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(sŭl`fər), nonmetallic chemical element; symbol S; at. no. 16; interval in which at. wt. ranges 32.059–32.076; m.p. 112.8°C; (rhombic), 119.0°C; (monoclinic), about 120°C; (amorphous); b.p. 444.674°C;; sp. gr. at 20°C;, 2.07 (rhombic), 1.957 (monoclinic), 1.92 (amorphous); valence −2, +4, or +6. Sulfur was known to the ancients; it is the brimstone of the Bible. It was first recognized as an element in 1777 by A. L. Lavoisier.

Properties and Compounds

Sulfur is found in Group 16 of the periodic tableperiodic table,
chart of the elements arranged according to the periodic law discovered by Dmitri I. Mendeleev and revised by Henry G. J. Moseley. In the periodic table the elements are arranged in columns and rows according to increasing atomic number (see the table entitled
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. It exhibits allotropyallotropy
[Gr.,=other form]. A chemical element is said to exhibit allotropy when it occurs in two or more forms in the same physical state; the forms are called allotropes.
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. Solid sulfur occurs principally in three forms, all of which are brittle, yellow in color, odorless, tasteless, and insoluble in water. Two of these solid forms are crystalline, composed of molecules containing eight sulfur atoms and having molecular weight 256.512 amu. Rhombic sulfur has orthorhombic crystalline structure and is stable below 95.5°C;; most sulfur is in this form. The monoclinic, or prismatic, form has long, needlelike, nearly transparent crystals; it is stable between 95.5°C; and its melting point but reverts to the rhombic form on standing at room temperature. Amorphous sulfur is a dark, noncrystalline, gumlike substance. It is often thought to be a supercooled liquid; it is formed by rapidly cooling molten sulfur, e.g., by pouring it into cold water. It slowly reverts to the rhombic form on standing. The crystalline forms are readily soluble in carbon disulfide, but the amorphous form is not. Many other forms of sulfur exist. Liquid sulfur is unusual in that its viscosity increases as it is heated. This property is thought to be due to the formation of long polymeric chains of sulfur molecules.

Sulfur is a chemically active element and forms many compounds, both by itself (sulfidessulfide,
chemical compound containing sulfur and one other element or sulfur and a radical. Sulfides may be salts or esters of hydrogen sulfide, H2S, or may be formed directly, e.g., by heating a metal with sulfur.
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) and in combination with other elements. It is part of many organic compounds, e.g., mercaptans (thiols) and thio compounds. It burns in air with a blue flame, forming sulfur dioxide, SO2.

Natural Occurrence and Processing

Sulfur is widely distributed in nature. It is found in many minerals and ores, e.g., iron pyrites, galena, cinnabar, zinc blende, gypsum, barite, and epsom salts and in mineral springs and other waters. It is found uncombined in some volcanic regions and in large underground deposits in Sicily and in Texas and Louisiana. Sulfur often occurs with coal, petroleum, and natural gas. Sulfur is found in meteorities, and deposits of it may be present near the lunar crater Aristarchus. The distinctive colors of Jupiter's moon Io are believed to result from forms of molten, solid, and gaseous sulfur. Sulfur is a component of all living cells. The amino acids cysteine, methionine, homocysteine, and taurine contain sulfur as do some common enzymes; it is a component of most proteins. Some forms of bacteria use hydrogen sulfide (H2S) in place of water in a rudimentary photosynthesislike process. Sulfur is absorbed by plants from soil as sulfate ions.

Sulfur is produced chiefly by the Frasch processFrasch process
[for Herman Frasch, the German-American chemist who devised it], process for the extraction of sulfur from subsurface deposits. Three pipes, one inside another, are sunk to the bottom of the sulfur bed.
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, although it is also produced by the Sicilian method and by other methods. In the Sicilian method the sulfur-bearing ores are piled in a mound and ignited. The heat produced by the burning melts some of the sulfur, which is collected and cast. This sulfur is impure and is usually purified by sublimation. Sulfur is also recovered from natural gas, coal, crude oil, and other sources, e.g., the flue dusts and gases from the refining of metal sulfide ores. Elemental sulfur is obtained in several forms, including flowers of sulfur, a fine crystalline powder, and roll sulfur (cast cakes or sticks).


