Brimstone (sulfur) - definition of Brimstone (sulfur) by The Free Dictionary Brimstone (sulfur) - definition of Brimstone (sulfur) by The Free Dictionary also sul·phur (sŭl′fər) n. 1. Symbol S A pale yellow nonmetallic element occurring widely in nature in several free, allotropic and crystal forms and combined in numerous sulfates and sulfides. It is used in black gunpowder, rubber vulcanization, the manufacture of insecticides and pharmaceuticals, and in the preparation of sulfur compounds such as hydrogen sulfide and sulfuric acid. Atomic number 16; atomic weight 32.066; melting point 115.21°C; boiling point 444.61°C; specific gravity at 20°C (rhombic) 2.07, (monoclinic) 2.00; valence 2, 4, 6. See Periodic Table . 2. Any of various butterflies of the subfamily Coliadinae of the family Pieridae, having yellow or orange wings often marked with black. tr.v.

sul·fured, sul·fur·ing, sul·furs also sul·phured or sul·phur·ing or sul·phurs To treat with sulfur or a compound of sulfur. [Middle English, from Anglo-Norman sulfre, from Latin sulfur.] sulfur (ˈsʌl fər) n. 1. Also, esp. Brit., sulphur. a nonmetallic element, ordinarily a flammable yellow solid, of widespread occurrence in combined form, as in sulfide and sulfate compounds and cellular protein: used esp. in making gunpowder and matches, in medicine, and in vulcanizing rubber.Symbol: S; at. wt.: 32.064; at. no.: 16; sp. gr.: 2.07 at 20° C. [1300–50; Middle English < Latin sulp(h)ur, sulfur brimstone, sulfur] sul·fur also sul·phur (sŭl′fər) Symbol S A pale-yellow, brittle nonmetallic element that occurs widely in nature, especially in volcanic deposits, many common minerals, natural gas, and petroleum.

It is used to make gunpowder and fertilizer, to vulcanize rubber, and to produce sulfuric acid. Atomic number 16. See Periodic Table . ThesaurusAntonymsRelated WordsSynonymsLegend: Brimstone | Article about brimstone by The Free Dictionary Brimstone | Article about brimstone by The Free Dictionary see sulfur sulfur or sulphur , nonmetallic chemical element; symbol S; at. no. 16; interval in which at. wt. ranges 32.059–32.076; m.p. 112.8&degC; (rhombic), 119.0&degC; (monoclinic), about 120&degC; (amorphous); b.p. 444.674&degC;; sp. gr. at 20&degC;, 2.07 (rhombic), 1. ..... Click the link for more information. .

Brimstone (pop culture) Raining fire and panic upon the scurrying populace who woefully find themselves in his lumbering path, the 50-foot terror called Brimstone is the DC Comics-equivalent of Godzilla—a creature that exists solely to destroy. Brimstone blazed into the DC Universe in Legends #1 (1986), the miniseries in which the evil despot Darkseid of Apokolips attempts to eliminate earthly opposition by eroding the public's confidence in its superheroes. Darkseid's obsequious envoy, the sadistic Desaad, sabotaged a S.T.A.R. Labs nuclear generator to explode, as a result creating Brimstone, a 60,000-pound behemoth possessing some degree of sentience. Programmed with the conviction that it is a fallen angel, Brimstone cuts a swath across the DC Universe with its fiery sword and generation of intense heat. The Justice League of America, Firestorm the Nuclear Man, and the Suicide Squad have narrowly avoided meltdowns in their conflicts with Brimstone. The creature has on several occasions been eliminated, but is capable of regeneration—and repeated devastation.

Brimstone shambled onto television in the “Initiation” episode of the Cartoon Network's Justice League Unlimited (2004–present). In its TV incarnation, the creature was the out-of-control creation of Chinese inventor Chong-Mi, with Justice League recruits including Green Arrow and Captain Atom taking down the fiery beast. Brimstone (Gonopteryx rhamni), a butterfly of the family Pieridae. The wings are angular, with a span of up to 6 cm. The males are lemon yellow, and the females are greenish white. In the middle of each wing there is a small orange-yellow spot. The brimstone is found throughout Europe and the temperate zone of Asia. Soon after the butterflies emerge from the pupae at the end of summer, they go into hibernation. Flight is resumed in early spring. The caterpillar, which is green with a white stripe along its sides, feeds on buckthorn leaves.

