Sulfur is the tenth most abundant element by mass in the universe and the fifth most common onEarth. Though sometimes found in pure,native form, sulfur on Earth usually occurs assulfide andsulfate minerals. Being abundant in native form, sulfur was known in ancient times, being mentioned for its uses inancient India,ancient Greece,China, andancient Egypt. Historically and in literature sulfur is also calledbrimstone,[10] which means "burning stone".[11] Almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants fromnatural gas andpetroleum.[12][13] The greatest commercial use of the element is the production ofsulfuric acid for sulfate and phosphatefertilizers, and other chemical processes. Sulfur is used inmatches,insecticides, andfungicides. Many sulfur compounds are odoriferous, and the smells of odorized natural gas,skunk scent,bad breath,grapefruit, andgarlic are due toorganosulfur compounds.Hydrogen sulfide gives the characteristic odor to rotting eggs and other biological processes.
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 isoctasulfur, cyclo-S8. Thepoint group of cyclo-S8 is D4d and its dipole moment is 0 D.[14] Octasulfur is a soft, bright-yellow solid that is odorless.[a] It melts at 115.21 °C (239.38 °F),[b] and boils at 444.6 °C (832.3 °F).[10] At 95.2 °C (203.4 °F), below its melting temperature, cyclo-octasulfur begins slowly changing from α-octasulfur to the β-polymorph.[16] 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),[17] as it predominantly consists of the β-S8 molecules.[c] Between its melting and boiling temperatures, octasulfur changes its allotrope again, turning from β-octasulfur to γ-sulfur, again accompanied by a lower density but increasedviscosity due to the formation ofpolymers.[16] 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.
Thesublimation 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).[21]
Sulfur is insoluble in water but soluble incarbon disulfide and, to a lesser extent, in othernonpolar organic solvents, such asbenzene andtoluene. Sulfur has been found to be soluble in super-critical carbon dioxide.[22]
The second, fourth and sixthionization 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. Applyingcatalysts and/orsupply 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 ofsulfur 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 lesserelectronegativity, 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 metaliridium (yieldingiridium disulfide).[24] Some of those reactions require elevated temperatures.[25]
Sulfur forms over 30 solidallotropes, more than any other element.[26] Besides S8, several other rings are known.[27] 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.[28] Larger rings have been prepared, including S12 and S18.[29][30]
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 ahelical structure with eight atoms per turn. The long coiled polymeric molecules make the brownish substanceelastic, and in bulk it has the feel of crude rubber. This form ismetastable 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.
Sulfur has 23 knownisotopes, four of which are stable:32S (94.99%±0.26%),33S (0.75%±0.02%),34S (4.25%±0.24%), and36S (0.01%±0.01%).[1] Other than35S, with ahalf-life of 87.37 days, theradioactive isotopes of sulfur have half-lives less than 3 hours.
The preponderance of32S is explained by its production in thealpha process (one of the main classes of nuclear fusion reactions) in exploding stars. Other stable sulfur isotopes are produced in the bypass processes related with34Ar[clarification needed], and their composition depends on the type of a stellar explosion. For example, proportionally more33S comes fromnovae than fromsupernovae.[31]
It has been found that the proportion of the two most abundant sulfur isotopes32S and34S varies in different samples by a surprising 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 oxygenfugacity, identify the activity of sulfate-reducing bacteria in the time of formation of the sample, or suggest the main sources of sulfur in ecosystems.[32] However, there are ongoing discussions over the real reason for the δ34S shifts, biological activity or postdeposit alteration.[33]
For example, whensulfide 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. Theδ13C and δ34S of coexistingcarbonate minerals and sulfides can be used to determine thepH and oxygen fugacity of the ore-bearing fluid during ore formation.
Scientists measure thesulfur isotopes ofminerals in rocks andsediments to study theredox conditions in past oceans.Sulfate-reducing bacteria in marine sediment fractionatesulfur isotopes as they take insulfate and producesulfide. Prior to the 2010s, it was thought that sulfate reduction could fractionatesulfur isotopes up to 46permil[34] and fractionation larger than 46 permil recorded in sediments must be due todisproportionation of sulfur compounds in the sediment. This view has changed since the 2010s as experiments showed thatsulfate-reducing bacteria can fractionate to 66 permil.[35] As substrates for disproportionation are limited by the product ofsulfate reduction, the isotopic effect of disproportionation should be less than 16 permil in most sedimentary settings.[36]
Inforest 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 inhydrologic studies. Differences in thenatural abundances can be used in systems where there is sufficient variation in the34S of ecosystem components.Rocky Mountain lakes thought to be dominated by atmospheric sources of sulfate have been found to have measurably different34S values than lakes believed to be dominated by watershed sources of sulfate.
