Sulfur

Sulfur or sulphur is a chemical element; it has symbol S and atomic number 16. It is abundant, multivalent and nonmetallic. Under normal conditions, sulfur atoms form cyclic octatomic molecules with the chemical formula S₈. Elemental sulfur is a bright yellow, crystalline solid at room temperature.
Sulfur is the tenth most abundant element by mass in the universe and the fifth most common on Earth. Though sometimes found in pure, native form, sulfur on Earth usually occurs as sulfide and sulfate minerals. Being abundant in native form, sulfur was known in ancient times, being mentioned for its uses in ancient India, ancient Greece, China, and ancient Egypt. Historically and in literature sulfur is also called brimstone, which means "burning stone". Almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants from natural gas and petroleum. The greatest commercial use of the element is the production of sulfuric acid for sulfate and phosphate fertilizers, and other chemical processes. Sulfur is used in matches, insecticides, and fungicides. Many sulfur compounds are odoriferous, and the smells of odorized natural gas, skunk scent, bad breath, grapefruit, and garlic are due to organosulfur compounds. Hydrogen sulfide gives the characteristic odor to rotting eggs and other biological processes.
Sulfur is an essential element for all life, almost always in the form of organosulfur compounds or metal sulfides. Amino acids and two vitamins 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 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.

Characteristics

Physical properties

Sulfur forms several polyatomic molecules. The best-known allotrope is octasulfur, cyclo-S₈. The point group of cyclo-S₈ is D_4d and its dipole moment is 0 D. Octasulfur is a soft, bright-yellow solid that is odorless. It melts at, and boils at. At, below its melting temperature, cyclo-octasulfur begins slowly changing from α-octasulfur to the β-polymorph. The structure of the S₈ ring is virtually unchanged by this phase transition, which affects the intermolecular interactions. Cooling molten sulfur freezes at, as it predominantly consists of the β-S₈ molecules. Between its melting and boiling temperatures, octasulfur changes its allotrope again, turning from β-octasulfur to γ-sulfur, again accompanied by a lower density but increased viscosity due to the formation of polymers. At higher temperatures, the viscosity decreases as depolymerization occurs. Molten sulfur assumes a dark red color above. 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 and, and occurs readily in boiling water at.
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.

Chemical properties

Under normal conditions, sulfur hydrolyzes very slowly to mainly form hydrogen sulfide and sulfuric acid:
The reaction involves adsorption of protons onto clusters, followed by disproportionation into the reaction products.
The second, fourth and sixth ionization energies of sulfur are 2252 kJ/mol, 4556 kJ/mol and 8495.8 kJ/mol, respectively. The composition of reaction products of sulfur with oxidants 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, formation of sulfur trioxide requires a temperature of and presence of a catalyst.
In reactions with elements of lesser electronegativity, it reacts as an oxidant and forms sulfides, where it has oxidation state −2.
Sulfur reacts with nearly all other elements except noble gases, even with the notoriously unreactive metal iridium. Some of those reactions require elevated temperatures.

Allotropes

Sulfur forms over 30 solid allotropes, more than any other element. Besides S₈, several other rings are known. Removing one atom from the crown gives S₇, which is of a deeper yellow than S₈. HPLC analysis of "elemental sulfur" reveals an equilibrium mixture of mainly S₈, but with S₇ and small amounts of S₆. Larger rings have been prepared, including S₁₂ and S₁₈.
Amorphous or "plastic" sulfur is produced by rapid cooling of molten sulfur—for example, by pouring it into cold water. X-ray crystallography studies show that the amorphous form may have a helical structure with eight atoms per turn. The long coiled polymeric molecules make the brownish substance elastic, and in bulk it has the feel of crude rubber. This form is metastable at room temperature and gradually reverts to the crystalline molecular allotrope, which is no longer elastic. This process happens over a matter of hours to days, but can be rapidly catalyzed.

