Sulfur

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

Sulfur, 16S
Sulfur
Alternative nameSulphur (British spelling)
Allotropessee Allotropes of sulfur
AppearanceLemon yellow sintered microcrystals
Standard atomic weight Ar°(S)
Sulfur in the periodic 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
kJ/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
at T (K) 375 408 449 508 591 717
Atomic properties
Discovery
before 2000 BCE[6]
Recognized as an element byAntoine 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 spelled sulphur in

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,

organosulfur compounds. Hydrogen sulfide
gives the characteristic odor to rotting eggs and other biological processes.

Sulfur is an

macronutrient
for all living organisms.

Characteristics

Physical properties

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 is

polymorph.[13] The structure of the S8 ring 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-S8 molecules.[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 increased viscosity due to the formation of polymers.[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.

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

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:

1⁄2 S
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.[19]

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).[20] Some of those reactions need elevated temperatures.[21]

Allotropes

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

Sulfur forms over 30 solid allotropes, 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. HPLC analysis of "elemental sulfur" reveals an equilibrium mixture of mainly S8, but with S7 and small amounts of S6.[24] Larger rings have been prepared, including S12 and S18.[25][26]

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

Main article: Isotopes of sulfur

Sulfur has 23 known

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

supernovae.[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 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.[30] However, there are ongoing discussions about what is the real reason of the δ34S shifts, biological activity or postdeposital alteration.[31]

For example, when

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

permil[32] 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.[33] 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.[34]

In most

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

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.[36]

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.
Sulfur extraction, East Java
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.[37] 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.[38] The distinctive colors of Jupiter's volcanic moon Io are attributed to various forms of molten, solid, and gaseous sulfur.[39]

It is the fifth most common element by mass in the Earth. Elemental sulfur can be found near

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.[40] Historically, Sicily was a major source of sulfur in the Industrial Revolution.[41] 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.[42]

Native sulfur is synthesised by

sulfate minerals such as gypsum in salt domes.[43][44] 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.[45] 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

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]

Compounds

Main article: Sulfur compounds

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

Electron transfer reactions

Lapis lazuli owes its blue color to a trisulfur radical anion (S−
3)

Sulfur polycations, S82+, S42+ and S162+ are produced when sulfur is reacted with oxidising agents in a strongly acidic solution.[46] 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.[13]

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:

4 Na + S8 → 2 Na2S4

Some of these dianions dissociate to give radical anions, such as S3− gives the blue color of the rock lapis lazuli.

Two parallel sulfur chains grown inside a single-wall carbon nanotube (CNT, a). Zig-zag (b) and straight (c) S chains inside double-wall CNTs[47]

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.[48] Ultimately, reduction of sulfur produces sulfide salts:

16 Na + S8 → 8 Na2S

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]

H2S ⇌ 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).

Combustion

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 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.[49] 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.[50] Bromine also oxidizes sulfur to form sulfur dibromide and disulfur dibromide.[50]

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.[51] Geoscientists also study the isotopes of metal sulfides in rocks and sediment to study environmental conditions in the Earth's past.[52]

Organic compounds

Main article:
Organosulfur compounds

Some of the main classes of sulfur-containing organic compounds include the following:[53]

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.[54]

Organosulfur compounds are responsible for some of the unpleasant odors of decaying organic matter. They are widely known as the

Sulfur mustard, a potent vesicant, was used in World War I as a disabling agent.[55]

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
.

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 was known in ancient times and is referred to in the

Melos. He mentions its use for fumigation, medicine, and bleaching cloth.[59]

A natural form of sulfur known as shiliuhuang (石硫黄) was known in China since the 6th century BC and found in

black powder, which is a mixture of potassium nitrate (KNO
3), charcoal, and sulfur.[61]

Sulfur
Brimstone
Alchemical signs for sulfur, or the combustible elements, and brimstone, an older/archaic name for sulfur[62]

Indian alchemists, practitioners of the "science of chemicals" (

rasaśāstra
tradition, sulfur is called "the smelly" (गन्धक, gandhaka).