The radioactive35S is formed incosmic ray spallation of the atmospheric40Ar. 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 reaction35Cl +n →35S +p is used, irradiatingpotassium chloride with neutrons.[37] The isotope35S is used in various sulfur-containing compounds as aradioactive tracer for many biological studies, for example, theHershey-Chase experiment.
Because of the weakbeta activity of35S, its compounds are relatively safe as long as they are not ingested or absorbed by the body.[38]
Sulfur vat from which railroad cars are loaded, Freeport Sulphur Co., Hoskins Mound, Texas (1943)Most of the yellow and orange hues ofIo are due to elemental sulfur and sulfur compounds deposited by activevolcanoes.Sulfur extraction, East JavaA man carrying sulfur blocks fromKawah Ijen, a volcano in East Java, Indonesia, 2009
32S is created inside massive stars, at a depth where the temperature exceeds 2.5 billion K, by thefusion of one nucleus of silicon plus one nucleus of helium.[39] As this nuclear reaction is part of thealpha process that produces elements in abundance, sulfur is the 10thmost common element in the universe.
Sulfur, usually as sulfide, is present in many types ofmeteorites.Ordinary chondrites contain on average 2.1% sulfur, andcarbonaceous chondrites may contain as much as 6.6%. It is normally present astroilite (FeS), but there are exceptions, with carbonaceous chondrites containing free sulfur, sulfates and other sulfur compounds.[40] The distinctive colors ofJupiter'svolcanic moonIo are attributed to various forms of molten, solid, and gaseous sulfur.[41] In July 2024, elemental sulfur was accidentally discovered to exist onMars after theCuriosity rover drove over and crushed a rock, revealing sulfur crystals inside it.[42]
Sulfur is the fifth most common element by mass in the Earth. Elemental sulfur can be found nearhot springs andvolcanic regions in many parts of the world, especially along thePacific 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).[43] Historically,Sicily was a major source of sulfur in theIndustrial Revolution.[44] Lakes of molten sulfur up to about 200 m (660 ft) in diameter have been found on the sea floor, associated withsubmarine volcanoes, at depths where the boiling point of water is higher than the melting point of sulfur.[45]
Native sulfur is synthesized byanaerobic bacteria acting onsulfate minerals such asgypsum insalt domes.[46][47] Significant deposits in salt domes occur along the coast of theGulf of Mexico, and inevaporites 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.[48] Such sources have become of secondary commercial importance, and most are no longer worked but commercial production is still carried out in theOsiek mine in Poland.
Common naturally occurring sulfur compounds include thesulfide minerals, such aspyrite (iron sulfide),cinnabar (mercury sulfide),galena (lead sulfide),sphalerite (zinc sulfide), andstibnite (antimony sulfide); and thesulfate minerals, such asgypsum (calcium sulfate),alunite (potassium aluminium sulfate), andbarite (barium sulfate). On Earth, just as upon Jupiter's moon Io, elemental sulfur occurs naturally in volcanic emissions, including emissions fromhydrothermal vents.
Sulfur polycations,S2+8,S2+4 andS2+19 are produced when sulfur is reacted with oxidizing agents in a strongly acidic solution.[49] The colored solutions produced by dissolving sulfur inoleum 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 andS2+19 is red.[16]
Reduction of sulfur gives variouspolysulfides with the formulaS2− x, many of which have been obtained in crystalline form. Illustrative is the production ofsodium tetrasulfide:
4 Na + S8 → 2 Na2S4
Some of these dianions dissociate to giveradical anions. For instance,S−3 gives the blue color of the rocklapis lazuli.
