Sulphur: Difference between revisions
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Common [[oxidation state]]s of sulfur include −2, +2, +4 and +6. Sulfur forms stable compounds with all elements except the [[noble gas]]es. Sulfur in the solid state ordinarily exists as cyclic crown-shaped S<sub>8</sub> [[molecule]]s. | Common [[oxidation state]]s of sulfur include −2, +2, +4 and +6. Sulfur forms stable compounds with all elements except the [[noble gas]]es. Sulfur in the solid state ordinarily exists as cyclic crown-shaped S<sub>8</sub> [[molecule]]s. | ||
{{Image|Sulfur crystals.jpg|right|175px|Sulfur crystals}} | |||
The [[crystallography]] of sulfur is complex. Depending on the specific conditions, the sulfur [[allotrope]]s form several distinct [[crystal structure]]s, with [[rhombic]] and [[monoclinic]] S<sub>8</sub> best known. | The [[crystallography]] of sulfur is complex. Depending on the specific conditions, the sulfur [[allotrope]]s form several distinct [[crystal structure]]s, with [[rhombic]] and [[monoclinic]] S<sub>8</sub> best known. | ||
Revision as of 02:10, 11 June 2009
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Sulfur, or sulphur, is a non-metallic chemical element.
Sulfur has the symbol S and an atomic weight of 32.065. It is a yellowish crystalline solid in its elemental form and it is an element essential for life.
The vast majority of the 66,000,000 metric tons of sulfur produced worldwide in 2006 was by-product sulfur recovered from petroleum refining and natural gas processing plants by the Claus process.[1]
Sulfur is widely used in the manufacture of sulfuric acid (H2SO4} and various fertilizers.
Characteristics
At room temperature, sulfur is a soft, bright-yellow solid. Elemental sulfur has only a faint odor, similar to that of matches.
Sulfur burns with a blue flame that emits sulfur dioxide, notable for its peculiar suffocating odor due to dissolving in the mucosa to form dilute sulfurous acid.
Sulfur itself is insoluble in water, but soluble in carbon disulfide and to a lesser extent in other non-polar organic solvents such as benzene and toluene.
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.
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 of sulfur is that its viscosity in its molten state, unlike most other liquids, increases above temperatures of 200 °C due to the formation of polymers. The molten sulfur assumes a dark red color above this temperature. At higher temperatures, however, the viscosity is decreased as depolymerization occurs.
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.