Sulfur

Sulfur, 16S
Sulfur-sample.jpg
Sulfur
Alternative namesulphur (British spelling)
Appearancelemon yellow sintered microcrystals
Standard atomic weight Ar, std(S)[32.05932.076] conventional: 32.06
Sulfur in the periodic table
HydrogenHelium
LithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeon
SodiumMagnesiumAluminiumSiliconPhosphorusSulfurChlorineArgon
PotassiumCalciumScandiumTitaniumVanadiumChromiumManganeseIronCobaltNickelCopperZincGalliumGermaniumArsenicSeleniumBromineKrypton
RubidiumStrontiumYttriumZirconiumNiobiumMolybdenumTechnetiumRutheniumRhodiumPalladiumSilverCadmiumIndiumTinAntimonyTelluriumIodineXenon
CaesiumBariumLanthanumCeriumPraseodymiumNeodymiumPromethiumSamariumEuropiumGadoliniumTerbiumDysprosiumHolmiumErbiumThuliumYtterbiumLutetiumHafniumTantalumTungstenRheniumOsmiumIridiumPlatinumGoldMercury (element)ThalliumLeadBismuthPoloniumAstatineRadon
FranciumRadiumActiniumThoriumProtactiniumUraniumNeptuniumPlutoniumAmericiumCuriumBerkeliumCaliforniumEinsteiniumFermiumMendeleviumNobeliumLawrenciumRutherfordiumDubniumSeaborgiumBohriumHassiumMeitneriumDarmstadtiumRoentgeniumCoperniciumNihoniumFleroviumMoscoviumLivermoriumTennessineOganesson
O

S

Se
phosphorussulfurchlorine
Atomic number (Z)16
Groupgroup 16 (chalcogens)
Periodperiod 3
Blockp-block
Element category  Reactive nonmetal
Electron configuration[Ne] 3s2 3p4
Electrons per shell2, 8, 6
Physical properties
Phase at STPsolid
Melting point388.36 K ​(115.21 °C, ​239.38 °F)
Boiling point717.8 K ​(444.6 °C, ​832.3 °F)
Density (near r.t.)alpha: 2.07 g/cm3
beta: 1.96 g/cm3
gamma: 1.92 g/cm3
when liquid (at m.p.)1.819 g/cm3
Critical point1314 K, 20.7 MPa
Heat of fusionmono: 1.727 kJ/mol
Heat of vaporizationmono: 45 kJ/mol
Molar heat capacity22.75 J/(mol·K)
Vapor pressure
P (Pa)1101001 k10 k100 k
at T (K)375408449508591717
Atomic properties
Oxidation states−2, −1, 0, +1, +2, +3, +4, +5, +6 (a strongly acidic oxide)
ElectronegativityPauling scale: 2.58
Ionization energies
  • 1st: 999.6 kJ/mol
  • 2nd: 2252 kJ/mol
  • 3rd: 3357 kJ/mol
  • (more)
Covalent radius105±3 pm
Van der Waals radius180 pm
Color lines in a spectral range
Spectral lines of sulfur
Other properties
Natural occurrenceprimordial
Crystal structureorthorhombic
Orthorhombic crystal structure for sulfur
Thermal conductivity0.205 W/(m·K) (amorphous)
Electrical resistivity2×1015  Ω·m (at 20 °C) (amorphous)
Magnetic orderingdiamagnetic[1]
Magnetic susceptibility(α) −15.5·10−6 cm3/mol (298 K)[2]
Bulk modulus7.7 GPa
Mohs hardness2.0
CAS Number7704-34-9
History
DiscoveryChinese[3] (before 2000 BCE)
Recognized as an element byAntoine Lavoisier (1777)
Main isotopes of sulfur
Iso­topeAbun­danceHalf-life (t1/2)Decay modePro­duct
32S94.99%stable
33S0.75%stable
34S4.25%stable
35Strace87.37 dβ35Cl
36S0.01%stable
| references

Sulfur (in British English, sulphur) is a chemical element with the symbol S and atomic number 16. It is abundant, multivalent, and nonmetallic. Under normal conditions, sulfur atoms form cyclic octatomic molecules with a chemical formula S8. Elemental sulfur is a bright yellow, crystalline solid at room temperature.

Sulfur is the tenth most common element by mass in the universe, and the fifth most common on Earth. Though sometimes found in pure, native 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 Egypt. In the Bible, sulfur is called brimstone,[4] which means "burning stone".[5] Today, almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants from natural gas and petroleum. The greatest commercial use of the element is the production of sulfuric acid for sulfate and phosphate fertilizers, and other chemical processes. The element sulfur is used in matches, insecticides, and fungicides. Many sulfur compounds are odoriferous, and the smells of odorized natural gas, skunk scent, grapefruit, and garlic are due to organosulfur compounds. Hydrogen sulfide gives the characteristic odor to rotting eggs and other biological processes.

