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= Sulfur =
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Introduction
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Sulfur (American spelling and the preferred IUPAC name) or sulphur
(Commonwealth spelling) is a chemical element; it has symbol S and
atomic number 16. It is abundant, multivalent and nonmetallic. Under
normal conditions, sulfur atoms form cyclic octatomic molecules with
the chemical formula S8. Elemental sulfur is a bright yellow,
crystalline solid at room temperature.
Sulfur is the tenth most abundant element by mass in the universe and
the fifth most common on Earth. Though sometimes found in pure, native
form, sulfur on Earth usually occurs as sulfide and sulfate minerals.
Being abundant in native form, sulfur was known in ancient times,
being mentioned for its uses in ancient India, ancient Greece, China,
and ancient Egypt. Historically and in literature sulfur is also
called brimstone, which means "burning stone". Almost all elemental
sulfur is produced as a byproduct of removing sulfur-containing
contaminants from natural gas and petroleum. The greatest commercial
use of the element is the production of sulfuric acid for sulfate and
phosphate fertilizers, and other chemical processes. Sulfur is used in
matches, insecticides, and fungicides. Many sulfur compounds are
odoriferous, and the smells of odorized natural gas, skunk scent, bad
breath, grapefruit, and garlic are due to organosulfur compounds.
Hydrogen sulfide gives the characteristic odor to rotting eggs and
other biological processes.
Sulfur is an essential element for all life, almost always in the form
of organosulfur compounds or metal sulfides. Amino acids (two
proteinogenic: cysteine and methionine, and many other non-coded:
cystine, taurine, etc.) and two vitamins (biotin and thiamine) are
organosulfur compounds crucial for life. Many cofactors also contain
sulfur, including glutathione, and iron-sulfur proteins. Disulfides,
S-S bonds, confer mechanical strength and insolubility of the (among
others) protein keratin, found in outer skin, hair, and feathers.
Sulfur is one of the core chemical elements needed for biochemical
functioning and is an elemental macronutrient for all living
organisms.
Physical properties
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Sulfur forms several polyatomic molecules. The best-known allotrope is
octasulfur, cyclo-S8. The point group of cyclo-S8 is D4d and its
dipole moment is 0 D. Octasulfur is a soft, bright-yellow solid that
is odorless. It melts at 115.21 C,{{efn|Sulfur's melting point at
115.21°C was determined by two laboratories of the US Department of
Energy (Jefferson Lab and Los Alamos National Lab).
Greenwood and Earnshaw say that at fast heating for microcrystalline
α-S8 the melting point is 115.1 C.}} and boils at 444.6 C. At 95.2 C,
below its melting temperature, cyclo-octasulfur begins slowly changing
from α-octasulfur to the β-polymorph. The structure of the S8 ring is
virtually unchanged by this phase transition, which affects the
intermolecular interactions. Cooling molten sulfur freezes at 119.6 C,
as it predominantly consists of the β-S8 molecules. 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. At higher temperatures, the viscosity decreases as
depolymerization occurs. Molten sulfur assumes a dark red color above
200 C. The density of sulfur is about 2 g/cm3, depending on the
allotrope; all of the stable allotropes are excellent electrical
insulators.
The sublimation of sulfur becomes noticeable more or less between 20 C
and 50 C, and occurs readily in boiling water at 100 C.
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
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Under normal conditions, sulfur hydrolyzes very slowly to mainly form
hydrogen sulfide and sulfuric acid:
The reaction involves adsorption of protons onto clusters, followed
by disproportionation into the reaction products.
The second, fourth and sixth ionization energies of sulfur are 2252
kJ/mol, 4556 kJ/mol and 8495.8 kJ/mol, respectively. The composition
of reaction products of sulfur with oxidants (and its oxidation state)
depends on whether releasing of reaction energy overcomes these
thresholds. Applying catalysts and/or supply of external energy may
vary sulfur's oxidation state and the composition of reaction
products. While reaction between sulfur and oxygen under normal
conditions gives sulfur dioxide (oxidation state +4), formation of
sulfur trioxide (oxidation state +6) requires a temperature of 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 except noble gases, even
with the notoriously unreactive metal iridium (yielding iridium
disulfide). Some of those reactions require elevated temperatures.