Isotopes

Sulfur has 23 known isotopes, four of which are stable: ³²S, ³³S, ³⁴S, and ³⁶S. 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 in exploding stars. Other stable sulfur isotopes are produced in the bypass processes related with ³⁴Ar, and their composition depends on the type of a stellar explosion. For example, proportionally more ³³S comes from novae than from supernovae.
It has been found that the proportion of the two most abundant sulfur isotopes ³²S and ³⁴S varies in different samples by a surprising large amount. Determination of the isotope ratio 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. However, there are ongoing discussions over the real reason for the δ³⁴S shifts, biological activity or postdeposit alteration.
For example, when sulfide minerals are precipitated, isotopic equilibration 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 and fractionation larger than 46 permil recorded in sediments must be due to disproportionation of sulfur compounds in the sediment. This view has changed since the 2010s as experiments showed that sulfate-reducing bacteria can fractionate to 66 permil. As substrates for disproportionation are limited by the product of sulfate reduction, the isotopic effect of disproportionation should be less than 16 permil in most sedimentary settings.
In forest ecosystems, sulfate is derived mostly from the atmosphere; weathering of ore minerals and evaporites contribute some sulfur. Sulfur with a distinctive isotopic composition has been used to identify pollution sources, and enriched sulfur has been added as a tracer in hydrologic studies. Differences in the natural abundances can be used in systems where there is sufficient variation in the ³⁴S of ecosystem components. Rocky Mountain lakes thought to be dominated by atmospheric sources of sulfate have been found to have measurably different ³⁴S values than lakes believed to be dominated by watershed sources of sulfate.
The radioactive ³⁵S is formed in cosmic ray spallation of the atmospheric ⁴⁰Ar. This fact may be used to verify the presence of recent 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. The isotope ³⁵S is used in various sulfur-containing compounds as a radioactive tracer for many biological studies, for example, the Hershey-Chase experiment.
Because of the weak beta activity of ³⁵S, its compounds are relatively safe as long as they are not ingested or absorbed by the body.

Natural occurrence

³²S is created inside massive stars, at a depth where the temperature exceeds 2.5 billion K, by the fusion of one nucleus of silicon plus one nucleus of helium. As this nuclear reaction is part of the alpha process that produces elements in abundance, sulfur is the 10th most common element in the universe.
Sulfur, usually as sulfide, is present in many types of meteorites. Ordinary chondrites contain on average 2.1% sulfur, and carbonaceous chondrites may contain as much as 6.6%. It is normally present as troilite, but there are exceptions, with carbonaceous chondrites containing free sulfur, sulfates and other sulfur compounds. The distinctive colors of Jupiter's volcanic moon Io are attributed to various forms of molten, solid, and gaseous sulfur. In July 2024, elemental sulfur was accidentally discovered to exist on Mars after the Curiosity rover drove over and crushed a rock, revealing sulfur crystals inside it.
Sulfur is the fifth most common element by mass in the Earth. Elemental sulfur can be found near hot springs and volcanic regions in many parts of the world, especially along the Pacific Ring of Fire; such volcanic deposits are mined in Indonesia, Chile, and Japan. These deposits are polycrystalline, with the largest documented single crystal measuring. Historically, Sicily was a major source of sulfur in the Industrial Revolution. Lakes of molten sulfur up to about in diameter have been found on the sea floor, associated with submarine volcanoes, at depths where the boiling point of water is higher than the melting point of sulfur.
Native sulfur is synthesized by anaerobic bacteria acting on sulfate minerals such as gypsum in salt domes. Significant deposits in salt domes occur along the coast of the Gulf of Mexico, and in evaporites in eastern Europe and western Asia. Native sulfur may be produced by geological processes alone. Fossil-based sulfur deposits from salt domes were once the basis for commercial production in the United States, Russia, Turkmenistan, and Ukraine. Such sources have become of secondary commercial importance, and most are no longer worked but commercial production is still carried out in the Osiek mine in Poland.
Common naturally occurring sulfur compounds include the sulfide minerals, such as pyrite, cinnabar, galena, sphalerite, and stibnite ; and the sulfate minerals, such as gypsum, alunite, and barite. On Earth, just as upon Jupiter's moon Io, elemental sulfur occurs naturally in volcanic emissions, including emissions from hydrothermal vents.
The main industrial source of sulfur has become petroleum and natural gas.