Early

eczema, and acne. The mechanism of action is unknown—though elemental sulfur does oxidize slowly to sulfurous acid, which is (through the action of sulfite) a mild reducing and antibacterial agent.[64][65][66]

Modern times

Today sulfur is known to have antifungal, antibacterial, and keratolytic activity; in the past it was used against acne vulgaris, rosacea, seborrheic dermatitis, dandruff, pityriasis versicolor, scabies, and warts.[67] This 1881 advertisement baselessly claims efficacy against rheumatism, gout, baldness, and graying of hair.

Sulfur appears in a column of fixed (non-acidic) alkali in a chemical table of 1718.[68] Antoine Lavoisier used sulfur in combustion experiments, writing of some of these in 1777.[69]

Sulfur deposits in

Sulfur Crisis of 1840, when King Ferdinand II gave a monopoly of the sulfur industry to a French firm, violating an earlier 1816 trade agreement with Britain. A peaceful solution was eventually negotiated by France.[70][71]

In 1867, elemental sulfur was discovered in underground deposits in Louisiana and Texas. The highly successful Frasch process was developed to extract this resource.[72]

In the late 18th century, furniture makers used molten sulfur to produce decorative inlays.[73] 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.[45]

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.[74]

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

Spelling and etymology

Sulfur is derived from the Latin word sulpur, which was Hellenized to sulphur in the erroneous belief that the Latin word came from Greek. This spelling was later reinterpreted as representing an /f/ sound and resulted in the spelling sulfur, which appears in Latin toward the end of the Classical period. The true Ancient Greek word for sulfur, θεῖον, theîon (from earlier θέειον, théeion), is the source of the international chemical prefix thio-. The Modern Standard Greek word for sulfur is θείο, theío.

In 12th-century Anglo-French, it was sulfre. In the 14th century, the erroneously Hellenized Latin -ph- was restored in Middle English sulphre. By the 15th century, both full Latin spelling variants sulfur and sulphur became common in English. The parallel f~ph spellings continued in Britain until the 19th century, when the word was standardized as sulphur.[75] On the other hand, sulfur was the form chosen in the United States, whereas Canada uses both.

The IUPAC adopted the spelling sulfur in 1990 or 1971, depending on the source cited,[76] as did the Nomenclature Committee of the Royal Society of Chemistry in 1992, restoring the spelling sulfur to Britain.[77] Oxford Dictionaries note that "in chemistry and other technical uses ... the -f- spelling is now the standard form for this and related words in British as well as US contexts, and is increasingly used in general contexts as well."[78]

Production

Sicilian kiln used to obtain sulfur from volcanic rock (diagram from a 1906 chemistry book)
Traditional sulfur mining at Ijen 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; pyrite has also been a source of sulfur.[79] In volcanic regions in Sicily, in ancient times, it was found on the surface of the Earth, and the "Sicilian process" was used: sulfur deposits were piled and stacked in brick kilns built on sloping hillsides, with airspaces between them. Then, some sulfur was pulverized, spread over the stacked ore and ignited, causing the free sulfur to melt down the hills. Eventually the surface-borne deposits played out, and miners excavated veins that ultimately dotted the Sicilian landscape with labyrinthine mines. Mining was unmechanized and labor-intensive, with pickmen freeing the ore from the rock, and mine-boys or carusi carrying baskets of ore to the surface, often through a mile or more of tunnels. Once the ore was at the surface, it was reduced and extracted in smelting ovens. The conditions in Sicilian sulfur mines were horrific, prompting Booker T. Washington to write "I am not prepared just now to say to what extent I believe in a physical hell in the next world, but a sulfur mine in Sicily is about the nearest thing to hell that I expect to see in this life."[80] Sulfur is still mined from surface deposits in poorer nations with volcanoes, such as Indonesia, and worker conditions have not improved much since Booker T. Washington's days.[81]

Elemental sulfur was extracted from salt 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 the Frasch process.[45] 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.[82][83]

Sulfur recovered from hydrocarbons in Alberta, stockpiled for shipment in North Vancouver, British Columbia

Today, sulfur is produced from petroleum,

Organosulfur compounds, undesirable impurities in petroleum, may be upgraded by subjecting them to hydrodesulfurization, which cleaves the C–S bonds:[82][83]

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:[82][83]

