|Name, Symbol, Number||Sulfur, S, 16|
|Group, Period, Block||16 (VIA), 3 , p|
|Density, Hardness||1960 kg/m3, 2|
|Atomic weight||32.065 amu|
|Atomic radius (calc.)||100 (88) pm|
|Covalent radius||102 pm|
|van der Waals radius||180 pm|
|Electron configuration||[Ne]3s2 3p4|
|e- 's per energy level||2, 8, 6|
|Oxidation states (Oxide)||±2,4,6 (strong acid)|
|State of matter||solid|
|Melting point||388.36 K (239.38 °F)|
|Boiling point||717.87 K (832.5 °F)|
|Molar volume||15.53 ×10-3 m3/mol|
|Heat of vaporization||no data|
|Heat of fusion||1.7175 kJ/mol|
|Vapor pressure||2.65 E-20 Pa at 388 K|
|Speed of sound||__ m/s at 293.15 K|
|Electronegativity||2.58 (Pauling scale)|
|Specific heat capacity||710 J/(kg*K)|
|Electrical conductivity||5.0 E-22 106/m ohm|
|Thermal conductivity||0.269 W/(m*K)|
|1st ionization potential||999.6 kJ/mol|
|2nd ionization potential||2252 kJ/mol|
|3rd ionization potential||3357 kJ/mol|
|4th ionization potential||4556 kJ/mol|
|5th ionization potential||7004.3 kJ/mol|
|6th ionization potential||8495.8 kJ/mol|
|Most Stable Isotopes|
|SI units & STP are used except where noted.|
Notable Characteristics This non-metal is pale yellow in appearance, soft, light, with a distinct odor when allied with hydrogen (rotten egg smell). It burns with a blue flame that emits a peculiar suffocating odor. Sulfur is insoluble in water but soluble in carbon disulfide. Common oxidation states of sulfur include -2, +2, +4 and +6. In all states, solid, liquid, and gaseous, sulfur has allotropic forms, whose relationships are not completely understood. Crystalline sulfur can be shown to form an 8 membered sulfur ring, S8.
Polymeric sulfur nitride[?] has metallic properties even though it doesn't contain any metal atoms. This compound also has unusual electrical and optical properties. Amorphous or "plastic" sulfur is produced through fast cooling crystalline sulfur. X-ray studies show that the amorphous form may have an eight atom per spiral helical structure
Sulfur can be obtained in two crystalline modifications, in orthorhombic octahedra, or in monoclinic prisms, the former of which is the more stable at ordinary temperatures. Applications It is used for many industrial processes such as the production of sulfuric acid (H2SO4) for batteries, the production of gunpowder, and the vulcanization of rubber. Sulfur is used as a fungicide[?], and in the manufacture of phosphate fertilizers. Sulfites are used to bleach papers and dried fruits. Sulfur also finds use in matches and fireworks. Sodium or ammonium thiosulfate are used as photographic fixing agents. Epsom salts, magnesium sulfate, can be used as a laxative, as a bath additive, as an exfoliant[?], or a magnesium supplement in plant nutrition. Biological Role The amino acids cysteine, methionine, homocysteine, and taurine contain sulfur, as do some common enzymes, making sulfur a necessary component of all living cells. Disulfide bonds between polypeptides are very important in protein assembly and structure. Some forms of bacteria use hydrogen sulfide (H2S) in the place of water as the electron doner in a primitive photosynthesis-like process. Sulfur is absorbed by plants from soil as sulfate ion. Inorganic sulfur forms a part of iron-sulfur clusters, and sulfur is the bridging ligand in the CuA site of cytochrome c oxidase. History Sulfur (Sanskrit, sulvere; Latin sulpur) was known in ancient times and was called brimstone in the Biblical story of Pentateuch (Genesis). Homer mentioned "pest-averting sulfur" in the 9th century BC and in 424 BC, the tribe of Bootier destroyed the walls of a city by burning 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. Through experimentation, alchemists knew that the element mercury can be combined with sulfur. In the late 1770s, Antoine Lavoisier helped convince the scientific community that sulfur was an element and not a compound. Occurrence Sulfur occurs naturally in large quantities compounded to other elements in sulfides (example: pyrites) and sulfates (example: gypsum). It is found in its free form near hot springs and volcanic regions and in ores like cinnabar, galena, sphalerite and stibnite[?]. This element is also found in small amounts in coal and petroleum, which produce sulfur dioxide when burned. Fuel standards increasingly require sulfur to be extracted from fossil fuels because sulfur dioxide combines with water droplets to produce acid rain. This extracted sulfur is then refined and represents a large portion of sulfur production. It is also mined along the US Gulf coast, by pumping hot water into sulfur containing deposits (such as salt domes) which melts the sulfur. The molten sulfur is then pumped to the surface.
Through its major derivative, sulfuric acid, 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 industrial and fertilizer complexes. Sulfuric acid production is the major end use for sulfur, and consumption of sulfuric acid has been regarded as one of the best indexes of a nation's industrial development. More sulfuric acid is produced in the United States every year than any other chemical.
The distinctive colors of Jupiter's volcanic moon Io, are from various forms of multen, 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. Compounds Many of the unpleasant odors of organic matter are based on sulfur-containing compounds such as hydrogen sulfide, which has the characteristic smell of rotten eggs. Dissolved in water, hydrogen sulfide is acidic (pKa1 = 7.00, pKa2 = 12.92) and will react with metals to form a series of metal sulfides. Natural metal sulfides are found, especially those of iron. Iron sulfides are called iron pyrites, the so called fool's gold. Interestingly, pyrites can show semiconductor properties. (http://home.earthlink.net/~lenyr/iposc.htm) Galena, a naturally occurring lead sulfide (as the detector in a "cat's hair" rectifier) was of course the original semiconductor discovered.
Some important compounds of sulfur include:
When sulfide minerals are precipitated, isotopic equilibration among solids and liquid may cause small differences in the dS-34 values of co-genetic minerals. The differences between minerals can be used to estimate the temperature of equilibration. The dC-13 and dS-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 contributes 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 hydologic studies. Differences in the natural abundances can also be used in systems where there is sufficient variation in the S-34 of ecosystem components. Rocky Mountain[?] lakes thought to be dominated by atmospheric sources of sulfate have been found to have different dS-34 values from lakes believed to be dominated by watershed sources of sulfate. Precautions Carbon disulfide, hydrogen sulfide, and sulfur dioxide should all be handled with care. In addition to being quite toxic (more toxic than cyanide), sulfur dioxide reacts with atmospheric water to produce acid rain. In high concentration this element can kill quickly by preventing respiration. Sulfur quickly deadens the sense of smell so potential victims may be unaware of its presence. Spelling Sulfur is traditionally spelled "sulphur" in British English, but IUPAC has adopted the spelling "sulfur", as has the Royal Society of Chemistry Nomenclature Committee[?]. Increasingly "sulfur" is being used in British English instead.