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Encyclopedia > Sulfer
phosphorussulfurchlorine
O
S
Se  
 
 
Image:S-TableImage.png
Full table
General
Name, Symbol, Number sulfur, S, 16
Chemical series nonmetals
Group, Period, Block 16 (VIA), 3 , p
Density, Hardness 1960 kg/m3, 2
Appearance Lemon yellow at STP
Atomic properties
Atomic weight 32.065 amu
Atomic radius (calc.) 100 pm (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) -1,2,4,6 (strong acid)
Crystal structure orthorhombic
Physical properties
State of matter solid
Melting point 388.36 K (239.38 F)
Boiling point 717.87 K (832.5 F)
Molar volume 15.53 10-6 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
Miscellaneous
Electronegativity 2.58 (Pauling scale)
Specific heat capacity 710 J/(kg*K)
Electrical conductivity 5.0 E-22 106/(mohm)
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
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
SI units & STP are used except where noted.


Sulfur (or Sulphur; see spelling below) is the chemical element in the periodic table that has the symbol S and atomic number 16. An abundant tasteless odorless multivalent non-metal, sulfur is best known as yellow crystals and occurs as sulfide and sulfate minerals and even in its native form (especially in volcanic regions). It is an essential element in all living organisms and is needed in several amino acids and hence in many proteins. It is primarily used in fertilizers but is also widely used in gunpowder, laxatives, matches, insecticides and fungicides.

Contents

Notable characteristics

This non-metal is pale yellow in appearance, soft, light, with a distinct odor when allied with hydrogen (rotten egg smell; it should be noted that, contrary to popular belief, the element itself does not have such an odor). It burns with a blue flame that emits a peculiar suffocating odor (sulfur dioxide, SO2). Sulfur is insoluble in water but soluble in carbon disulfide. Common oxidation states of sulfur include -2, +2, +4 and +6. Elemental sulfur forms stable compounds with all elements except the noble gases.


Sulfur ordinarily exists as molecules in the solid state, mainly adopting a cyclic crown-shaped S8. The S8 structure is easily represented on paper by writing two W letters of different sizes, with their ends connected on each side: each corner represents a S atom. Removing one atom from the crown gives S7, which is responsible for the bright yellow color associated with "sulfur". Many different rings have been prepared, e.g. S18 and S12.



S8 crystallizes in several ways, such as "rhombic" and "monoclinic" crystalline forms, depending on the specific conditions. Of course the other rings also can crystallize in multiple "habits". Because of its diverse rings sizes and crystallization patterns, elemental sulfur is probably the most complicated element. By contrast, its lighter neighbor, oxygen only exists in two states of chemical significance - O2 and O3 (ozone, whose structure resembles that of SO2). The heavier analogue of sulfur, selenium, forms a few rings but is more stable as a gray polymer.



The most noteworthy property of sulfur is that in its molten state, its viscosity increases with temperature due to the formation of polymers. The viscosity of most liquids decrease with increasing temperature. "Amorphous" sulfur, which is elastic, is obtained by pouring molten sulfur into ice water. At room temperature, the polymeric form of sulfur loses its elasticity and converts into rhombic sulfur within a matter of hours or few days.


Applications

It is used for many industrial processes such as the production of sulfuric acid (H2SO4) for batteries and detergents, 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 donor 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.


The massive burning of coal by industry and power plants liberates huge amounts of sulfur dioxide, which reacts with atmospheric water and oxygen to produce sulfuric acid. By changing the pH of soil and freshwater bodies, the resulting acid rain has led to substantial damage to the natural environment in some regions.


History

Enlarge
Sulfur crystal

Sulfur (Sanskrit, sulvere; Latin sulpur) was known in ancient times, and is referred to in the Biblical story of Pentateuch (Genesis). English translations of this commonly refer to sulfur as "brimstone", giving rise to the name of 'Fire and brimstone' sermons, which are sermons where hell and eternal damnation for sinners is stressed. It is from this part of the bible that hell is thought to smell of sulfur. The word itself is almost certainly from the Arabic "sufra" meaning yellow, from the bright color of the naturally-occurring form. If there is indeed a Sanskrit word 'sulvere', this can only be a borrowing. 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 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. 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. 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

Enlarge
Sulfur

Sulfur occurs naturally in large quantities compounded to other elements in sulfides (example: pyrite) and sulfates (example: gypsum). It is found in its free form near hot springs and volcanic regions (hence the name brimstone, from being found at the brim of craters) 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 the Frasch process, which involves pumping a mixture of compressed air and superheated water into sulfur containing deposits (such as salt domes). The hot water melts the sulfur, and the pressure of the air drives the molten sulfur 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 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.


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 pyrite, the so called fool's gold. Interestingly, pyrite can show semiconductor properties.[1] (http://home.earthlink.net/~lenyr/iposc.htm) Galena, a naturally occurring lead sulfide, was the first semiconductor discovered, and found a use as a signal rectifier in the early "cat's whisker" (crystal radios)


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 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, however, and gradually reverts back to crystalline form.


Other important compounds of sulfur include:

  • sodium dithionite, Na2S2O4, a powerful reducing agent.
  • sulfurous acid, H2SO3, created by dissolving SO2 in water. Sulfurous acid and the corresponding sulfites are fairly strong reducing agents. Other compounds derived from SO2 include the pyrosulfite ion (S2O52-).
  • The thiosulfates (S2O32-). Thiosulfates are used in photographic fixing, are oxidizing agents, and ammonium thiosulfate is being investigated as a cyanide replacement in leaching gold.[2] (http://doccopper.tripod.com/gold/AltLixiv.html)
  • Compounds of dithionic acid (H2S2O6)
  • The polythionic acids, (H2SnO6), where n can range from 3 to 80.
  • The sulfates, the salts of sulfuric acid. Epsom salts are magnesium sulfate.
  • Sulfides are simple compounds of sulfur with some other chemical element.
  • Sulfuric acid reacting with SO3 in equimolar ratios forms pyrosulfuric acid.
  • peroxymonosulfuric acid and peroxydisulfuric acids, made from the action of SO3 on concentrated H2O2, and H2SO4 on concentrated H2O2 respectively.
  • tetrasulfur tetranitride S4N4.
  • Thiocyanates are compounds containing the thiocyanate ion, SCN-
  • thiocyanogen, (SCN)2.
  • A thioether is a molecule with the form R-S-R', where R and R' are organic groups. These are the sulfur equivalents of ethers.
  • A thiol (also known as a mercaptan) is a molecule with an -SH functional group. These are the sulfur equivalents of alcohols.
  • A thiolate ion has an -S- functional group attached. These are the sulfur equivalent of alkoxide ions.

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. Sulfur-35 is formed from cosmic ray spallation of argon-40 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 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.


Although sulfur dioxide is sufficiently safe to be used as a food additive in small amounts, in higher atmospheric concentration it reacts with water in the lungs to form sulfurous acid there; this causes immediate bleeding, the lungs fill up with blood and suffocation results. In creatures without lungs such as insects or plants, it otherwise prevents respiration.


Hydrogen sulfide is quite toxic (more toxic than cyanide). Although very smelly 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, India, and New Zealand, but sulfur in the United States, while both spellings are used in Australia and Canada. The IUPAC has adopted the spelling "sulfur", as has the Royal Society of Chemistry Nomenclature Committee.


See also

References

External links


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