Elemental sulfur is used in black gunpowdergunpowder,
explosive mixture; its most common formula, called "black powder," is a combination of saltpeter, sulfur, and carbon in the form of charcoal. Historically, the relative amounts of the components have varied.
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, matches, and fireworks; in the vulcanizationvulcanization
, treatment of rubber to give it certain qualities, e.g., strength, elasticity, and resistance to solvents, and to render it impervious to moderate heat and cold.
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 of rubber; as a fungicide, insecticide, and fumigant; in the manufacture of phosphate fertilizers; and in the treatment of certain skin diseases. The principal use of sulfur, however, is in the preparation of its compounds. The most important sulfur compound is sulfuric acidsulfuric acid,
chemical compound, H2SO4, colorless, odorless, extremely corrosive, oily liquid. It is sometimes called oil of vitriol. Concentrated Sulfuric Acid

When heated, the pure 100% acid loses sulfur trioxide gas, SO3
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. Other important compounds include sulfur dioxide, used as a bleaching agent, disinfectant, and refrigerant; sodium bisulfite, used in paper manufacture; carbon disulfide, an important organic solvent; hydrogen sulfide, sulfur trioxide, and thionyl chloride, used as reagents in chemistry; Epsom saltsEpsom salts,
common name for magnesium sulfate heptahydrate, MgSO4·7H2O, a water-soluble bitter-tasting compound that occurs as white or colorless needle-shaped crystals.
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 (magnesium sulfate), used as a laxative, bath additive, exfoliant, and magnesium supplement in plant nutrition; the numerous other sulfatesulfate,
chemical compound containing the sulfate (SO4) radical. Sulfates are salts or esters of sulfuric acid, H2SO4, formed by replacing one or both of the hydrogens with a metal (e.g., sodium) or a radical (e.g., ammonium or ethyl).
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 compounds; and sulfa drugssulfa drug,
any of a class of synthetic chemical substances derived from sulfanilamide, or para-aminobenzenesulfonamide. Sulfa drugs are used to treat bacterial infections, although they have largely been replaced for this purpose by antibiotics; some are also used in the
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S, a chemical element in Group VI of Mendeleev’s periodic system. Atomic number, 16; atomic weight, 32.06. Sulfur occurs in nature as a mixture of the four stable isotopes 32S (95.02 percent), 33S (0.75 percent), 34S (4.21 percent), and 36S (0.02 percent). The artificial radioisotopes 31S (half-life, 2.4 sec), 35S (half-life, 87.1 days), and 37S (half-life, 5.04 min) have also been obtained.

Historical survey. Sulfur, both in the native state and in compound form, has been known since antiquity. It is mentioned in the Bible and in the poems of Homer. It was a component of the smoke used in religious rites, and the odor of burning sulfur was thought to ward off evil spirits. Sulfur has long been a necessary component of the incendiary mixtures used in war, for example, Greek Fire (tenth century A.D). It came into use in China around the eighth century as a material for fireworks. Skin diseases have long been treated with sulfur and sulfur compounds. Arabic alchemy was responsible for the hypothesis that sulfur (“spirit” of combustibility) and mercury (“spirit” of metallicity) were components of all metals. Sulfur was established as an element by A. L. Lavoisier, who included it in his list of nonmetal-lic simple substances in 1789. In 1822, E. Mitscherlich discovered sulfur’s allotropy.