Allotropes
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Sulfur forms over 30 solid allotropes, more than any other element.
Besides S8, several other rings are known. 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. Larger rings have been prepared,
including S12 and S18.
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 it has the
feel of crude rubber. This form is metastable at room temperature and
gradually reverts to the crystalline molecular allotrope, which is no
longer elastic. This process happens over a matter of hours to days,
but can be rapidly catalyzed.
Isotopes
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Sulfur has 23 known isotopes, four of which are stable: 32S (), 33S
(), 34S (), and 36S (). Other than 35S, with a half-life of 87 days,
the radioactive isotopes of sulfur have half-lives less than 3 hours.
The preponderance of 32S is explained by its production in the alpha
process (one of the main classes of nuclear fusion reactions) in
exploding stars. Other stable sulfur isotopes are produced in the
bypass processes related with 34Ar, and their composition depends on a
type of a stellar explosion. For example, proportionally more 33S
comes from novae than from supernovae.
On the planet Earth the sulfur isotopic composition was determined by
the Sun. Though it was assumed that the distribution of different
sulfur isotopes would be more or less equal, it has been found that
proportions of the two most abundant sulfur isotopes 32S and 34S
varies in different samples. Assaying of the isotope ratio (δ34S) in
the samples suggests 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. However, there are ongoing discussions over the real
reason for the δ34S shifts, biological activity or postdeposit
alteration.
For example, when sulfide minerals are precipitated, isotopic
equilibration among solids and liquid may cause small differences in
the δ34S values of co-genetic minerals. The differences between
minerals can be used to estimate the temperature of equilibration. The
δ13C and δ34S of coexisting carbonate minerals and sulfides can be
used to determine the pH and oxygen fugacity of the ore-bearing fluid
during ore formation.
Scientists measure the sulfur isotopes of minerals in rocks and
sediments to study the redox conditions in past oceans.
Sulfate-reducing bacteria in marine sediment fractionate sulfur
isotopes as they take in sulfate and produce sulfide. Prior to the
2010s, it was thought that sulfate reduction could fractionate sulfur
isotopes up to 46 permil 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 showed that sulfate-reducing bacteria can fractionate to
66 permil. 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.
In forest ecosystems, sulfate is derived mostly from the atmosphere;
weathering of ore minerals and evaporites contribute some sulfur.
Sulfur with a distinctive isotopic composition has been used to
identify pollution sources, and enriched sulfur has been added as a
tracer in hydrologic studies. Differences in the natural abundances
can be used in systems where there is sufficient variation in the 34S
of ecosystem components. Rocky Mountain lakes thought to be dominated
by atmospheric sources of sulfate have been found to have measurably
different 34S values than lakes believed to be dominated by watershed
sources of sulfate.
The radioactive 35S is formed in cosmic ray spallation of the
atmospheric 40Ar. This fact may be used to verify the presence of
recent (up to 1 year) atmospheric sediments in various materials. This
isotope may be obtained artificially by different ways. In practice,
the reaction 35Cl + n → 35S + p is used by irradiating potassium
chloride with neutrons. The isotope 35S is used in various
sulfur-containing compounds as a radioactive tracer for many
biological studies, for example, the Hershey-Chase experiment.
Because of the weak beta activity of 35S, its compounds are relatively
safe as long as they are not ingested or absorbed by the body.
Natural occurrence
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32S is created inside massive stars, at a depth where the temperature
exceeds 2.5 billion K, by the fusion of one nucleus of silicon plus
one nucleus of helium. 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. The
distinctive colors of Jupiter's volcanic moon Io are attributed to
various forms of molten, solid, and gaseous sulfur. In July 2024,
elemental sulfur was accidentally discovered to exist on Mars after
the Curiosity rover drove over and crushed a rock, revealing sulfur
crystals inside it.
