Sulfur

From New World Encyclopedia
16 phosphorussulfurchlorine
O

S

Se
S-TableImage.png
periodic table
General
Name, Symbol, Number sulfur, S, 16
Chemical series nonmetals
Group, Period, Block 16, 3, p
Appearance lemon yellow
Sulfur.jpg
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) 2.07 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)
Critical point 1314 K, 20.7 MPa
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

Sulfur or sulphur (see spelling below) (chemical symbol S, atomic number 16) is a yellow crystalline solid at ordinary temperatures and pressures. It is tasteless and odorless and is classified as a nonmetal. It forms stable compounds with all elements except the noble gases. Abundant in nature, it can be found as the pure element or as sulfide and sulfate minerals.

Sulfur (incorporated into certain compounds) is an essential element for living organisms. For instance, it is part of the structure of two amino acids. It is primarily used in fertilizers, but it also useful for making gunpowder, matches, insecticides, and fungicides.

Occurrence

Sulfur crystalites at Waiotapu hot springs, New Zealand.

Elemental sulfur can be found near hot springs and volcanic regions in many parts of the world, especially along the Pacific "Ring of Fire"—a zone of frequent earthquakes and volcanic eruptions encircling the Pacific Ocean. 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. (Evaporites are mineral sediments that are left behind after evaporation of the water they were once dissolved in.) The sulfur in these deposits is believed to come from the action of anaerobic bacteria on sulfate minerals, especially gypsum. Such deposits form 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 led to a glut on the market, and huge stockpiles of sulfur can be seen throughout Alberta.

Common sulfur compounds in nature include:

  • metal sulfides, such as pyrite (iron sulfide), cinnabar (mercury sulfide), galena (lead sulfide), sphalerite (zinc sulfide), and stibnite (antimony sulfide);
  • metal sulfates, such as gypsum (calcium sulfate), alunite (potassium aluminum sulfate), and barite (barium sulfate).

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

History

Sulfur crystal

Sulfur (Sanskrit, sulvere; Latin sulpur) was known in ancient times, and is referred to in several books of the Bible, including the book of Genesis. It has been suggested that the word may have been derived from the Arabic sufra, meaning yellow, which is the color of the naturally occurring form of the element.

In the eighth century B.C.E., Homer mentioned "pest-averting sulfur"; and in 424 B.C.E., the tribe of Boeotia destroyed the walls of a city by burning a mixture of coal, sulfur, and tar under them. Sometime in the twelfth century, the Chinese invented gunpowder, which is a mixture of potassium nitrate (KNO3), carbon, and sulfur. Early alchemists gave sulfur its own alchemical symbol—a triangle at the top of a cross. In the late 1770s, Antoine Lavoisier helped convince the scientific community that sulfur was an element, 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, it was removed by the Frasch process—the underground sulfur was melted by pumping superheated steam through pipes set in the ground, and the molten sulfur was pumped out.

Notable characteristics

A piece of sulfur melts to a blood-red liquid. When burned, it emits a blue flame.

In the periodic table, sulfur is located in group 16 (formerly group 6A), between oxygen and selenium. It is thus a member of the oxygen family of elements, also called the chalcogens. In addition, it lies between phosphorus and chlorine in period 3. It can combine with other elements in different proportions, and it is therefore described as being multivalent.

Elemental sulfur has a faint odor, similiar to that of matches. The common belief that sulfur smells like rotten eggs is actually an association with the smell of hydrogen sulfide (H2S).

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.

The structure of the S8 molecule

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 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.

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 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.

Applications

Sulfur powder.

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.

Biological role

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.

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.

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, 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 (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 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

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 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 thios (θιοσ), which does not bear resemblance to the Latin word.

Fire and brimstone

In the Bible, sulfur is referred to as "brimstone." The book of Genesis, in particular, mentions that God punished evildoers in Sodom and Gomorrah by raining "brimstone and fire" upon them. Accordingly, a "fire and brimstone" sermon is one in which listeners are reminded of the fate of eternal damnation that awaits the unrepentant sinner. Also, hell is implied as having the smell of sulfur, although, as mentioned above, sulfur is odorless. The "smell of sulfur" usually refers to the odor of hydrogen sulfide, which is produced by rotten eggs.

See also

References
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External links

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