brimstone Brimstone (sulfur) | Article about Brimstone (sulfur) by The Free Dictionary Brimstone (sulfur) | Article about Brimstone (sulfur) by The Free Dictionary or sulphur (sŭl`fər), nonmetallic chemical element; symbol S; at. no. 16; interval in which at. wt. ranges 32.059–32.076; m.p. 112.8&degC; (rhombic), 119.0&degC; (monoclinic), about 120&degC; (amorphous); b.p. 444.674&degC;; sp. gr. at 20&degC;, 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 table periodic 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 ..... Click the link for more information. . It exhibits allotropy allotropy [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. ..... Click the link for more information. . 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&degC;; most sulfur is in this form.

The monoclinic, or prismatic, form has long, needlelike, nearly transparent crystals; it is stable between 95.5&degC; 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 ( sulfides sulfide, 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. ..... Click the link for more information.

) 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 process Frasch 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. ..... Click the link for more information. , 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).

Uses Elemental sulfur is used in black gunpowder gunpowder, 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. ..... Click the link for more information. , matches, and fireworks; in the vulcanization vulcanization , 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. ..... Click the link for more information. 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 acid sulfuric 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 ..... Click the link for more information. .

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 salts Epsom salts, common name for magnesium sulfate heptahydrate, MgSO4·7H2O, a water-soluble bitter-tasting compound that occurs as white or colorless needle-shaped crystals. ..... Click the link for more information. (magnesium sulfate), used as a laxative, bath additive, exfoliant, and magnesium supplement in plant nutrition; the numerous other sulfate sulfate, 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). ..... Click the link for more information.

compounds; and sulfa drugs sulfa 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 ..... Click the link for more information. . Sulfur 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. I. K. M ALINA 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. L. I. B [′səl·fər] (chemistry) 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.

(mineralogy) A yellow orthorhombic mineral occurring in crystals, masses, or layers, and existing in several allotropic forms; the native form of the element. sulphur (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 Brimstone | Define Brimstone at Dictionary.com brimstone before 1150; Middle English brinston, etc., late Old English brynstān. See burn 1, stone Related forms Examples from the Web for brimstone Expand Contemporary Examples The problem: Pawlenty comes off like a Good Samaritan at a time when the religious right wants fire and brimstone. Netanyahu Swims Against Iranian Diplomatic Current Ali Gharib October 1, 2013 Creators David Benioff and D.B. Weiss brought the fire and brimstone this season, and for this, deserve the shiny hardware. Jim Spurling, Fisherman Albert Walter Tolman The recipe was brimstone and treacle, but the brimstone predominated, and was the more operative ingredient. Flowers of Freethought George W. Foote At last he was stemming the strong tide-rip off brimstone Point. Jim Spurling, Fisherman Albert Walter Tolman British Dictionary definitions for brimstone Expand an obsolete name for sulphur 2.

a common yellow butterfly, Gonepteryx rhamni, of N temperate regions of the Old World: family Pieridae 3. (archaic) a scolding nagging woman; virago Word Origin Old English brynstān; related to Old Norse brennistein; see burn1, stone Collins English Dictionary - Complete & Unabridged 2012 Digital Edition © William Collins Sons & Co. Ltd. 1979, 1986 © HarperCollins Publishers 1998, 2000, 2003, 2005, 2006, 2007, 2009, 2012 Word Origin and History for brimstone Expand n. Old English brynstan, from brin- stem of brinnen "to burn" (see burn (v.)) + stan (see stone (n.)). In Middle English the first element also recorded as brem-, brom-, brum-, bren-, brin-, bron-, brun-, bern-, born-, burn-, burned-, and burnt-. Formerly "the mineral sulfur," now restricted to biblical usage.

