Sulfur

16 phosphorus ← sulfur → chlorine O ↑ S ↓ Se periodic table
General
Name, Symbol, Number	sulfur, S, 16
Chemical series	nonmetals
Group, Period, Block	16, 3, p
Appearance	lemon yellow
Atomic mass	32.065 (5) g/mol
Electron configuration	[ Ne] 3s2 3p4
Electrons per shell	2, 8, 6
Physical properties
Phase	solid
Density (near r.t.)	(alpha) 2.08 g/cm³
Density (near r.t.)	(beta) 1.96 g/cm³
Density (near r.t.)	(gamma) 1.92 g/cm³
Liquid density at m.p.	1.819 g/cm³
Melting point	388.36 K (115.21 ° C, 239.38 ° F)
Boiling point	717.8 K (444.6 ° C, 832.3 ° F)
Heat of fusion	(mono) 1.727 kJ/mol
Heat of vaporization	(mono) 45 kJ/mol
Heat capacity	(25 °C) 22.75 J/(mol·K)
Vapor pressure P/Pa 1 10 100 1 k 10 k 100 k at T/K 375 408 449 508 591 717
Atomic properties
Crystal structure	orthorhombic
Oxidation states	−1, ±2, 4, 6 (strongly acidic oxide)
Electronegativity	2.58 ( Pauling scale)
Ionization energies ( more)	1st: 999.6 kJ/mol
	2nd: 2252 kJ/mol
	3rd: 3357 kJ/mol
Atomic radius	100 pm
Atomic radius (calc.)	88 pm
Covalent radius	102 pm
Van der Waals radius	180 pm
Miscellaneous
Magnetic ordering	no data
Electrical resistivity	(20 °C) (amorphous) 2×1015 Ω·m
Thermal conductivity	(300 K) (amorphous) 0.205 W/(m·K)
Bulk modulus	7.7 GPa
Mohs hardness	2.0
CAS registry number	7704-34-9
Notable isotopes
Main article: Isotopes of sulfur iso NA half-life DM DE ( MeV) DP 32S 95.02% S is stable with 16 neutrons 33S 0.75% S is stable with 17 neutrons 34S 4.21% S is stable with 18 neutrons 35S syn 87.32 d β- 0.167 35Cl 36S 0.02% S is stable with 20 neutrons
References

Sulfur or sulphur (see spelling below) is the chemical element in the periodic table that has the symbol S and atomic number 16. It is an abundant, tasteless, odorless, multivalent non-metal. Sulfur, in its native form, is a yellow crystaline solid. In nature, it can be found as the pure element or as sulfide and sulfate minerals. It is an essential element for life and is found in two amino acids. Its commercial uses are primarily in fertilizers but it is also widely used in gunpowder, matches, insecticides and fungicides.

Notable characteristics

At room temperature, sulfur is a soft bright yellow solid. Although sulfur is infamous for its smell - frequently compared to rotten eggs - the odor is actually characteristic of hydrogen sulfide (H₂S); elemental sulfur is odorless. It burns with a blue flame that emits 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 organic solvents such as benzene. Common oxidation states of sulfur include −2, +2, +4 and +6. Sulfur forms stable compounds with all elements except the noble gases.

Sulfur in the solid state ordinarily exists as cyclic crown-shaped S₈ molecules. Sulfur has many allotropes besides S₈. Removing one atom from the crown gives S₇, which is responsible for sulfur's distinctive yellow colour. Many other rings have been prepared, including S₁₂ and S₁₈. By contrast, its lighter neighbor oxygen only exists in two states of chemical significance: O₂ and O₃. Selenium, the heavier analogue of sulfur can form rings but is more often found as a polymer chain.

The crystallography of sulfur is complex. Depending on the specific conditions, the sulfur allotropes form several distinct crystal structures, with rhombic and monoclinic S₈ best known.