Two parallel sulfur chains grown inside a single-wallcarbon nanotube (CNT, a). Zig-zag (b) and straight (c) S chains inside double-wall CNTs[50]
This reaction highlights a distinctive property of sulfur: its ability tocatenate (bind to itself by formation of chains).Protonation of these polysulfide anions produces thepolysulfanes, H2Sx, wherex = 2, 3, and 4.[51] Ultimately, reduction of sulfur produces sulfide salts:
Treatment of sulfur with hydrogen giveshydrogen sulfide. When dissolved in water, hydrogen sulfide is mildly acidic:[10]
H2S ⇌ HS− + H+
Hydrogen sulfide gas and the hydrosulfide anion are extremely toxic to mammals, due to their inhibition of the oxygen-carrying capacity ofhemoglobin and certaincytochromes in a manner analogous tocyanide andazide (see below, underprecautions).
Sulfur reacts with many metals. Electropositive metals give polysulfide salts. Copper, zinc, and silver are attacked by sulfur; seetarnishing. Although manymetal sulfides are known, most are prepared by high temperature reactions of the elements.[54] Geoscientists also study the isotopes of metal sulfides in rocks and sediment to study environmental conditions in the Earth's past.[55]
Some of the main classes of sulfur-containing organic compounds include the following:[56]
Thiols or mercaptans (so called because they capture mercury aschelators) are the sulfur analogs ofalcohols; treatment of thiols with base givesthiolate ions.
Sulfoxides andsulfones 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 issulfolane.
Compounds with carbon–sulfur multiple bonds are uncommon, an exception beingcarbon disulfide, a volatile colorless liquid that is structurally similar to carbon dioxide. It is used as a reagent to make the polymerrayon and many organosulfur compounds.[57] Unlikecarbon monoxide,carbon monosulfide is stable only as an extremely dilute gas, found between solar systems.[58]
Organosulfur compounds are responsible for some of the unpleasant odors of decaying organic matter. They are widely known as theodorant in domestic natural gas, garlic odor, and skunk spray, as well as a component ofbad breath odor. Not all organic sulfur compounds smell unpleasant at all concentrations: the sulfur-containingmonoterpenoidgrapefruit mercaptan in small concentrations is the characteristic scent of grapefruit, but has a generic thiol odor at larger concentrations.Sulfur mustard, a potentvesicant, wasused in World War I as a disabling agent.[59]
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 naturalrubber, elemental sulfur is heated with the rubber to the point that chemical reactions formdisulfide bridges betweenisoprene units of the polymer. This process, patented in 1843,[citation needed] made rubber a major industrial product, especially in automobile tires. Because of the heat and sulfur, the process was namedvulcanization, after the Roman god of the forge andvolcanism.
Pharmaceutical container for sulfur from the first half of the 20th century. From theMuseo del Objeto del Objeto collection
According to theEbers Papyrus, a sulfur ointment was used in ancientEgypt to treat granular eyelids. Sulfur was used forfumigation in preclassicalGreece;[60] this is mentioned in theOdyssey.[61]Pliny the Elder discusses sulfur in hisNatural History, saying that its best-known source is the island ofMelos. He mentions its use for fumigation, medicine, and bleaching cloth.[62]
A natural form of sulfur known asshiliuhuang was known in China since the 6th century BC and found inHanzhong.[63] By the 3rd century, the Chinese had discovered that sulfur could be extracted frompyrite.[63] ChineseDaoists were interested in sulfur's flammability and its reactivity with certain metals, yet its earliest practical uses were found intraditional Chinese medicine.[63] TheWujing Zongyao of 1044 AD described formulas for Chineseblack powder, which is a mixture ofpotassium nitrate,charcoal, and sulfur.[64]
Sulfur
Brimstone
Alchemical signs for sulfur, or thecombustible elements, and brimstone, an older/archaic name for sulfur[65]
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.[66] In therasaśāstra tradition, sulfur is called "the smelly" (गन्धक,gandhaka).EarlyEuropeanalchemists gave sulfur analchemical symbol of a triangle atop a cross (🜍). The variation known as brimstone has a symbol combining atwo-barred cross atop alemniscate (🜏). In traditional skin treatment, elemental sulfur was used (mainly in creams) to alleviate such conditions asscabies,ringworm,psoriasis,eczema, andacne. The mechanism of action is unknown—though elemental sulfur does oxidize slowly to sulfurous acid, a mild reducing and antibacterial agent.[67][68][69]
Sulfur has antifungal, antibacterial, andkeratolytic activity; in the past it was used against acne vulgaris, rosacea, seborrheic dermatitis, dandruff, pityriasis versicolor, scabies, and warts.[70] This 1881 advertisement baselessly claims efficacy against rheumatism, gout, baldness, and graying of hair.
