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Average: 4/5
Sulfur microcrystals. Source: Ben Mills.
March 25, 2007, 12:00 am
November 1, 2011, 12:18 pm
C Michael Hogan
Stephen C. Nodvin
| Topics: |  Physics & Chemistry Uses Of Chemicals Minerals & Mining Chemical Engineering Air Pollution & Air Quality Environmental Chemistry Biogeochemistry Extremophile |
| Previous Element: Phosphorus Next Element: Chlorine |
|
|
| Physical Properties | ||
|---|---|---|
| Color | pale yellow | |
| Phase at Room Temp. | solid | |
| Density (g/cm3) | 2.069 | |
| Hardness (Mohs) | 2 | |
| Melting Point (K) | 386 | |
| Boiling Point (K) | 717.9 | |
| Heat of Fusion (kJ/mol) | 1.2 | |
| Heat of Vaporization (kJ/mol) | --- | |
| Heat of Atomization (kJ/mol) | 279 | |
| Thermal Conductivity (J/m sec K) | 0.27 | |
| Electrical Conductivity (/mohm cm) | 0 | |
| Source | pyrite ore | |
| Atomic Properties | ||
| Electron Configuration | [Ne]3s23p4 | |
| Number of Isotopes (total, natural) | 25, 4 | |
| Electron Affinity (kJ/mol) | 200.4144 | |
| First Ionization Energy (kJ/mol) | 999.6 | |
| Second Ionization Energy (kJ/mol) | 2251 | |
| Third Ionization Energy (kJ/mol) | 3360.6 | |
| Electronegativity | 2.58 | |
| Polarizability (Å3) | 2.9 | |
| Atomic Weight | 32.06 | |
| Atomic Volume (cm3/mol) | 15.5 | |
| Ionic Radius2- (pm) | 170 | |
| Ionic Radius1- (pm) | --- | |
| Atomic Radius (pm) | 103 | |
| Ionic Radius1+ (pm) | --- | |
| Ionic Radius2+ (pm) | --- | |
| Ionic Radius3+ (pm) | --- | |
| Common Oxidation Numbers | -2; +2,4,6 | |
| Other Oxid. Numbers | -1; +1,3,5 | |
| Abundance | ||
| In Earth's Crust (mg/kg) | 3.50×102 | |
| In Earth's Ocean (mg/L) | 9.05×102 | |
| In Human Body (%) | 0.20% | |
| Regulatory / Health | ||
| CAS Number | 7704-34-9 | |
| OSHA Permissible Exposure | No limits | |
| OSHA PEL Vacated 1989 | No limits | |
| NIOSH Recommended Exposure | No limits | |
| Sources: Mineral Information Institute Jefferson Accelerator Laboratory EnvironmentalChemistry.com |
||
Copper sulfate crystal, one of the many sulfur compounds.
Metallic sulfides form an important class of compounds, since many heavy metals such as zinc, copper, nickel, cobalt and molybdenum occur naturally in large quantities as the sulfide. These compounds are generally semiconducting and are somewhat resistant to degradation by water or weak acids. In the natural environment, such metal sulfides are typically produced by reaction of hydrogen sulfide with metal salts. Lead sulfide is a mineral that was used early as a semiconductor.
Standing within a gentoo penguin colony on King George Island, Antarctica, Dr. C. Michael Hogan served a term as Editor in-Chief of the Encyclopedia of Earth which ended in 2012. In addition to authoring a number of papers for the Encyclopedia of Earth, he is a physicist who has published over 1220 peer reviewed articles in other journals and government monographs in the fields of molecular biology, quantum spinwaves, atmospheric physics, biogeochemistry, hydrological modeling, species populat ... (Full Bio)
| Previous Element: Phosphorus Next Element: Chlorine |
|
|
| Physical Properties | ||
|---|---|---|
| Color | pale yellow | |
| Phase at Room Temp. | solid | |
| Density (g/cm3) | 2.069 | |
| Hardness (Mohs) | 2 | |
| Melting Point (K) | 386 | |
| Boiling Point (K) | 717.9 | |
| Heat of Fusion (kJ/mol) | 1.2 | |
| Heat of Vaporization (kJ/mol) | --- | |
| Heat of Atomization (kJ/mol) | 279 | |
| Thermal Conductivity (J/m sec K) | 0.27 | |
| Electrical Conductivity (/mohm cm) | 0 | |
| Source | pyrite ore | |
| Atomic Properties | ||
| Electron Configuration | [Ne]3s23p4 | |
| Number of Isotopes (total, natural) | 25, 4 | |
| Electron Affinity (kJ/mol) | 200.4144 | |
| First Ionization Energy (kJ/mol) | 999.6 | |
| Second Ionization Energy (kJ/mol) | 2251 | |
| Third Ionization Energy (kJ/mol) | 3360.6 | |
| Electronegativity | 2.58 | |
| Polarizability (Å3) | 2.9 | |
| Atomic Weight | 32.06 | |
| Atomic Volume (cm3/mol) | 15.5 | |
| Ionic Radius2- (pm) | 170 | |
| Ionic Radius1- (pm) | --- | |
| Atomic Radius (pm) | 103 | |
| Ionic Radius1+ (pm) | --- | |
| Ionic Radius2+ (pm) | --- | |
| Ionic Radius3+ (pm) | --- | |
| Common Oxidation Numbers | -2; +2,4,6 | |
| Other Oxid. Numbers | -1; +1,3,5 | |
| Abundance | ||
| In Earth's Crust (mg/kg) | 3.50×102 | |
| In Earth's Ocean (mg/L) | 9.05×102 | |
| In Human Body (%) | 0.20% | |
| Regulatory / Health | ||
| CAS Number | 7704-34-9 | |
| OSHA Permissible Exposure | No limits | |
| OSHA PEL Vacated 1989 | No limits | |
| NIOSH Recommended Exposure | No limits | |
| Sources: Mineral Information Institute Jefferson Accelerator Laboratory EnvironmentalChemistry.com |
||
At room temperature, sulfur is a soft, bright-yellow solid with a faint odor, similar to that of a burning match; a strong sulfurous odor is usually attributed to the presence of hydrogen sulfide or related compounds. Sulfur is an electrical insulator, with a melting point slightly above 100 °C; sulfur is readily subject to sublimation.
