Sulfur

Sulfur, also spelled as sulphur,[1] 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,[2] which means "burning stone".Cite error: Invalid `<ref>` tag; refs with no name must have content Almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants from natural gas and petroleum.[3][4] 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.

## Contents

* 1 Characteristics
* 1.1 Physical properties
* 1.2 Chemical properties
* 1.3 Allotropes
* 1.4 Isotopes
* 1.5 Natural occurrence
* 2 Compounds
* 2.1 Electron transfer reactions
* 2.2 Hydrogenation
* 2.3 Combustion
* 2.4 Halogenation
* 2.5 Pseudohalides
* 2.6 Metal sulfides
* 2.7 Organic compounds
* 3 History
* 3.1 Antiquity
* 3.2 In religion and alchemy
* 3.2.1 In religion
* 3.2.2 Sulfur and lightning
* 3.2.3 In alchemy
* 3.2.4 Black sulfur
* 3.3 In medicine
* 3.4 Medieval and Early Modern period
* 3.4.1 In warfare
* 3.4.2 Torture and executions
* 3.4.3 Scientific and Chemical Revolutions
* 3.4.4 Other uses
* 3.5 Modern times
* 3.6 Spelling and etymology
* 4 Production
* 5 Applications
* 5.1 Sulfuric acid
* 5.2 Other important sulfur chemistry
* 5.3 Fertilizer
* 5.4 Fungicide and pesticide
* 5.5 Pharmaceuticals
* 5.6 Batteries
* 6 Biological role
* 6.1 Transferring sulfur between inorganic and biomolecules
* 6.2 Protein and organic metabolites
* 6.3 Metalloproteins and inorganic cofactors
* 6.4 Sulfate
* 6.5 Deficiency
* 7 Precautions
* 7.1 Toxicity of sulfur compounds
* 8 Notes
* 9 See also
* 10 References
* 11 Further reading
* 12 External links

## 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 octasulfur, cyclo-S8. The point group of cyclo-S8 is D4d and its dipole moment is 0 D.[5] Octasulfur is a soft, bright-yellow solid that is odorless.[lower-alpha 1] It melts at 115.21 °C (239.38 °F),[lower-alpha 2] and boils at 444.6 °C (832.3 °F).[2] At 95.2 °C (203.4 °F), below its melting temperature, cyclo-octasulfur begins slowly changing from α-octasulfur to the β-polymorph.[7] 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 (247.3 °F),[8] as it predominantly consists of the β-S8 molecules.[lower-alpha 3] 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.[7] 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.

The sublimation of sulfur becomes noticeable more or less between 20 °C (68 °F) and 50 °C (122 °F), and occurs readily in boiling water at 100 °C (212 °F).[12]

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. Sulfur is also soluble in supercritical carbon dioxide.[13]

### Chemical properties

Under normal conditions, sulfur hydrolyzes very slowly to mainly form hydrogen sulfide and sulfuric acid:

​1⁄2 S8 \+ 4 H2O → 3 H2S \+ H2SO4

Top: Liquid hydrogen sulfide inside a test tube. Bottom: A bottle of sulfuric acid.

The reaction involves adsorption of protons onto S8 clusters, followed by disproportionation into the reaction products.[14]

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 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 except noble gases, even with the notoriously unreactive metal iridium (yielding iridium disulfide).[15] Some of those reactions require elevated temperatures.[16]

### Allotropes

Main article: Chemistry:Allotropes of sulfur

The structure of the cyclooctasulfur molecule, S8

Sulfur forms over 30 solid allotropes, more than any other element.[17] Besides S8, several other rings are known.[18] 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.[19] Larger rings have been prepared, including S12 and S18.[20][21]

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

Main article: Physics:Isotopes of sulfur

Sulfur has 23 known isotopes, four of which are stable: 32S (94.99%±0.26%), 33S (0.75%±0.02%), 34S (4.25%±0.24%), and 36S (0.01%±0.01%).[22] Other than 35S, with a half-life of 87.37 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 [clarification needed], and their composition depends on the type of a stellar explosion. For example, proportionally more 33S comes from novae than from supernovae.[23]

It has been found that the proportion of the two most abundant sulfur isotopes 32S and 34S varies in different samples by a surprisingly large amount. Determination of the isotope ratio (δ34S) in the samples indicates 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.[24] However, there are ongoing discussions over the real reason for the δ34S shifts, biological activity or postdeposit alteration.[25]

For example, when sulfide minerals are precipitated, isotopic equilibration between solid 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[26] 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.[27] 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.[28]

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 (less than a year old) atmospheric sediments in various materials. This isotope may be obtained artificially in different ways. In practice, the reaction 35Cl \+ n → 35S + p is used, irradiating potassium chloride with neutrons.[29] 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.[30]

### 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 A man carrying sulfur blocks from Kawah Ijen, a volcano in East Java, Indonesia, 2009

32S is produced by stellar nucleosynthesis inside massive stars, at a depth where the temperature exceeds 2.5×109 kelvins (4.50×109 °F), by the fusion of one nucleus of silicon plus one nucleus of helium.[31] Because this nuclear reaction is part of the alpha process, which produces highly abundant elements, sulfur is the 10th most common element in the universe.