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.[84] Another way of storing sulfur is as a binder for concrete, the resulting product having some desirable properties (see sulfur concrete).[85]

The world production of sulfur in 2011 amounted to 69 million tonnes (Mt), with more than 15 countries contributing more than 1 Mt each. Countries producing more than 5 Mt are China (9.6), the United States (8.8), Canada (7.1) and Russia (7.1).[86] Production has been slowly increasing from 1900 to 2010; the price was unstable in the 1980s and around 2010.[87]

Applications

Sulfuric acid

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

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

In 2010, the United States produced more sulfuric acid than any other inorganic industrial chemical.[87] 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.[45]

Other important sulfur chemistry

Sulfur reacts directly with methane to give

sodium lauryl sulfate) are sulfate derivatives. Calcium sulfate, gypsum (CaSO4·2H2O) is mined on the scale of 100 million tonnes each year for use in Portland cement
and fertilizers.

When silver-based photography was widespread, sodium and ammonium thiosulfate were widely used as "fixing agents". Sulfur is a component of gunpowder ("black powder").

Fertilizer

atmospheric oxygen and soil bacteria can oxidize and convert elemental sulfur to soluble derivatives, which can then be used by microorganisms and plants. Sulfur improves the efficiency of other essential plant nutrients, particularly nitrogen and phosphorus.[88] Biologically produced sulfur particles are naturally hydrophilic due to a biopolymer
coating and are easier to disperse over the land in a spray of diluted slurry, resulting in a faster uptake by plants.

The plants requirement for sulfur equals or exceeds the requirement for phosphorus. It is an essential nutrient for plant growth, root nodule formation of legumes, and immunity and defense systems. Sulfur deficiency has become widespread in many countries in Europe.[89][90][91] Because atmospheric inputs of sulfur continue to decrease, the deficit in the sulfur input/output is likely to increase unless sulfur fertilizers are used. Atmospheric inputs of sulfur decrease because of actions taken to limit acid rains.[92][88]

Fungicide and pesticide

Sulfur candle originally sold for home fumigation

Elemental sulfur is one of the oldest fungicides and

organically farmed apple production against the main disease apple scab
under colder conditions. Biosulfur (biologically produced elemental sulfur with hydrophilic characteristics) can also be used for these applications.

Standard-formulation dusting sulfur is applied to crops with a sulfur duster or from a dusting plane. Wettable sulfur is the commercial name for dusting sulfur formulated with additional ingredients to make it water miscible.[85][93] It has similar applications and is used as a fungicide against mildew and other mold-related problems with plants and soil.

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

A diluted solution of

dermatoses and parasites
.

Sulfur candles of almost pure sulfur were burned to

fumigate
structures and wine barrels, but are now considered too toxic for residences.

Pharmaceuticals

Main article:
Sulfur (pharmacy)

Sulfur (specifically

acne rosacea, and seborrhoeic dermatitis.[95]

Many drugs contain sulfur.

cephalosporins and monobactams contain sulfur.[53]

Batteries

Due to their high energy density and the availability of sulfur, there is ongoing research in creating rechargeable lithium–sulfur batteries. Until now, carbonate electrolytes have caused failures in such batteries after a single cycle. In February 2022, researchers at Drexel University have not only created a prototypical battery that lasted 4000 recharge cycles, but also found the first monoclinic gamma sulfur that remained stable below 95 degrees Celsius.[97]

Biological role

Sulfur is an essential component of all living cells. It is the eighth most abundant element in the human body by weight,[98] about equal in abundance to potassium, and slightly greater than sodium and chlorine.[99] A 70 kg (150 lb) human body contains about 140 grams (4.9 oz) of sulfur.[100] The main dietary source of sulfur for humans is sulfur-containing amino-acids,[101] which can be found in plant and animal proteins.[102]

Transferring sulfur between inorganic and biomolecules

See also: Sulfur cycle and Sulfur metabolism

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).[103] Another contributor, who continued to study it was Selman Waksman.[104] Primitive bacteria that live around deep ocean volcanic vents oxidize hydrogen sulfide for their nutrition, as discovered by Robert Ballard.[10]