The Lord reynede vpon Sodom and Gomor brenstoon and fier. [Wycliff's rendition (1382) of Gen. xix:24] The Old Norse cognate compound brennusteinn meant "amber," as does German Bernstein. Online Etymology Dictionary, © 2010 Douglas Harper brimstone in the Bible Expand an inflammable mineral substance found in quantities on the shores of the Dead Sea. The cities of the plain were destroyed by a rain of fire and brimstone (Gen. 19:24, 25). In Isa. 34:9 allusion is made to the destruction of these cities. This word figuratively denotes destruction or punishment (Job 18:15; Isa. 30:33; 34:9; Ps. 11:6; Ezek. 38:22). It is used to express the idea of excruciating torment in Rev. 14:10; 19:20; 20:10. Easton's 1897 Bible Dictionary What does sulfur mean?

definition, meaning and pronunciation (Free English Language Dictionary) atomic number 16 ; sulfur; sulphur ; S Hypernyms ("sulfur" is a kind of...): chemical element ; element (any of the more than 100 known substances (of which 92 occur naturally) that cannot be separated into simpler substances and that singly or in combination constitute all matter) Hyponyms (each of the following is a kind of "sulfur"): brimstone ; native sulfur ; native sulphur (an old name for sulfur) Holonyms ("sulfur" is a substance of...): sulfide ; sulphide (a compound of sulphur and some other element that is more electropositive) oil of vitriol ; sulfuric acid ; sulphuric acid ; vitriol ((H2SO4) a highly corrosive acid made from sulfur dioxide; widely used in the chemical industry) • SULFUR (verb) sulfur (S) | chemical element | Britannica.com chemical element Alternative Titles: brimstone, S, sulphur Related Topics matter Sulfur (S), also spelled sulphur, nonmetallic chemical element belonging to the oxygen group (Group 16 [VIa] of

 the periodic table), one of the most reactive of the elements.

Pure sulfur is a tasteless, odourless, brittle solid that is pale yellow in colour, a poor conductor of electricity, and insoluble in water. It reacts with all metals except gold and platinum , forming sulfides ; it also forms compounds with several nonmetallic elements. Millions of tons of sulfur are produced each year, mostly for the manufacture of sulfuric acid, which is widely used in industry. Sulfur crystals from Sicily (greatly enlarged) Courtesy of the Illinois State Museum; photograph, John H. Gerard/EB Inc. In cosmic abundance, sulfur ranks ninth among the elements, accounting for only one atom of every 20,000–30,000. Sulfur occurs in the uncombined state as well as in combination with other elements in rocks and minerals that are widely distributed, although it is classified among the minor constituents of Earth ’s crust, in which its proportion is estimated to be between 0.03 and 0.06 percent. On the basis of the finding that certain meteorites contain about 12 percent sulfur, it has been suggested that deeper layers of Earth contain a much larger proportion.

Seawater contains about 0.09 percent sulfur in the form of sulfate. In underground deposits of very pure sulfur that are present in domelike geologic structures, the sulfur is believed to have been formed by the action of bacteria upon the mineral anhydrite , in which sulfur is combined with oxygen and calcium . Deposits of sulfur in volcanic regions probably originated from gaseous hydrogen sulfide generated below the surface of Earth and transformed into sulfur by reaction with the oxygen in the air. Boiling pots of molten sulfur on the slope of the Nikko volcano near the Mariana Islands. Major funding for this expedition was provided by NOAA Ocean Exploration Program and NOAA Vents Program; video clips edited by Bill Chadwick, Oregon State University/NOAA One arm of the Jason remotely operated vehicle breaking through the thin crust on a deposit of … Major funding for this expedition was provided by NOAA Ocean Exploration Program and NOAA Vents Program; video clips edited by Bill Chadwick, Oregon State University/NOAA History boron (B) The history of sulfur is part of antiquity.

The name itself probably found its way into Latin from the language of the Oscans, an ancient people who inhabited the region including Vesuvius, where sulfur deposits are widespread. Prehistoric humans used sulfur as a pigment for cave painting; one of the first recorded instances of the art of medication is in the use of sulfur as a tonic. organosulfur compound: The sulfur atom The combustion of sulfur had a role in Egyptian religious ceremonials as early as 4,000 years ago. “Fire and brimstone” references in the Bible are related to sulfur, suggesting that “hell’s fires” are fuelled by sulfur. The beginnings of practical and industrial uses of sulfur are credited to the Egyptians, who used sulfur dioxide for bleaching cotton as early as 1600 bce. Greek mythology includes sulfur chemistry: Homer tells of Odysseus’ use of sulfur dioxide to fumigate a chamber in which he had slain his wife’s suitors. The use of sulfur in explosives and fire displays dates to about 500 bc in China, and flame-producing agents used in warfare (Greek fire) were prepared with sulfur in the Middle Ages.