Sulfur, usually in the sulfide oxidation state, 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.[32] The distinctive colors of Jupiter's volcanic moon Io are attributed to various forms of molten, solid, and gaseous sulfur.[33] 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.[34]

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 22 cm × 16 cm × 11 cm (8.7 in × 6.3 in × 4.3 in).[35] Historically, Sicily was a major source of sulfur in the Industrial Revolution.[36] 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.[37]

Native sulfur is synthesized by anaerobic bacteria acting on sulfate minerals such as gypsum in salt domes.[38][39] 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.[40] 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.

Petroleum and natural gas are the main industrial sources of sulfur.[3]

## 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, S2+
8, S2+
4 and S2+
19 are produced when sulfur is reacted with oxidizing agents in a strongly acidic solution.[41] 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. S2+
8 is deep blue, S2+
4 is yellow and S2+
19 is red.[7]

When sulfur is dissolved in oleum, a dark blue solution containing S82+ is obtained

Reduction of sulfur gives various polysulfides with the formula S2−x, many of which have been obtained in crystalline form. Illustrative is the production of sodium tetrasulfide:

4 Na + S
8 → 2 Na

Some of these dianions dissociate to give radical anions. For instance, S−
3 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[42]

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.[43] 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:[2]

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.[44] 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.[45] Bromine also oxidizes sulfur to form sulfur dibromide and disulfur dibromide.[45]

### 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.[46] Geoscientists also study the isotopes of metal sulfides in rocks and sediment to study environmental conditions in the Earth's past.[47]

### Organic compounds

Main article: Chemistry:Organosulfur compounds

* Illustrative organosulfur compounds
* (L)-Cysteine, an amino acid containing a thiol group

* Thiamine or vitamin B1

* Biotin or vitamin B7

* Penicillin, an antibiotic ("R" is the variable group)

* Allicin, a chemical compound in garlic

* Diphenyl disulfide, a representative disulfide

* Dibenzothiophene, a component of crude oil

* Perfluorooctanesulfonic acid (PFOS), a surfactant

Some of the main classes of sulfur-containing organic compounds include the following:[48]

* 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.[49] Unlike carbon monoxide, carbon monosulfide is stable only as an extremely dilute gas, found between solar systems.[50]

Organosulfur compounds are responsible for some of the unpleasant odors of decaying organic matter. They are widely known as the odorant in domestic natural gas, garlic odor, and skunk spray, as well as a component of bad breath odor. Not all organic sulfur compounds smell unpleasant at all concentrations: the sulfur-containing monoterpenoid grapefruit mercaptan in small concentrations is the characteristic scent of grapefruit, but has a generic thiol odor at larger concentrations. Sulfur mustard, a potent vesicant, was used in World War I as a disabling agent.[51]

Sulfur–sulfur bonds are a structural component used to stiffen rubber, similar to the disulfide bridges that rigidify proteins (see biological below). In the most common type of industrial "curing" or hardening and strengthening of natural rubber, elemental sulfur is heated with the rubber to the point that chemical reactions form disulfide bridges between isoprene units of the polymer. This process, patented in 1843, [52] 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

Roman brooch in the form of a panther (copper alloy, 100-300 AD). The spots are made of niello, a black mixture of metal and sulfur.

According to the Ebers Papyrus, a sulfur ointment was used in ancient Egypt to treat granular eyelids. There are ancient Egyptian necklaces, dated between the Ptolemaic period and the early Roman period, with beads made of sulfur, shaped like flowers and cow heads.[53][54]

Sulfur is one of the ingredients of niello, a dark material used to decorate metallic items. Pliny the Elder mentioned, in the first century AD, that the Romans used a type of niello inherited from the Egyptians, who, according to him, made a dark pigment to depict Anubis on silver vessels. He lists a recipe where copper and silver are mixed with sulfur and heated to produce this material.[55]

In Roman jewelry, sulfur was used as a filler in hollow gold items, especially in gold rings. This practice was widespread enough to be included by Artemidorus in the section of the Oneirocritica about gold finger-rings, where he mentions rings "that are hollow and filled with brimstone". Modern analysis has confirmed this, detecting sulfur in ancient Roman gold artifacts.[56]

In the 2nd century BC, Cato the Elder gave a recipe to protect vines against caterpillars, which contains sulfur as an ingredient, in his De Agri Cultura, an early example of its use as pest control.[57]

Pliny the Elder further discussed sulfur in his Natural History, saying that its best-known source is the island of Melos. He listed four types of sulfur: one used for medicinal purposes, one used by fullers, one used to bleach wool and one used for lamp wicks. He noted its use in Ancient Rome in religious ceremonies to purify homes by fumigation (something also briefly evoked in a line of Ovid's Ars Amatoria about purifying houses with eggs and sulfur).[58][59] The use of sulfur for fumigation is very ancient and was already a practice in preclassical Greece;[60] this is mentioned in the Odyssey.[61]

Several Roman authors of the second half of the first century, such as Martial and Statius, reference a type of street trade seemingly common at the time, where peddlers in the streets would exchange pieces of sulfur or sulfur "matches" (actually some kind of sulfured wooden splints used as a fire starter, rather than friction matches) for broken glass.[62][63][64][65]

The Persian army used sulfur as a chemical weapon during the siege of Dura-Europos, in the 3rd century. They burned pitch and sulfur in tunnels, creating a toxic cloud of sulfur dioxide which killed the Roman soldiers in the tunnels.[66]