Sulfur appears in a column of fixed (non-acidic)alkali in a chemical table of 1718.[71]Antoine Lavoisier used sulfur in combustion experiments, writing of some of these in 1777.[72]
Sulfur deposits inSicily were the dominant source for more than a century. By the late 18th century, about 2,000 tonnes per year of sulfur were imported intoMarseille, France, for the production ofsulfuric acid for use in theLeblanc process. Inindustrializing Britain, with the repeal oftariffs 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 theSulfur Crisis of 1840, whenKing 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.[73][74]
In 1867, elemental sulfur was discovered in underground deposits inLouisiana andTexas. The highly successfulFrasch process was developed to extract this resource.[75]
In the late 18th century,furniture makers used molten sulfur to producedecorative inlays.[76] 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.[48]
Since the advent of thecontact process, the majority of sulfur is used to make sulfuric acid for a wide range of uses, particularly fertilizer.[77]
In recent times, the main source of sulfur has becomepetroleum andnatural gas. This is due to the requirement to remove sulfur from fuels in order to preventacid rain, and has resulted in a surplus of sulfur.[12]
Sulfur is derived from the Latin wordsulpur, which wasHellenized tosulphur 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 spellingsulfur, which appears in Latin toward the end of theClassical period. The true Ancient Greek word for sulfur,θεῖον,theîon (from earlierθέειον,théeion), is the source of the international chemical prefixthio-. The Modern Standard Greek word for sulfur is θείο,theío.
In 12th-centuryAnglo-French, it wassulfre. In the 14th century, the erroneously Hellenized Latin-ph- was restored in Middle Englishsulphre. By the 15th century, both full Latin spelling variantssulfur andsulphur became common in English. The parallelf~ph spellings continued in Britain until the 19th century, when the word was standardized assulphur.[78] On the other hand,sulfur was the form eventually chosen in the United States, though multiple place names (such asWhite Sulphur Springs) use-ph-. Canada uses both spellings.
IUPAC adopted the spellingsulfur in 1990[79][80] as did the Nomenclature Committee of theRoyal Society of Chemistry in 1992, restoring the spellingsulfur to Britain.[81]Oxford Dictionaries note that "in chemistry and other technical uses ... the-f- spelling is now the standard form for this and related words in British as well as US contexts, and is increasingly used in general contexts as well."[82]
Sicilian kiln used to obtain sulfur from volcanic rock (diagram from a 1906 chemistry book)Traditional sulfur mining atIjen 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.[83] In volcanic regions inSicily, 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 orcarusi 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 inSicilian sulfur mines were horrific, promptingBooker 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."[84] Sulfur is still mined from surface deposits in poorer nations with volcanoes, such asIndonesia, and problems with working conditions still exist.[85]
Elemental sulfur was extracted fromsalt 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 theFrasch process.[48] 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.[86][87]
Sulfur recovered from hydrocarbons inAlberta, stockpiled for shipment inNorth Vancouver, British Columbia
The resulting hydrogen sulfide from this process, and also as it occurs in natural gas, is converted into elemental sulfur by theClaus process, which entails oxidation of some hydrogen sulfide to sulfur dioxide and then thecomproportionation of the two:[86][87]
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 theAthabasca Oil Sands, stockpiles of elemental sulfur from this process exist throughoutAlberta, Canada.[88] Another way of storing sulfur is as abinder for concrete, the resulting product having some desirable properties (seesulfur concrete).[89]
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 areChina (9.6), theUnited States (8.8),Canada (7.1) andRussia (7.1).[90] Production has been slowly increasing from 1900 to 2010; the price was unstable in the 1980s and around 2010.[91]
Elemental sulfur is used mainly as a precursor to other chemicals. Approximately 85% (1989) is converted tosulfuric acid (H2SO4):
S8 + 12 O2 + 8 H2O → 8 H2SO4
Sulfuric acid production in 2000
In 2010, the United States produced more sulfuric acid than any other inorganic industrial chemical.[91] 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.[48]
When silver-basedphotography was widespread, sodium and ammoniumthiosulfate were widely used as "fixing agents". Sulfur is a component ofgunpowder ("black powder").