Distribution in nature. Sulfur is a very common chemical element (clarke, 4.7 × 10-2). It is encountered in the free state (native sulfur) and in the form of compounds, for example, sulfides, polysulfides, and sulfates. Seawater contains sulfates of sodium, magnesium, and calcium. Endogenic processes have produced more than 200 known minerals of sulfur. More than 150 sulfur minerals, mostly sulfates, are formed in the biosphere; here, processes wherein sulfides are oxidized to sulfates, which in turn are reduced to secondary H2 S and sulfides, are common. These reactions occur with the participation of microorganisms. Many processes of the biosphere lead to the concentration of sulfur; thus, sulfur is accumulated in humus, coal, and petroleum, as well as in the sea (8.9 x 10-2 percent) and groundwater and in lakes and salt marshes. The sulfur content in clays and shales is six times greater than in the earth’s crust as a whole; it is 200 times greater in gypsum and tens of times greater in subsurface sulfate waters. The sulfur cycle in the biosphere involves the deposit of sulfur on land with atmospheric precipitation and the subsequent return of the element to the sea with runoff. In the earth’s geological past, sulfur was supplied mainly by products of volcanic eruptions containing SO2 and H2 S. Man’s industrial activities have accelerated the migration of sulfur, and the oxidation of sulfides has been intensified.

Physical and chemical properties. Sulfur is a crystalline substance that is stable in the form of two allotropie modifications. Orthorhombic α-S has a lemon-yellow color, a density of 2.07 g/cm3, and a melting point of 112.8°C; it is stable below 95.6°C. Monoclinic β-S has a honey-yellow color, a density of 1.96 g/cm3, and a melting point of 119.3°C; it is stable between 95.6°C and the melting point. Both these forms are composed of cyclic S8 molecules for which the S—S bond energy is 225.7 kilojoules (kJ) per mole.

Upon melting, sulfur is converted into a mobile yellow liquid, which turns brown above 160°C and becomes a viscous, dark brown mass at about 190°C. The viscosity decreases above 190°C, and at 300°C sulfur again becomes a flowing liquid. This change in viscosity reflects the change in molecular structure: at 160°C, the S8 rings begin to rupture, becoming open chains; further heating above 190°C decreases the average length of these chains.

When molten sulfur heated to 250°–300°C is poured in a thin stream into cold water, a yellow-brown elastic mass (plastic sulfur) is obtained. The mass is only partially soluble in carbon disulfide, and an unconsolidated powder remains as a residue. The modification that is soluble in CS2 is called λ-S, and the insoluble, μ -S. At room temperature, both these modifications are converted to the stable, brittle α-S. The boiling point of sulfur is 444.6°C, one of the standard points on the international temperature scale. Sulfur vapors at the boiling point include, in addition to S8 molecules, S6, S4, and S2 molecules. Upon further heating, the larger molecules decompose, and at 900°C only S2 molecules remain. At approximately 1500°C, these molecules dissociate to a significant extent into atoms. Upon freezing strongly heated sulfur vapors by liquid nitrogen, a purple modification composed of S2 molecules is obtained, which is stable below – 80°C.

Sulfur is a poor conductor of heat and electricity. It is practically insoluble in water but is readily soluble in anhydrous ammonia, carbon disulfide, and a number of organic solvents, including phenol, benzene, and dichloroethane.

The electronic configuration of the outer electrons of the sulfur atom is 3s23p4. In compounds, sulfur has oxidation states of-2, +4, and +6.

Sulfur is chemically reactive, especially upon heating, and combines with almost all the elements, with the exception of N2, I2, Au, Pt, and the inert gases. At temperatures above 300° C, sulfur reacts with O2 in the air to form sulfur dioxide (SO2) and sulfur trioxide (SO3), from which, respectively, sulfurous and sulfuric acids and sulfite and sulfate salts are obtained. Even in the cold, sulfur combines vigorously with F2; upon heating, it reacts with Cl2. With bromine, sulfur forms only S2 Br2; the iodides of sulfur are unstable. Upon heating to 150°–200°C, the reversible reaction of sulfur with H2 begins, resulting in the production of hydrogen sulfide. Sulfur also forms compounds with hydrogen in which there is more than one atom of sulfur; the general formula for these compounds, known as sulfanes, is H2 Sx. Numerous organosulfur compounds are known.

