# Sulfur

Sulfur
16S
O

S

Se
phosphorussulfurchlorine
Sulfur in the periodic table
Appearance
lemon yellow sintered microcrystals

Spectral lines of sulfur
General properties
Name, symbol, numbersulfur, S, 16
Pronunciation SUL-fər
Element categorypolyatomic nonmetal
Group, period, block16 (chalcogens), 3, p
Standard atomic weight32.066(1)
Electron configuration[Ne] 3s2 3p4
2, 8, 6
Physical properties
Phasesolid
Density (near r.t.)(alpha) 2.07 g·cm−3
Density (near r.t.)(beta) 1.96 g·cm−3
Density (near r.t.)(gamma) 1.92 g·cm−3
Liquid density at m.p.1.819 g·cm−3
Melting point388.36 K, 115.21 °C, 239.38 °F
Boiling point717.8 K, 444.6 °C, 832.3 °F
Critical point1314 K, 20.7 MPa
Heat of fusion(mono) 1.727 kJ·mol−1
Heat of vaporization(mono) 45 kJ·mol−1
Molar heat capacity22.75 J·mol−1·K−1
Vapor pressure
 P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 375 408 449 508 591 717
Atomic properties
Oxidation states6, 5, 4, 3, 2, 1, -1, -2
(strongly acidic oxide)
Electronegativity2.58 (Pauling scale)
Ionization energies
(more)
1st: 999.6 kJ·mol−1
2nd: 2252 kJ·mol−1
3rd: 3357 kJ·mol−1
Miscellanea
Crystal structureorthorhombic
Magnetic orderingdiamagnetic[1]
Electrical resistivity(20 °C) (amorphous)
2×1015 Ω·m
Thermal conductivity(amorphous)
0.205 W·m−1·K−1
Bulk modulus7.7 GPa
Mohs hardness2.0
CAS registry number7704-34-9
History
DiscoveryChinese[2] (Before 2000BC)
Recognized as an element byAntoine Lavoisier (1777)
Most stable isotopes
Main article: Isotopes of sulfur
isoNAhalf-lifeDMDE (MeV)DP
32S95.02%32S is stable with 16 neutrons
33S0.75%33S is stable with 17 neutrons
34S4.21%34S is stable with 18 neutrons
35Ssyn87.32 dβ0.16735Cl
36S0.02%36S is stable with 20 neutrons

(Redirected from Sulfer)
Sulfur
16S
O

S

Se
phosphorussulfurchlorine
Sulfur in the periodic table
Appearance
lemon yellow sintered microcrystals

Spectral lines of sulfur
General properties
Name, symbol, numbersulfur, S, 16
Pronunciation SUL-fər
Element categorypolyatomic nonmetal
Group, period, block16 (chalcogens), 3, p
Standard atomic weight32.066(1)
Electron configuration[Ne] 3s2 3p4
2, 8, 6
Physical properties
Phasesolid
Density (near r.t.)(alpha) 2.07 g·cm−3
Density (near r.t.)(beta) 1.96 g·cm−3
Density (near r.t.)(gamma) 1.92 g·cm−3
Liquid density at m.p.1.819 g·cm−3
Melting point388.36 K, 115.21 °C, 239.38 °F
Boiling point717.8 K, 444.6 °C, 832.3 °F
Critical point1314 K, 20.7 MPa
Heat of fusion(mono) 1.727 kJ·mol−1
Heat of vaporization(mono) 45 kJ·mol−1
Molar heat capacity22.75 J·mol−1·K−1
Vapor pressure
 P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 375 408 449 508 591 717
Atomic properties
Oxidation states6, 5, 4, 3, 2, 1, -1, -2
(strongly acidic oxide)
Electronegativity2.58 (Pauling scale)
Ionization energies
(more)
1st: 999.6 kJ·mol−1
2nd: 2252 kJ·mol−1
3rd: 3357 kJ·mol−1
Miscellanea
Crystal structureorthorhombic
Magnetic orderingdiamagnetic[1]
Electrical resistivity(20 °C) (amorphous)
2×1015 Ω·m
Thermal conductivity(amorphous)
0.205 W·m−1·K−1
Bulk modulus7.7 GPa
Mohs hardness2.0
CAS registry number7704-34-9
History
DiscoveryChinese[2] (Before 2000BC)
Recognized as an element byAntoine Lavoisier (1777)
Most stable isotopes
Main article: Isotopes of sulfur
isoNAhalf-lifeDMDE (MeV)DP
32S95.02%32S is stable with 16 neutrons
33S0.75%33S is stable with 17 neutrons
34S4.21%34S is stable with 18 neutrons
35Ssyn87.32 dβ0.16735Cl
36S0.02%36S is stable with 20 neutrons

Sulfur or sulphur is a chemical element with the symbol S and atomic number 16. It is an abundant, multivalent non-metal. Under normal conditions, sulfur atoms form cyclic octatomic molecules with chemical formula S8. Elemental sulfur is a bright yellow crystalline solid when at room temperature. Chemically, sulfur can react as either an oxidant or reducing agent. It oxidizes most metals and several nonmetals, including carbon, which leads to its negative charge in most organosulfur compounds, but it reduces several strong oxidants, such as oxygen and fluorine.