A noteworthy property is that the viscosity of molten sulfur, unlike most other liquids, increases with temperature due to the formation of polymer chains. However, after a certain temperature is reached, the viscosity is reduced because there is enough energy to break the chains.

Amorphous or "plastic" sulfur can be produced through the rapid cooling of molten sulfur. X-ray crystallography studies show that the amorphous form may have a helical structure with eight atoms per turn. This form is metastable at room temperature and gradually reverts back to crystalline form. This process happens within a matter of hours to days but can be rapidly catalyzed.

Applications

Sulfur has many industrial uses. Through its major derivative, sulfuric acid (H₂SO₄), sulfur ranks as one of the more important elements used as an industrial raw material. It is of prime importance to every sector of the world's economies.

Sulfuric acid production is the major end use for sulfur, and consumption of sulfuric acid has been regarded as one of the best indices of a nation's industrial development. More sulfuric acid is produced in the United States every year than any other industrial chemical.

Sulfur is also used in batteries, detergents, the vulcanization of rubber, fungicides, and in the manufacture of phosphate fertilizers. Sulfites are used to bleach paper and as a preservative in wine and dried fruit. Because of its flammable nature, sulfur also finds use in matches, gunpowder, and fireworks. Sodium or ammonium thiosulfate are used as photographic fixing agents. Magnesium sulfate, better known as Epsom salts, can be used as a laxative, a bath additive, an exfoliant, or a magnesium supplement for plants.

In the late 1700s, furniture makers used molten sulfur to produce decorative inlays in their craft. Because of the sulfur dioxide produced during the process of melting sulfur, the craft of sulfur inlays was soon abandoned.

Biological role

The amino acids cysteine and methionine contain sulphur, as do all polypeptides, proteins, and enzymes which contain these amino acids. This makes sulphur a necessary component of all living cells. Disulfide bonds between polypeptides are very important in protein assembly and structure. Homocysteine and taurine are also sulphur containing amino acids but are not coded for by DNA nor are they part of the primary structure of proteins. Some forms of bacteria use hydrogen sulfide (H₂S) in the place of water as the electron donor in a primitive photosynthesis-like process. Sulfur is absorbed by plants from soil as the sulphate ion. Inorganic sulphur forms a part of iron-sulphur clusters, and sulphur is the bridging ligand in the Cu_A site of cytochrome c oxidase. Sulphur is an important component of coenzyme A

Environmental Impact

The burning of coal and petroleum by industry and power plants liberates huge amounts of sulfur dioxide (SO₂) which reacts with atmospheric water and oxygen to produce sulfuric acid. This sulfuric acid is a component of acid rain, which lowers the pH of soil and freshwater bodies, resulting in substantial damage to the natural environment and chemical weathering of statues and architecture. Fuel standards increasingly require sulfur to be extracted from fossil fuels to prevent the formation of acid rain. This extracted sulfur is then refined and represents a large portion of sulfur production.

History

Sulfur ( Sanskrit, sulvere; Latin sulpur) was known in ancient times, and is referred to in the Biblical Pentateuch ( Genesis). The word itself is almost certainly from the Arabic sufra meaning yellow, from the bright colour of the naturally-occurring form.

English translations of the Bible commonly refer to sulfur as "brimstone", giving rise to the name of 'Fire and brimstone' sermons, which sinners are reminded of their fate of eternal damnation. It is from this part of the Bible that hell is implied to "smell of sulfur", although as mentioned above sulfur is in fact odorless. The "smell of sulfur" usually refers to the odor of hydrogen sulfide, e.g. from rotten eggs. Burning sulfur produces sulfur dioxide, the smell associated with burnt matches.