A natural form of sulfur known as shiliuhuang (Chinese: 石硫黃; pinyin: shí liú huáng) was known in China since the 6th century BC and found in Hanzhong.[67] By the 3rd century, the Chinese had discovered that sulfur could be extracted from pyrite.[67] 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.[67]

### In religion and alchemy

#### In religion

Sulfur has a long history of religious associations and spiritual symbolism. In Antiquity, it was associated with the divine: used in fumigation to purify and prevent contagion, it was considered divine incense, and its name in Ancient Greek, theion (θεῖον) resembled a form of the word theios (θεῖος), divine.[68] In the Iliad, Achilles is portrayed cleansing a cup with sulfur before making a libation to Zeus.[69]

In the Christian Bible, sulfur is consistently associated with God's wrath, purification of evil and the punishment of sinners. The most well-known example is in the Book of Genesis, with the account of God's destruction of the cities of Sodom and Gomorrah by "brimstone and fire from the LORD out of heaven".[70] Likewise, in Revelation, the Devil is cast in a lake of fire and brimstone (hence the idiom). Subsequently, sulfur's role as hellfire made it a symbol of Hell, its powerful, unpleasant smell becoming a feature of the Devil and demons in folklore and popular imagination.

#### Sulfur and lightning

Sulfur’s role as a symbol of divine wrath and punishment may come from its association with lightning. For centuries, lightning was thought to cause a sulfur smell when it struck, and was therefore linked with it, including in religious contexts.[71] In the 12th century, the Byzantine scholar John Tzetzes discussed the different associations around the word theion in his Chiliades, mentioning different meanings such as sulfur, incense (which he links to the breath of the immortals) and thunder, as well as it referring to "bewildering" and "amazing" things, the stars, souls, and angels. He explains that thunder is divine fire, and that it gives off a strong smell resembling sulfur.[72] Lightning being sulfurous is a very ancient belief, mentioned by ancient Roman authors and present in the works of Homer, where descriptions of the divine lightning of Zeus mention sulfur several times (in the Iliad, it strikes in front of the horses of Diomedes, causing a smell of sulfur and flames, at another point, Hector is compared to an oak uprooted by Zeus' lightning, smelling of sulfur, while in the Odyssey, it strikes a ship, filling it with sulfurous smoke).[73][74][75][76] The 18th century theologian Jonathan Edwards considered lightning to be "a string of brimstone".[77]

Lightning does not actually contain sulfur or produce sulfur odorants. Ozone would later be identified, in the 19th century, as the element behind the smell of lightning, but even then, people unfamiliar with ozone were said to still commonly misidentify the smell as brimstone. In a somewhat ironic twist, after the discovery of ozone, a popular Victorian belief started to attribute health benefits to sea air due to its smell, which was thought to be caused by a high ozone content. That smell is in fact caused by dimethyl sulfide, a sulfur compound.[78][79]

#### In alchemy

Sulfur has also been a major ingredient in alchemy since its early days, different branches of Indian, Muslim and European alchemists ascribing it esoteric symbolism based on its strong bond with mercury and the interactions between the two elements. Greco-Alexandrian practitioners of Hellenistic alchemy like Mary the Jewess and Zosimos of Panopolis were interested in sulfur and sulfur compounds, and mentioned them in their writings.[80] Indian alchemists, practitioners of the "science of chemicals" (Sanskrit: रसशास्त्र, romanized: rasaśāstra), wrote extensively about the use of sulfur in alchemical operations with mercury, from the eighth century AD onwards.[81] In the rasaśāstra tradition, sulfur is called "the smelly" (गन्धक, gandhaka), and is thought to represent the feminine principle, the menstrual blood of the goddess Gauri (Shakti), while mercury is its male counterpart, the semen of Shiva.[82][83] Alchemists in Europe and the Islamic world (basing themselves off the "sulfur-mercury theory of metals" from Arabic alchemical texts of the ninth century) had the opposite imagery in their traditions, with sulfur representing the masculine principle and mercury the feminine one, mixing together to form various metals. This theory of metals was very influential during the Middle Ages, and persisted until the eighteenth century, when Lavoisier proved that metals were distinct elements.[84][85][86]

In the sixteenth century, Paracelsus added Salt as a third element to the Sulfur-Mercury dyad, making them the three basic pillars of alchemy according to his theory: the Tria Prima.[87] Like most alchemical terms, the word "Sulfur" came to mean a lot of different things in alchemical works beyond its standard definition, representing concepts such as the active, masculine or fixing principle, combustion, or the soul. There were multiple "sulfurs", as the word was used as a Deckname, a secret codename for other substances.[88][89][90]

#### Black sulfur

Sulfur

Black Sulfur

Alchemical signs for sulfur, or the combustible elements, and black sulfur, an impure type of sulfur also known as horse brimstone (sulfur caballinum) [91][92]