Pliny the Elder in 50 ce reported a number of individual uses of sulfur and ironically was himself killed, in all probability by sulfur fumes, at the time of the great Vesuvius eruption (79 ce). Sulfur was regarded by the alchemists as the principle of combustibility. Lavoisier recognized it as an element in 1777, although it was considered by some to be a compound of hydrogen and oxygen; its elemental nature was established by the French chemists Joseph Gay-Lussac and Louis Thenard. Natural occurrence and distribution Periodic Table of the Elements Many important metal ores are compounds of sulfur, either sulfides or sulfates. Some important examples are galena (lead sulfide, PbS), blende (zinc sulfide, ZnS), pyrite (iron disulfide, FeS2), chalcopyrite (copper iron sulfide, CuFeS2), gypsum (calcium sulfate dihydrate, CaSO4∙2H2O) and barite (barium sulfate, BaSO4).

The sulfide ores are valued chiefly for their metal content, although a process developed in the 18th century for making sulfuric acid utilized sulfur dioxide obtained by burning pyrite. Coal, petroleum , and natural gas contain sulfur compounds. Allotropy Ringling Bros. Folds Its Tent In sulfur, allotropy arises from two sources: (1) the different modes of bonding atoms into a single molecule and (2) packing of polyatomic sulfur molecules into different crystalline and amorphous forms. Some 30 allotropic forms of sulfur have been reported, but some of these probably represent mixtures. Only eight of the 30 seem to be unique; five contain rings of sulfur atoms and the others contain chains. In the rhombohedral allotrope, designated ρ-sulfur, the molecules are composed of rings of six sulfur atoms. This form is prepared by treating sodium thiosulfate with cold, concentrated hydrochloric acid , extracting the residue with toluene , and evaporating the solution to give hexagonal crystals. ρ-sulfur is unstable, eventually reverting to orthorhombic sulfur (α-sulfur).

A second general allotropic class of sulfur is that of the eight-membered ring molecules, three crystalline forms of which have been well characterized. One is the orthorhombic (often improperly called rhombic) form, α-sulfur. It is stable at temperatures below 96 °C. Another of the crystalline S8 ring allotropes is the monoclinic or β-form, in which two of the axes of the crystal are perpendicular, but the third forms an oblique angle with the first two. There are still some uncertainties concerning its structure; this modification is stable from 96 °C to the melting point , 118.9 °C. A second monoclinic cyclooctasulfur allotrope is the γ-form, unstable at all temperatures, quickly transforming to α-sulfur. An orthorhombic modification, S12 ring molecules, and still another unstable S10 ring allotrope are reported. The latter reverts to polymeric sulfur and S8. At temperatures above 96 °C, the α-allotrope changes into the β-allotrope.

If enough time is allowed for this transition to occur completely, further heating causes melting to occur at 118.9 °C; but if the α-form is heated so rapidly that the transformation to β-form does not have time to occur, the α-form melts at 112.8 °C. Connect with Britannica Facebook Twitter YouTube Instagram Pinterest Just above its melting point, sulfur is a yellow, transparent, mobile liquid. Upon further heating, the viscosity of the liquid decreases gradually to a minimum at about 157 °C, but then rapidly increases, reaching a maximum value at about 187 °C; between this temperature and the boiling point of 444.6 °C, the viscosity decreases. The colour also changes, deepening from yellow through dark red, and, finally, to black at about 250 °C. The variations in both colour and viscosity are considered to result from changes in the molecular structure.

A decrease in viscosity as temperature increases is typical of liquids, but the increase in the viscosity of sulfur above 157 °C probably is caused by rupturing of the eight-membered rings of sulfur atoms to form reactive S8 units that join together in long chains containing many thousands of atoms. The liquid then assumes the high viscosity characteristic of such structures. At a sufficiently high temperature, all of the cyclic molecules are broken, and the length of the chains reaches a maximum. Beyond that temperature, the chains break down into small fragments. Upon vaporization, cyclic molecules (S8 and S6) are formed again; at about 900 °C, S2 is the predominant form; finally, monatomic sulfur is formed at temperatures above 1,800 °C. Commercial production Elemental sulfur is found in volcanic regions as a deposit formed by the emission of hydrogen sulfide, followed by aerial oxidation to the element. Underground deposits of sulfur associated with salt domes in limestone rock provide a substantial portion of the world’s supply of the element. These domes are located in the Louisiana swamplands of the United States and offshore in the Gulf of Mexico .