Sulfur occurs naturally as the pure element (native sulfur) and as sulfide and sulfate minerals. Elemental sulfur crystals are commonly sought after by mineral collectors for their distinct, brightly colored polyhedron shapes. Being abundant in native form, sulfur was known in ancient times, mentioned for its uses in ancient India, ancient Greece, China and Egypt. Fumes from burning sulfur were used as fumigants, and sulfur-containing medicinal mixtures were used as balms and antiparasitics. Sulfur is referred to in the Bible as brimstone (burn stone) in English, with this name still used in several nonscientific tomes.[3] It was needed to make the best quality of black gunpowder. In 1777, Antoine Lavoisier helped convince the scientific community that sulfur was a basic element rather than a compound.

Elemental sulfur was once extracted from salt domes where it sometimes occurs in nearly pure form, but this method has been obsolete since the late 20th century. Today, almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants from natural gas and petroleum. The element's commercial uses are primarily in fertilizers, because of the relatively high requirement of plants for it, and in the manufacture of sulfuric acid, a primary industrial chemical. Other well-known uses for the element are in matches, insecticides and fungicides. Many sulfur compounds are odoriferous, and the smell of odorized natural gas, skunk scent, grapefruit, and garlic is due to sulfur compounds. Hydrogen sulfide produced by living organisms imparts the characteristic odor to rotting eggs and other biological processes.

Sulfur is an essential element for all life, and is widely used in biochemical processes. In metabolic reactions, sulfur compounds serve as both fuels (electron donors) and respiratory (oxygen-alternative) materials (electron acceptors). Sulfur in organic form is present in the vitamins biotin and thiamine, the latter being named for the Greek word for sulfur. Sulfur is an important part of many enzymes and in antioxidant molecules like glutathione and thioredoxin. Organically bonded sulfur is a component of all proteins, as the amino acids cysteine and methionine. Disulfide bonds are largely responsible for the mechanical strength and insolubility of the protein keratin, found in outer skin, hair, and feathers, and the element contributes to their pungent odor when burned.

## Spelling and etymology

Sulfur or sulphur comes via Old French from Latin sulphur, which in turn is apparently formed on a root meaning "to burn".[4] The element was traditionally spelled sulphur in the United Kingdom (since the 14th century),[5] and most of the Commonwealth (including Australia, India, Malaysia, South Africa), Hong Kong, the Caribbean and Ireland. Sulfur is used in the United States, while both spellings are used in Canada and the Philippines.[5]

However, the IUPAC adopted the spelling sulfur in 1990, as did the Royal Society of Chemistry Nomenclature Committee in 1992.[6] The Qualifications and Curriculum Authority for England and Wales recommended its use in 2000,[7] and it now appears in GCSE exams.[8] The Oxford Dictionaries note that "In chemistry... the -f- spelling is now the standard form in all related words in the field in both British and US contexts."[9]

In Latin, the word is variously written sulpur, sulphur, and sulfur (the Oxford Latin Dictionary lists the spellings in this order). It is an original Latin name and not a Classical Greek loan, so the ph variant does not denote the Greek letter φ (phi). Sulfur in Greek is thion (θείον), whence comes the prefix thio-. The simplification of the Latin word's p or ph to an f appears to have taken place towards the end of the classical period.[10][11]

## Characteristics

When burned, sulfur melts to a blood-red liquid and emits a blue flame that is best observed in the dark.

### Physical properties

Sulfur forms polyatomic molecules with different chemical formulas, with the best-known allotrope being octasulfur, cyclo-S8. Octasulfur is a soft, bright-yellow solid with only a faint odor, similar to that of matches.[12] It melts at 115.21 °C, boils at 444.6 °C and sublimes easily.[3] At 95.2 °C, below its melting temperature, cyclo-octasulfur changes from α-octasulfur to the β-polymorph.[13] 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.[13] At even higher temperatures, however, the viscosity decreases as depolymerization occurs. Molten sulfur assumes a dark red color above 200 °C. The density of sulfur is about 2 g·cm−3, depending on the allotrope; all of its stable allotropes are excellent electrical insulators.

### Chemical properties

Sulfur burns with a blue flame concomitant with formation of sulfur dioxide, notable for its peculiar suffocating 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 the second ionization energies of sulfur are 999.6 and 2252 kJ·mol−1, 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−1, the magnitude of the figures caused by electron transfer between orbitals; these states are only stable with strong oxidants as fluorine, oxygen, and chlorine.

### Allotropes

The structure of the cyclooctasulfur molecule, S8.

Sulfur forms over 30 solid allotropes, more than any other element.[14] Besides S8, several other rings are known.[15] Removing one atom from the crown gives S7, which is more deeply yellow than S8. HPLC analysis of "elemental sulfur" reveals an equilibrium mixture of mainly S8, but with S7 and small amounts of S6.[16] Larger rings have been prepared, including S12 and S18.[17][18]

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 25 known isotopes, four of which are stable: 32S (95.02%), 33S (0.75%), 34S (4.21%), and 36S (0.02%). 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 170 minutes.

When sulfide minerals are precipitated, isotopic equilibration among solids and liquid may cause small differences in the δS-34 values of co-genetic minerals. The differences between minerals can be used to estimate the temperature of equilibration. The δC-13 and δS-34 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 different δ34S values from lakes believed to be dominated by watershed sources of sulfate.