Early European alchemists gave sulfur an alchemical symbol of a triangle atop a cross (🜍). Another type of sulfur mentioned in alchemy, black sulfur (sulfur nigrum), also had its own symbol combining a two-barred cross atop a lemniscate (🜏). This symbol was later used by Anton LaVey as a satanist symbol, the "Leviathan cross", as he associated it with the brimstone of Hell.[93] This has led to some confusion and misconceptions about its original alchemical meaning: in alchemical texts, black sulfur was not a synonym for "brimstone" or an alternative symbol for sulfur as one of the Three Primes, but was instead considered a distinct type of sulfur, listed separately. Black sulfur, also known as grey sulfur (sulfur griseum) or "horse brimstone" (sulfur caballinum), was a grey, impure form of sulfur, usually the dregs of sulfur purification, which was used for medicinal purposes despite often containing arsenic.[94][95][96] Iron shavings or hammerscale were sometimes added to it in its manufacture.[97][98] The "horse" in the name is thought to reference its veterinary use: low quality sulfur, only fit to treat horses.[99] In the 19th century, it was sometimes also sold under the name "sulphur of ivy" (a corruption of sulfur vivum) and used on hops.[100]

### In medicine

Sulfur has antifungal, antibacterial, and keratolytic activity; in the past it was used against acne vulgaris, rosacea, seborrheic dermatitis, dandruff, pityriasis versicolor, scabies, and warts.[101] This 1881 advertisement baselessly claims efficacy against rheumatism, gout, baldness, and graying of hair.

The main use of sulfur for therapeutic purposes has been, throughout history, to cure people and domestic animals of a wide range of skin conditions and parasites. To this day, it remains a popular treatment to alleviate such conditions as scabies, ringworm, psoriasis, eczema, and acne. The exact mechanism of action is unknown, but it is thought to react with the cysteine in the skin to produce hydrogen sulfide, which has antifungal, antioxidant and keratolytic properties, as well as converting to pentathiolic acid, which is antifungal, when applied to the skin.[102][103][104][105]

In 1696, Haarlem oil (colloquially known as "Dutch drops"), a sulfur-based potion, was created in Haarlem. Initially sold as a kidney diuretic, it acquired a reputation for being a panacea, and was popularized in other countries by sailors, among whom it was widely used. It is still commercialized three centuries later, and is now often used as a health supplement for horses.[106][107]

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

In the 19th century, it was common to give children a mixture of sulfur and treacle as a remedy. The sulfur in it was meant to act as a spring tonic and a laxative.[108][109] Reportedly foul-tasting, "brimstone and treacle" was considered by many an unpleasant childhood experience, and it is mentioned as such in a number of cultural references, including Mary Poppins, the works of Charles Dickens, and the controversial 1976 British play of the same name.[110][111][112]

Madam C. J. Walker, who became the first female millionnaire by selling hair products for Black women in the early 1900s, notably used sulfur as an ingredient in her products.[113]

Sulfur was also used in fumigations against the plague. In the early 20th century, during the Third Plague Pandemic, devices like the Clayton machine, used in harbors across the world, or the Sulfurozador in Latin America, were used to spread sulfur dioxide gas to kill insects, rats and pathogens.[114][115]

### Medieval and Early Modern period

#### In warfare

Sulfur has been hypothesized to have been part of the now-lost recipe for Greek fire, a powerful naval incendiary weapon used by the Byzantine Empire between the 7th and 15th century, but this remains debated. Regardless, it was present in Byzantine incendiary weapons, as the 12th century Byzantine princess and historian Anna Komnene mentioned a flammable mixture of pine resin and sulfur, blown through reed pipes to burn the enemy, in her Alexiad.[116][117]

Sulfur was a tool of medieval siege warfare. Raymond of Aguilers described incendiary mixtures containing pitch, wax, sulfur and tow being thrown at the Crusaders during the siege of Jerusalem in 1099, and William of Tyre mentioned darts set on fire with sulfur, paste and oil.[118]

In the 10th century, princess Olga of Kiev assieged the city of Iskorosten, seeking revenge against the Drevlians for the murder of her husband. Olga told her enemies they would only need to pay a small tribute if they surrendered: three pigeons and three sparrows from each household. They complied, and at night, she ordered her army to release the birds with pieces of hot sulfur tied to their legs. The birds flew back to their previous nests, setting them aflame and burning down the city.[119]

One of the most significant historical applications of sulfur has been its role in the manufacture of black powder, which is a mixture of saltpeter, charcoal and sulfur. Black powder is thought to have been accidentally discovered by Chinese alchemists using saltpeter and sulfur as materials in their experiments to make gold or prepare an elixir of life. The Zhenyuan Miaodao Yaolue, a Taoist alchemical text from the Late Tang era, warned of a flammable mixture of saltpeter, realgar, sulfur, and honey burning the alchemists preparing it. In China, black powder was first used in fireworks for amusement and religious purposes, as well as in construction works, to blast rocks in order to open roads or waterways; no mention of it being used on the batterfield is made before the 10th century. In weapons, it was first used in incendiary projectiles such as fire arrows and grenades, later joined by the fire lance, an early firearm. The Wujing Zongyao of 1044 described formulas for Chinese black powder and listed a number of Chinese gunpowder weapons, accompanied by illustrations.[120][121][122]

#### Torture and executions

The use of sulfur has been reported in some Early Modern executions by burning. In 1553, in Geneva, Spanish theologian Michael Servetus was condemned to be burned at the stake for heresy, wearing a wreath of straw and leaves sprayed with sulfur.[123] In France, some of the people convicted of offenses punishable by burning were given a sulfur shirt ("chemise soufrée") to wear at the stake. This kind of garment has been mentioned in convictions for witchcraft, as with Urbain Grandier in 1634,[124] and convictions for homosexuality, as with Jean Diot and Bruno Lenoir in 1750.[125]