Where deposits of sulfur are located in salt domes, as they are along the coast of the Gulf of Mexico, the element was recovered by the Frasch process , named after German-born U.S. chemist Herman Frasch . Ordinary underground mining procedures were inapplicable since highly poisonous hydrogen sulfide gas accompanies the element in the domes. Beginning in 1894, the Frasch process, which takes advantage of the low melting point of sulfur (112 °C), made sulfur of a high purity (up to 99.9 percent pure) available in large quantities and helped establish sulfur as an important basic chemical commodity. Wells were drilled from 60 to 600 m (200 to 2,000 feet) into the sulfur formation and then lined with a 15-cm (6-inch) pipe in which an air pipe and a water pipe of smaller diameter were concentrically placed. Superheated water, injected into the circular space between the three- and six-inch pipes, penetrated the cap rock through holes on the bottom of the pipe. As the sulfur melted, it settled to the bottom of the deposit.

From there it was pumped to the surface by applying air pressure through the central pipe. Several such wells operated under the ocean floor in the Gulf of Mexico. The sulfur was collected in reservoirs, or sumps, and from there transferred to vats or bins to solidify for storage and stockpiling. Vats contained as much as 300,000 tons of sulfur. Frasch-process sulfur produced at the Gulf Coast salt domes constituted the major source of U.S. sulfur production and dominated the world market until approximately 1970. Thereafter, non-Frasch sources such as the purification of sour (high sulfur-content) petroleum, the refining of natural gas, and improved methods for obtaining sulfur from metal sulfides gained a greater share of the market. The Frasch process is still used today in Poland and Russia. Britannica Lists & Quizzes Editor Picks: Exploring 10 Types of Basketball Movies About 9,000,000 tons of sulfur are recovered in the United States each year from natural gas, petroleum refinery gases, pyrites, and smelter gases from the processing of copper , zinc , and lead ores.

In most cases sulfur is separated from other gases as hydrogen sulfide and then converted to elemental sulfur by the Claus process, which involves the partial burning of hydrogen sulfide to sulfur dioxide, with subsequent reaction between the two to yield sulfur. Another important source is the sulfur dioxide emitted into the atmosphere by coal-fired steam power plants. In the early 1970s techniques to collect this sulfur dioxide and convert it into usable sulfur were developed. Trending Topics Opium Wars A few of the non-Frasch processes for sulfur production may be mentioned. Sulfur-bearing rock is piled into mounds. Shafts are bored vertically and fires set at the top of the shafts. The burning sulfur provides sufficient heat to melt the elemental sulfur in the rock layers below, and it flows out at the bottom of the pile. This is an old process, still used to some extent in Sicily . The product is of low purity and must be refined by distillation. The air pollution in the area of the process is so great that its operation is limited to certain times of the year when prevailing winds will carry the fumes away from populated areas. Rock bearing sulfur is treated with superheated water in retorts, melting the sulfur, which flows out.

This process is a modification of the Frasch method. Sulfates (such as gypsum or barite) may be treated with carbon at high temperatures, forming the metal sulfides CaS or BaS (the Chance-Claus process). The metal sulfides can be treated with acid, generating hydrogen sulfide, which in turn can be burned to give elemental sulfur. Tremendous tonnages of sulfur are available from smelter operations and from power production by combustion of fossil and sour petroleum fuels, some of which contain as much as 4 percent sulfur. Thus, generation of electrical power and heat represent a major source of atmospheric pollution by sulfur dioxide. Unfortunately, recovery and purification of sulfur dioxide from stack gases are expensive operations. Wherever such metals as lead, zinc, copper, cadmium , or nickel (among others) are processed, much of the sulfuric acid needed in the metallurgical operations may be obtained on the site by converting sulfur dioxide, produced by roasting the ores, to sulfur trioxide, SO3, and thence to sulfuric acid. Sulfur available in bulk from commercial production usually is more than 99 percent pure, and some grades contain 99.9 percent sulfur.