### Natural occurrence

Most of the yellow and orange hues of Io are due to elemental sulfur and sulfur compounds, produced by active volcanoes.
Native sulfur crystals
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.[19] As this 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.[20] The distinctive colors of Jupiter's volcanic moon Io are attributed to various forms of molten, solid and gaseous sulfur.[21]

On 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. Such deposits are polycrystalline, with the largest documented single crystal measuring 22×16×11 cm.[22] Historically, Sicily was a large source of sulfur in the Industrial Revolution.[23]

Significant deposits of elemental sulfur, believed to have been (and are still being) synthesised by anaerobic bacteria on sulfate minerals like gypsum, exist in salt domes 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 have until recently been the basis for commercial production in the United States, Russia, Turkmenistan, and Ukraine.[24] 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 sulfates, 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.

## Production

Sulfur may be found by itself and historically was usually obtained in this way, while pyrite has been a source of sulfur via sulfuric acid.[25] 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 sulphur mine in Sicily is about the nearest thing to hell that I expect to see in this life.".[26]

Today's sulfur production is as a side product of other industrial processes such as oil refining; in these processes, sulfur often occurs as undesired or detrimental compounds that are extracted and converted to elemental sulfur. As a mineral, native sulfur under salt domes is thought to be a fossil mineral resource, produced by the action of ancient bacteria on sulfate deposits. It was removed from such salt-dome mines mainly by the Frasch process.[24] 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. However, 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.[27][28]

Sulfur recovered from hydrocarbons in Alberta, stockpiled for shipment in North Vancouver, B.C.

Today, sulfur is produced from petroleum, natural gas, and related fossil resources, from which it is obtained mainly as hydrogen sulfide. Organosulfur compounds, undesirable impurities in petroleum, may be upgraded by subjecting them to hydrodesulfurization, which cleaves the C–S bonds:[27][28]

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 the Claus process. This process entails oxidation of some hydrogen sulfide to sulfur dioxide and then the comproportionation of the two:[27][28]

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 the Athabasca Oil Sands, stockpiles of elemental sulfur from this process now exist throughout Alberta, Canada.[29] Another way of storing sulfur is as a binder for concrete, the resulting product having many desirable properties (see sulfur concrete).[30]

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), US (8.8), Canada (7.1) and Russia (7.1).[31] While the production has been slowly increasing from 1900 to 2010, the price was much less stable, especially in the 1980s and around 2010.[32]

## Compounds

Common oxidation states of sulfur range from −2 to +6. Sulfur forms stable compounds with all elements except the noble gases.

### Sulfides

Treatment of sulfur with hydrogen gives hydrogen sulfide. When dissolved in water, hydrogen sulfide is mildly acidic:[3]

H2S $\overrightarrow{\leftarrow}$ HS + H+

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).

Reduction of elemental sulfur gives polysulfides, which consist of chains of sulfur atoms terminated with S centers:

2 Na + S8 → Na2S8

This reaction highlights arguably the single most distinctive property of sulfur: its ability to catenate (bind to itself by formation of chains). Protonation of these polysulfide anions gives the polysulfanes, H2Sx where x = 2, 3, and 4.[33] Ultimately reduction of sulfur gives sulfide salts:

16 Na + S8 → 8 Na2S

The interconversion of these species is exploited in the sodium-sulfur battery. The radical anion S3 gives the blue color of the mineral lapis lazuli.

Lapis lazuli owes its blue color to a sulfur radical.

With very strong oxidants, S8 can be oxidized, for example, to give bicyclic S82+.

### Oxides and oxyanions

The principal sulfur oxides are obtained by burning sulfur:

S + O2 → SO2
2 SO2 + O2 → 2 SO3

Other oxides are known, e.g. sulfur monoxide and disulfur mono- and dioxides, but they are unstable.

The sulfur oxides form numerous oxyanions with the formula SOn2–. Sulfur dioxide and sulfites (SO2−
3
) are related to the unstable sulfurous acid (H2SO3). Sulfur trioxide and sulfates (SO2−
4
) are related to sulfuric acid. Sulfuric acid and SO3 combine to give oleum, a solution of pyrosulfuric acid (H2S2O7) in sulfuric acid.

Peroxides convert sulfur into unstable oxides such as S8O, a sulfoxide. Peroxymonosulfuric acid (H2SO5) and peroxydisulfuric acids (H2S2O8), made from the action of SO3 on concentrated H2O2, and H2SO4 on concentrated H2O2 respectively.

The sulfate anion, SO2−
4

Thiosulfate salts (S
2
O2−
3
), sometimes referred as "hyposulfites", used in photographic fixing (HYPO) and as reducing agents, feature sulfur in two oxidation states. Sodium dithionite (Na
2
S
2
O
4
), contains the more highly reducing dithionite anion (S
2
O2−
4
). Sodium dithionate (Na2S2O6) contains the dithionate anion (S2O62-) and is the first member of the polythionic acids (H2SnO6), where n can range from 3 to many. Thiosulfurous acid (HS-S(=O)-OH) is formed in trace amounts when hydrogen sulfide and sulfur dioxide gases are mixed at room temperature, but its salts (thiosulfites) are unknown.