Sulfur was also part of the French punishment for regicide. The hand that held the assassination weapon would be scorched with flaming sulfur; the body of the culprit would be torn with pincers, and the wounds covered with molten lead, boiling oil, pitch, hot resin, wax and sulfur, before being quartered by horses. Such was the fate of Ravaillac for murdering Henry IV of France in 1610, and of Robert-François Damiens for his failed assassination attempt targeting Louis XV in 1757.[126][127]

A sulfur-based torture device was used in 18th century Brittany by the local authorities. "Escarpins soufrés" or "chaussons soufrés" (meaning "sulfured slippers") were a pair of shoes, usually made of leather, smeared with sulfur to burn the victim's feet during interrogation.[128]

#### Scientific and Chemical Revolutions

Guericke's sulfur globe experiments.

In 1660, German scientist Otto von Guericke studied static electricity by building a device comprising a large, rotating sulfur globe, now regarded as the first electrostatic generator.[129]

Sulfur appeared in the 1718 "affinity table" of the French chemist Étienne-François Geoffroy, a document that would become influential in chemistry works of the 18th century. Also present in that table was the so-called "Sulfur Principle": at the time, many chemists were convinced that sulfur was not a standalone element but instead contained multiple substances, including a distinct flammable one. This was supported by the phlogiston theory of combustion, which had emerged few decades prior; Geoffroy would later consider the sulfurous principle in his table to be phlogiston. Widely accepted for a century, phlogiston theory was debunked by the work of Antoine Lavoisier on combustion and oxygen.[130][131]

Antoine Lavoisier used sulfur in his combustion experiments, writing of some of these in 1777.[132] In his 1789 Traité Élémentaire de Chimie, considered to be the first modern chemistry textbook, Lavoisier placed sulfur as its own chemical element in a "table of simple substances".[133]

#### Other uses

Sulfur fumigation could be used to disinfect wine barrels and prevent wine spoilage: In the mid-15th century, under the reign of Frederick III, several imperial decrees forbade the use of sulfur in wine barrels in Germany (thus implying that it was an attested practice). The ban was relaxed by his successor, Maximilian I. The practice then spread to France; in the 17th century, French wines were known to be sulfured when exported, French white wines being particularly affected. The Netherlands, importer of French wine, had their own sulfuring process using "allumettes hollandaises" (Dutch wicks), and were said to excessively sulfur the wine. That type of heavily sulfured wine was called "vin muet" in French, and "stomme" in Dutch, which both translate to "silent wine". Mixing some stomme with wine would make the wine sweeter, as well as making it last longer. However, as sulfuring was known to alter the taste and to potentially transfer impurities like arsenic to the wine, there were some health concerns around it. The Netherlands therefore decided to ban stomme in 1613, but the practice nevertheless perdured.[134]

In the late 18th century, furniture makers of German and Swiss origin in Pennsylvania, Maryland, Virginia and North Carolina used molten sulfur to produce decorative inlays. This type of inlay was mostly used on chests, schranks and clocks, but a wide range of surviving examples can be found, including tables, cupboards, cradles and many more. Similarly, a number of 18th century Lutheran tombstones, most of which with epitaphs in German, have been found in Maryland with a similar sulfur inlay. The origin of this technique is unclear, but it may have been inspired by the German use of marzipan sulfur molds (Schwefelform). Molten sulfur was poured in casts with designs carved, then solidified into that shape, creating a mold that could be used to make decorative marzipan pieces. The designs in the molds could be very detailed and ornate. The practice lasted until the 1950s in Central Europe, and some of the molds can still be found in German and Swiss museums.[135][136][137][138]

Molten sulfur was sometimes used in construction from the 18th century on, as an easier and cheaper alternative to molten lead for purposes such as anchoring bolts into stone or concrete.[139][140] There are mentions of an 18th century French practice in La Rochelle and Île de Ré, where the stones in some structures would be secured together using sheep or ox tibia bones encased in molten sulfur to anchor them, the use of other materials such as metal being deemed unsuitable due to seaside corrosion.[141]

### Modern times

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.[142][143]

In 1867, elemental sulfur was discovered in underground deposits in Louisiana and Texas. The highly successful Frasch process was developed to extract this resource, and the United States replaced Sicily as the main producer of sulfur worldwide in the early 20th century.[144]

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.[145]

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.[3]

### 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 Ancient Greek word for sulfur, θεῖον, theîon (from earlier θέειον, théeion), is the source of the international chemical prefix thio-.

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.[146] In the U.S., Webster's Dictionary (1828) lists only sulphur,[147] and sulfur only becomes the primary entry in 1961. Nonetheless, the form sulfur was widely adopted in the United States during the 19th century.[148] Canada uses both spellings.