Sulfur is the fifth most common element by mass in the Earth.
Elemental sulfur can be found near hot springs and volcanic regions in
many parts of the world, especially along the Pacific Ring of Fire;
such volcanic deposits are mined in Indonesia, Chile, and Japan. These
deposits are polycrystalline, with the largest documented single
crystal measuring . Historically, Sicily was a major source of sulfur
in the Industrial Revolution. Lakes of molten sulfur up to about 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.
Native sulfur is synthesized by anaerobic bacteria acting on sulfate
minerals such as gypsum in salt domes. 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. Such sources have become of
secondary commercial importance, and most are no longer worked but
commercial production is still carried out in the Osiek mine in
Poland.
Common naturally occurring sulfur compounds include the sulfide
minerals, such as pyrite (iron sulfide), cinnabar (mercury sulfide),
galena (lead sulfide), sphalerite (zinc sulfide), and stibnite
(antimony sulfide); and the sulfate minerals, such as gypsum (calcium
sulfate), alunite (potassium aluminium sulfate), and barite (barium
sulfate). On Earth, just as upon Jupiter's moon Io, elemental sulfur
occurs naturally in volcanic emissions, including emissions from
hydrothermal vents.
The main industrial source of sulfur has become petroleum and natural
gas.
Compounds
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Common oxidation states of sulfur range from −2 to +6. Sulfur forms
stable compounds with all elements except the noble gases.
Electron transfer reactions
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Sulfur polycations, , and are produced when sulfur is reacted with
oxidizing agents in a strongly acidic solution. 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. is
deep blue, is yellow and is red.
Reduction of sulfur gives various polysulfides with the formula , many
of which have been obtained in crystalline form. Illustrative is the
production of sodium tetrasulfide:
Some of these dianions dissociate to give radical anions. For
instance, gives the blue color of the rock lapis lazuli.
This reaction highlights a distinctive property of sulfur: its ability
to catenate (bind to itself by formation of chains). Protonation of
these polysulfide anions produces the polysulfanes, H2S'x', where 'x'
= 2, 3, and 4. Ultimately, reduction of sulfur produces sulfide salts:
The interconversion of these species is exploited in the sodium-sulfur
battery.
Hydrogenation
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Treatment of sulfur with hydrogen gives hydrogen sulfide. When
dissolved in water, hydrogen sulfide is mildly acidic:
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
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The two principal sulfur oxides are obtained by burning sulfur:
Many other sulfur oxides are observed including the sulfur-rich oxides
include sulfur monoxide, disulfur monoxide, disulfur dioxides, and
higher oxides containing peroxo groups.
Halogenation
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Sulfur reacts with fluorine to give the highly reactive sulfur
tetrafluoride and the highly inert sulfur hexafluoride. 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.
Bromine also oxidizes sulfur to form sulfur dibromide and disulfur
dibromide.
Pseudohalides
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Sulfur oxidizes cyanide and sulfite to give thiocyanate and
thiosulfate, respectively.
Metal sulfides
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Sulfur reacts with many metals. Electropositive metals give
polysulfide salts. Copper, zinc, and silver are attacked by sulfur;
see tarnishing. Although many metal sulfides are known, most are
prepared by high temperature reactions of the elements. Geoscientists
also study the isotopes of metal sulfides in rocks and sediment to
study environmental conditions in the Earth's past.