### Halides and oxyhalides

The two main sulfur fluorides are sulfur hexafluoride, a dense gas used as nonreactive and nontoxic propellant, and sulfur tetrafluoride, a rarely used organic reagent that is highly toxic.[34] Their chlorinated analogs are sulfur dichloride and sulfur monochloride. Sulfuryl chloride and chlorosulfuric acid are derivatives of sulfuric acid; thionyl chloride (SOCl2) is a common reagent in organic synthesis.[35]

### Pnictides

An important S–N compound is the cage tetrasulfur tetranitride (S4N4). Heating this compound gives polymeric sulfur nitride ((SN)x), which has metallic properties even though it does not contain any metal atoms. Thiocyanates contain the SCN group. Oxidation of thiocyanate gives thiocyanogen, (SCN)2 with the connectivity NCS-SCN. Phosphorus sulfides are numerous, the most important commercially being the cages P4S10 and P4S3.[36][37]

### Metal sulfides

The principal ores of copper, zinc, nickel, cobalt, molybdenum, and other metals are sulfides. These materials tend to be dark-colored semiconductors that are not readily attacked by water or even many acids. They are formed, both geochemically and in the laboratory, by the reaction of hydrogen sulfide with metal salts. The mineral galena (PbS) was the first demonstrated semiconductor and found a use as a signal rectifier in the cat's whiskers of early crystal radios. The iron sulfide called pyrite, the so-called "fool's gold," has the formula FeS2.[38] The upgrading of these ores, usually by roasting, is costly and environmentally hazardous. Sulfur corrodes many metals via the process called tarnishing.

### Organic compounds

Some of the main classes of sulfur-containing organic compounds include the following:[39]

Compounds with carbon–sulfur bonds are uncommon with the notable exception of 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 only stable as a dilute gas, as in the interstellar medium.[40]

Organosulfur compounds are responsible for the some of the unpleasant odors of decaying organic matter. They are used in the odoration of natural gas and cause the odor of garlic and skunk spray. Not all organic sulfur compounds smell unpleasant at all concentrations: the sulfur-containing monoterpenoid grapefruit mercaptan in small concentrations is responsible for 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.[41]

Sulfur-sulfur bonds are a structural component to stiffen rubber, in a way similar to the biological role of disulfide bridges to 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, allowed rubber to become a major industrial product, especially automobile tires. Because of the heat and sulfur, the process was named vulcanization, after the Roman god of the forge and volcanism.

## History

### Antiquity

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

Being abundantly available in native form, sulfur (Latin sulphur) was known in ancient times and is referred to in the Torah (Genesis). English translations of the Bible commonly referred to burning sulfur as "brimstone", giving rise to the name of '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 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;[42] this is mentioned in the Odyssey.[43] 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.[44]

A natural form of sulfur known as shiliuhuang was known in China since the 6th century BC and found in Hanzhong.[45] By the 3rd century, the Chinese discovered that sulfur could be extracted from pyrite.[45] 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.[45] A Song Dynasty military treatise of 1044 AD described different formulas for Chinese black powder, which is a mixture of potassium nitrate (KNO
3
), charcoal, and sulfur.

Indian alchemists, practitioners of "the science of mercury" (sanskrit rasaśāstra, रसशास्त्र), wrote extensively about the use of sulfur in alchemical operations with mercury, from the eighth century AD onwards.[46] In the rasaśāstra tradition, sulfur is called "the smelly" (sanskrit gandhaka, गन्धक).

Early European alchemists gave sulfur its own alchemical symbol, a triangle at the top of a cross. In traditional skin treatment before the modern era of scientific medicine, elemental sulfur was used, mainly in creams, to alleviate conditions such as scabies, ringworm, psoriasis, eczema, and acne. The mechanism of action is unknown—though elemental sulfur does oxidize slowly to sulfurous acid, which in turn (through the action of sulfite) acts as a mild reducing and antibacterial agent.[47][48][49]

### Modern times

Sicilian kiln used to obtain sulfur from volcanic rock.

In 1777, Antoine Lavoisier helped convince the scientific community that sulfur was an element, not a compound. Sulfur deposits in Sicily were the dominant supply source for over half a century. Approximately 2000 tons per year of sulfur were imported into Marseilles, France for the production of sulphuric acid via the Leblanc process by the late 18th century. In industrializing Britain, with the repeal of tarrifs on 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 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 negotiated solution was eventually mediated by France.[50][51]

In 1867, sulfur was discovered in underground deposits in Louisiana and Texas. The highly successful Frasch process was developed to extract this resource.[52]

In the late 18th century, furniture makers used molten sulfur to produce decorative inlays in their craft. Because of the sulfur dioxide produced during the process of melting sulfur, the craft of sulfur inlays was soon abandoned. Molten sulfur is sometimes still used for setting steel bolts into drilled concrete holes where high shock resistance is desired for floor-mounted equipment attachment points. Pure powdered sulfur was used as a medicinal tonic and laxative.[24] With the advent of the contact process, the majority of sulfur today is used to make sulfuric acid for a wide range of uses, particularly fertilizer.[53]

## Applications

### Sulfuric acid

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

2 S + 3 O2 + 2 H2O → 2 H2SO4

With sulfuric acid being of central importance to the world's economies, its production and consumption is an indicator of a nation's industrial development.[54] For example with 32.5 million tonnes in 2010, the United States produces more sulfuric acid every year than any other inorganic industrial chemical.[32] 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.[24]

Sulfuric acid production in 2000

### Other large-scale sulfur chemicals

Sulfur reacts directly with methane to give carbon disulfide, which is used to manufacture cellophane and rayon.[24] One of the direct uses of sulfur is in vulcanization of rubber, where polysulfides crosslink organic polymers. Sulfites are heavily used to bleach paper and as preservatives in dried fruit. Many surfactants and detergents, e.g. sodium lauryl sulfate, are produced are sulfate derivatives. Calcium sulfate, gypsum, (CaSO4·2H2O) is mined on the scale of 100 million tons 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.