Sulfur oxidizers can use as energy sources reduced sulfur compounds, including hydrogen sulfide, elemental sulfur, sulfite, thiosulfate, and various polythionates (e.g., tetrathionate).[105] They depend on enzymes such as sulfur oxygenase and sulfite oxidase to oxidize sulfur to sulfate. Some lithotrophs can even use the energy contained in sulfur compounds to produce sugars, a process known as chemosynthesis. Some bacteria and archaea use hydrogen sulfide in place of water as the electron donor in chemosynthesis, a process similar to photosynthesis that produces sugars and uses oxygen as the electron acceptor. Sulfur-based chemosynthesis may be simplifiedly compared with photosynthesis:

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

There are bacteria combining these two ways of nutrition:

giant tube worm.[107]

There are

sulfate-reducing bacteria
, that, by contrast, "breathe sulfate" instead of oxygen. They use organic compounds or molecular hydrogen as the energy source. They use sulfur as the electron acceptor, and reduce various oxidized sulfur compounds back into sulfide, often into hydrogen sulfide. They can grow on other partially oxidized sulfur compounds (e.g. thiosulfates, thionates, polysulfides, sulfites).

There are studies pointing that many deposits of native sulfur in places that were the bottom of the ancient oceans have biological origin.[108][109][110] These studies indicate that this native sulfur have been obtained through biological activity, but what is responsible for that (sulfur-oxidizing bacteria or sulfate-reducing bacteria) is still unknown for sure.

Sulfur is absorbed by plants roots from soil as sulfate and transported as a phosphate ester. Sulfate is reduced to sulfide via sulfite before it is incorporated into cysteine and other organosulfur compounds.[111]

SO42− → SO32− → H2S → cysteine (thiol) → methionine (thioether)

While the plants' role in transferring sulfur to animals by food chains is more or less understood, the role of sulfur bacteria is just getting investigated.[112][113]

Protein and organic metabolites

In all forms of life, most of the sulfur is contained in two proteinogenic amino acids (cysteine and methionine), thus the element is present in all proteins that contain these amino acids, as well as in respective peptides.[114] Some of the sulfur is comprised in certain metabolites—many of which are cofactors—and sulfated polysaccharides of connective tissue (chondroitin sulfates, heparin).

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

Proteins, to execute their biological function, need to have specific space geometry. Formation of this geometry is performed in a process called protein folding, and is provided by intra- and inter-molecular bonds. The process has several stages. While at premier stages a polypeptide chain folds due to hydrogen bonds, at later stages folding is provided (apart from hydrogen bonds) by covalent bonds between 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 in insulin. As the bond energy of a covalent disulfide bridge is higher than the energy of a coordinate bond or hydrophobic interaction, higher disulfide bridges content leads to higher energy needed for protein denaturation. 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.[115] Within cytoplasm cysteine residues of proteins are saved in reduced state (i.e. in -SH form) by thioredoxins.[116]

This property manifests in following examples. Lysozyme is stable enough to be applied as a drug.[117] Feathers and hair have relative strength, and consisting in them keratin is considered indigestible by most organisms. However, there are fungi and bacteria containing keratinase, 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 are

primary structure of proteins, take part in various locations of mammalian physiology.[119][120] Two of the 13 classical vitamins, biotin and thiamine, contain sulfur, and serve as cofactors to several enzymes.[121][122]
In intracellular chemistry, sulfur operates as a carrier of reducing hydrogen and its electrons for cellular repair of oxidation. Reduced glutathione, a sulfur-containing tripeptide, is a reducing agent through its sulfhydryl (-SH) moiety derived from cysteine.

Methanogenesis, the route to most of the world's methane, is a multistep biochemical transformation of carbon dioxide. This conversion requires several organosulfur cofactors. These include coenzyme M, CH3SCH2CH2SO3−, the immediate precursor to methane.[123]

Metalloproteins and inorganic cofactors

Metalloproteins—in which the active site is a transition metal ion (or metal-sulfide cluster) often coordinated by sulfur atoms of cysteine residues

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.[126]

Sulfur is also present in molybdenum cofactor.[127]

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

Sulfate

Main article: Sulfation § Sulfation in biology

Deficiency

In humans methionine is an essential amino acid; cysteine is conditionally essential and may be synthesized from non-essential serine (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,[128] but the method is still potentially dangerous.[129]