Organic compounds
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File:L-Cystein - L-Cysteine.svg |('L')-cysteine, an amino acid
containing a thiol group
File:Methionin - Methionine.svg|Methionine, an amino acid containing a
thioether
File:Thiamin.svg|Thiamine or vitamin B1
File:Biotin_structure.svg|Biotin or vitamin B7
File:Penicillin core.svg|Penicillin, an antibiotic ("R" is the
variable group)
File:Allicin skeletal.svg|Allicin, a chemical compound in garlic
File:Diphenyl disulfide.svg|Diphenyl disulfide, a representative
disulfide
File:Dibenzothiophen - Dibenzothiophene.svg|Dibenzothiophene, a
component of crude oil
File:Perfluorooctanesulfonic acid
structure.svg|Perfluorooctanesulfonic acid (PFOS), a surfactant
Some of the main classes of sulfur-containing organic compounds
include the following:
* Thiols or mercaptans (so called because they capture mercury as
chelators) are the sulfur analogs of alcohols; treatment of thiols
with base gives thiolate ions.
* Thioethers are the sulfur analogs of ethers.
* Sulfonium ions have three groups attached to a cationic sulfur
center. Dimethylsulfoniopropionate (DMSP) is one such compound,
important in the marine organic sulfur cycle.
* Sulfoxides and sulfones are thioethers with one and two oxygen atoms
attached to the sulfur atom, respectively. The simplest sulfoxide,
dimethyl sulfoxide, is a common solvent; a common sulfone is
sulfolane.
* Sulfonic acids are used in many detergents.
Compounds with carbon-sulfur multiple bonds are uncommon, an exception
being carbon disulfide, a volatile colorless liquid that is
structurally similar to carbon dioxide. It is used as a reagent to
make the polymer rayon and many organosulfur compounds. Unlike carbon
monoxide, carbon monosulfide is stable only as an extremely dilute
gas, found between solar systems.
Organosulfur compounds are responsible for some of the unpleasant
odors of decaying organic matter. They are widely known as the odorant
in domestic natural gas, garlic odor, and skunk spray, as well as a
component of bad breath odor. Not all organic sulfur compounds smell
unpleasant at all concentrations: the sulfur-containing monoterpenoid
grapefruit mercaptan in small concentrations is the characteristic
scent of grapefruit, but has a generic thiol odor at larger
concentrations. Sulfur mustard, a potent vesicant, was used in World
War I as a disabling agent.
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.
Antiquity
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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; this is mentioned in the 'Odyssey'. Pliny the
Elder discusses sulfur his 'Natural History', saying that its
best-known source is the island of Melos. He mentions its use for
fumigation, medicine, and bleaching cloth.
A natural form of sulfur known as was known in China since the 6th
century BC and found in Hanzhong. By the 3rd century, the Chinese had
discovered that sulfur could be extracted from pyrite. 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. The 'Wujing Zongyao' of 1044 AD
described formulas for Chinese black powder, which is a mixture of
potassium nitrate, charcoal, and sulfur.
English translations of the Christian Bible commonly referred to
burning sulfur as "brimstone", giving rise to the term
"fire-and-brimstone" sermons, in which listeners are reminded of the
fate of eternal damnation that await the unbelieving and unrepentant.
Hell is implied to smell of sulfur.
Indian alchemists, practitioners of the "science of chemicals" (),
wrote extensively about the use of sulfur in alchemical operations
with mercury, from the eighth century AD onwards. In the rasa shastra
tradition, sulfur is called "the smelly" (, ).
Early European alchemists gave sulfur an alchemical symbol of a
triangle atop a cross (🜍). The variation known as brimstone has a
symbol combining a two-barred cross atop a lemniscate (🜏). In
traditional skin treatment, elemental sulfur was used (mainly in
creams) to alleviate such conditions as scabies, ringworm, psoriasis,
eczema, and acne. The mechanism of action is unknown--though elemental
sulfur does oxidize slowly to sulfurous acid, a mild reducing and
antibacterial agent.
Modern times
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Sulfur appears in a column of fixed (non-acidic) alkali in a chemical
table of 1718. Antoine Lavoisier used sulfur in combustion
experiments, writing of some of these in 1777.
Sulfur deposits in Sicily were the dominant source for more than a
century. By the late 18th century, about 2,000 tonnes per year of
sulfur were imported into Marseille, France, for the production of
sulfuric acid for use in the Leblanc process. In industrializing
Britain, with the repeal of tariffs on salt in 1824, demand for sulfur
from Sicily surged. The increasing British control and exploitation of
the mining, refining, and transportation of 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
solution was eventually negotiated by France.
