Sulfur
Chemical element, symbol S and atomic number 16 / From Wikipedia, the free encyclopedia
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Sulfur (also spelled sulphur in British English) 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 abundant 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,[7] which means "burning stone".[8] Today, almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants from natural gas and petroleum.[9][10] 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.
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.[11] Octasulfur is a soft, bright-yellow solid that is odorless, but impure samples have an odor similar to that of matches.[12] It melts at 115.21 °C (239.38 °F), boils at 444.6 °C (832.3 °F)[7] and sublimes more or less between 20 °C (68 °F) and 50 °C (122 °F).[13] At 95.2 °C (203.4 °F), below its melting temperature, cyclo-octasulfur changes from α-octasulfur to the β-polymorph.[14] 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.[14] 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.
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 properties
Under normal conditions, sulfur hydrolyzes very slowly to mainly form hydrogen sulfide and sulfuric acid:
8 + 4 H
2O → 3 H
2S + H
2SO
4
The reaction involves adsorption of protons onto S
8 clusters, followed by disproportionation into the reaction products.[15]
The second, fourth and sixth ionization energies of 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. Applying catalysts and / or supply of outer energy may 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 of sulfur trioxide (oxidation state +6) requires temperature 400–600 °C (750–1,100 °F) 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 with the exception of the noble gases, even with the notoriously unreactive metal iridium (yielding iridium disulfide).[16] Some of those reactions need elevated temperatures.[17]
Allotropes
Sulfur forms over 30 solid allotropes, more than any other element.[18] Besides S8, several other rings are known.[19] 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.[20] Larger rings have been prepared, including S12 and S18.[21][22]
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 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
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%).[23][24] 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. 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.[25]
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 isotopes 32S and 34S 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.[26] However, there are ongoing discussions about what is the real reason of the δ34S shifts, biological activity or postdeposital alteration.[27]
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 the oceans in the past. Sulfate-reducing bacteria in marine sediment fractionate sulfur isotopes as they take in sulfate and produce sulfide. Prior to 2010s, it was thought that sulfate reduction could fractionate sulfur isotopes up to 46 permil[28] 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 show that sulfate-reducing bacteria can fractionate to 66 permil.[29] 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.[30]
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 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 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 reaction 35Cl + n → 35S + p is used by irradiating potassium chloride with neutrons.[31] 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.[32]
Natural occurrence
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.[33] 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.[34] The distinctive colors of Jupiter's volcanic moon Io are attributed to various forms of molten, solid, and gaseous sulfur.[35]
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.[36] Historically, Sicily was a major source of sulfur in the Industrial Revolution.[37] Lakes of molten sulfur up to about 200 m (660 ft) 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.[38]
Native sulfur is synthesised by anaerobic bacteria acting on sulfate minerals such as gypsum in salt domes.[39][40] 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.[41] 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.
The main industrial source of sulfur is now petroleum and natural gas.[9]
Common oxidation states of sulfur range from −2 to +6. Sulfur forms stable compounds with all elements except the noble gases.
Electron transfer reactions
Sulfur polycations, S82+, S42+ and S162+ are produced when sulfur is reacted with oxidising agents in a strongly acidic solution.[42] The colored solutions produced by dissolving sulfur in oleum were first reported as early as 1804 by C.F. Bucholz, but the cause of the color and the structure of the polycations involved was only determined in the late 1960s. S82+ is deep blue, S42+ is yellow and S162+ is red.[14]
Reduction of sulfur gives various polysulfides with the formula Sx2−, 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, such as S3− gives the blue color of the rock lapis lazuli.
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.[44] Ultimately, reduction of sulfur produces sulfide salts:
The interconversion of these species is exploited in the sodium–sulfur battery.
Hydrogenation
Treatment of sulfur with hydrogen gives hydrogen sulfide. When dissolved in water, hydrogen sulfide is mildly acidic:[7]
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).
Combustion
The two principal sulfur oxides are obtained by burning sulfur:
Many other sulfur oxides are observed including the sulfur-rich oxides include sulfur monoxide, disulfur monoxide, disulfur dioxides, and higher oxides containing peroxo groups.
Halogenation
Sulfur reacts with fluorine to give the highly reactive sulfur tetrafluoride and the highly inert sulfur hexafluoride.[45] 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.[46] Bromine also oxidizes sulfur to form sulfur dibromide and disulfur dibromide.[46]
Pseudohalides
Sulfur oxidizes cyanide and sulfite to give thiocyanate and thiosulfate, respectively.
Metal sulfides
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.[47] Geoscientists also study the isotopes of metal sulfides in rocks and sediment to study environmental conditions in the Earth's past.[48]
Organic compounds
- (L)-cysteine, an amino acid containing a thiol group
- Methionine, an amino acid containing a thioether
- Thiamine or vitamin B1
- Biotin or vitamin B7
- Penicillin, an antibiotic ("R" is the variable group)
- Allicin, a chemical compound in garlic
- Diphenyl disulfide, a representative disulfide
- Dibenzothiophene, a component of crude oil
- Perfluorooctanesulfonic acid (PFOS), a surfactant
Some of the main classes of sulfur-containing organic compounds include the following:[49]
- 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. Unlike carbon monoxide, carbon monosulfide is stable only as an extremely dilute gas, found between solar systems.[50]
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.[51]
Sulfur–sulfur bonds are a structural component used to stiffen rubber, similar to the disulfide bridges that rigidify proteins (see biological below). In the most common type of industrial "curing" or hardening and strengthening of natural rubber, elemental sulfur is heated with the rubber to the point that chemical reactions form disulfide bridges between isoprene units of the polymer. This process, patented in 1843, made rubber a major industrial product, especially in automobile tires. Because of the heat and sulfur, the process was named vulcanization, after the Roman god of the forge and volcanism.