|Name, Symbol, Number|| sulfur, S, 16
|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|
|Density (near r.t.)||(alpha) 2.07 g·cm−3|
|Density (near r.t.)||(beta) 1.96 g·cm−3|
|Density (near r.t.)||(gamma) 1.92 g·cm−3|
|Liquid density at m.p.||1.819 g·cm−3|
|Melting point|| 388.36 K|
(115.21 °C, 239.38 °F)
|Boiling point|| 717.8 K|
(444.6 °C, 832.3 °F)
|Critical point||1314 K, 20.7 MPa|
|Heat of fusion||(mono) 1.727 kJ·mol−1|
|Heat of vaporization||(mono) 45 kJ·mol−1|
|Heat capacity||(25 °C) 22.75 J·mol−1·K−1|
|Oxidation states|| −1, ±2, 4, 6|
(strongly acidic oxide)
|Electronegativity||2.58 (Pauling scale)|
| Ionization energies|
|1st: 999.6 kJ·mol−1|
|2nd: 2252 kJ·mol−1|
|3rd: 3357 kJ·mol−1|
|Atomic radius||100 pm|
|Atomic radius (calc.)||88 pm|
|Covalent radius||102 pm|
|Van der Waals radius||180 pm|
|Magnetic ordering||no data|
|Electrical resistivity|| (20 °C) (amorphous)|
|Thermal conductivity|| (300 K) (amorphous)|
|Bulk modulus||7.7 GPa|
|CAS registry number||7704-34-9|
Sulfur or sulphur (IPA: /ˈsʌlfə(ɹ)/, 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 (Chalcogen). Sulfur, in its native form, is a yellow crystalline 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, Cysteine and Methionine. Its commercial uses are primarily in fertilizers, but it is also widely used in gunpowder, matches, insecticides, and fungicides.
At room temperature, sulfur is a soft bright yellow solid. Although sulfur is blamed for the smell—, e.g. of rotten eggs— elemental sulfur has only the faintest odor (the odor associated with rotten eggs is actually due to hydrogen sulfide and organic sulfur compounds). 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 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.
A noteworthy property of sulfur is that its viscosity in the molten sulfur, unlike most other liquids, increases with temperature due to the formation of 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.
Sulfur has many industrial uses. Through its major derivative, sulfuric acid (H2SO4), sulfur ranks as one of the more 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, or a magnesium supplement for plants. Sulfur is used as the light-generating medium in the rare lighting fixtures known as sulfur lamps.
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.
The amino acids cysteine and methionine contain sulfur, as do all polypeptides, proteins, and enzymes which contain these amino acids. This makes sulfur a necessary component of all living cells. Disulfide bonds between polypeptides are very important in protein assembly and structure. Homocysteine and taurine are also sulfur 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 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). Inorganic sulfur forms a part of iron-sulfur clusters, and sulfur is the bridging ligand in the CuA site of cytochrome c oxidase. Sulfur is an important component of coenzyme A.
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.
Sulfur (Sanskrit, sulvari; Latin sulfur or sulpur) was known in ancient times, and is referred to in the Biblical book of 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 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 gunpowder, 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.
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.
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, 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 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). 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.
Sulfur is extracted by mainly 2 processes. The first of is the Sicialian proceces. It was used in ancient times to get sulfur from rocks present in volcanic regions. The process was first used in Sicily. In this process, the sulfur deposits are piled and stacked in brick kilns. The kilns are built on sloping hillsides with airspaces between them. Then powdered sulfur is put on top of the sulfur deposit and lit on fire. The sulfur burns and evolves heat. This melts the sulfur deposits and molten sulfur flows down the sloping hillsides and is collescted in wooden buckets. The second process if the Frasch process. In this three coencentric pipes are used. The outer with superheated water and the innermost with hot compressed air. The water melts the sufur and the air creates foam and prssure. The sulfur foam is then expelled out through the middle pipe. This process produces sulfur of 99.5% purity and needs no further purification, while the sicillian process sulfur must be purified by distillation.
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. 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.
- Sodium dithionite, Na2S2O4, is the highly reducing dianion derived from hyposulfurous/dithionous acid.
- Sodium dithionate (Na2S2O6).
- Polythionic acids (H2SnO6), 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 (Na2Sx)
- 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 alcohol.
- 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.
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.
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.
The element has traditionally been spelled sulphur in the United Kingdom, Ireland, Hong Kong 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. This spelling has begun to replace its variant in official use, unlike aluminum, which is not commonly used outside North America, and which IUPAC rejected in 1990 in favor of aluminium.
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).
Fire and brimstoneEdit
Christian countries often associate sulfur, (in English usually under its ancient name, brimstone) with Hell and divine wrath, mostly due to the phrase "fire and brimstone", which occurs in the Bible in descriptions of Hell and divine punishment. "Fire and brimstone" sermons are those used by some preachers to compel belief by depictions of the horrors of Hell and its punishments. A joke among scientists has used those descriptions of Hell to conclude that, whereas Heaven's temperature would be a scorching 525 degrees Celsius (because it is bathed in a light of the sun... sevenfold as the light of seven days) Hell can be no hotter than the boiling point of brimstone (a mere 444.6 degrees Celsius), and thus cannot be as hot as Heaven. 
|This page uses content from Wikipedia. The original article was at Sulfur. The list of authors can be seen in the page history. As with Chemistry, the text of Wikipedia is available under the GNU Free Documentation License.|