cofactors to several enzymes.[123][124] In intracellular chemistry, sulfur operates as a carrier of reducing hydrogen and its electrons for cellular repair of oxidation. Reducedglutathione,a sulfur-containing tripeptide, is a reducing agent through its sulfhydryl (–SH) moiety derived fromcysteine.

Methanogenesis,the route to most of the world's methane, is a multistep biochemical transformation ofcarbon dioxide.This conversion requires several organosulfur cofactors. These includecoenzyme M,CH3SCH2CH2SO−3,the immediate precursor tomethane.[125]

Metalloproteins and inorganic cofactors

[edit]

Metalloproteins—in which the active site is a transition metal ion (or metal-sulfide cluster) often coordinated by sulfur atoms of cysteine residues[126]—are essential components of enzymes involved in electron transfer processes. Examples includeplastocyanin(Cu2+) andnitrous oxide reductase(Cu–S). The function of these enzymes is dependent on the fact that the transition metal ion can undergoredox reactions.Other examples include many zinc proteins,[127]as well asiron–sulfur clusters.Most pervasive are theferrodoxins,which serve as electron shuttles in cells. In bacteria, the importantnitrogenaseenzymes contain an Fe–Mo–S cluster and is acatalystthat performs the important function ofnitrogen 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.[128]

Sulfur is also present inmolybdenum cofactor.[129]

Easiness of electron flow in a cluster provides catalytic effect of a respective enzyme.

Sulfate

[edit]
Main article:Sulfation § Sulfation in biology

Deficiency

[edit]

In humansmethionineis anessential amino acid;cysteineis conditionally essential and may be synthesized from non-essentialserine(sulfur donor would be methionine in this case). Dietary deficiency rarely happens in common conditions. Artificial methionine deficiency is attempted to apply in cancer treatment,[130]but the method is still potentially dangerous.[131]

Isolated sulfite oxidase deficiencyis a rare, fatal genetic disease preventing production ofsulfite oxidase,needed to metabolize sulfites to sulfates.[132]

Precautions

[edit]
Sulfur
Hazards
GHSlabelling:
GHS07: Exclamation markGHS02: Flammable
Warning
H315[133]
NFPA 704(fire diamond)
[134]
NFPA 704 four-colored diamondHealth 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g. chloroformFlammability 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g. canola oilInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
2
1
0
Effect of acid rain on a forest, Jizera Mountains, Czech Republic

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,[135]and cases of life-threatening metabolic acidosis have been reported after patients deliberately consumed sulfur as a folk remedy.[136][137]

Toxicity of sulfur compounds

[edit]

When sulfur burns in air, it producessulfur dioxide.In water, this gas produces sulfurous acid and sulfites; sulfites are antioxidants that inhibit growth of aerobic bacteria and a usefulfood additivein small amounts. At high concentrations these acids harm thelungs,eyes,or othertissues.[138]In organisms without lungs such as insects, sulfite in high concentration preventsrespiration.[139]

Sulfur trioxide(made by catalysis from sulfur dioxide) andsulfuric acidare 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.[140]

The burning ofcoaland/orpetroleumby industry andpower plantsgenerates sulfur dioxide (SO2) that reacts with atmospheric water and oxygen to producesulfurous acid(H2SO3).[141]These acids are components ofacid rain,lowering thepHofsoiland freshwater bodies, sometimes resulting in substantial damage to theenvironmentandchemical weatheringof statues and structures. Fuel standards increasingly require that fuel producers extract sulfur fromfossil fuelsto prevent acid rain formation. This extracted and refined sulfur represents a large portion of sulfur production. In coal-fired power plants,flue gasesare sometimes purified. More modern power plants that usesynthesis gasextract the sulfur before they burn the gas.

Hydrogen sulfideis about one-half astoxicashydrogen cyanide,and intoxicates by the same mechanism (inhibition of the respiratory enzymecytochrome oxidase),[142]though hydrogen sulfide is less likely to cause sudden poisonings from small inhaled amounts (near itspermissible exposure limit(PEL) of 20 ppm) because of its disagreeable odor.[143]However, its presence in ambient air at concentration over 100–150 ppm quickly deadens the sense of smell,[144]and a victim may breathe increasing quantities without noticing until severe symptoms cause death. Dissolvedsulfideandhydrosulfidesalts are toxic by the same mechanism.

Notes

[edit]
  1. ^But impure samples have an odor similar to that ofmatches.A strong odor called "smell of sulfur" actually is given off by several sulfur compounds, such ashydrogen sulfideandorganosulfurcompounds.
  2. ^A figure of sulfur's melting point in 115.21°C had been determined by two laboratories of the US Department of Energy (Jefferson Lab and Los Alamos National Lab).[12] Greenwood and Earnshaw say that at fast heating for microcrystalline α-S8the melting point in 115.1 °C (239.2 °F) is specified.[7]
  3. ^Historically, it was a rather difficult task to find the exact melting point of sulfur.[15]When heated sluggishly, the factual melting point may lie within the range from 114.6 °C (238.3 °F) or even lower, to 120.4 °C (248.7 °F)[7](factors, that interfere a definite melting point, is apolymerlike natureof sulfur[16]and a large number of allotropes.[17]) Melting point may be presented as a temperature interval, depending on the allotrope composition of a sample at a moment of melting.

See also

[edit]

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Further reading

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Wikiquote has quotations related toSulfur.

Sigel, Astrid; Freisinger, Eva; Sigel, Roland K.O., eds. (2020).Transition Metals and Sulfur: A Strong Relationship for Life.Guest Editors Martha E Sosa Torres and Peter M.H.Kroneck. Berlin/Boston: de Gruyter. pp. xlv+455.ISBN978-3-11-058889-7.

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