Sulfur belongs to a nonmetallic chemical element (pure product: yellow crystalline solid) under the symbol S. It can actively react with many other elements. It exists in various kinds of forms and compound such as sulfide and sulfate minerals which can be found everywhere around the universe and earth. It is also a key element for all life as the major component of amino acids, vitamins and many other cofactors. Sulfur has applications in various kinds of fields. For example, one of its biggest applications is for the production of sulfuric acid for sulfate and phosphate fertilizers. It is also used for the manufacturing of insecticides, fungicides, and bactericides. In pharmaceutical, it can be used for the manufacturing of many kinds of sulfur-containing antibiotics.
Can cause eye irritation; may rarely irritate skin. If recovered sulfur, refer to hydrogen sulfide.*
Elemental sulfur is recovered from its ore deposits found throughout the world. It is obtained commercially by the Frasch process, recovery from wells sunk into salt domes. Heated water under pressure is forced into the underground deposits to melt sulfur. Liquid sulfur is then brought to the surface. Sulfur is recovered by distillation. Often the ore is concentrated by froth flotation.
Elemental sulfur also is recovered as a by-product in processing natural gas and petroleum. Refining operations of natural gas and petroleum crude produce hydrogen sulfide, which also may occur naturally. Hydrogen sulfide is separated from hydrocarbon gases by absorption in an aqueous solution of alkaline solvent such as monoethanol amine. Hydrogen sulfide is concentrated in this solvent and gas is stripped out and oxidized by air at high temperature in the presence of a catalyst (Claus process).
Elemental sulfur also may be obtained by smelting sulfide ores with a reducing agent, such as coke or natural gas, or by reduction of sulfur dioxide.
A pale yellow crystalline solid with a faint odor of rotten eggs. Insoluble in water. A fire and explosion risk above 450° F. Transported as a yellow to red liquid. Handled at elevated temperature (typically 290°F) to prevent solidification and makes transfers easier. Hot enough that plastic or rubber may melt or lose strength. Causes thermal burns to skin on contact. Cools rapidly and solidifies if released. Equipment designed to protect against ordinary chemical exposure is ineffective against the thermal hazard. Exercise caution walking on the surface of a spill to avoid breakthrough into pockets of molten sulfur below the crust. Do not attempt to remove sulfur impregnated clothing because of the danger of tearing flesh if a burn has resulted. May be irritatin to skin, eyes and mucous membranes. Used in sulfuric acid production, petroleum refining, and pulp and paper manufacturing.
Sulfur forms two oxides, sulfur dioxide, SO2, and the trioxide, SO3. It burns in oxygen at about 250°C or in air above 260°C, forming sulfur dioxide. In excess oxygen the trioxide is obtained.
Sulfur reacts with hydrogen at 260 to 350°C forming hydrogen sulfide. The reaction is slow at this temperature and does not go to completion. The reaction is catalyzed by activated alumina.
Reactions with excess chlorine or fluorine yield sulfur tetrachloride, SCl4, or hexafluoride, SF6. These reactions occur under cold conditions.
Sulfur reacts with sulfur dioxide in an electric discharge to form disulfuroxide, S2O.
Sulfur reacts with aqueous sulfide to form polysulfides: S + Na2S → Na2S2
With aqueous solution of sulfite the product is thiosulfate:
S + SO32– → S2O32–
Thiosulfate also is obtained by heating sulfur with powdered sulfite:
S + Na2SO3 → Na2S2O3
When heated with alkali cyanide, thiocyanate salt is obtained:
S + KCN → KSCN
A similar reaction occurs in the aqueous phase in which thiocyanate is obtained by evaporation and crystallization.
Sulfur combines with alkali metals, copper, silver, and mercury on cold contact with the solid, forming sulfides. Reactions with magnesium, zinc, and cadmium occur to a small degree at ordinary temperatures, but rapidly on heating. Sulfur reacts with phosphorus, arsenic, antimony, bismuth, and silicon at their melting points and with other elements at elevated temperatures forming binary sulfides. Sulfides of tellurium, gold, platinum, and iridium are difficult to obtain even at elevated temperatures. Sulfur does not react with inert gases, nitrogen, and iodine.
Elemental sulfur is used for vulcanizing rubber; making black gunpowder; as a soil conditioner; as a fungicide; preparing a number of metal sulfides; and producing carbon disulfide. It also is used in matches; bleaching wood pulp, straw, silk, and wool; and in synthesis of many dyes. Pharmaceutical grade precipitated and sublimed sulfurs are used as scabicides and as antiseptics in lotions and ointments.
Important sulfur compounds include sulfuric acid, sulfur dioxide, hydrogen 890 SULFUR sulfide, sulfur trioxide, and a number of metal sulfides and metal oxo- salts such as sulfates, bisulfates, and sulfites. Numerous organic compounds contain sulfur, such as mercaptans, thiophenes, thiophenols, sulfate esters, sulfones, and carbon disulfide.
Sulfur was known to the alchemists from ancient times as brimstone. Lavoisier in 1772 proved sulfur to be an element. The element derived its name from both the Sanskrit and Latin names Sulvere and Sulfurium, respectively. Sulfur is widely distributed in nature, in earth's crust, ocean, meteorites, the moon, sun, and certain stars. It also is found in volcanic gases, natural gases, petroleum crudes, and hot springs. It is found in practically all plant and animal life. Most natural sulfur is in iron sulfides in the deep earth mantle. The abundance of sulfur in earth’s crust is about 350 mg/kg. Its average concentration in seawater is estimated to be about 0.09%. Sulfur occurs in earth’s crust as elemental sulfur (often found in the vicinity of volcanoes), sulfides, and sulfates. The most important sulfur-containing ores are iron pyrite, FeS2; chalcopyrite, CuFeS2; sphalerite, ZnS; galena, PbS; cinnabar HgS; gypsum CaSO4•2H2O; anhydrite CaSO4; kieserite, MgSO4•H2O; celestite, SrSO4; barite, BaSO4; and. stibnite, Sb2S3.