Upon heating, sulfur reacts with metals, forming the corresponding sulfides and polysulfides. At 800°–900°C, sulfur vapor reacts with carbon, forming carbon disulfide (CS2). Compounds of sulfur with nitrogen (N4 S4 and N2 S5) can be obtained only through indirect methods.

Production. Elemental sulfur is produced from native sulfur, as well as by the oxidation of hydrogen sulfide and reduction of sulfur dioxide. Sources of hydrogen sulfide for the production process include natural gas, coke-oven gas, and gas generated in the cracking of petroleum. Of the many methods developed for treating hydrogen sulfide, two have the greatest importance. In the first, H2 S is extracted from gases by a solution of sodium monohydrogen thioarsenate:

Na2 HAsS2 O2 + H2 S = Na2 HAsS3 O + H2 O

Free sulfur is then precipitated by passing air through the solution:

NaHAsS3 O + ½O2 = Na2 HAsS2 O2 + S

In the second method, hydrogen sulfide is separated from gases in concentrated form. Then, most of it is oxidized by atmospheric oxygen to sulfur and, partially, SO2. Upon cooling, the H2 S and gases that have been formed (SO2, N2, CO2) enter two consecutively arranged converters, where in the presence of a catalyst (activated bauxite or specially prepared alumina gel), the following reaction takes place:

2H2 S + SO2 = 3S + 2H2 O

The basis for the production of sulfur from sulfur dioxide is the reduction of sulfur by coal or natural hydrocarbon gases. This production process is sometimes combined with the treatment of pyrite ores.

In 1972, world production (exclusive of socialist countries) of elemental sulfur totaled 32.0 million tons, with most obtained from native ores. During the 1970’s, methods of producing sulfur from H2 S have acquired greater importance in connection with the discovery of large deposits of natural gas containing hydrogen sulfide.

Types. Sulfur melted directly from sulfur ores is called lump sulfur, while that obtained from H2 S and SO2 is called gaseous lump sulfur. Lump sulfur purified by distillation is called refined sulfur. Sulfur condensed from sulfur vapors at a temperature above the melting point and then poured into molds is called roll sulfur. Sulfur formed as a fine powder in the condensation of sulfur at a temperature below the melting point on the walls of a condensing chamber is known as flowers of sulfur. Sulfur in a very highly dispersed state is called colloidal sulfur.

Use. Sulfur is used primarily in the production of sulfuric acid. In the paper industry it is used in the production of sulfite pulp, and in agriculture, for combating plant diseases, especially those attacking grapevines and cotton plants. Sulfur finds use in the rubber industry as a vulcanizing agent. It is also used in the production of dyes, luminescent substances, black (hunting) gunpowder, and matches.


Sulfur is used in medicine because of its capacity to react with organic matter of organisms to form sulfides and penta-thionic acid. These substances must be present before certain keratolytic, antimicrobial, and antiparasitic processes can occur. Sulfur is a component of Wilkinson ointment and other preparations used in treating scabies. Purified and precipitated sulfur is used in ointments and dusting powders for treating certain skin diseases (seborrhea, psoriasis), in a powder used against helminthic invasions (enterobiasis), and in solutions for the fever therapy of progressive paralysis.

Sulfur in organisms. Sulfur in the form of organic and inorganic compounds is always present in all living organisms and is an important biogenic element. The average content of sulfur, as a percentage of dry weight, is 1.2 percent in marine plants, 0.3 percent in terrestrial plants, 0.5–2 percent in marine animals, and 0.5 percent in terrestrial animals. The biological role of sulfur derives from its inclusion in compounds that are widely distributed in the animal world. These compounds include amino acids (methionine, cysteine) and, hence, proteins and peptides, coenzymes (coenzyme A, lipoic acid), vitamins (biotin, thiamine), and glutathione. The mercapto groups (—SH) of cysteine residues play an important role in the structure and catalytic activity of many enzymes. By forming disulfide bonds (—S—S—) within and between individual polypeptide chains, mercapto groups help support the structure of protein molecules. In animals, sulfur is also found in the form of organic sulfates and such sulfonic acids as chondroitinsulfuric acid (in cartilage and bone), taurocholic acid (in bile), heparin, and taurine. In certain proteins containing iron, for example, ferredoxins, sulfur is found in the form of an acid-labile sulfide. Sulfur is capable of forming energy-rich bonds in high-energy compounds.