Homer mentioned "pest-averting sulfur" in the 9th century BC and in 424 BC, the tribe of Boeotia destroyed the walls of a city by burning a mixture of coal, sulfur, and tar under them. Sometime in the 12th century, the Chinese invented gun powder which is a mixture of potassium nitrate (KNO₃), carbon, and sulfur. Early alchemists gave sulfur its own alchemical symbol which was a triangle at the top of a cross. In the late 1770s, Antoine Lavoisier helped convince the scientific community that sulfur was an element and not a compound. In 1867, sulfur was discovered in underground deposits in Louisiana and Texas. The overlying layer of earth was quicksand, prohibiting ordinary mining operations. Therefore the Frasch process was utilized.

Occurrence

Elemental sulfur can be found near hot springs and volcanic regions in many parts of the world, especially along the Pacific Ring of Fire. These occurrences are the basis for the traditional name brimstone, since sulfur could be found near the brims of volcanic craters. Such volcanic deposits are currently exploited in Indonesia, Chile, and Japan.

Significant desposits of elemental sulfur also exist in salt domes along the coast of the Gulf of Mexico, and in evaporites in eastern Europe and western Asia. The sulfur in these deposits is believed to come from the action of anaerobic bacteria on sulfate minerals, especially gypsum. Such deposits are the basis for commercial production in the United States, Poland, Russia, Turkmenistan, and Ukraine.

Sulfur extracted from oil, gas and the Athabasca Oil Sands has become a glut on the market, with huge stockpiles of sulfur in existence throughout Alberta.

Common naturally-occurring sulfur compounds include the metal sulfides, such as pyrite (iron sulfide), cinnabar (mercury sulfide), galena ( lead sulfide), sphalerite (zinc sulfide) and stibnite (antimony sulfide); and the metal sulfates, such as gypsum (calcium sulfate), alunite (potassium aluminium sulfate), and barite (barium sulfate). Hydrogen sulfide is the gas responsible for the odor of rotten eggs. It occurs naturally in volcanic emissions, such as from hydrothermal vents, and from bacterial action on decaying sulfur-containing organic matter.

The distinctive colors of Jupiter's volcanic moon, Io, are from various forms of molten, solid and gaseous sulfur. There is also a dark area near the Lunar crater Aristarchus that may be a sulfur deposit. Sulfur is also present in many types of meteorites.

see also Category:Sulfide minerals, Category:Sulfate minerals

Compounds

Hydrogen sulfide has the characteristic smell of rotten eggs. Dissolved in water, hydrogen sulfide is acidic and will react with metals to form a series of metal sulfides. Natural metal sulfides are common, especially those of iron. Iron sulfide is called pyrite, the so called fool's gold. Interestingly, pyrite can show semiconductor properties. [1] Galena, a naturally occurring lead sulfide, was the first semiconductor discovered, and found a use as a signal rectifier in the "cat's whiskers" of early crystal radios.

Many of the unpleasant odors of organic matter are based on sulfur-containing compounds such as methyl and ethyl mercaptan used to scent natural gas so that leaks are easily detectable. The odor of garlic and "skunk stink" are also caused by sulfur-containing organic compounds. However, not all organic sulfur compounds smell unpleasant; for example, margin, a sulfur containing terpene is responsible for the characteristic scent of grapefruit.

Polymeric sulfur nitride has metallic properties even though it does not contain any metal atoms. This compound also has unusual electrical and optical properties. This polymer can be made from tetrasulfur tetranitride S₄N₄.

Phosphorus sulfides are important in synthesis. For example, P₄S₁₀ and its derivatives Lawesson's reagent and naphthalen-1,8-diyl 1,3,2,4-dithiadiphosphetane 2,4-disulfide are used to replace oxygen from some organic molecules with sulfur.