Isolated sulfite oxidase deficiency is a rare, fatal genetic disease preventing production of sulfite oxidase, needed to metabolize sulfites to sulfates.[130]

Precautions

Sulfur
Hazards
GHS labelling:
GHS07: Exclamation mark GHS02: Flammable
Warning
H315[131]
NFPA 704 (fire diamond)
[132]
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,[133] and cases of life-threatening metabolic acidosis have been reported after patients deliberately consumed sulfur as a folk remedy.[134][135]

Toxicity of sulfur compounds

When sulfur burns in air, it produces

tissues.[136] In organisms without lungs such as insects, sulfite in high concentration prevents respiration.[137]

Sulfur trioxide (made by catalysis from sulfur dioxide) and sulfuric acid are similarly highly acidic and corrosive in the presence of water. Concentrated sulfuric acid is a strong dehydrating agent that can strip available water molecules and water components from sugar and organic tissue.[138]

The burning of

synthesis gas
extract the sulfur before they burn the gas.

hydrosulfide
salts are toxic by the same mechanism.

Notes

  1. organosulfur
    compounds.
  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-S8 the 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 a polymerlike nature of 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

References

  1. ^ "Standard Atomic Weights: Sulfur". CIAAW. 2009.
  2. ISSN 1365-3075
    .
  3. ^
    ISBN 978-1-62708-155-9
    .
  4. ISBN 0-8493-0486-5
    .
  5. ISBN 0-8493-0464-4
    .
  6. ^ "Sulfur History". Georgiagulfsulfur.com. Retrieved 12 February 2022.
  7. ^
    ISBN 0-7506-3365-4
    .
  8. ^ Chisholm, Hugh, ed. (1911). "Brimstone" . Encyclopædia Britannica. Vol. 4 (11th ed.). Cambridge University Press. p. 571.
  9. ^ a b c d Laurence Knight (19 July 2014). "Sulphur surplus: Up to our necks in a diabolical element". BBC.
  10. ^ a b "Sulfur". Elements. BBC. 11 October 2014.. Download here.
  11. doi:10.1107/S0108270187088152
    .
  12. ^ "Sulfur | S (Element) - PubChem". pubchem.ncbi.nlm.nih.gov. Retrieved 15 April 2024.
  13. ^
    ISBN 978-0-08-037941-8
    .
  14. ISBN 978-0-07-149999-6
    .
  15. ^ "7.5: Changes of State". Chemistry LibreTexts. 3 October 2013. Retrieved 15 April 2024.
  16. PMID 15473808
    .
  17. ISSN 2296-6463
    .
  18. ISSN 0019-7866
    .
  19. ISSN 0007-0599
    . Retrieved 23 June 2022.
  20. doi:10.1021/ic50060a047. Archived from the original
    (PDF) on 12 April 2019. Retrieved 19 January 2019.
  21. ISBN 978-0-12-352651-9
    .
  22. ISBN 978-3-540-40191-9
    .
  23. ISBN 978-3-540-11345-4
    .
  24. doi:10.1021/ja00382a050
    .
  25. doi:10.1021/cr60230a004
    .
  26. doi:10.1021/cr60301a003
    .
  27. ^ Sulfur. Commission on Isotopic Abundances and Atomic Weights
  28. ISBN 1-4398-5511-0
    .
  29. ^ "Searching for the Origins of Presolar Grains". Energy.gov. Retrieved 4 February 2023.
  30. ISBN 9780128243602
    .
  31. ^ "NASA Astrobiology". astrobiology.nasa.gov. Retrieved 4 February 2023.
  32. doi:10.1016/0304-4203(80)90063-8
    .
  33. S2CID 1248182
    .
  34. S2CID 258600480
    .
  35. ^ Kim, Ik Soo; Kwak, Seung Im; Park, Ul Jae; Bang, Hong Sik; Han, Hyun Soo (1 July 2005). Production of Sulfur-35 by the Cation Exchange Process. 2005 autumn meeting of the KNS, Busan (Korea, Republic of), 27–28 Oct 2005.