The IUPAC adopted the spelling sulfur in 1990.[1][149] UK institutions soon followed: the Royal Society of Chemistry in 1992;[150] the British Standards Institute in 1993; the UK national school curriculum in 2000[151][152] (there was some controversy).[153] It is now the more common spelling in British books.[148] The World Health Organization has mandated sulfur since at least 1993.[154] The 1989 edition of the Oxford English Dictionary gave sulphur as the headword, and listed sulfur as a U.S. spelling;[155] the current edition gives sulfur first,[156] which indicates that it is the usual British spelling.[157]

## 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.[158] 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."[159] Sulfur is still mined from surface deposits in poorer nations with volcanoes, such as Indonesia, and problems with working conditions still exist.[160]

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.[40] 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.[161][162]

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

Since then, sulfur has typically been produced from petroleum, natural gas, and related fossil resources, from which it is obtained mainly as hydrogen sulfide.[3] Organosulfur compounds, undesirable impurities in petroleum, may be upgraded by subjecting them to hydrodesulfurization, which cleaves the C–S bonds:[161][162]

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, which entails oxidation of some hydrogen sulfide to sulfur dioxide and then the comproportionation of the two:[161][162]

3 O2 \+ 2 H2S → 2 SO2 \+ 2 H2O

SO2 \+ 2 H2S → 3 S + 2 H2O

Production and price (US market) of elemental sulfur

Due to the high sulfur content of the Athabasca Oil Sands, stockpiles of elemental sulfur from this process exist throughout Alberta, Canada.[163] Another way of storing sulfur is as a binder for concrete, the resulting product having some desirable properties (see sulfur concrete).[164]

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).[165] Production has been slowly increasing from 1900 to 2010; the price was unstable in the 1980s and around 2010.[166]

## Applications

### Sulfuric acid

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

8 \+ 12 O
2 \+ 8 H
2O → 8 H
2SO

Sulfuric acid production in 2000

In 2010, the United States produced more sulfuric acid than any other inorganic industrial chemical.[166] 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.[40]

### Other important sulfur chemistry

Sulfur reacts directly with methane to give carbon disulfide, which is used to manufacture cellophane and rayon.[40] 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, abiotic processes involving atmospheric oxygen as well as biological processes involving soil bacteria can oxidize and convert elemental sulfur to soluble compounds, which can in turn be used by microorganisms and plants. Sulfur improves the efficiency of other essential plant nutrients, particularly nitrogen and phosphorus.[167] 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 plant 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.[168][169][170] 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.[171][167]

### Fungicide and pesticide

Sulfur candle originally sold for home fumigation

Elemental sulfur is one of the oldest fungicides and pesticides. "Dusting sulfur", elemental sulfur in powdered form, is a common fungicide for grapes, strawberry, many vegetables and several other crops. It has a good efficacy against a wide range of powdery mildew diseases as well as black spot. In organic production, sulfur is the most important fungicide. It is the only fungicide used in organically farmed apple production against the main disease apple scab under colder conditions. Biosulfur (biologically produced elemental sulfur with hydrophilic characteristics) can 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.[164][172] 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

Main article: Medicine:Sulfur (pharmacy)

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.[173] 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.[174]

Many drugs contain sulfur.[175] 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.[48]

### 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.[176]

## Biological role

Sulfur is an essential component of all living cells. It is the eighth most abundant element in the human body by weight,[177] about equal in abundance to potassium, and slightly greater than sodium and chlorine.[178] A 70 kg (150 lb) human body contains about 140 grams (4.9 oz) of sulfur.[179] The main dietary source of sulfur for humans is sulfur-containing amino acids,[180] which can be found in plant and animal proteins.[181]

### Transferring sulfur between inorganic and biomolecules

See also: Biology: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).[182] Another contributor, who continued to study it was Selman Waksman.[183] Primitive bacteria that live around deep ocean volcanic vents oxidize hydrogen sulfide for their nutrition, as discovered by Robert Ballard.[4]

Sulfur oxidizers can use as energy sources reduced sulfur compounds, including hydrogen sulfide, elemental sulfur, sulfite, thiosulfate, and various polythionates (e.g., tetrathionate).[184] 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: green sulfur bacteria and purple sulfur bacteria.[185] 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.[186]

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.[187][188][189] 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.[190]

SO2−
4 → SO2−
3 → H
2S → 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.[191][192]

### 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.[193] Some of the sulfur is present in certain metabolites—many of which are cofactors—and sulfated polysaccharides of connective tissue (chondroitin sulfates, heparin).

Schematic representation of disulfide bridges (in yellow) between two protein helices

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.[194] Within the cytoplasm, disulfide bonds may instead be reduced (i.e. in -SH form) to their constituent cysteine residues by thioredoxins.[195]

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.[196] 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.[197][198] Two of the 13 classical vitamins, biotin and thiamine, contain sulfur, and serve as cofactors to several enzymes.[199][200] 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. Sulfur atoms may be catenated on the sulfhydryl groups of cysteine or glutathione to form a hydropersulfide (–SSH), which act as signaling molecules and donors of sulfur atoms in the biosynthesis of iron-sulfur clusters and some sulfur containing metabolites. Polysulfides and elemental sulfur are also known to exist in both prokaryotes and eukaryotes.[201]

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, CH
3SCH
2CH
2SO−
3, the immediate precursor to methane.[202]

### 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[203]—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,[204] 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.[205]

Sulfur is also present in molybdenum cofactor.[206]

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

### Sulfate

### 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.[citation needed]

Isolated sulfite oxidase deficiency is a rare, fatal genetic disease caused by mutations to sulfite oxidase, which is needed to metabolize sulfites to sulfates.[207]

## Precautions

Sulfur  Hazards

GHS pictograms |
GHS Signal word | Warning

GHS hazard statements

| H315[208]
NFPA 704 (fire diamond)  |

[209]

Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Infobox references

Effect of acid rain on a forest, thumb 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,[210] and cases of life-threatening metabolic acidosis have been reported after patients deliberately consumed sulfur as a folk remedy.[211][212]

### 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.[213] In organisms without lungs such as insects, sulfite in high concentration prevents respiration.[214]

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.[215]

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).[216] 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),[217] 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.[218] However, its presence in ambient air at concentration over 100–150 ppm quickly deadens the sense of smell,[219] 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.