Inorganic compounds of sulfur in higher animals occur in small quantities, primarily as sulfates (in blood and urine), but also as thiocyanates (in saliva, gastric juice, milk, urine). Marine organisms have higher contents of inorganic sulfur compounds than freshwater and terrestrial organisms. In plants and many microorganisms, sulfate (SO42-), along with phosphate and nitrate, serves as a vital source of mineral nutrition. Before incorporation into organic compounds, sulfur undergoes changes in valence and is then converted into the organic form in its lowest oxidation state; in this way, sulfur participates extensively in the oxidation-reduction reactions in cells. In cells, sulfates, reacting with adenosinetriphosphate (ATP), are converted to the active form, adenylyl sulfate:

The enzyme that catalyzes this reaction is sulfurylase

(ATP:sulfate—adenylyltransferase), which is widely distributed in nature. In this activated form, the sulfonyl group undergoes further conversion, being transferred to another acceptor or reduced.

Animals assimilate sulfur as components of organic compounds. Autotrophic organisms obtain all the sulfur found in their cells from inorganic compounds, mainly in the form of sulfates. Higher plants, algae, fungi, and bacteria are able to autotrophically assimilate sulfur. (A special protein has been isolated from bacteria cultures that transfers the sulfate through the cell membrane from the medium to the cell.) Desulfurizing bacteria and sulfur bacteria play an important role in the sulfur cycle in nature. Many of the sulfur deposits currently being worked are of biogenic origin. Sulfur is a component of antibiotics (penicillins, cephalosporins), and sulfur compounds are used as substances that protect plants against ionizing radiation.



Spravochnik sernokislotchika, 2nd ed. Edited by K. M. Malin. Moscow, 1971.
Prirodnaiasera. Edited by M. A. Menkovskii. Moscow, 1972.
Nekrasov, B. V. Osnovy obshchei khimii, 3rd ed., vol. 1. Moscow, 1973.
Remy, H. Kurs neorganicheskoi khimii, vol. I. Moscow, 1972. (Translated from German.)
Young, L., and G. Maw. Melabolizm soedinenii sery. Moscow, 1961. (Translated from English.)
Gorizonty biokhimii. Moscow, 1964. (Translated from English.)
Biokhimiia rastenii. Moscow, 1968. Chapter 19. (Translated from English.)
Torchinskii, Iu. M. Sul’fgidril’nye i disul’fidnye gruppy belkov. Moscow, 1971.
Dagley, S., and D. Nicholson. Metabolicheskie puti. Moscow, 1973. (Translated from English.)


A nonmetallic element in group 16, symbol S, atomic number 16, atomic weight 32.06, existing in a crystalline or amorphous form and in four stable isotopes; used as a chemical intermediate and fungicide, and in rubber vulcanization.
A yellow orthorhombic mineral occurring in crystals, masses, or layers, and existing in several allotropic forms; the native form of the element.


(US), sulfur
a. an allotropic nonmetallic element, occurring free in volcanic regions and in combined state in gypsum, pyrite, and galena. The stable yellow rhombic form converts on heating to monoclinic needles. It is used in the production of sulphuric acid, in the vulcanization of rubber, and in fungicides. Symbol: S; atomic no.: 16; atomic wt.: 32.066; valency: 2, 4, or 6; relative density: 2.07 (rhombic), 1.957 (monoclinic); melting pt.: 115.22°C (rhombic), 119.0°C (monoclinic); boiling pt.: 444.674°C
b. (as modifier): sulphur springs