Sulfur

This article is about the chemical element. For other meanings of "sulfur" or "sulphur", see sulphur (disambiguation)

Sulfur, ₁₆S

Sulfur
Alternative name	Sulphur (British spelling)
Allotropes	see Allotropes of sulfur
Appearance	Lemon yellow sintered microcrystals
Standard atomic weight Ar°(S)
	[32.059, 32.076] 32.06±0.02 (abridged)
Sulfur in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson O ↑ S ↓ Se phosphorus ← sulfur → chlorine
Atomic number (Z)	16
Group	group 16 (chalcogens)
Period	period 3
Block	p-block
Electron configuration	[Ne] 3s2 3p4
Electrons per shell	2, 8, 6
Physical properties
Phase at STP	solid
Melting point	alpha (α-S8): 388.36 K ​(115.21 °C, ​239.38 °F)
Boiling point	717.8 K ​(444.6 °C, ​832.3 °F)
Density (near r.t.)	alpha (α-S8): 2.07 g/cm3 beta (β-S8): 1.96 g/cm3 gamma (γ-S8): 1.92 g/cm3
when liquid (at m.p.)	1.819 g/cm3
Critical point	1314 K, 20.7 MPa
Heat of fusion	beta (β-S8): 1.727 kJ/mol
Heat of vaporization	beta (β-S8): 45 kJ/mol
Molar heat capacity	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
Oxidation states	−2, −1, 0, +1, +2, +3, +4, +5, +6 (a strongly acidic oxide)
Electronegativity	Pauling scale: 2.58
Ionization energies	1st: 999.6 kJ/mol 2nd: 2252 kJ/mol 3rd: 3357 kJ/mol (more)
Covalent radius	105±3 pm
Van der Waals radius	180 pm
Spectral lines of sulfur
Other properties
Natural occurrence	primordial
Crystal structure	alpha (α-S8): ​orthorhombic (oF128)
Lattice constants	a = 1.0460 nm b = 1.2861 nm c = 2.4481 nm (at 20 °C)
Crystal structure	beta (β-S8): ​monoclinic (mP48)
Lattice constants	a = 1.0923 nm b = 1.0851 nm c = 1.0787 nm β = 95.905° (at 20 °C)
Thermal conductivity	0.205 W/(m⋅K) (amorphous)
Electrical resistivity	2×1015  Ω⋅m (at 20 °C) (amorphous)
Magnetic ordering	diamagnetic
Molar magnetic susceptibility	alpha (α-S8): −15.5×10−6 cm3/mol (298 K)
Bulk modulus	7.7 GPa
Mohs hardness	2.0
CAS Number	7704-34-9
History
Discovery	before 2000 BCE
Recognized as an element by	Antoine Lavoisier (1777)
Isotopes of sulfur
Main isotopes Decay abun­dance half-life (t1/2) mode pro­duct 32S 94.8% stable 33S 0.760% stable 34S 4.37% stable 35S trace 87.37 d β− 35Cl 36S 0.02% stable 34S abundances vary greatly (between 3.96 and 4.77 percent) in natural samples.

Sulfur or sulphur (IPA: /ˈsʌlfə(ɹ)/, see spelling below) is the chemical element that has the symbol S and atomic number 16. It is an abundant, tasteless, multivalent non-metal. Sulfur, in its native form, is a yellow crystalline solid, with a vile rotten smelling odor. 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, cysteine and methionine. Its commercial uses are primarily in fertilizers, but it is also widely used in gunpowder, matches, insecticides and fungicides.

Notable characteristics

Sulfur melts to a blood-red liquid. When burned, it emits a blue flame.

At room temperature, sulfur is a soft bright yellow solid. Elemental sulfur has only a faint odor similar to that of matches. The odor associated with rotten eggs is from hydrogen sulfide (H₂S) and organic sulfur compounds. Sulfur 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 nonpolar organic solvents such as benzene and toluene. 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 color. Many other rings have been prepared, including S₁₂ and S₁₈. By contrast, its lighter neighbor oxygen only exists in two states of allotropic significance: O₂ and O₃. Selenium, the heavier analogue of sulfur can form rings but is more often found as a polymer chain.

The structure of the cyclooctasulfur molecule, S₈

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 of sulfur is that its viscosity in its molten state, unlike most other liquids, increases above temperatures of 200°C due to the formation of polymer chains. The molten sulfur also becomes dark red in colour above this temperature due to the presence of free valences on terminal atoms of the polymer chains. However, after a specific 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 most important industrial raw materials. 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 is used as photographic fixing agents. Magnesium sulfate, better known as Epsom salts, can be used as a laxative, a bath additive, an exfoliant, a magnesium supplement for plants, or a dessicant. Sulfur is used as the light-generating medium in the rare lighting fixtures known as sulfur lamps. Elemental sulfur crystals are commonly sought after by rock collectors for their brightly colored polyhedron shapes.

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

Sulfur is an essential component of all living cells.

Sulfur may also serve as chemical food source for some primitive organisms: some forms of bacteria use hydrogen sulfide (H₂S) in the place of water as the electron donor in a primitive photosynthesis-like process. Inorganic sulfur forms a part of iron-sulfur clusters, and sulfur is the bridging ligand in the Cu_A site of cytochrome c oxidase, a basic substance involved in utilization of oxygen by all aerobic life.

