Spectral lines of sulfur
|Name, symbol||sulfur, S|
|Appearance||lemon yellow sintered microcrystals|
|Sulfur in the periodic table|
|Standard atomic weight||32.06(1)|
|Element category||polyatomic nonmetal|
|Group, block||group 16 (chalcogens), p-block|
|Electron configuration||[Ne] 3s2 3p4|
|per shell||2, 8, 6|
|Melting point||388.36 K (115.21 °C, 239.38 °F)|
|Boiling point||717.8 K (444.6 °C, 832.3 °F)|
|Density near r.t.||
alpha: 2.07 g·cm−3
beta: 1.96 g·cm−3
gamma: 1.92 g·cm−3
|liquid, at m.p.||1.819 g·cm−3|
|Critical point||1314 K, 20.7 MPa|
|Heat of fusion||mono: 1.727 kJ·mol−1|
|Heat of vaporization||mono: 45 kJ·mol−1|
|Molar heat capacity||22.75 J·mol−1·K−1|
|Oxidation states||6, 5, 4, 3, 2, 1, −1, −2 (a strongly acidic oxide)|
|Electronegativity||Pauling scale: 2.58|
1st: 999.6 kJ·mol−1
2nd: 2252 kJ·mol−1
3rd: 3357 kJ·mol−1
|Covalent radius||105±3 pm|
|Van der Waals radius||180 pm|
|Thermal conductivity||0.205 W·m−1·K−1 (amorphous)|
|Electrical resistivity||at 20 °C: 2×1015 Ω·m (amorphous)|
|Bulk modulus||7.7 GPa|
|Discovery||Chinese (Before 2000 BCE)|
|Recognized as an element by||Antoine Lavoisier (1777)|
|Most stable isotopes|
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. 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 largest commercial use (after mostly being converted to sulfuric acid) is to produce sulfate and phosphate fertilizers, because of the relatively high requirement of plants for sulfur and phosphorus. Sulfuric acid is also a primary industrial chemical outside fertilizer manufacture. 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 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 1
- Physical properties 2.1
- Chemical properties 2.2
- Allotropes 2.3
- Isotopes 2.4
- Natural occurrence 2.5
- Production 3
- Sulfur polycations 4.1
- Sulfides 4.2
- Oxides, oxoacids and oxoanions 4.3
- Halides and oxyhalides 4.4
- Pnictides 4.5
- Metal sulfides 4.6
- Organic compounds 4.7
- Antiquity 5.1
- Modern times 5.2
- Sulfuric acid 6.1
- Other large-scale sulfur chemicals 6.2
- Fertilizer 6.3
- Fine chemicals 6.4
- Fungicide and pesticide 6.5
- Bactericide in winemaking and food preservation 6.6
Biological role 7
- Protein and organic cofactors 7.1
- Metalloproteins and inorganic cofactors 7.2
- Sulfur metabolism and the sulfur cycle 7.3
- Precautions 8
- See also 9
- References 10
- External links 11
Spelling and etymology
Sulfur is historically a Latin word. The original Latin spelling was sulpur, but this was Hellenized to sulphur; the form sulfur appears toward the end of the Classical period. (The true Greek word for sulfur, θεῖον, is the source of the international chemical prefix thio-.) In 12th-century Anglo-French, it was sulfre; in the 14th century the Latin ph was restored, for sulphre; and by the 15th century the full Latin spelling was restored, for sulfur, sulphur. The parallel f~ph spellings continued in Britain until the 19th century, when the word was standardized as sulphur. Sulfur was the form chosen in the United States, while Canada uses both. However, the 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. 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."
The Latin word also continues in the Romance languages: French soufre, Italian zolfo (from solfo), Spanish azufre (from açufre, from earlier çufre), Portuguese enxofre (from xofre). The Spanish and Portuguese forms are prefixed with the Arabic article, despite not being Arabic words. The root has been traced back to reconstructed proto-Indo-European *swépl̥ (genitive *sulplós), a nominal derivative of *swelp 'to burn', a lineage also preserved in the Germanic languages, where it is found for example as modern German Schwefel, Dutch zwavel, and Swedish svavel, and as Old English swefl.
Sulfur forms polyatomic molecules with different chemical formulas, with the best-known allotrope being 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 with only a faint odor, similar to that of matches. It melts at 115.21 °C, boils at 444.6 °C and sublimes easily. At 95.2 °C, below its melting temperature, cyclo-octasulfur changes from α-octasulfur to the β-polymorph. 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. 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.
Sulfur burns with a blue flame concomitant with formation of 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.
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 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. 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 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.