Liquamat (Galderma ); Sastid (Stiefel); Sulfur Soap (Stiefel).
Murphy, Clabaugh & Gilchrist [J Res Nat Bur Stand 64A 355 1960] have obtained sulfur of about 99.999% purity by the following procedure: Roll sulfur was melted and filtered through a coarse-porosity glass filter funnel into a 2L round-bottomed Pyrex flask with two necks. Conc H2SO4 (300mL) was added to the sulfur (2.5kg), and the mixture was heated to 150o, stirring continuously for 2hours. Over the next 6hours, conc HNO3 was added in about 2mL portions at 10-15minutes intervals to the heated mixture. It was then allowed to cool to room temperature and the acid was poured off. The sulfur was rinsed several times with distilled water, then remelted, cooled, and rinsed several times with distilled water again, this process being repeated four or five times to remove most of the acid entrapped in the sulfur. An air-cooled reflux tube (ca 40cm long) was attached to one of the necks of the flask, and a gas delivery tube (the lower end about 2.5cm above the bottom of the flask) was inserted into the other. While the sulfur was boiled under reflux, a stream of helium or N2 was passed through to remove any water, HNO3 or H2SO4, as vapours. After 4hours, the sulfur was cooled so that the reflux tube could be replaced by a bent air-cooled condenser. The sulfur was then distilled, rejecting the first and the final 100mL portions, and transferred in 200mL portions to 400mL glass cylinder ampoules (which were placed on their sides during solidification). After adding about 80mL of water, displacing the air with N2, the ampoule was cooled, and the water was titrated with 0.02M NaOH, the process being repeated until the acid content was negligible. Finally, entrapped water was removed by alternate evacuation to 10mm Hg and refilling with N2 while the sulfur was kept molten. The ampoules were then sealed. Other purifications include crystallisation from CS2 (which is less satisfactory because the sulfur retains appreciable amounts of organic material), *benzene or *benzene/acetone, followed by melting and degassing. It has also been boiled with 1% MgO, then decanted, and dried under a vacuum at 40o for 2days over P2O5. [For the purification of S6, “recrystallised S8” and “Bacon-Fanelli sulfur” see Bartlett et al. J Am Chem Soc 83 103, 109 1961.]
SULFUR reacts violently with strong oxidizing agents causing fire and explosion hazards [Handling Chemicals Safely 1980 p. 871]. Reacts with iron to give pyrophoric compounds. Attacks copper, silver and mercury. Reacts with bromine trifluoride, even at 10°C [Mellor 2:113. 1946-47]. Ignites in fluorine gas at ordinary temperatures [Mellor 2:11-13 1946-47]. Reacts to incandescence with heated with thorium [Mellor 7:208 1946-47]. Can react with ammonia to form explosive sulfur nitride. Reacts with calcium phosphide incandescently at about 300°C. Reacts violently with phosphorus trioxide [Chem. Eng. News 27:2144 1949]. Mixtures with ammonium nitrate or with metal powders can be exploded by shock [Kirk and Othmer 8:644]. Combinations of finely divided sulfur with finely divided bromates, chlorates, or iodates of barium, calcium, magnesium, potassium, sodium, or zinc can explode with heat, friction, percussion, and sometimes light [Mellor 2 Supp.1:763. 1956]. A mixture with barium carbide heated to 150°C becomes incandescent. Reacts incandescently with calcium carbide or strontium carbide at 500°C. Attacks heated lithium, or heated selenium carbide with incandescence [Mellor 5:862 1946-47]. Reacts explosively if warmed with powdered zinc [Mellor 4:476. 1946-47]. Reacts vigorously with tin [Mellor 7:328. 1946-47]. A mixture with potassium nitrate and arsenic trisulfide is a known pyrotechnic formulation [Ellern 1968 p. 135]. Mixtures with any perchlorate can explode on impact [ACS 146:211-212]. A mixture of damp sulfur and calcium hypochlorite produces a brilliant crimson flash with scatter of molten sulfur [Chem. Eng. News 46(28):9 1968]. Takes fire spontaneously in chlorine dioxide and may produce an explosion [Mellor 2:289 (1946-47)]. Ignites if heated with chromic anhydride ignite and can explode, [Mellor 10:102 (1946-47)]. Even small percentages of hydrocarbons in contact with molten sulfur generate hydrogen sulfide and carbon disulfide, which may accumulate in explosive concentrations. Sulfur reacts with Group I metal nitrides to form flammable mixtures, evolving flammable and toxic NH3 and H2S gasses if water is present. (Mellor, 1940, Vol. 8, 99).
Sulfur, S, is a nonmetallic element that exists in a crystalline or amorphous form and in four stable isotopes. Sulfur melts at temperatures rangingfrom 112.8°C (234 °F) for the rhombic form to 120.0°C(248 °F) for amorphous sulfur,and all forms boil at 444.7°C (835°F). Sulfur occurs as free sulfur in many volcanic areas and is often associated with gypsum and limestone. It is used as a chemical intermediate and fungicide and in the vulcanization of rubber.
Air & Water Reactions
Flammable. Insoluble in water.