### Fertilizer

Sulfur is increasingly used as a component of fertilizers. The most important form of sulfur for fertilizer is the mineral calcium sulfate. Elemental sulfur is hydrophobic (that is, it is not soluble in water) and, therefore, cannot be directly utilized by plants. Over time, soil bacteria can convert it to soluble derivatives, which can then be utilized by plants. Sulfur improves the use efficiency of other essential plant nutrients, particularly nitrogen and phosphorus.[55] Biologically produced sulfur particles are naturally hydrophilic due to a biopolymer coating. This sulfur is, therefore, easier to disperse over the land (via spraying as a diluted slurry), and results in a faster release.

Plant requirements for sulfur are equal to or exceed those for phosphorus. It is one of the major nutrients essential for plant growth, root nodule formation of legumes and plants protection mechanisms. Sulfur deficiency has become widespread in many countries in Europe.[56][57][58] Because atmospheric inputs of sulfur continue to decrease, the deficit in the sulfur input/output is likely to increase, unless sulfur fertilizers are used.

### Fine chemicals

A molecular model of the pesticide malathion.

Organosulfur compounds are used in pharmaceuticals, dyestuffs, and agrochemicals. Many drugs contain sulfur, early examples being antibacterial sulfonamides, known as sulfa drugs. Sulfur is a part of many bacterial defense molecules. Most β-lactam antibiotics, including the penicillins, cephalosporins and monolactams contain sulfur.[39]

Magnesium sulfate, known as Epsom salts when in hydrated crystal form, can be used as a laxative, a bath additive, an exfoliant, magnesium supplement for plants, or (when in dehydrated form) as a desiccant.

### Fungicide and pesticide

Sulfur candle originally sold for home fumigation

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 be used well 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.[59][60] 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 use is to dust clothing or limbs with sulfur powder.

Diluted solutions of lime sulfur (made by combinding calcium hydroxide with elemental sulfur in water), are used as a dip for pets to destroy ringworm (fungus), mange and other dermatoses and parasites. Sulfur candles consist of almost pure sulfur in blocks or pellets that are burned to fumigate structures. It is no longer used in the home due to the toxicity of the products of combustion.

### Bactericide in winemaking and food preservation

Small amounts of sulfur dioxide gas addition (or equivalent potassium metabisulfite addition) to fermented wine to produce traces of sulfurous acid (produced when SO2 reacts with water) and its sulfite salts in the mixture, has been called "the most powerful tool in winemaking."[61] After the yeast-fermentation stage in winemaking, sulfites absorb oxygen and inhibit aerobic bacterial growth that otherwise would turn ethanol into acetic acid, souring the wine. Without this preservative step, indefinite refrigeration of the product before consumption is usually required. Similar methods go back into antiquity but modern historical mentions of the practice go to the fifteenth century. The practice is used by large industrial wine producers and small organic wine producers alike.

Sulfur dioxide and various sulfites have been used for their antioxidant antibacterial preservative properties in many other parts of the food industry also. The practice has declined since reports of an allergy-like reaction of some persons to sulfites in foods.

## Biological role

### Protein and organic cofactors

Sulfur is an essential component of all living cells. It is the seventh or eighth most abundant element in the human body by weight, being about as common as potassium, and a little more common than sodium or chlorine. A 70 kg human body contains about 140 grams of sulfur.

In plants and animals, the amino acids cysteine and methionine contain most of the sulfur. The element is thus present in all polypeptides, proteins, and enzymes that contain these amino acids. In humans, methionine is an essential amino acid that must be ingested. However, save for the vitamins biotin and thiamine, cysteine and all sulfur-containing compounds in the human body can be synthesized from methionine. The enzyme sulfite oxidase is needed for the metabolism of methionine and cysteine in humans and animals.

Disulfide bonds (S-S bonds) formed between cysteine residues in peptide chains are very important in protein assembly and structure. These covalent bonds between peptide chains confer extra toughness and rigidity.[62] For example, the high strength of feathers and hair is in part due to their high content of S-S bonds and their high content of cysteine and sulfur. Eggs are high in sulfur because large amounts of the element are necessary for feather formation, and the characteristic odor of rotting eggs is due to hydrogen sulfide. The high disulfide bond content of hair and feathers contributes to their indigestibility and to their characteristic disagreeable odor when burned.