Amino acids synthesized byliving organisms such asmethionine andcysteine containorganosulfur groups (thioester andthiol respectively). Theantioxidantglutathione protecting many living organisms againstfree radicals andoxidative stress also contains organic sulfur. Somecrops such asonion andgarlic also produce differentorganosulfur compounds such assyn-propanethial-S-oxide responsible of lacrymal irritation (onions), ordiallyl disulfide andallicin (garlic).Sulfates, commonly found insoils andgroundwaters are often a sufficient natural source of sulfur for plants and bacteria.Atmospheric deposition ofsulfur 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 alimiting nutrient for plants andmicroorganisms (seeLiebig's barrel). However, in some circumstances, soils can be depleted insulfate, e.g. if this later is leached bymeteoric 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 offertilizers. The most important form of sulfur for fertilizer iscalcium sulfate, commonly found in nature as the mineralgypsum (CaSO4·2H2O). Elemental sulfur ishydrophobic (not soluble in water) and cannot be used directly by plants. Elemental sulfur (ES) is sometimes mixed withbentonite to amend depleted soils for crops with high requirement in organo-sulfur. Over time,oxidationabiotic processes withatmosphericoxygen andsoil bacteria canoxidize and convert elemental sulfur to soluble derivatives, which can then be used by microorganisms and plants. Sulfur improves the efficiency of other essential plant nutrients, particularlynitrogen and phosphorus.[92] Biologically produced sulfur particles are naturallyhydrophilic due to abiopolymer 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 forphosphorus. It is anessential nutrient for plant growth,root nodule formation of legumes, and immunity and defense systems. Sulfur deficiency has become widespread in many countries in Europe.[93][94][95] 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 limitacid rains.[96][92]
Elemental sulfur is one of the oldest fungicides andpesticides. "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 ofpowdery mildew diseases as well as black spot. In organic production, sulfur is the most important fungicide. It is the only fungicide used inorganically farmed apple production against the main diseaseapple 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 orfrom a dusting plane. Wettable sulfur is the commercial name for dusting sulfur formulated with additional ingredients to make it watermiscible.[89][97] It has similar applications and is used as afungicide againstmildew and other mold-related problems with plants and soil.
Elemental sulfur powder is used as an "organic" (i.e., "green")insecticide (actually anacaricide) againstticks andmites. A common method of application is dusting the clothing or limbs with sulfur powder.
Sulfur (specificallyoctasulfur, S8) is used in pharmaceutical skin preparations for the treatment ofacne and other conditions. It acts as akeratolytic agent and also kills bacteria, fungi,scabies mites, and other parasites.[98] Precipitated sulfur and colloidal sulfur are used, in form oflotions, creams, powders, soaps, and bath additives, for the treatment ofacne vulgaris,acne rosacea, andseborrhoeic dermatitis.[99]
Due to their high energy density and the availability of sulfur, there is ongoing research in creating rechargeablelithium–sulfur batteries. Until now, carbonate electrolytes have caused failures in such batteries after a single cycle. In February 2022, researchers atDrexel 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.[101]
Sulfur is an essential component of all livingcells. It is the eighth most abundant element in the human body by weight,[102] about equal in abundance topotassium, and slightly greater thansodium andchlorine.[103] A 70 kg (150 lb) human body contains about 140 grams (4.9 oz) of sulfur.[104] The main dietary source of sulfur for humans issulfur-containing amino acids,[105] which can be found in plant and animal proteins.[106]
Transferring sulfur between inorganic and biomolecules
In the 1880s, while studyingBeggiatoa (a bacterium living in a sulfur rich environment),Sergei Winogradsky found that it oxidizedhydrogen sulfide (H2S) as an energy source, forming intracellular sulfur droplets. Winogradsky referred to this form of metabolism as inorgoxidation (oxidation of inorganic compounds).[107] Another contributor, who continued to study it wasSelman Waksman.[108] Primitive bacteria that live around deep oceanvolcanic vents oxidize hydrogen sulfide for their nutrition, as discovered byRobert Ballard.[13]
Sulfur oxidizers can use as energy sources reduced sulfur compounds, including hydrogen sulfide, elemental sulfur,sulfite,thiosulfate, and variouspolythionates (e.g.,tetrathionate).[109] They depend on enzymes such assulfur oxygenase andsulfite oxidase to oxidize sulfur to sulfate. Somelithotrophs can even use the energy contained in sulfur compounds to produce sugars, a process known aschemosynthesis. Somebacteria andarchaea use hydrogen sulfide in place of water as theelectron donor in chemosynthesis, a process similar tophotosynthesis that produces sugars and uses oxygen as theelectron 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 andpurple sulfur bacteria.[110] Also sulfur-oxidizing bacteria can go into symbiosis with larger organisms, enabling the later to use hydrogen sulfide as food to be oxidized. Example: thegiant tube worm.[111]