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 3 hours.
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.
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. 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. The distinctive colors of Jupiter's volcanic moon Io are attributed to various forms of molten, solid and gaseous sulfur.
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. Historically, Sicily was a large source of sulfur in the Industrial Revolution.
Native sulfur is synthesised 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 have until recently been the basis for commercial production in the United States, Russia, Turkmenistan, and Ukraine. 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.
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. 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.".
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. 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.
Today, sulfur is produced from petroleum, hydrodesulfurization, which cleaves the C–S bonds:
- 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:
- 3 O2 + 2 H2S → 2 SO2 + 2 H2O
- SO2 + 2 H2S → 3 S + 2 H2O
Owing to the high sulfur content of the Athabasca Oil Sands, stockpiles of elemental sulfur from this process now exist throughout Alberta, Canada. Another way of storing sulfur is as a binder for concrete, the resulting product having many 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), US (8.8), Canada (7.1) and Russia (7.1). While the production has been slowly increasing from 1900 to 2010, the price was much less stable, especially in the 1980s and around 2010.
Sulfur polycations, S82+, S42+ and S192+ are produced when sulfur is reacted with mild oxidising 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. S82+ is deep blue, S42+ is yellow and S192+ is red.
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).
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. Ultimately reduction of sulfur gives sulfide salts:
- 16 Na + S8 → 8 Na2S
Oxides, oxoacids and oxoanions
The principal sulfur oxides are obtained by burning sulfur:
- S + O2 → SO2
- 2 SO2 + O2 → 2 SO3
Sulfur forms a number of sulfur oxoacids, some of which cannot be isolated and are only known through their salts. 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.
Thiosulfate salts (S
3), sometimes referred as "hyposulfites", used in photographic fixing (HYPO) and as reducing agents, feature sulfur in two oxidation states. Sodium dithionite (Na
4), contains the more highly reducing dithionite anion (S
Halides and oxyhalides
The two main sulfur fluorides are
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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.
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.
When sulfur burns in air, it produces 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.
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.
- SO42– → SO32– → H2S → cysteine → methionine
 Sulfur is absorbed by
The so-called flatus) and decomposition products.
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 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 sulfur cycle was the first of the Selman Waksman until the 1950s.
Sulfur metabolism and the sulfur cycle
 Inorganic sulfur forms a part of
Metalloproteins and inorganic cofactors
coenzyme M, CH3SCH2CH2SO3–, the immediate precursor to methane.
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. 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.
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. 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.
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.
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.
Protein and organic cofactors
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.
Small amounts of 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.
Bactericide in winemaking and food preservation
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.
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 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 apple scab under colder conditions. Biosulfur (biologically produced elemental sulfur with hydrophilic characteristics) can be used well for these applications.
Fungicide and pesticide
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.
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.
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. Because atmospheric inputs of sulfur continue to decrease, the deficit in the sulfur input/output is likely to increase, unless sulfur fertilizers are used.
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. 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.
Sulfur reacts directly with methane to give 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.
Other large-scale sulfur chemicals
Because of its importance, sulfuric acid was considered an excellent indicator of a country's industrial well-being. For example with 32.5 million tonnes in 2010, the United States produces more sulfuric acid every year 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.
- 2 S + 3 O2 + 2 H2O → 2 H2SO4
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. 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.
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 sulfuric acid via the Leblanc process by the late 18th century. In industrializing Britain, with the repeal of tariffs 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.
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.
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. In the rasaśāstra tradition, sulfur is called "the smelly" (sanskrit gandhaka, गन्धक).
A natural form of sulfur known as shiliuhuang was known in China since the 6th century BC and found in Hanzhong. By the 3rd century, the Chinese 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. 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.
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; this is mentioned in the Odyssey. 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.
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.
Organosulfur compounds are responsible for 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.
Compounds with carbon–sulfur bonds are uncommon with the notable exception of carbon monoxide, carbon monosulfide is only stable as a dilute gas, as in the interstellar medium.
- Thiols or mercaptans (as they are mercury capturers as chelators) are the sulfur analogs of alcohols; treatment of thiols with base gives thiolate ions.
- Thioethers are the sulfur analogs of ethers.
- 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.
Some of the main classes of sulfur-containing organic compounds include the following:
Allicin, the active ingredient in garlic
Diphenyl disulfide, a representative disulfide
Perfluorooctanesulfonic acid, a controversial surfactant
Dibenzothiophene, a component of crude oil
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. The upgrading of these ores, usually by roasting, is costly and environmentally hazardous. Sulfur corrodes many metals via the process called tarnishing.
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.