Homocysteine and taurine are other sulfur-containing acids that are similar in structure, but not coded by DNA, and are not part of the primary structure of proteins. Many important cellular enzymes use prosthetic groups ending with -SH moieties to handle reactions involving acyl-containing biochemicals: two common examples from basic metabolism are coenzyme A and alpha-lipoic acid.[62] Two of the 13 classical vitamins, biotin and thiamine contain sulfur, with the latter being named for its sulfur content. Sulfur plays an important part, 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. The thioredoxins, a class of small protein essential to all known life, using neighboring pairs of reduced cysteines to act as general protein reducing agents, to similar effect.

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, CH3SCH2CH2SO3, the immediate precursor to methane.[63]

### Metalloproteins and inorganic cofactors

Inorganic sulfur forms a part of iron-sulfur clusters as well as many copper, nickel, and iron proteins. Most pervasive are the ferrodoxins, which serve as electron shuttles in cells. In bacteria, the important nitrogenase enzymes contains an Fe-Mo-S cluster, 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.[64]

### Sulfur metabolism and the sulfur cycle

The sulfur cycle was the first of the biogeochemical cycles to be discovered. 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). He continued to study it together with Selman Waksman until the 1950s.

Sulfur oxidizers can use as energy sources reduced sulfur compounds, including hydrogen sulfide, elemental sulfur, sulfite, thiosulfate, and various polythionates (e.g., tetrathionate).[65] 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 utilizes oxygen as the electron acceptor. The photosynthetic green sulfur bacteria and purple sulfur bacteria and some lithotrophs use elemental oxygen to carry out such oxidization of hydrogen sulfide to produce elemental sulfur (S0), oxidation state = 0. Primitive bacteria that live around deep ocean volcanic vents oxidize hydrogen sulfide in this way with oxygen; see giant tube worm for an example of large organisms that use hydrogen sulfide (via bacteria) as food to be oxidized.

The so-called sulfate-reducing bacteria, 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 a number of other partially oxidized sulfur compounds (e.g. thiosulfates, thionates, polysulfides, sulfites). The hydrogen sulfide produced by these bacteria is responsible for some of the smell of intestinal gases (flatus) and decomposition products.

Sulfur is absorbed by plants via the roots from soil as the sulfate and transported as a phosphate ester. Sulfate is reduced to sulfide via sulfite before it is incorporated into cysteine and other organosulfur compounds.[66]

SO42– → SO32– → H2S → cysteine → methionine

## Precautions

Effect of acid rain on a forest, Jizera Mountains, Czech Republic

Elemental sulfur is non-toxic, as generally are the soluble sulfate salts, such as Epsom salts. Soluble sulfate salts are poorly absorbed and laxative. However, when injected parenterally, they are freely filtered by the kidneys and eliminated with very little toxicity in multi-gram amounts.

When sulfur burns in air, it produces sulfur dioxide. In water, this gas produces sulfurous acid and sulfites, which are antioxidants that inhibit growth of aerobic bacteria and allow its use as a food additive in small amounts. However, at high concentrations these acids harm the lungs, eyes or other tissues. In organisms without lungs such as insects or plants, it otherwise prevents respiration in high concentrations. Sulfur trioxide (made by catalysis from sulfur dioxide) and sulfuric acid are similarly highly corrosive, due to the strong acids that form on contact with water.

The burning of coal and/or petroleum by industry and power plants generates sulfur dioxide (SO2), which reacts with atmospheric water and oxygen to produce sulfuric acid (H2SO4) and sulfurous acid (H2SO3). These acids are components of acid rain, which lower 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 as toxic as hydrogen cyanide, and kills by the same mechanism, though hydrogen sulfide is less likely to cause surprise poisonings from small inhaled amounts, because of its disagreeable warning odor. Though pungent at first, however, hydrogen sulfide quickly deadens the sense of smell—so a victim may breathe increasing quantities and be unaware of its presence until severe symptoms occur, which can quickly lead to death. Dissolved sulfide and hydrosulfide salts are also toxic by the same mechanism.

