iso	NA	half-life	DM	DE (MeV)	DP
³²S	95.02%	³²S is stable with 16 neutrons
³³S	0.75%	³³S is stable with 17 neutrons
³⁴S	4.21%	³⁴S is stable with 18 neutrons
³⁵S	syn	87.32 d	β⁻	0.167	³⁵Cl
³⁶S	0.02%	³⁶S is stable with 20 neutrons

Sulfur (US English) or sulphur (

/ˈsʌlfər/ sul-fər; see spelling below) is the chemical element with atomic number 16. In the periodic table it is represented by the symbol S. It is an abundant, multivalent non-metal. Under normal conditions, sulfur atoms form cyclic octatomic molecules with chemical formula S₈. Elemental sulfur is a bright yellow crystalline solid when at room temperature. Chemically, sulfur can react as either an oxidant or reducing agent. It oxidizes most metals and several nonmetals, including carbon, which leads to its negative charge in most organosulfur compounds, but it reduces several strong oxidants, such as oxygen and fluorine. It is also the lightest element to easily produce stable exceptions to the octet rule.
In nature, sulfur can be found as the pure element and as sulfide and sulfate minerals. Elemental sulfur crystals are commonly sought after by mineral collectors for their brightly colored polyhedron shapes. Being abundant in native form, sulfur was known in ancient times, mentioned for its uses in ancient Greece, China and Egypt. Sulfur fumes were used as fumigants, and sulfur-containing medicinal mixtures were used as balms and antiparasitics. Sulfur is referenced in the Bible as brimstone in English, with this name still used in several nonscientific terms.^[2] Sulfur was considered important enough to receive its own alchemical symbol. It was needed to make the best quality of black gunpowder, and the bright yellow powder was hypothesized by alchemists to contain some of the properties of gold, which they sought to synthesize from it. In 1777, Antoine Lavoisier helped convince the scientific community that sulfur was a basic element, rather than a compound.
Elemental sulfur was once extracted from salt domes where it sometimes occurs in nearly pure form, but this method has been obsolete since the late 20th century. Today, almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants from natural gas and petroleum. The element's commercial uses are primarily in fertilizers, because of the relatively high requirement of plants for it, and in the manufacture of sulfuric acid, a primary industrial chemical. Other well-known uses for the element are in matches, insecticides and fungicides. Many sulfur compounds are odiferous, and the smell of odorized natural gas, skunk scent, grapefruit, and garlic is due to sulfur compounds. Hydrogen sulfide produced by living organisms imparts the characteristic odor to rotting eggs and other biological processes.
Sulfur is an essential element for all life, and is widely used in biochemical processes. In metabolic reactions, sulfur compounds serve as both fuels and respiratory (oxygen-replacing) materials for simple organisms. Sulfur in organic form is present in the vitamins biotin and thiamine, the latter being named for the Greek word for sulfur. Sulfur is an important part of many enzymes and in antioxidant molecules like glutathione and thioredoxin. Organically bonded sulfur is a component of all proteins, as the amino acids cysteine and methionine. Disulfide bonds are largely responsible for the mechanical strength and insolubility of the protein keratin, found in outer skin, hair, and feathers, and the element contributes to their pungent odor when burned.

Characteristics

When burned, sulfur melts to a blood-red liquid and emits a blue flame which is best observed in the dark.

Physical

Sulfur forms polyatomic molecules with different chemical formulas, with the best-known allotrope being octasulfur, cyclo-S₈. Octasulfur is a soft, bright-yellow solid with only a faint odor, similar to that of matches.^[3] It melts at 115.21 °C, boils at 444.6 °C and sublimes easily.^[2] At 95.2 °C, below its melting temperature, cyclo-octasulfur changes from α-octasulfur to the β-polymorph.^[4] The structure of the S₈ ring is virtually unchanged by this phase change, which affects the intermolecular interactions. Between its melting and boiling temperatures, octasulfur changes its allotrope again, turning from β-octasulfur to γ-sulfur, again accompanied by a lower density but increased viscosity due to the formation of polymers.^[4] At even higher temperatures, however, the viscosity decreases as depolymerization occurs. Molten sulfur assumes a dark red color above 200 °C. The density of sulfur is about 2 g·cm⁻³, depending on the allotrope; all of its stable allotropes are excellent electrical insulators.