## Notes

1. ↑ But impure samples have an odor similar to that of matches. A strong odor called "smell of sulfur" actually is given off by several sulfur compounds, such as hydrogen sulfide and organosulfur compounds.
2. ↑ 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).[6] Greenwood and Earnshaw say that at fast heating for microcrystalline α-S8 the melting point is 115.1 °C (239.2 °F).[2]
3. ↑ Historically, it was rather difficult to find the exact melting point of sulfur.[9] When heated slowly, the melting point may range from 114.6 °C (238.3 °F) to 120.4 °C (248.7 °F)[2] (factors that interfere with a definite melting point, are the polymerlike nature of sulfur[10] and a large number of allotropes.[11]) Melting point may be presented as a temperature range, depending on the allotropic composition of a sample at the time of melting.

## See also

* Blue lava
* Stratospheric sulfur aerosols
* Sulfur assimilation
* Sulfur isotope biogeochemistry
* Ultra-low-sulfur diesel

## References

1. ↑ 1.0 1.1 "So long sulphur". Nature Chemistry 1 (5): 333. 4 August 2009. doi:10.1038/nchem.301. PMID 21378874. Bibcode: 2009NatCh...1Q.333..
2. ↑ 2.0 2.1 2.2 2.3 2.4 Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. ISBN 0-7506-3365-4.
3. ↑ 3.0 3.1 3.2 3.3 Laurence Knight (Jul 19, 2014). "Sulphur surplus: Up to our necks in a diabolical element". BBC. https://www.bbc.com/news/magazine-28369829.
4. ↑ 4.0 4.1 "Sulfur". BBC. Oct 11, 2014. https://www.bbc.co.uk/programmes/p027xnp2. . Download here.
5. ↑ Rettig, S. J.; Trotter, J. (15 December 1987). "Refinement of the structure of orthorhombic sulfur, α-S8". Acta Crystallographica Section C 43 (12): 2260–2262. doi:10.1107/S0108270187088152. Bibcode: 1987AcCrC..43.2260R.
6. ↑ "Sulfur | S (Element) - PubChem". https://pubchem.ncbi.nlm.nih.gov/element/Sulfur.
7. ↑ 7.0 7.1 7.2 Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. pp. 645–665. ISBN 978-0-08-037941-8.
8. ↑ Poling, Bruce E.; Prausnitz, John M.; O'Connell, John P. (2000-11-27) (in en). The Properties of Gases and Liquids 5E. McGraw Hill Professional. ISBN 978-0-07-149999-6. https://books.google.com/books?id=9tGclC3ZRX0C.
9. ↑ "7.5: Changes of State" (in en). 2013-10-03. https://chem.libretexts.org/Bookshelves/General_Chemistry/Chem1_(Lower)/07:_Solids_and_Liquids/7.05:_Changes_of_State.
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19. ↑ Tebbe, Fred N.; Wasserman, E.; Peet, William G.; Vatvars, Arturs; Hayman, Alan C. (1982). "Composition of Elemental Sulfur in Solution: Equilibrium of S6, S7, and S8 at Ambient Temperatures". Journal of the American Chemical Society 104 (18): 4971–4972. doi:10.1021/ja00382a050. Bibcode: 1982JAChS.104.4971T.S6,+S7,+and+S8+at+Ambient+Temperatures&rft.jtitle=Journal+of+the+American+Chemical+Society&rft.aulast=Tebbe&rft.aufirst=Fred+N.&rft.au=Tebbe, Fred+N.&rft.au=Wasserman, E.&rft.au=Peet, William+G.&rft.au=Vatvars, Arturs&rft.au=Hayman, Alan+C.&rft.date=1982&rft.volume=104&rft.issue=18&rft.pages=4971–4972&rft_id=info:doi/10.1021/ja00382a050&rft_id=info:bibcode/1982JAChS.104.4971T&rfr_id=info:sid/en.wikibooks.org:Chemistry:Sulfur">
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29. ↑ Kim, Ik Soo; Kwak, Seung Im; Park, Ul Jae; Bang, Hong Sik; Han, Hyun Soo (2005-07-01). "Production of Sulfur-35 by the Cation Exchange Process". 2005 autumn meeting of the KNS, Busan (Korea, Republic of), 27–28 Oct 2005.
30. ↑ "Sulfur-35 (35 S) safety information and specific handling precautions". https://ehs.yale.edu/sites/default/files/files/radioisotope-s35.pdf.
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35. ↑ Rickwood, P. C. (1981). "The largest crystals". American Mineralogist 66: 885–907. http://www.minsocam.org/ammin/AM66/AM66_885.pdf.
36. ↑ Kutney, Gerald (2007). Sulfur: history, technology, applications & industry. Toronto: ChemTec. p. 43. ISBN 978-1-895198-37-9. OCLC 79256100.
37. ↑ de Ronde, C. E. J.; Chadwick, W. W. Jr.; Ditchburn, R. G.; Embley, R. W.; Tunnicliffe, V.; Baker, E. T.; Walker, S. L.; Ferrini, V. L. et al. (2015). "Molten Sulfur Lakes of Intraoceanic Arc Volcanoes". Volcanic Lakes. Springer. pp. 261–288. doi:10.1007/978-3-642-36833-2. ISBN 978-3-642-36832-5. Bibcode: 2015vola.book.....R.
38. ↑ Klein, Cornelis; Hurlbut, Cornelius S. Jr. (1985). Manual of Mineralogy (20th ed.). Wiley. pp. 265–66. ISBN 0-471-80580-7.
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40. ↑ 40.0 40.1 40.2 40.3 Nehb, Wolfgang; Vydra, Karel (2006). "Sulfur". Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH Verlag. doi:10.1002/14356007.a25_507.pub2. ISBN 978-3-527-30673-2.
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46. ↑ Vaughan, D. J.; Craig, J. R. (1978). Mineral Chemistry of Metal Sulfides. Cambridge: Cambridge University Press. ISBN 0-521-21489-0.
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## Further reading