There aresulfate-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 ofthe ancient oceans have biological origin.[112][113][114] 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 byplantsroots from soil assulfate and transported as a phosphate ester. Sulfate is reduced to sulfide via sulfite before it is incorporated intocysteine and other organosulfur compounds.[115]
While the plants' role in transferring sulfur to animals byfood chains is more or less understood, the role of sulfur bacteria is just getting investigated.[116][117]
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 ofprotein folding, which is facilitated by a variety of intra- and inter-molecular bonds. While much of the folding is driven by the formation ofhydrogen bonds,covalent bonding of cysteine residues into disulfide bridges imposes constraints that stabilize particular conformations while preventing others from forming. As thebond energy of a covalent disulfide bridge is higher than the energy of acoordinate bond or hydrophobic interaction, greater numbers of disulfide bridges lead to higher energies required for proteindenaturation. Disulfide bonds often serve to stabilize protein structures in the more oxidizing conditions of the extracellular environment.[119] Within thecytoplasm, disulfide bonds may instead be reduced (i.e. in -SH form) to their constituent cysteine residues bythioredoxins.[120]
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 arecoenzyme A andalpha-lipoic acid.[121] Cysteine-related metaboliteshomocysteine andtaurine are other sulfur-containing amino acids that are similar in structure, but not coded byDNA, and are not part of theprimary structure of proteins, take part in various locations of mammalian physiology.[122][123] Two of the 13 classical vitamins,biotin andthiamine, contain sulfur, and serve as cofactors to several enzymes.[124][125]In intracellular chemistry, sulfur operates as a carrier of reducing hydrogen and its electrons for cellular repair of oxidation. Reducedglutathione, a sulfur-containing tripeptide, is a reducing agent through its sulfhydryl (–SH) moiety derived fromcysteine.
Methanogenesis, the route to most of the world's methane, is a multistep biochemical transformation ofcarbon dioxide. This conversion requires several organosulfur cofactors. These includecoenzyme M,CH3SCH2CH2SO−3, the immediate precursor tomethane.[126]
Metalloproteins—in which the active site is a transition metal ion (or metal-sulfide cluster) often coordinated by sulfur atoms of cysteine residues[127]—are essential components of enzymes involved in electron transfer processes. Examples includeplastocyanin (Cu2+) andnitrous oxide reductase (Cu–S). The function of these enzymes is dependent on the fact that the transition metal ion can undergoredox reactions. Other examples include many zinc proteins,[128] as well asiron–sulfur clusters. Most pervasive are theferrodoxins, which serve as electron shuttles in cells. In bacteria, the importantnitrogenase enzymes contain an Fe–Mo–S cluster and is acatalyst that performs the important function ofnitrogen 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.[129]
In humansmethionine is anessential amino acid;cysteine is conditionally essential and may be synthesized from non-essentialserine 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 tosulfite oxidase, which is needed to metabolize sulfites to sulfates.[131]
Effect of acid rain on a forest, Jizera Mountains, Czech Republic
Though elemental sulfur is only minimally absorbed through the skin and is of low toxicity to humans, inhalation of sulfur dust or contact with eyes or skin may cause irritation. Excessive ingestion of sulfur can cause a burning sensation or diarrhea,[134] and cases of life-threatening metabolic acidosis have been reported after patients deliberately consumed sulfur as a folk remedy.[135][136]
When sulfur burns in air, it producessulfur dioxide. In water, this gas produces sulfurous acid and sulfites; sulfites are antioxidants that inhibit growth of aerobic bacteria and a usefulfood additive in small amounts. At high concentrations these acids harm thelungs,eyes, or othertissues.[137] In organisms without lungs such as insects, sulfite in high concentration preventsrespiration.[138]
Sulfur trioxide (made by catalysis from sulfur dioxide) andsulfuric 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.[139]
The burning ofcoal and/orpetroleum by industry andpower plants generates sulfur dioxide (SO2) that reacts with atmospheric water and oxygen to producesulfurous acid (H2SO3).[140] These acids are components ofacid rain, lowering thepH ofsoil and freshwater bodies, sometimes resulting in substantial damage to theenvironment andchemical weathering of statues and structures. Fuel standards increasingly require that fuel producers extract sulfur fromfossil 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 usesynthesis gas extract the sulfur before they burn the gas.