In 1867, elemental sulfur was discovered in underground deposits in
Louisiana and Texas. The highly successful Frasch process was
developed to extract this resource.
In the late 18th century, furniture makers used molten sulfur to
produce decorative inlays. 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.
Since the advent of the contact process, the majority of sulfur is
used to make sulfuric acid for a wide range of uses, particularly
fertilizer.
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.
Spelling and etymology
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'Sulfur' is derived from the Latin word ', which was Hellenized to '
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 ', which appears in Latin toward the end of
the Classical period. The true Ancient Greek word for sulfur, ,
'theîon' (from earlier , '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 '. In the 14th century, the
erroneously Hellenized Latin ' was restored in Middle English '. 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'. On the other hand, 'sulfur' was the form
eventually chosen in the United States, though multiple place names
(such as White Sulphur Springs) use '-ph-'. Canada uses both
spellings.
IUPAC adopted the spelling 'sulfur' in 1990 as did the Nomenclature
Committee of the Royal Society of Chemistry in 1992, restoring the
spelling 'sulfur' to Britain. 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."
Production
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Sulfur may be found by itself and historically was usually obtained in
this form; pyrite has also been a source of sulfur. 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." Sulfur is still
mined from surface deposits in poorer nations with volcanoes, such as
Indonesia, and problems with working conditions still exist.
Elemental sulfur was extracted from salt domes (where it sometimes
occurs in nearly pure form) until the late 20th century, when it
became a side product of other industrial processes such as in oil
refining, in which sulfur is undesirable. 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. 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 had
significant use anywhere in the world since 2002.
Since then, sulfur has typically been 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:
The resulting hydrogen sulfide from this process, and also as it
occurs in natural gas, is converted into elemental sulfur by the Claus
process, which entails oxidation of some hydrogen sulfide to sulfur
dioxide and then the comproportionation of the two:
Due to the high sulfur content of the Athabasca Oil Sands, stockpiles
of elemental sulfur from this process exist throughout Alberta,
Canada. Another way of storing sulfur is as a binder for concrete, the
resulting product having some desirable properties (see sulfur
concrete).
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). Production has been slowly
increasing from 1900 to 2010; the price was unstable in the 1980s and
around 2010.
Sulfuric acid
===============
Elemental sulfur is used mainly as a precursor to other chemicals.
Approximately 85% (1989) is converted to sulfuric acid (H2SO4):
In 2010, the United States produced more sulfuric acid than any other
inorganic industrial chemical. 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.
Other important sulfur chemistry
==================================
Sulfur reacts directly with methane to give carbon disulfide, which is
used to manufacture cellophane and rayon. One of the uses of elemental
sulfur is in vulcanization of rubber, where polysulfide chains
crosslink organic polymers. Large quantities of sulfites are used to
bleach paper and to preserve dried fruit. Many surfactants and
detergents (e.g. 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
============
Amino acids synthesized by living organisms such as methionine and
cysteine contain organosulfur groups (thioester and thiol
respectively). The antioxidant glutathione protecting many living
organisms against free radicals and oxidative stress also contains
organic sulfur. Some crops such as onion and garlic also produce
different organosulfur compounds such as 'syn'-propanethial-'S'-oxide
responsible of lacrymal irritation (onions), or diallyl disulfide and
allicin (garlic). Sulfates, commonly found in soils and groundwaters
are often a sufficient natural source of sulfur for plants and
bacteria. Atmospheric deposition of sulfur dioxide (SO2) is also a
common artificial source (coal combustion) of sulfur for the soils.