  36. ^ "Sulfur-35 (35 S) safety information and specific handling precautions" (PDF). Yale Environmental Health & Safety.
  37. ^ Cameron, A. G. W. (1957). "Stellar Evolution, Nuclear Astrophysics, and Nucleogenesis" (PDF). CRL-41.
  38. ISBN 978-0-908678-84-6
    .
  39. S2CID 44208045
    .
  40. ^ Rickwood, P. C. (1981). "The largest crystals" (PDF). American Mineralogist. 66: 885–907.
  41. OCLC 79256100
    .
  42. S2CID 199492543
    .
  43. ISBN 0-471-80580-7
    .
  44. ^ "Sulphur: Mineral information, data and localities". www.mindat.org.
  45. ^
    ISBN 978-3-527-30673-2
    .
  46. ^ Shriver, Atkins. Inorganic Chemistry, Fifth Edition. W. H. Freeman and Company, New York, 2010; pp 416
  47. PMID 23851903
    .
  48. ^ Brauer, G., ed. (1963). Handbook of Preparative Inorganic Chemistry. Vol. 1 (2nd ed.). New York: Academic Press. p. 421.
  49. doi:10.1002/0471264180.os041.28
    .
  50. ^
    ISBN 978-3527306732
    .
  51. ISBN 0-521-21489-0
    .
  52. ISSN 2296-6846
    .
  53. ^
    ISBN 0-471-95512-4
    .
  54. doi:10.1086/182072
    .
  55. OCLC 697506461
    .
  56. ^ "Sulfur in the Bible (14 instances)". bible.knowing-jesus.com. Retrieved 19 May 2022.
  57. ISBN 978-3-540-78593-4
    .
  58. ^ Odyssey, book 22, lines 480–495. www.perseus.tufts.edu. Retrieved on 16 August 2012.
  59. ISBN 0-19-814687-6
    , pp. 247–249.
  60. ^
    S2CID 144187385
    .
  61. OCLC 489677531
    .
  62. ISBN 0-486-20162-7
    .
  63. ISBN 978-0-226-89499-7
    .
  64. PMID 2450900
    .
  65. PMID 2365870
    .
  66. PMID 15303787
    .
  67. PMID 15303787
    .
  68. ISBN 978-0-521-56672-8
    .
  69. ISBN 978-0-8122-1649-3
    .
  70. ISBN 9780191542619
    . Retrieved 7 February 2013.
  71. S2CID 145807900
    .
  72. doi:10.1002/1521-3781(200110)35:5<324::AID-CIUZ324>3.0.CO;2-9
    .
  73. S2CID 250882259
    .
  74. OCLC 62805047
    .
  75. ^ "sulphur". Oxford English Dictionary (Online ed.). Oxford University Press. (Subscription or participating institution membership required.)
  76. PMID 21378874
    .
  77. doi:10.1039/AN9911601094
    .
  78. ^ "sulphur – definition of sulphur in English". Oxford Dictionaries. Archived from the original on 20 November 2016. Retrieved 19 November 2016.
  79. OCLC 74650396
    .
  80. ^ Washington, Booker T. (1912). The Man Farthest Down: A Record of Observation and Study in Europe. Doubleday, Page. p. 214.
  81. ^ McElvaney, Kevin (25 February 2015). "The Men Who Mine Volcanos". The Atlantic. Retrieved 26 February 2015.
  82. ^
    doi:10.1002/ep.670210312
    .
  83. ^
    doi:10.1002/ciuz.200800461
    .
  84. ISBN 978-0-8412-0713-4
    .
  85. ^
    OCLC 531718953
    .
  86. ^ Apodaca, Lori E. (2012) "Sulfur". Mineral Commodity Summaries. United States Geological Survey.
  87. ^ a b Apodaca, Lori E. "Mineral Yearbook 2010: Sulfur" (PDF). United States Geological Survey.
  88. ^ a b "FAQ – The Sulphur Institute". sulphurinstitute.org. The Sulphur Institute. 2020. Retrieved 27 February 2020.
  89. doi:10.1006/jcrs.1998.0241
    .
  90. S2CID 44208638
    .
  91. S2CID 42207099
    .
  92. ^ Glossary, United States: NASA Earth Observatory, acid rain, archived from the original on 13 December 2011, retrieved 15 February 2013