Hydrogen sulfide is about one-half astoxic ashydrogen cyanide, and intoxicates by the same mechanism (inhibition of the respiratory enzymecytochrome oxidase),[141] though hydrogen sulfide is less likely to cause sudden poisonings from small inhaled amounts (near itspermissible exposure limit (PEL) of 20 ppm) because of its disagreeable odor.[142] However, its presence in ambient air at concentration over 100–150 ppm quickly deadens the sense of smell,[143] and a victim may breathe increasing quantities without noticing until severe symptoms cause death. Dissolvedsulfide andhydrosulfide salts are toxic by the same mechanism.
^But impure samples have an odor similar to that ofmatches. A strong odor called "smell of sulfur" actually is given off by several sulfur compounds, such ashydrogen sulfide andorganosulfur compounds.
^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).[15]Greenwood and Earnshaw say that at fast heating for microcrystalline α-S8 the melting point is 115.1 °C (239.2 °F).[10]
^Historically, it was rather difficult to find the exact melting point of sulfur.[18] When heated slowly, the melting point may range from 114.6 °C (238.3 °F) to 120.4 °C (248.7 °F)[10] (factors that interfere with a definite melting point, are thepolymerlike nature of sulfur[19] and a large number of allotropes.[20]) Melting point may be presented as a temperature range, depending on the allotropic composition of a sample at the time of melting.
^abArblaster, John W. (2018).Selected Values of the Crystallographic Properties of Elements. Materials Park, Ohio: ASM International.ISBN978-1-62708-155-9.
^Wadsley, M. W. "Inorganic Substances in Super-Critical Fluids". Doctor of Philosophy Thesis, Department of Chemical Engineering, Monash University, Clayton, Victoria, Australia, November, 1995.
^Steudel, Ralf; Eckert, Bodo (2003).Solid Sulfur Allotropes Sulfur Allotropes. Topics in Current Chemistry. Vol. 230. pp. 1–80.doi:10.1007/b12110.ISBN978-3-540-40191-9.
^Tebbe, Fred N.; Wasserman, E.; Peet, William G.; Vatvars, Arturs; Hayman, Alan C. (1982). "Composition of Elemental Sulfur in Solution: Equilibrium ofS 6, S7, and S8 at Ambient Temperatures".Journal of the American Chemical Society.104 (18):4971–4972.Bibcode:1982JAChS.104.4971T.doi:10.1021/ja00382a050.
^Meyer, Beat (1964). "Solid Allotropes of Sulfur".Chemical Reviews.64 (4):429–451.doi:10.1021/cr60230a004.
^Pliny the Elder on science and technology, John F. Healy, Oxford University Press, 1999,ISBN0-19-814687-6, pp. 247–249.
^abcZhang, Yunming (1986). "The History of Science Society: Ancient Chinese Sulfur Manufacturing Processes".Isis.77 (3): 487.doi:10.1086/354207.S2CID144187385.
^Needham, Joseph;Yates, Robin (1994).Science and Civilisation in China, Volume 5: Chemistry and Chemical Technology, Part 6, Military Technology: Missiles and Sieges. Cambridge: Cambridge University Press. p. 120.ISBN9780521327275.OCLC489677531.