Other important compounds of sulfur include:

Inorganic sulfur compounds:

• Sulfides (S^2-), a complex family of compounds usually derived from S^2-. Cadmium sulfide (CdS) is an example.
• Sulfites (SO₃^2-), the salts of sulfurous acid (H₂SO₃) which is generated by dissolving SO₂ in water. Sulfurous acid and the corresponding sulfites are fairly strong reducing agents. Other compounds derived from SO₂ include the pyrosulfite or metabisulfite ion (S₂O₅²⁻).
• Sulfates (SO₄^2-), the salts of sulfuric acid. Sulfuric acid also reacts with SO₃ in equimolar ratios to form pyrosulfuric acid (H₂S₂O₇).
• Thiosulfates (sometimes refered as thiosulfites or "hyposulfites") (S₂O₃²⁻). Thiosulfates are used in photographic fixing (HYPO) as reducing agents. Ammonium thiosulfate is being investigated as a cyanide replacement in leaching gold. [2]
• Sodium dithionite, Na₂S₂O₄, is the highly reducing dianion derived from hyposulfurous/dithionous acid.
• Sodium dithionate (Na₂S₂O₆).
• Polythionic acids (H₂S_nO₆), where n can range from 3 to 80.
• Peroxymonosulfuric acid (H₂SO₅) and peroxydisulfuric acids (H₂S₂O₈), made from the action of SO₃ on concentrated H₂O₂, and H₂SO₄ on concentrated H₂O₂ respectively.
• Sodium polysulfides (Na₂S_x)
• Sulfur hexafluoride, SF₆, a dense gas at ambient conditions, is used as nonreactive and nontoxic propellant
• Sulfur nitrides are chain and cyclic compounds containing only S and N. Tetrasulfur tetranitride S₄N₄ is an example.
• Thiocyanates contain the SCN^- group. Oxidation of thiocyanoate gives thiocyanogen, (SCN)₂ with the connectivity NCS-SCN.

Organic sulfur compounds (where R, R', and R are organic groups such as CH₃):

• Thioethers have the form R-S-R′. These compounds are the sulfur equivalents of ethers.
• Sulfonium ions have the formula RR'S-'R'", i.e. where three groups are attached to the cationic sulfur center. Dimethylsulfoniopropionate ( DMSP; (CH₃ )₂S⁺CH₂CH₂COO^-) is a sulfonium ion, which is important in the marine organic sulfur cycle.
• Thiols (also known as mercaptans) have the form R-SH. These are the sulfur equivalents of alcohols.
• Thiolates ions s have the form R-S^-. Such anions arise upon treatment of thiols with base.
• Sulfoxides have the form R-S(=O)-R&prime. A common sulfoxide is DMSO.
• Sulfones have the form R-S(=O)₂-R&prime. A common sulfone is sulfolane C₄H₈SO₂.

Isotopes

Sulfur has 18 isotopes, of which four are stable: ³²S (95.02%), ³³S (0.75%), ³⁴S (4.21%), and ³⁶S (0.02%). Other than ³⁵S, the radioactive isotopes of sulfur are all short lived. ³⁵S is formed from cosmic ray spallation of ⁴⁰Ar in the atmosphere. It has a half-life of 87 days.

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 co-existing 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 also 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 also 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 δS-34 values from lakes believed to be dominated by watershed sources of sulfate.

Precautions

Carbon disulfide, carbon oxysulfide, hydrogen sulfide, and sulfur dioxide should all be handled with care.

Although sulfur dioxide is sufficiently safe to be used as a food additive in small amounts, at high concentrations it reacts with moisture to form sulfurous acid which in sufficient quantities may harm the lungs, eyes or other tissues. In creatures without lungs such as insects or plants, it otherwise prevents respiration.

Hydrogen sulfide is quite toxic (more toxic than cyanide). Although very pungent at first, it quickly deadens the sense of smell, so potential victims may be unaware of its presence until it is too late.

Spelling

The element has traditionally been spelled sulphur in the United Kingdom, Ireland, and India, but sulfur in the United States and Canada, while both spellings are used in Australia and New Zealand. The IUPAC adopted the spelling "sulfur" in 1990, as did the Royal Society of Chemistry Nomenclature Committee in 1992. This spelling has begun to replace its variant in educated circles, unlike aluminum, which did not stick outside the US and Canada.