Chemical

Sulfur burns with a blue flame concomitant with formation of sulfur dioxide, notable for its peculiar suffocating odor. 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. The first and the second ionization energies of sulfur are 999.6 and 2252 kJ·mol⁻¹, respectively. Despite such figures, S²⁺ is rare, S^{4, 6+} being more common. The fourth and sixth ionization energies are 4556 and 8495.8 kJ·mol⁻¹, the magnitude of the figures caused by electron transfer between orbitals; these states are only stable with strong oxidants as fluorine, oxygen, and chlorine.

Allotropes

The structure of the cyclooctasulfur molecule, S₈.

Main article: Allotropes of sulfur

Sulfur forms more than 30 solid allotropes, more than any other element.^[5] Besides S₈, several other rings are known.^[6] Removing one atom from the crown gives S₇, which is more deeply yellow than S₈. HPLC analysis of "elemental sulfur" reveals an equilibrium mixture of mainly S₈, but with S₇ and small amounts of S₆.^[7] Larger rings have been prepared, including S₁₂ and S₁₈.^[8]^[9]
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 cause the brownish substance to be elastic, and in bulk this form has the feel of crude rubber. This form is metastable at room temperature and gradually reverts to crystalline molecular allotrope, which is no longer elastic. This process happens within a matter of hours to days, but can be rapidly catalyzed.

Isotopes

Main article: Isotopes of sulfur

Sulfur has 25 known isotopes, four of which are stable: ³²S (95.02%), ³³S (0.75%), ³⁴S (4.21%), and ³⁶S (0.02%). Other than ³⁵S, with a half-life of 87 days and formed in cosmic ray spallation of ⁴⁰Ar, the radioactive isotopes of sulfur have half-lives less than 170 minutes.
When sulfide minerals are precipitated, isotopic equilibration among solids and liquid may cause small differences in the δS-34 values of co-genetic minerals. The differences between minerals can be used to estimate the temperature of equilibration. The δC-13 and δS-34 of coexisting carbonates and sulfides can be used to determine the pH and oxygen fugacity of the ore-bearing fluid during ore formation.
In most forest ecosystems, sulfate is derived mostly from the atmosphere; weathering of ore minerals and evaporites contribute some sulfur. Sulfur with a distinctive isotopic composition has been used to identify pollution sources, and enriched sulfur has been added as a tracer in hydrologic studies. Differences in the natural abundances can be used in systems where there is sufficient variation in the ³⁴S of ecosystem components. Rocky Mountain lakes thought to be dominated by atmospheric sources of sulfate have been found to have different δ³⁴S values from lakes believed to be dominated by watershed sources of sulfate.

Natural occurrence

Most of the yellow and orange hues of Io are due to elemental sulfur and sulfur compounds, produced by active volcanoes.