Molten sulfur increases in viscosity as temperature increases, in vivid contrast with most other elements in their liquid form up to 200°C due to the formation of polymers. The molten sulfur form assumes a dark red color above this threshhold temperature. At yet higher temperatures, viscosity decreases, with depolymerization occurring.
Burning Sulfur produces sulfur dioxide gas, emitting a blue flame in the process. Sulfur dioxide is noted for its pungent suffocating odor. Sulfur is insoluble in water, but soluble in carbon disulfide, somewhat soluble in other non-polar organic solvents such as the aromatics benzene and toluene. Solid state Sulfur characteristically exists as cyclic crown-shaped S8 molecules. The crystallography of sulfur is a complex subject, since sulfur allotropes form several crystal structures, with both rhombic and monoclinic S8 forms.
Sulfur was known in ancient times, and was identified by the Chinese as early as the sixth century BC, and extracted from pyrite ore by them in the third century BC. It was used by the Egyptians as a cosmetic agent. Within the Bible, in the Book of Genesis, the text terms burning sulfur as brimstone. Pliny notes the place of most abundant ancient occurrence as the Aegean island of Milos.
In 1777, Lavoisier in 1777 established that sulfur is an element rather than a compound. Major deposits of sulfur were found in Texas and Louisianna in the latter nineteenth century, and the Frasch process was subsequently developed to extract sulfur from these ores.
Substantial deposits of elemental sulfur occur in salt domes along the coastal zone of the Gulf of Mexico, as well as in evaporite form in eastern Europe and western Asia. These occurrences are thought to derive anaerobic bacteria acting on sulfate minerals, but native sulfur may be also simply be produced by geological processes. However, fossil-based sulfur deposits from salt domes have, until recently, been the main source of commercial production in the USA, Poland, Russia, Turkmenistan and Ukraine.
Elemental sulfur may also occur near hot springs and volcanic regions, especially along the Pacific Rim. These deposits are currently mined in Indonesia, Chile and Japan. Sulfur deposits are typically polycrystalline, with single crystal occurrences sometimes exceeding 3000 cubic centimeters.
Sulfur is a good example of a material that can be produced from industrial wastes. There are large quantities of sulfur recovered from scrubbed flue gas in combusion of fossil fuels in western countries.
There are a wide variety of sulfur compounds, among them are chiefly:
The only stable oxides are sulfur dioxide and sulfur trioxide, both of which are produced by combusion of sulfur. There are a number of thiosulfate salts that are widely used in industry, especially for photographic developing.
Copper sulfate crystal, one of the many sulfur compounds.
Metallic sulfides form an important class of compounds, since many heavy metals such as zinc, copper, nickel, cobalt and molybdenum occur naturally in large quantities as the sulfide. These compounds are generally semiconducting and are somewhat resistant to degradation by water or weak acids. In the natural environment, such metal sulfides are typically produced by reaction of hydrogen sulfide with metal salts. Lead sulfide is a mineral that was used early as a semiconductor.
Arguably the most important sulfide compound is the gas (at standard temperature and pressure) hydrogen sulfide, which is most typically produced by reaction of elemental sulfur with hydrogen gas. This sulfide gas is slightly acidic when bubbled through water. When sulfur is reduced it yields a variety of polysulfide compounds; these molecules are chains of sulfur atoms terminated with an S-. Ultimate reduction of sulfur leads to sulfur salts such as calcium sulfide and sodium sulfide.