For research purposes, the proportion of impurities has been reduced to as little as one part in 10,000,000 by the application of procedures such as zone melting, column chromatography, electrolysis, or fractional distillation. China, Canada, Germany and Japan led the world in sulfur production in the early 21st century. Uses of sulfur Sulfur is so widely used in industrial processes that its consumption often is regarded as a reliable indicator of industrial activity and the state of the national economy. Approximately six-sevenths of all the sulfur produced is converted into sulfuric acid , for which the largest single use is in the manufacture of fertilizers (phosphates and ammonium sulfate). Other important uses include the production of pigments, detergents, fibres, petroleum products, sheet metal, explosives, and storage batteries; hundreds of other applications are known. Sulfur not converted to sulfuric acid is used in making paper , insecticides, fungicides, dyestuffs, and numerous other products.

Compounds Sulfur forms compounds in oxidation states −2 (sulfide, S2−), +4 (sulfite, SO32−), and +6 (sulfate, SO42−). It combines with nearly all elements. An unusual feature of some sulfur compounds results from the fact that sulfur is second only to carbon in exhibiting catenation—i.e., the bonding of an atom to another identical atom. This allows sulfur atoms to form ring systems and chain structures. The more significant sulfur compounds and compound groups are as follows. One of the most familiar sulfur compounds is hydrogen sulfide , also known as sulfureted hydrogen, or stinkdamp, H2S, the colourless, extremely poisonous gas responsible for the characteristic odour of rotten eggs. It is produced naturally by the decay of organic substances containing sulfur and is often present in vapours from volcanoes and mineral waters. Large amounts of hydrogen sulfide are obtained in the removal of sulfur from petroleum. It was formerly used extensively in chemical laboratories as an analytical reagent. All the metals except gold and platinum combine with sulfur to form inorganic sulfides.

Such sulfides are ionic compounds containing the negatively charged sulfide ion S2−; these compounds may be considered as salts of hydrogen sulfide. Some inorganic sulfides are important ores of such metals as iron, nickel, copper, cobalt , zinc, and lead. Several oxides are formed by sulfur and oxygen; the most important is the heavy, colourless, poisonous gas sulfur dioxide , SO2. It is used primarily as a precursor of sulfur trioxide, SO3, and thence sulfuric acid, H2SO4. It is also utilized as a bleach and an industrial reducing agent. Other noteworthy applications include its use in food preservation and for fruit ripening. (See also sulfur oxide .) Sulfur forms a wide variety of compounds with halogen elements. In combination with chlorine it yields sulfur chlorides such as disulfur dichloride, S2Cl2, a corrosive, golden-yellow liquid used in the manufacture of chemical products. It reacts with ethylene to produce mustard gas, and with unsaturated acids derived from fats it forms oily products that are basic components of lubricants.

With fluorine , sulfur forms sulfur fluorides, the most useful of which is sulfur hexafluoride , SF6, a gas employed as an insulator in various electrical devices. Sulfur also forms oxyhalides, in which the sulfur atom is bonded to both oxygen and halogen atoms. When such compounds are named, the term thionyl is used to designate those containing the SO unit and the term sulfuryl for those with SO2. Thionyl chloride , SOCl2, is a dense, toxic, volatile liquid used in organic chemistry to convert carboxylic acids and alcohols into chlorine-containing compounds. Sulfuryl chloride, SO2Cl2, is a liquid of similar physical properties utilized in the preparation of certain compounds that contain sulfur, chlorine, or both. Sulfur forms some 16 oxygen-bearing acids. Only four or five of them, however, have been prepared in the pure state. These acids, particularly sulfurous acid and sulfuric acid, are of considerable importance to the chemical industry. Sulfurous acid, H2SO3, is produced when sulfur dioxide is added to water.

Its most important salt is sodium sulfite, Na2SO3, a reducing agent employed in the manufacture of paper pulp, in photography, and in the removal of oxygen from boiler feedwater. Sulfuric acid is one of the most valuable of all chemicals. Prepared commercially by the reaction of water with sulfur trioxide, the compound is used in manufacturing fertilizers, pigments, dyes, drugs, explosives, detergents, and inorganic salts and esters. The organic compounds of sulfur constitute a diverse and important subdivision of organic substances. Some examples include the sulfur-containing amino acids (e.g., cysteine , methionine , and taurine), which are key components of hormones, enzymes, and coenzymes. Significant, too, are the synthetic organic sulfur compounds, among them numerous pharmaceuticals (sulfa drugs, dermatological agents), insecticides, solvents, and agents such as those used in preparing rubber and rayon .

Question: Brimstone is an old word for which chemical element?
Answer: Sulphur