## References

1. ^
2. ^ "Sulfur History". Georgiagulfsulfur.com. Retrieved 2008-09-12.
3. ^ a b c Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.), Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.
4. ^ Online Etymology Dictionary entry for sulfur. Retrieved 2011-08-18.
5. ^ a b Michie, C. A.; Langslow, D. R. (1988). "Sulphur or sulfur? A tale of two spellings". British Medical Journal 297 (6664): 1697–1699. doi:10.1136/bmj.297.6664.1697.
6. ^ McNaught, Alan (1991). "Journal style update". The Analyst 116 (11): 1094. Bibcode:1991Ana...116.1094M. doi:10.1039/AN9911601094.
7. ^ Sulphur, Worldwidewords
8. ^ "General Certificate of Secondary Education (Science A Chemistry". Foundation Tier and Higher Tier. March 2010. Retrieved 2011-02-27.
9. ^ "Ask Oxford". Retrieved 2010-11-12.
10. ^ "Sulphuricum Sulphur". Vanderkrogt.net. Retrieved 2011-02-27.
11. ^ Kelly, Donovan P. (1995). "Sulfur and its Doppelgänger". Archives of Microbiology 163 (3): 157–158. doi:10.1007/BF00305347.
12. ^ A strong odor called "smell of sulfur" actually is given off by several sulfur compounds, such as hydrogen sulfide and organosulfur compounds.
13. ^ a b Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. pp. 645–662. ISBN 0080379419.
14. ^ Steudel, Ralf; Eckert, Bodo (2003). "Solid Sulfur Allotropes Sulfur Allotropes". Topics in Current Chemistry. Topics in Current Chemistry 230: 1–80. doi:10.1007/b12110. ISBN 978-3-540-40191-9.
15. ^ Steudel, R. (1982). "Homocyclic Sulfur Molecules". Topics in Current Chemistry. Topics in Current Chemistry 102: 149–176. doi:10.1007/3-540-11345-2_10. ISBN 978-3-540-11345-4.
16. ^ Tebbe, Fred N.; Wasserman, E.; Peet, William G.; Vatvars, Arturs; Hayman, Alan C. (1982). "Composition of Elemental Sulfur in Solution: Equilibrium of S
6
, S7, and S8 at Ambient Temperatures". Journal of the American Chemical Society 104 (18): 4971–4972. doi:10.1021/ja00382a050.
17. ^ Meyer, Beat (1964). "Solid Allotropes of Sulfur". Chemical Reviews 64 (4): 429–451. doi:10.1021/cr60230a004.
18. ^ Meyer, Beat (1976). "Elemental sulfur". Chemical Reviews 76 (3): 367–388. doi:10.1021/cr60301a003.
19. ^ Cameron, A. G. W. (1957). "Stellar Evolution, Nuclear Astrophysics, and Nucleogenesis". CRL-41.
20. ^ Mason, B. (1962). Meteorites. New York: John Wiley & Sons. p. 160. ISBN 0-908678-84-3.
21. ^ Lopes, Rosaly M. C.; Williams, David A. (2005). "Io after Galileo". Reports on Progress in Physics 68 (2): 303–340. Bibcode:2005RPPh...68..303L. doi:10.1088/0034-4885/68/2/R02.
22. ^ Rickwood, P. C. (1981). "The largest crystals". American Mineralogist 66: 885–907.
23. ^ Kutney, Gerald (2007). Sulfur: history, technology, applications & industry. Toronto: ChemTec Publications. p. 43. ISBN 978-1-895198-37-9. OCLC 79256100.
24. Nehb, Wolfgang; Vydra, Karel (2006). Sulfur. "Ullmann's Encyclopedia of Industrial Chemistry". Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH Verlag. doi:10.1002/14356007.a25_507.pub2. ISBN 3-527-30673-0.
25. ^ Riegel, Emil; Kent, James (2007). Kent and Riegel's handbook of industrial chemistry and biotechnology, Volume 1. New York: Springer. p. 1171. ISBN 978-0-387-27842-1. OCLC 74650396.
26. ^ Washington, Booker T. The Man Farthest Down: A Record of Observation and Study in Europe.
27. ^ a b c Eow, John S. (2002). "Recovery of sulfur from sour acid gas: A review of the technology". Environmental Progress 21 (3): 143–162. doi:10.1002/ep.670210312.
28. ^ a b c Schreiner, Bernhard (2008). "Der Claus-Prozess. Reich an Jahren und bedeutender denn je". Chemie in unserer Zeit 42 (6): 378–392. doi:10.1002/ciuz.200800461.
29. ^ Hyndman, A. W.; Liu, J. K.; Denney, D. W. (1982). "Sulfur Recovery from Oil Sands". Sulfur: New Sources and Uses. ACS Symposium Series 183. pp. 69–82. doi:10.1021/bk-1982-0183.ch005. ISBN 0-8412-0713-5.
30. ^ Mohamed, Abdel-Mohsen; El-Gamal, Maisa (2010). Sulfur concrete for the construction industry: a sustainable development approach. Fort Lauderdale: J. Ross Publishing. p. 109. ISBN 978-1-60427-005-1. OCLC 531718953.
31. ^ Apodaca, Lori E. (2012) Sulfur. Mineral Commodity Summaries. USGS
32. ^ a b Apodaca, Lori E. "Mineral Yearbook 2010: Sulfur". United States Geological Survey.
33. ^ Handbook of Preparative Inorganic Chemistry, 2nd ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 421.
34. ^ Hasek, W. R. (1961), "1,1,1-Trifluoroheptane", Org. Synth. 41: 104
35. ^ Rutenberg, M. W; Horning, E. C. (1950), "1-Methyl-3-ethyloxindole", Org. Synth. 30: 62
36. ^ Heal, H. G. (1980). The Inorganic Heterocyclic Chemistry of Sulfur, Nitrogen, and Phosphorus. London: Academic Press. ISBN 0-12-335680-6.