Native sulfur crystals

Native sulfur crystals, Iran

Sulfur crystals, khanegiran, Iran

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

³²S is created inside massive stars, at a depth where the temperature exceeds 2.5×10⁹ K, by the fusion of one nucleus of silicon plus one nucleus of helium.^[10] As this is part of the alpha process that produces elements in abundance, sulfur is the 10th most common element in the universe.
Sulfur, usually as sulfide, is present in many types of meteorites. Ordinary chondrites contain on average 2.1% sulfur, and carbonaceous chondrites may contain as much as 6.6%. It is normally present as troilite (FeS), but there are exceptions, with carbonaceous chondrites containing free sulfur, sulfates and other sulfur compounds.^[11] The distinctive colors of Jupiter's volcanic moon Io are attributed to various forms of molten, solid and gaseous sulfur.^[12]
On Earth, elemental sulfur can be found near hot springs and volcanic regions in many parts of the world, especially along the Pacific Ring of Fire; such volcanic deposits are currently mined in Indonesia, Chile, and Japan. Such deposits are polycrystalline, with the largest documented single crystal measuring 22×16×11 cm.^[13] Historically, Sicily was a large source of sulfur in the Industrial Revolution.^[14]
Significant deposits of elemental sulfur, believed to have been (and are still being) synthesised by anaerobic bacteria on sulfate minerals like gypsum, exist in salt domes along the coast of the Gulf of Mexico, and in evaporites in eastern Europe and western Asia. Native sulfur may be produced by geological processes alone. Fossil-based sulfur deposits from salt domes have until recently been the basis for commercial production in the United States, Poland, Russia, Turkmenistan, and Ukraine.^[15] Such sources are now of secondary commercial importance, and most are no longer worked.
Common naturally-occurring sulfur compounds include the sulfide minerals, such as pyrite (iron sulfide), cinnabar (mercury sulfide), galena (lead sulfide), sphalerite (zinc sulfide) and stibnite (antimony sulfide); and the sulfates, such as gypsum (calcium sulfate), alunite (potassium aluminium sulfate), and barite (barium sulfate). On Earth, just as upon Jupiter's moon Io, elemental sulfur occurs naturally in volcanic emissions, including emissions from hydrothermal vents.

Production

Sulfur may be found by itself and historically was usually obtained in this way, while pyrite has been a source of sulfur via sulfuric acid.^[16] The Sicilian process was used in ancient times to obtain sulfur from rocks present in volcanic regions of Sicily: sulfur deposits were piled and stacked in brick kilns built on sloping hillsides, with airspaces between them. Then, powdered sulfur was put on top of the deposit and ignited, causing the deposits to melt down the hills. Today's sulfur production is as a side product of other industrial processes such as oil refining; in these processes, sulfur often occurs as undesired or detrimental compounds that are extracted and converted to elemental sulfur. As a mineral, native sulfur under salt domes is thought to be a fossil mineral resource, produced by the action of ancient bacteria on sulfate deposits. It was removed from such salt-dome mines mainly by the Frasch process.^[15] 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 which required no further purification. However, due to a limited number of such sulfur deposits and the high cost of working them, this process for mining sulfur has not been employed in a major way anywhere in the world since 2002.^[17]^[18]

Sulfur recovered from hydrocarbons in Alberta, stockpiled for shipment in North Vancouver, B.C.

Today, sulfur is produced from petroleum, natural gas, and related fossil resources, from which it is obtained mainly as hydrogen sulfide. Organosulfur compounds, undesirable impurities in petroleum, may be upgraded by subjecting them to hydrodesulfurization, which cleaves the C–S bonds:^[17]^[18]

R-S-R + 2 H₂ → 2 RH + H₂S

The resulting hydrogen sulfide from this process, and also as it occurs in natural gas, is converted into elemental sulfur by the Claus process. This process entails oxidation of some hydrogen sulfide to sulfur dioxide and then the comproportionation of the two:^[17]^[18]

3 O₂ + 2 H₂S → 2 SO₂ + 2 H₂O

SO₂ + 2 H₂S → 3 S + 2 H₂O

Owing to the high sulfur content of the Athabasca Oil Sands, stockpiles of elemental sulfur from this process now exist throughout Alberta, Canada.^[19] Another way of storing sulfur is as a binder for concrete, the resulting product having many desirable properties.^[20] The price of sulfur increased from 2007 to 2008, and decreased thereafter.^[21]

Compounds

Sulfides

Treatment of sulfur with hydrogen gives hydrogen sulfide. When dissolved in water, hydrogen sulfide is mildly acidic:^[2]

H₂S $\overrightarrow{\leftarrow}$ HS^– + H⁺

Hydrogen sulfide gas and the dissolved sulfide and hydrosulfide anions are extremely toxic to mammals, due to their inhibition of the oxygen-carrying capacity of hemoglobin and certain cytochromes in a manner analogous to cyanide and azide (see below, under precautions).
Reduction of elemental sulfur gives polysulfides, which consist of chains of sulfur atoms terminated with S^– centers:

2 Na + S₈ → Na₂S₈

This reaction highlights arguably the single most distinctive property of sulfur: its ability to catenate (bind to itself by formation of chains). Protonation of these polysulfide anions gives the polysulfanes, H₂S_x where x = 2, 3, and 4.^[22] Ultimately reduction of sulfur gives sulfide salts:

16 Na + S₈ → 8 Na₂S

The interconversion of these species is exploited in the sodium-sulfur battery. The radical anion S₃^– gives the blue color to the mineral lapis lazuli.

Lapis lazuli owes its blue color to a sulfur radical.

Elemental sulfur can be oxidized, for example, to give bicyclic S₈²⁺.

Oxides and oxyanions

The principal sulfur oxides are obtained by burning sulfur:

S + O₂ → SO₂

2 SO₂ + O₂ → 2 SO₃

Other oxides are known, e.g. sulfur monoxide and disulfur mono- and dioxides, but they are unstable.
The sulfur oxides form numerous oxyanions with the formula SO_n^2–. Sulfur dioxide and sulfites (SO2−
3) are related to the unstable sulfurous acid (H₂SO₃). Sulfur trioxide and sulfates (SO2−
4) are related to sulfuric acid. Sulfuric acid and SO₃ combine to give oleum, a solution of pyrosulfuric acid (H₂S₂O₇) in sulfuric acid.

Peroxydisulfuric acid

Peroxides convert sulfur into unstable such as S₈O, a sulfoxide. Peroxymonosulfuric acid (H₂SO₅) and peroxydisulfuric acids (H₂S₂O₈), made from the action of SO₃ on concentrated H₂O₂, and H₂SO₄ on concentrated H₂O₂ respectively.

The sulfate anion, SO2−
4

Thiosulfate salts (S₂O2−
3), sometimes referred as "hyposulfites", used in photographic fixing (HYPO) and as reducing agents, feature sulfur in two oxidation states. Sodium dithionite, (S₂O2−
4), contains the more highly reducing dithionite anion. Sodium dithionate (Na₂S₂O₆) is the first member of the polythionic acids (H₂S_nO₆), where n can range from 3 to many.

Halides and oxyhalides

The two main sulfur fluorides are sulfur hexafluoride, a dense gas used as nonreactive and nontoxic propellant, and sulfur tetrafluoride, a rarely used organic reagent that is highly toxic.^[23] Their chlorinated analogs are sulfur dichloride and sulfur monochloride. Sulfuryl chloride and chlorosulfuric acid are derivatives of sulfuric acid; thionyl chloride (SOCl₂) is a common reagent in organic synthesis.^[24]

Pnictides

The most important S–N compound is the cage tetrasulfur tetranitride (S₄N₄). Heating this compound gives polymeric sulfur nitride ((SN)_x), which has metallic properties even though it does not contain any metal atoms. Thiocyanates contain the SCN⁻ group. Oxidation of thiocyanate gives thiocyanogen, (SCN)₂ with the connectivity NCS-SCN. Phosphorus sulfides are numerous, the most important commercially being the cages P₄S₁₀ and P₄S₃.^[25]^[26]

Metal sulfides

Main article: Sulfide mineral

Many if not most minerals occur as sulfides. The principal ores of copper, zinc, nickel, cobalt, molybdenum and others are sulfides. These materials tend to be dark-colored semiconductors that are not readily attacked by water or even many acids. They are formed, both geochemically and in the laboratory, by the reaction of hydrogen sulfide with metal salts to form the metal sulfides. The mineral galena (PbS) was the first demonstrated semiconductor and found a use as a signal rectifier in the cat's whiskers of early crystal radios. The iron sulfide called pyrite, the so-called "fool's gold," has the formula FeS₂.^[27] The upgrading of these ores, usually by roasting, is costly and environmentally hazardous. Sulfur corrodes many metals via the process called tarnishing.