Under normal circumstances, in most agricultural soils, sulfur is not
a limiting nutrient for plants and microorganisms (see Liebig's
barrel). However, in some circumstances, soils can be depleted in
sulfate, e.g. if this later is leached by meteoric water (rain) or if
the requirements in sulfur for some types of crops are high. This
explains that sulfur is increasingly recognized and used as a
component of fertilizers. The most important form of sulfur for
fertilizer is calcium sulfate, commonly found in nature as the mineral
gypsum (CaSO4·2H2O). Elemental sulfur is hydrophobic (not soluble in
water) and cannot be used directly by plants. Elemental sulfur (ES) is
sometimes mixed with bentonite to amend depleted soils for crops with
high requirement in organo-sulfur. Over time, oxidation abiotic
processes with 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.
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. 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.
Fungicide and pesticide
=========================
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 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. 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 lime sulfur (made by combining calcium hydroxide
with elemental sulfur in water) is used as a dip for pets to destroy
ringworm (fungus), mange, and other 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
=================
Sulfur (specifically octasulfur, S8) is used in pharmaceutical skin
preparations for the treatment of acne and other conditions. It acts
as a keratolytic agent and also kills bacteria, fungi, scabies mites,
and other parasites. Precipitated sulfur and colloidal sulfur are
used, in form of lotions, creams, powders, soaps, and bath additives,
for the treatment of acne vulgaris, acne rosacea, and seborrhoeic
dermatitis.
Many drugs contain sulfur. Early examples include antibacterial
sulfonamides, known as 'sulfa drugs'. A more recent example is
mucolytic acetylcysteine. Sulfur is a part of many bacterial defense
molecules. Most β-lactam antibiotics, including the penicillins,
cephalosporins and monobactams contain sulfur.
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.
Biological role
======================================================================
Sulfur is an essential component of all living cells. It is the eighth
most abundant element in the human body by weight, about equal in
abundance to potassium, and slightly greater than sodium and chlorine.
A 70 kg human body contains about of sulfur. The main dietary source
of sulfur for humans is sulfur-containing amino acids, which can be
found in plant and animal proteins.
Transferring sulfur between inorganic and biomolecules
========================================================
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). Another
contributor, who continued to study it was Selman Waksman. Primitive
bacteria that live around deep ocean volcanic vents oxidize hydrogen
sulfide for their nutrition, as discovered by Robert Ballard.
Sulfur oxidizers can use as energy sources reduced sulfur compounds,
including hydrogen sulfide, elemental sulfur, sulfite, thiosulfate,
and various polythionates (e.g., tetrathionate). 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:
There are bacteria combining these two ways of nutrition: green sulfur
bacteria and purple sulfur bacteria. Also sulfur-oxidizing bacteria
can go into symbiosis with larger organisms, enabling the later to use
hydrogen sulfide as food to be oxidized. Example: the giant tube worm.
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. 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.
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.
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. Some of the
sulfur is present in certain metabolites--many of which are
cofactors--and sulfated polysaccharides of connective tissue
(chondroitin sulfates, heparin).
The functionality of a given protein is heavily dependent on its
structure. Proteins reach this structure through the process of
protein folding, which is facilitated by a variety of intra- and
inter-molecular bonds. While much of the folding is driven by the
formation of hydrogen bonds, covalent bonding of cysteine residues
into disulfide bridges imposes constraints that stabilize particular
conformations while preventing others from forming. As the bond energy
of a covalent disulfide bridge is higher than the energy of a
coordinate bond or hydrophobic interaction, greater numbers of
disulfide bridges lead to higher energies required for protein
denaturation. Disulfide bonds often serve to stabilize protein
structures in the more oxidizing conditions of the extracellular
environment. Within the cytoplasm, disulfide bonds may instead be
reduced (i.e. in -SH form) to their constituent cysteine residues by
thioredoxins.
Many important cellular enzymes use prosthetic groups ending with
sulfhydryl (-SH) moieties to handle reactions involving
acyl-containing biochemicals: two common examples from basic
metabolism are coenzyme A and alpha-lipoic acid. Cysteine-related
metabolites homocysteine and taurine are other sulfur-containing amino
acids that are similar in structure, but not coded by DNA, and are not
part of the primary structure of proteins, take part in various
locations of mammalian physiology. Two of the 13 classical vitamins,
biotin and thiamine, contain sulfur, and serve as cofactors to several
enzymes.