Sulfur is absorbed by plants via the roots from soil as the sulfate ion and reduced to sulfide before it is incorporated into cysteine and other organic sulfur compounds (sulfur assimilation).

In plants and animals the amino acids cysteine and methionine contain sulfur, as do all polypeptides, proteins, and enzymes which contain these amino acids. Homocysteine and taurine are other sulfur-containing acids which are similar in structure, but which are not coded for by DNA, and are not part of the primary structure of proteins. Glutathione is an important sulfur-containing tripeptide which plays a role in cells as a source of chemical reduction potential in the cell, through its sulfhydryl (-SH) moiety. Many important cellular enzymes use prosthetic groups ending with -SH moieties to handle reactions involving acyl-containing biochemicals: two common examples from basic metabolism are coenzyme A and alpha-lipoic acid.

Disulfide bonds (S-S bonds) formed between cysteine residues in peptide chains are very important in protein assembly and structure. These strong covalent bonds between peptide chains give proteins a great deal of extra toughness and resiliancy. For example, the high strength of feathers and hair is in part due to their high content of S-S bonds and their high content of cysteine and sulfur (eggs are high in sulfur because large amounts of the element are necessary for feather formation). The high disulfide content of hair and feathers also contributes to their indigestibility, and also their disagreeable odor when burned.

Environmental impact

The burning of coal and petroleum by industry and power plants creates massive 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 structures. 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 crystal from Agrigento, Sicily, Italy

Sulfur (Sanskrit, sulvari; Latin sulfur or sulpur) was known in ancient times, and is referred to in the Biblical Pentateuch (Genesis). The word itself probably is from the Arabic sufra meaning yellow, from the bright color of the naturally occurring form, although the Sanskrit name for sulfur, sulvari could also be interpreted as meaning "enemy of copper".

English translations of the Bible commonly refer to sulfur as "brimstone", giving rise to the name of 'Fire and brimstone' sermons, in which listeners are reminded of the fate of eternal damnation that awaits the nonbelieving and unrepented. 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 8th 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

Sulfur crystalites at Waiotapu hot springs, New Zealand

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

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 exploited in Indonesia, Chile, and Japan. Sicily is also famous for its sulphur mines.

Significant deposits 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, although apparently native sulfur may be produced by geological processes alone, without the aid of living organisms (see below). However, fossil-based sulfur deposits from salt domes are the basis for commercial production in the United States, Poland, Russia, Turkmenistan, and Ukraine.

Sulfur production through hydrodesulfurization of oil, gas, and the Athabasca Oil Sands has produced a surplus - huge stockpiles of sulfur now exist throughout Alberta, Canada.

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). 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.

Extraction

Sulfur is extracted by mainly two processes: the Sicilian process and the Frasch process. The Sicilian process, which was first used in Sicily, was utilized in ancient times to get sulfur from rocks present in volcanic regions. In this process, the sulfur deposits are piled and stacked in brick kilns built on sloping hillsides, and with airspaces between them. Then powdered sulfur is then put on top of the sulfur deposit and lit on fire. As the sulfur burns it produces heat, which melts the sulfur deposits, causing the molten sulfur to flow down the sloping hillsides. The molten sulfur can then be collected in wooden buckets.

The second process used to obtain sulfur is the Frasch process. In this method, three concentric pipes are used: The outermost pipe contains superheated water, which melts the sulfur, and the innermost pipe is filled with hot compressed air, which serves to create foam and pressure. The resulting sulfur foam is then expelled out through the middle pipe.

The Frasch process produces sulfur with a 99.5% purity content, and which needs no further purification. On the other hand, the sulfur produced by the Sicilian process must be purified by distillation.

Compounds

Sulfur powder

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, grapefruit mercaptan, a sulfur-containing monoterpenoid 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.

Inorganic sulfur compounds:

• Sulfides (S²⁻), a complex family of compounds usually derived from S²⁻. Cadmium sulfide (CdS) is an example.
• Sulfites (SO₃²⁻), 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 (SO₄²⁻), the salts of sulfuric acid. Sulfuric acid also reacts with SO₃ in equimolar ratios to form pyrosulfuric acid (H₂S₂O₇).
• Thiosulfates (sometimes referred to 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′. A common sulfoxide is DMSO.
• Sulfones have the form R-S(=O)₂-R′. A common sulfone is sulfolane C₄H₈SO₂.

See also Category: sulfur compounds and organosulfur chemistry

Isotopes

Sulfur has 18 isotopes, four of which 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 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 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, Hong Kong, the Commonwealth Caribbean and India, but sulfur in the United States, while both spellings are used in Australia, Canada and New Zealand. IUPAC adopted the spelling “sulfur” in 1990, as did the Royal Society of Chemistry Nomenclature Committee in 1992. The spelling of the term in non-official texts is gradually becoming uniform as sulfur.

The Latin name of the element is sulfur with an F. Since it is an original Latin name and not a Classical Greek loan, the fricative phoneme is indeed denoted with f rather than ph (which would denote the Greek letter φ). Sulfur in Greek is theion (θεῖον), from whence comes the prefix thio-.

See also

• Sulfur cycle
• Disulfide bond
• Sulfonium S⁺, S⁺R₃
• Ultra low sulphur diesel

External links

• Sulfur phase diagram.
• WebElements.com – Sulfur
• chemicalelements.com/sulfur