  93. ^ Every, Richard L.; et al. (20 August 1968). "Method for Preparation of Wettable Sulfur" (PDF). Retrieved 20 May 2010.
  94. ISBN 978-3-540-07738-1
    .
  95. ISBN 978-3-8047-2575-1
    .
  96. ISBN 978-3-030-25598-5
    , retrieved 8 March 2023
  97. S2CID 246704531
    .
  98. ^ "Sulphur and the Human Body". The Sulfur Institute. Retrieved 3 April 2021.
  99. ^ "What is the body made of?". New Scientist. Archived from the original on 3 November 2021. Retrieved 9 November 2021.
  100. ThoughtCo. Archived
    from the original on 13 April 2021. Retrieved 21 November 2021.
  101. PMID 11896744
    .
  102. PMID 23815141
    .
  103. PMID 22092289
    .
  104. PMID 19871997
    .
  105. doi:10.1111/j.1574-6968.1990.tb04103.x
    .
  106. PMID 18929068
    , retrieved 17 May 2022
  107. doi:10.1093/icb/34.1.79
    .
  108. PMID 17793426
    .
  109. S2CID 128827551
    .
  110. ^ "Petrography and mineralogy of the crystalline limestone of Fatha Formation from Mishraq area, Iraq". ResearchGate. Retrieved 15 April 2022.
  111. ISBN 978-3-8274-0103-8
    .
  112. PMID 6127737
    .
  113. PMID 28419734
    .
  114. ^ Gutiérrez-Preciado, A.; Romero, H.; Peimbert, M. (2010). "An Evolutionary Perspective on Amino Acids". Nature Education. 3 (9): 29.
  115. ISBN 978-0-8153-3218-3
    .
  116. PMID 11012661
    .
  117. ^ "Lysozyme". Drugs.com. Retrieved 19 May 2022.
  118. ISBN 978-1-57259-153-0
    .
  119. PMID 10448523
    .
  120. PMID 1731369
    .
  121. ^ "The Function of Biotin". www.chem.uwec.edu. Retrieved 10 June 2022.
  122. ^ Edwards, Katie A. "Thiamine Biochemistry". thiamine.dnr.cornell.edu. Retrieved 10 June 2022.
  123. PMID 9782487
    .
  124. PMID 24970223
    .
  125. PMID 12954327
    .
  126. ISBN 978-0-935702-73-6
    .
  127. PMID 16669776
    .
  128. S2CID 73470156
    .
  129. PMID 33673497
    .
  130. PMID 16234925
    .
  131. ^ "Sulfur 84683". S.
  132. ^ "Chemical Datasheet. SULFUR, MOLTEN".
  133. ^ "Sulfur General Fact Sheet". npic.orst.edu. Retrieved 2 September 2022.
  134. PMID 3718141
    .
  135. PMID 904661
    . Retrieved 2 September 2022.
  136. ^ "Sulfur Dioxide | Toxicological Profile | ATSDR". Centers for Disease Control and Prevention. 26 March 2014. Retrieved 24 October 2023.
  137. ^ "Sulfur General Fact Sheet". npic.orst.edu. Retrieved 20 January 2024.
  138. ^ Baker, Colin (1 March 2007). "The dehydration of sucrose". Education in Chemistry. Royal Society of Chemistry. Retrieved 14 June 2018.
  139. ^ "4.1: Fossil Fuels and Products of Combustion". Engineering LibreTexts. 28 February 2021. Retrieved 24 October 2023.
  140. ^ "Hydrogen Sulfide Toxicity: Practice Essentials, Pathophysiology, Etiology". Medscape. 30 March 2017 – via eMedicine.
  141. ^ Summers, Vincent (8 April 2017). "Hydrogen Sulfide or Hydrogen Cyanide: Which is More Dangerous?". Quirky Science. Retrieved 23 August 2022.
  142. ^ "Hydrogen Sulfide – Hazards | Occupational Safety and Health Administration". www.osha.gov. Retrieved 23 August 2022.

Further reading

Wikiquote has quotations related to Sulfur.

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.

ISBN 978-3-11-058889-7
.

External links
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