Sulfur is an essential element for all life, but almost always in the form of organosulfur compounds or metal sulfides. Three amino acids (cysteine, cystine, and methionine) and two vitamins (biotin and thiamine) are organosulfur compounds. Many cofactors also contain sulfur, including glutathione, thioredoxin, and iron–sulfur proteins. Disulfides, S–S bonds, confer mechanical strength and insolubility of the protein keratin, found in outer skin, hair, and feathers. Sulfur is one of the core chemical elements needed for biochemical functioning and is an elemental macronutrient for all living organisms.

Characteristics

When burned, sulfur melts to a blood-red liquid and emits a blue flame.

Physical properties

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.[6] Octasulfur is a soft, bright-yellow solid that is odorless, but impure samples have an odor similar to that of matches.[7] It melts at 115.21 °C (239.38 °F), boils at 444.6 °C (832.3 °F) and sublimes easily.[4] At 95.2 °C (203.4 °F), below its melting temperature, cyclo-octasulfur changes from α-octasulfur to the β-polymorph.[8] The structure of the S8 ring is virtually unchanged by this phase change, which affects the intermolecular interactions. Between its melting and boiling temperatures, octasulfur changes its allotrope again, turning from β-octasulfur to γ-sulfur, again accompanied by a lower density but increased viscosity due to the formation of polymers.[8] 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.

Chemical properties

Sulfur burns with a blue flame with formation of sulfur dioxide, which has a suffocating and irritating odor. Sulfur is insoluble in water but soluble in carbon disulfide and, to a lesser extent, in other nonpolar organic solvents, such as benzene and toluene. The first and second ionization energies of sulfur are 999.6 and 2252 kJ/mol, respectively. Despite such figures, the +2 oxidation state is rare, with +4 and +6 being more common. The fourth and sixth ionization energies are 4556 and 8495.8 kJ/mol, the magnitude of the figures caused by electron transfer between orbitals; these states are only stable with strong oxidants such as fluorine, oxygen, and chlorine.[citation needed] Sulfur reacts with nearly all other elements with the exception of the noble gases, even with the notoriously unreactive metal iridium (yielding iridium disulfide).[9] Some of those reactions need elevated temperatures.[10]

Allotropes

The structure of the cyclooctasulfur molecule, S8

Sulfur forms over 30 solid allotropes, more than any other element.[11] Besides S8, several other rings are known.[12] Removing one atom from the crown gives S7, which is more of a deep yellow than the S8. HPLC analysis of "elemental sulfur" reveals an equilibrium mixture of mainly S8, but with S7 and small amounts of S6.[13] Larger rings have been prepared, including S12 and S18.[14][15]

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 this form has the feel of crude rubber. This form is metastable at room temperature and gradually reverts to crystalline molecular allotrope, which is no longer elastic. This process happens within a matter of hours to days, but can be rapidly catalyzed.

Isotopes

Sulfur has 23 known isotopes, four of which are stable: 32S (94.99%±0.26%), 33S (0.75%±0.02%), 34S (4.25%±0.24%), and 36S (0.01%±0.01%).[16][17] Other than 35S, with a half-life of 87 days and formed in cosmic ray spallation of 40Ar, the radioactive isotopes of sulfur have half-lives less than 3 hours.

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.

In most forest ecosystems, sulfate is derived mostly from the atmosphere; weathering of ore minerals and evaporites contribute some sulfur. Sulfur with a distinctive isotopic composition has been used to identify pollution sources, and enriched sulfur has been added as a tracer in 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 characteristic 34S values from lakes believed to be dominated by watershed sources of sulfate.

Natural occurrence

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

32S is created inside massive stars, at a depth where the temperature exceeds 2.5×109 K, by the fusion of one nucleus of silicon plus one nucleus of helium.[18] 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.[19] The distinctive colors of Jupiter's volcanic moon Io are attributed to various forms of molten, solid, and gaseous sulfur.[20]

Sulfur occurs in fumaroles such as this one in Vulcano, Italy

It is the fifth most common element by mass in the Earth. Elemental sulfur can be found near hot springs and volcanic regions in many parts of the world, especially along the Pacific Ring of Fire; such volcanic deposits are currently mined in Indonesia, Chile, and Japan. These deposits are polycrystalline, with the largest documented single crystal measuring 22×16×11 cm.[21] Historically, Sicily was a major source of sulfur in the Industrial Revolution.[22]

Native sulfur is synthesised by anaerobic bacteria acting on sulfate minerals such as gypsum in salt domes.[23][24] 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 until recently the basis for commercial production in the United States, Russia, Turkmenistan, and Ukraine.[25] Currently, commercial production is still carried out in the Osiek mine in Poland. Such sources are now of secondary commercial importance, and most are no longer worked.

Common naturally occurring sulfur compounds include the 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.