The most important halide of sulfur is the hexafluoride; this substance is a dense gas at standard temperature and pressure; sulfur hexafluoride is emplyed widely as a tracer gas and as a nontoxic propellant. Sulfur hexafluoride was used by the EPA in the first tracer gas calibration study to confirm the validity of a line source model for roadway air pollutant dispersal. Sulfur tetrafluoride is a seldomly utilized reagent which is extremely toxic. The two chief sulfur chlorides are sulfur dichloride and sulfur monochloride. Important oxyhalides include sulfuryl chloride and chlorosulfuric acid, which are derivatives of sulfuric acid. Thionyl chloride is a chemical commonly employed in the synthesis of a variety of organic chemicals.
Commercial uses of sulfur include fertilizers, gunpowder, matches, insecticides, herbicides and fungicides; moreover, large quantities of sulfur are employed in the vulcanization of rubber. Vulcanization involves the process of adding sulfur to rubber in the presence of elevated temperature, in order to improve the stability of rubber at temperature extrema and to enhance the material's resistance to wear. Production of sulfuric acid is also an important aspect of industrial sulfur utilization. Elemental sulfur crystals are prized by mineral collectors for their brightly colored polyhedron shapes.
Fossil fuels, such as coal and oil, can contain sulfur from trace amounts to a few percent, and the combustion of these fuels results in emissions of sulfur dioxide. The only fossil fuel which can contain minimal sulfur is natural gas. Sulfur emissions by fossil fuel-fired facilities can be appreciably reduced (>95%) by the application of wet and dry desulfurization installations, based on either washing with a calcium hydroxide solution or reaction with calcium oxide.
The sulfur dioxide produced by petroleum, coal or natural gas combustion constitutes a major air pollutant worldwide. In the last three decades this pollutant has come under significant control in the USA, Canada, western Europe, Australia and New Zealand; however, China, India, Brazil and Russia constitute major sources of ongoing sulfur dioxide air pollution. Sulfur dioxide represents a significant cause of lung disease and human mortality in the world, especially in the four countries cited above as ongoing major polluters; there are an estimated 300,000 to 700,000 deaths per annum worldwide attributable to sulfur dioxide.
Sulfur is an essential element within all organism cells. Amino acids cysteine and methionine contain sulfur, as well as polypeptides, proteins, and enzymes that contain these amino acids. Disulfide bonds of cysteine residues in peptide chains are vital in protein snthesis. These covalent molecular bonds between peptide chains confer toughness and rigidity. The high strength of feathers and hair is in part due to their high content of disulfide bonds and the correspondingly elevated content of cysteine and sulfur. Bird eggs are high in sulfur since high disulfide levels are essental for competent feathers; moreover, the well-known smell of rotten egg is from hydrogen sulfide. The high disulfide bond content of hair and feathers contributes to their indigestibility, and also to their foul aroma when combusted.
Sulfate is taken up by the roots for many vascular plants. The maximum uptake rate is typically attained at sulfate levels of 0.1 mM. The uptake of sulfate through the root epidermis and its upward transport strictly regulated at the root zone. Sulfate is actively transported across the plasma membrane of the root cells, subsequently loadthence moving into the xylem and pumped to the shoot by the transpiration mechanism.
The uptake and transport of sulfate is energy dependent (driven by a proton gradient controlled by ATPases) via proton/sulfate co-transport. In the growing region of the upper plant, the sulfate is conveyed to the chloroplasts where it is reduced. The balance of sulfate in plant tissue is predominantly resident in the vacuole. Plants may take in certain quantities of sulfur dioxide from its air pollutant status in the atmosphere, and even metabolize low concentrations of this gas; however, at elevated concentrations sulfur dioxide gas is phytotoxic.
Sulfur is the energy source for a number of bacteria genera that utilize hydrogen sulfide rather than water as an electron donor in a primordial photosynthenthic type process for which oxygen is the electron receptor. both green and purple sulfur bacteria, as well as certain chemolithotrophs employ elemental oxygen to effect oxidization of hydrogen sulfide producing elemental sulfur with a zero oxidation state. Primitive extremophile bacteria found in deep ocean thermal vents oxidize hydrogen sulfide; for example, the giant tube worm is a macro-invertebrate organism making metabolic use of hydrogen sulfide via the intermediary of bacteria.
Sulfur bacteria breathe sulfate instead of oxygen. They use sulfur as the electron acceptor, and reduce various oxidized sulfur compounds back into sulfide, often into hydrogen sulfide. They also can grow on a number of other partially oxidized sulfur compounds (e.g. thiosulfates, thionates, polysulfides, sulfites). The hydrogen sulfide and related sulfur compounds produced by these bacteria causes the odor of some intestinal gases and decomposition products.
The chief forms and products of sulfur are:
Sulfur exhibits 25 well defined isotopes, four of which are stable: 32S (95.02%), 33S (0.75%), 34S (4.21%), and 36S (0.02%). The abundance of sulfur-32 is based on its primordial production from carbon-12, and successive fusion capture of five helium nuclei, in the alpha process of exploding type II supernovae.
Besides 35S, the radioactive isotopes of sulfur are all relatively short-lived. 35S is formed from cosmic ray spallation of 40A (argon) in the Earth's atmosphere.Sulfur-35 has a decay half life of 87 days; the next longest half life radioisotope is sulfur-38, with a half life of 170 minutes.
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