^Koch, Rudolf (1955).The book of signs : which contains all manner of symbols used from the earliest times to the Middle Ages by primitive peoples and early Christians. New York: Dover Publications.ISBN0-486-20162-7.{{cite book}}:ISBN / Date incompatibility (help)
^White, David Gordon (1996).The Alchemical Body — Siddha Traditions in Medieval India. Chicago: University of Chicago Press. pp. passim.ISBN978-0-226-89499-7.
^Lin, A. N.; Reimer, R. J.; Carter, D. M. (1988). "Sulfur revisited".Journal of the American Academy of Dermatology.18 (3):553–558.doi:10.1016/S0190-9622(88)70079-1.PMID2450900.
^Thomson, D. W. (April 1995). "Prelude to the Sulphur War of 1840: The Neapolitan Perspective".European History Quarterly.25 (2):163–180.doi:10.1177/026569149502500201.S2CID145807900.
^Mass, Jennifer L; Anderson, Mark J (2003). "Pennsylvania German sulfur-inlaid furniture: characterization, reproduction, and ageing phenomena of the inlays".Measurement Science and Technology.14 (9): 1598.doi:10.1088/0957-0233/14/9/311.ISSN0957-0233.S2CID250882259.
^Kogel, Jessica (2006).Industrial minerals & rocks: commodities, markets, and uses (7th ed.). Colorado: Littleton. p. 935.ISBN978-0-87335-233-8.OCLC62805047.
^International Union of Pure and Applied Chemistry: Inorganic Chemistry Division, Commission on Nomenclature of Inorganic Chemistry (1990).Nomenclature of Inorganic Chemistry, (Recommendations 1990). Oxford, UK: Blackwell Scientific Publications. pp. 39, 40, 41, 240, 247.
^abcSchreiner, Bernhard (2008). "Der Claus-Prozess. Reich an Jahren und bedeutender denn je".Chemie in unserer Zeit.42 (6):378–392.doi:10.1002/ciuz.200800461.
^Hyndman, A. W.; Liu, J. K.; Denney, D. W. (1982). "Sulfur Recovery from Oil Sands".Sulfur: New Sources and Uses. ACS Symposium Series. Vol. 183. pp. 69–82.doi:10.1021/bk-1982-0183.ch005.ISBN978-0-8412-0713-4.
^Zhao, F.; Hawkesford, M. J.; McGrath, S. P. (1999). "Sulphur Assimilation and Effects on Yield and Quality of Wheat".Journal of Cereal Science.30 (1):1–17.doi:10.1006/jcrs.1998.0241.
^Ceccotti, S. P. (1996). "Plant nutrient sulphur-a review of nutrient balance, environmental impact and fertilizers".Fertilizer Research.43 (1–3):117–125.doi:10.1007/BF00747690.S2CID42207099.
^Hagers Handbuch der Pharmazeutischen Praxis (in German). Vol. 6B (4th ed.). Berlin–Heidelberg–New York: Springer. 1978. pp. 672–9.ISBN978-3-540-07738-1.
^Arzneibuch-Kommentar. Wissenschaftliche Erläuterungen zum Europäischen Arzneibuch und zum Deutschen Arzneibuch [Pharmacopoeia Commentary. Scientific annotations to the European Pharmacopoeia and the German Pharmacopoeia] (in German) (23rd ed.). Stuttgart: Wissenschaftliche Verlagsgesellschaft. 2004. MonographieSchwefel zum äußerlichen Gebrauch [MonographSulfur for external use].ISBN978-3-8047-2575-1.
^Parcell, Stephen (February 2002). "Sulfur in human nutrition and applications in medicine".Alternative Medicine Review.7 (1):22–44.ISSN1089-5159.PMID11896744.
^Heldt, Hans-Walter (1996).Pflanzenbiochemie (in German). Heidelberg: Spektrum Akademischer Verlag. pp. 321–333.ISBN978-3-8274-0103-8.
^Kuenen, J. G.; Beudeker, R. F. (13 September 1982). "Microbiology of thiobacilli and other sulphur-oxidizing autotrophs, mixotrophs and heterotrophs".Philosophical Transactions of the Royal Society of London. Series B, Biological Sciences.298 (1093):473–497.Bibcode:1982RSPTB.298..473K.doi:10.1098/rstb.1982.0093.ISSN0962-8436.PMID6127737.
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.ISBN978-3-11-058889-7.
_NIST_ASD_emission_spectrum.png)
.svg)
.svg)