37. ^ Chivers, T. (2004). A Guide To Chalcogen-Nitrogen Chemistry. Singapore: World Scientific. ISBN 981-256-095-5.
38. ^ Vaughan, D. J.; Craig, J. R. "Mineral Chemistry of Metal Sulfides" Cambridge University Press, Cambridge (1978) ISBN 0-521-21489-0
39. ^ a b Cremlyn R. J.; "An Introduction to Organosulfur Chemistry" John Wiley and Sons: Chichester (1996). ISBN 0-471-95512-4.
40. ^ Wilson, R. W.; Penzias, A. A.; Wannier, P. G.; Linke, R. A. (March 15, 1976). "Isotopic abundances in interstellar carbon monosulfide". Astrophysical Journal 204: L135–L137. Bibcode:1976ApJ...204L.135W. doi:10.1086/182072.
41. ^ Banoub, Joseph (2011). Detection of Biological Agents for the Prevention of Bioterrorism. Dordrecht: Springer. p. 183. ISBN 978-90-481-9815-3. OCLC 697506461.
42. ^ Rapp, George Robert (2009-02-04). Archaeomineralogy. p. 242. ISBN 978-3-540-78593-4.
43. ^ Odyssey, book 22, lines 480–495 • www.perseus.tufts.edu. Retrieved on 2012-08-16.
44. ^ Pliny the Elder on science and technology, John F. Healy, Oxford University Press, 1999, ISBN 0-19-814687-6, pp. 247–249.
45. ^ a b c Zhang, Yunming (1986). "The History of Science Society: Ancient Chinese Sulfur Manufacturing Processes". Isis 77 (3): 487. doi:10.1086/354207.
46. ^ White, David Gordon (1996). The Alchemical Body — Siddha Traditions in Medieval India. Chicago: University of Chicago Press. pp. passim. ISBN 978-0-226-89499-7.
47. ^ Lin, A. N.; Reimer, R. J.; Carter, D. M. (1988). "Sulfur revisited". Journal of the American Academy of Dermatology 18 (3): 553–558. doi:10.1016/S0190-9622(88)70079-1. PMID 2450900.
48. ^ Maibach, HI; Surber, C.; Orkin, M. (1990). "Sulfur revisited". Journal of the American Academy of Dermatology 23 (1): 154–156. doi:10.1016/S0190-9622(08)81225-X. PMID 2365870.
49. ^ Gupta, A. K.; Nicol, K. (2004). "The use of sulfur in dermatology". Journal of drugs in dermatology : JDD 3 (4): 427–31. PMID 15303787.
50. ^ Riall, Lucy (1998). Sicily and the Unification of Italy: Liberal Policy and Local Power, 1859–1866. Oxford University Press. ISBN 9780191542619. Retrieved 2013-02-07.
51. ^ Thomson, D. W. (April 1995). "Prelude to the Sulphur War of 1840: The Neapolitan Perspective". European History Quarterly 25 (2): 163–180. doi:10.1177/026569149502500201.
52. ^ Botsch, Walter (2001). "Chemiker, Techniker, Unternehmer: Zum 150. Geburtstag von Hermann Frasch". Chemie in unserer Zeit (in German) 35 (5): 324–331. doi:10.1002/1521-3781(200110)35:5<324::AID-CIUZ324>3.0.CO;2-9.
53. ^ Kogel, Jessica (2006). Industrial minerals & rocks: commodities, markets, and uses (7th ed.). Colorado: Littleton. p. 935. ISBN 978-0-87335-233-8. OCLC 62805047.
54. ^ Sulfuric Acid Growth. Pafko.com. Retrieved on 2012-08-16.
55. ^ Sulfur as a fertilizer. Sulphurinstitute.org. Retrieved on 2012-08-16.
56. ^ Zhao, F.; Hawkesford, MJ; McGrath, SP (1999). "Sulphur Assimilation and Effects on Yield and Quality of Wheat". Journal of Cereal Science 30 (1): 1–17. doi:10.1006/jcrs.1998.0241.
57. ^ Blake-Kalff, M. M. A. (2000). Plant and Soil 225 (1/2): 95–107. doi:10.1023/A:1026503812267.
58. ^ Ceccotti, S. P. (1996). "Plant nutrient sulphur-a review of nutrient balance, environmental impact and fertilizers". Fertilizer Research 43 (1–3): 117–125. doi:10.1007/BF00747690.
59. ^ Mohamed, Abdel-Mohsen Onsy; El Gamal, M. M (2010-07-13). Sulfur Concrete for the Construction Industry: A Sustainable Development Approach. pp. 104–105. ISBN 978-1-60427-005-1.
60. ^ Every, Richard L., et al. (1968-08-20). "Method for Preparation of Wettable Sulfur". Retrieved 2010-05-20.
61. ^ Spencer, Benjamin Sulfur in wine demystified. intowine.com. Retrieved Oct 26, 2011.
62. ^ a b Nelson, D. L.; Cox, M. M. (2000). Lehninger, Principles of Biochemistry (3rd ed.). New York: Worth Publishing. ISBN 1-57259-153-6.
63. ^ Thauer, R. K. (1998). "Biochemistry of methanogenesis: a tribute to Marjory Stephenson:1998 Marjory Stephenson Prize Lecture". Microbiology 144 (9): 2377–2406. doi:10.1099/00221287-144-9-2377. PMID 9782487.
64. ^ Lippard, S. J.; Berg, J. M. (1994). Principles of Bioinorganic Chemistry. University Science Books. ISBN 0-935702-73-3.
65. ^ Pronk JT, Meulenberg R, Hazeu W, Bos P, Kuenen JG (1990). "Oxidation of reduced inorganic sulphur compounds by acidophilic thiobacilli". FEMS Microbiology letters 75 (2–3): 293–306. doi:10.1111/j.1574-6968.1990.tb04103.x.
66. ^ Heldt, Hans-Walter (1996). Pflanzenbiochemie. Heidelberg: Spektrum Akademischer Verlag. pp. 321–333. ISBN 3-8274-0103-8.