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, , the immediate precursor to methane.
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--are essential components of enzymes involved in
electron transfer processes. Examples include plastocyanin (Cu2+) and
nitrous oxide reductase (Cu-S). The function of these enzymes is
dependent on the fact that the transition metal ion can undergo redox
reactions. Other examples include many zinc proteins, as well as
iron-sulfur clusters. Most pervasive are the ferrodoxins, which serve
as electron shuttles in cells. In bacteria, the important nitrogenase
enzymes contain an Fe-Mo-S cluster and 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.
Sulfur is also present in molybdenum cofactor.
:thumb
Deficiency
============
In humans methionine is an essential amino acid; cysteine is
conditionally essential and may be synthesized from non-essential
serine via sulfur salvaged from methionine. Sulfur deficiency is
uncommon due to the ubiquity of cysteine and methionine in food.
Isolated sulfite oxidase deficiency is a rare, fatal genetic disease
caused by mutations to sulfite oxidase, which is needed to metabolize
sulfites to sulfates.
Precautions
======================================================================
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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, and cases of
life-threatening metabolic acidosis have been reported after patients
deliberately consumed sulfur as a folk remedy.
Toxicity of sulfur compounds
==============================
When sulfur burns in air, it produces sulfur dioxide. In water, this
gas produces sulfurous acid and sulfites; sulfites are antioxidants
that inhibit growth of aerobic bacteria and a useful food additive in
small amounts. At high concentrations these acids harm the lungs,
eyes, or other tissues. In organisms without lungs such as insects,
sulfite in high concentration prevents respiration.
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.
The burning of coal and/or petroleum by industry and power plants
generates sulfur dioxide (SO2) that reacts with atmospheric water and
oxygen to produce sulfurous acid (H2SO3). These acids are components
of acid rain, lowering 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 about one-half as toxic as hydrogen cyanide, and
intoxicates by the same mechanism (inhibition of the respiratory
enzyme cytochrome oxidase), though hydrogen sulfide is less likely to
cause sudden poisonings from small inhaled amounts (near its
permissible exposure limit (PEL) of 20 ppm) because of its
disagreeable odor. However, its presence in ambient air at
concentration over 100-150 ppm quickly deadens the sense of smell, and
a victim may breathe increasing quantities without noticing until
severe symptoms cause death. Dissolved sulfide and hydrosulfide salts
are toxic by the same mechanism.
See also
======================================================================
*Blue lava
*Stratospheric sulfur aerosols
*Sulfur assimilation
*Sulfur isotope biogeochemistry
*Ultra-low-sulfur diesel
External links
======================================================================
*[
http://www.periodicvideos.com/videos/016.htm Sulfur] at 'The
Periodic Table of Videos' (University of Nottingham)
*[
https://physics.nist.gov/PhysRefData/Handbook/Tables/sulfurtable1.htm
Atomic Data for Sulfur], NIST Physical Measurement Laboratory
*[
https://library.tedankara.k12.tr/chemistry/vol2/allotropy/z129.htm
Sulfur phase diagram] , Introduction to Chemistry for Ages 13-17
*[
https://www.swisseduc.ch/stromboli/perm/vulcano/sulphur-vulcano-en.html
Crystalline, liquid and polymerization of sulfur on Vulcano Island,
Italy]
*[
https://extoxnet.orst.edu/pips/sulfur.htm Sulfur and its use as a
pesticide]
*[
https://www.sulphurinstitute.org/ The Sulphur Institute]
*[
https://web.archive.org/web/20140714221502/http://www.nutrientstewardship.com/partners/products-and-services/sulfur-institute
Nutrient Stewardship and The Sulphur Institute]
License
=========
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License URL:
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Original Article:
http://en.wikipedia.org/wiki/Sulfur