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

* Sulfur at The Periodic Table of Videos (University of Nottingham)
* Atomic Data for Sulfur, NIST Physical Measurement Laboratory
* Sulfur phase diagram , Introduction to Chemistry for Ages 13–17
* Crystalline, liquid and polymerization of sulfur on Vulcano Island, Italy
* Sulfur and its use as a pesticide
* The Sulphur Institute
* Nutrient Stewardship and The Sulphur Institute

* v
* t
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Periodic table

|  | 1 | 2 |  | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 | 11 | 12 | 13 | 14 | 15 | 16 | 17 | 18

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1 | H |  | He
2 | Li | Be |  | B | C | N | O | F | Ne
3 | Na | Mg |  | Al | Si | P | S | Cl | Ar
4 | K | Ca |  | Sc | Ti | V | Cr | Mn | Fe | Co | Ni | Cu | Zn | Ga | Ge | As | Se | Br | Kr
5 | Rb | Sr |  | Y | Zr | Nb | Mo | Tc | Ru | Rh | Pd | Ag | Cd | In | Sn | Sb | Te | I | Xe
6 | Cs | Ba | La | Ce | Pr | Nd | Pm | Sm | Eu | Gd | Tb | Dy | Ho | Er | Tm | Yb | Lu | Hf | Ta | W | Re | Os | Ir | Pt | Au | Hg | Tl | Pb | Bi | Po | At | Rn
7 | Fr | Ra | Ac | Th | Pa | U | Np | Pu | Am | Cm | Bk | Cf | Es | Fm | Md | No | Lr | Rf | Db | Sg | Bh | Hs | Mt | Ds | Rg | Cn | Nh | Fl | Mc | Lv | Ts | Og

| Alkali metal | Alkaline earth metal | Lan­thanide | Actinide | Transition metal | Post-​transition metal | Metalloid | Reactive nonmetal | Noble gas

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Sulfur compounds

* Al2S3
* As2S2
* As2S3
* As2S5
* As4S4
* Au2S
* Au2S3
* B2S3
* BaS
* BeS
* Bi2S3
* Br2S
* Br2S2
* CS2
* C3S2
* C6S6
* CaS
* CdS
* CeS
* SCl2
* S2Cl2
* CoS
* Cr2S3
* CuS
* D2S
* Dy2S3
* Er2S3
* EuS
* SF4
* S2F10
* SF6
* FeS2
* GaS
* H2S
* HfS2
* HgS
* InS
* I2S
* LaS
* LiS
* MgS
* MoS3
* NiS
* SO2
* SO3
* P4S7
* PbS
* PbS2
* PtS
* ReS2
* SrS
* TlS
* SV
* SeS2
* S2U
* WS2
* Sb2S3
* Sb2S5
* Sm2S3
* Y2S3
* Ag2SO4
* SOBr2
* CSTe
* C2H4S
* C2H6S3
* C4H4S
* CaSO4
* C32H66S2
* CuFeS2
* H2SO4
* H2SO3
* F2OS
* NaHS
* K2SO3
* O3S3Sb4
* Yb2(SO4)3
* AlKO8S2
* CHCl3S
* KSCN
* CdSO3
* PSCl3
* SOCl2
* Cs2O4S
* Re2S7
* Na2S
* K2S
* H2S2O7
* H2SO5
* NH5S
* HgSO4
* K2SO4
* RaSO4
* SnSO4
* SrSO4
* Zr(SO4)2
* Ti(SO4)2
* Tm2(SO4)3
* AlNa(SO4)2
* Er2(SO4)3
* Eu2(SO4)3
* CHNS
* Co(SCN)2
* C2H3SN
* PSI3
* ZrS2
* SiS
* CSSe

Chemical formulas

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* Sulfur
* Period 3 elements
* Chalcogens
* Chemical elements (sorted alphabetically)
* Chemical elements (sorted by number)

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