Sulphur:

Symbol S, tasteless, odourless, light yellow nonmetallic element. Sulphur is in group 16 (or VIa) of the periodic table. Its atomic number is 16, and its atomic weight is 32.064.

Also called brimstone, sulphur has been known since prehistoric times and is mentioned in the Bible and classical records. Because of its flammability, alchemists regarded sulphur as essential in combustion.

Properties

All forms of sulphur are insoluble in water, but the crystalline forms are soluble in carbon disulphide. When ordinary sulphur melts, it forms a straw-coloured liquid that turns darker with additional heating and then finally boils. When molten sulphur is slowly cooled, its physical properties change in accordance with the temperature, pressure, and method of crust formation. Sulphur thus exists in a variety of forms called allotropic modifications, which consist of the liquids Sl, and Sµ, and several solid varieties, of which the most familiar are rhombic sulphur and monoclinic sulphur (See Crystal). The most stable variety of the element is rhombic sulphur, a yellow, crystalline solid with a density of 2.06 g/cu cm at 20° C (68° F). Rhombic sulphur is slightly soluble in alcohol and ether, moderately soluble in oils and extremely soluble in carbon disulphide. When kept at temperatures above 94.5° C (202.1° F) but below 120° C (248° F) the rhombic form changes into monoclinic sulphur consisting of elongated, transparent, needlelike structures with a density of 1.96 g/cu cm at 20° C (68° F). The temperature at which rhombic and monoclinic sulphur are in equilibrium, 94.5° C (202.1° F), is known as the transition temperature. When ordinary rhombic sulphur is melted at 115.21° C (239.38° F), it forms the mobile, pale yellow liquid Sl, which becomes dark and viscous at 160° C (320° F) to form Sµ. If sulphur is heated almost to its boiling point of 444.6° C (832.3° F) and is then poured rapidly into cold water, it does not have time to crystallize into the rhombic or monoclinic state, but forms a transparent, sticky, elastic substance known as amorphous, or plastic, sulphur, which consists for the most part of supercooled Sµ.

Sulphur has valences of two, four, and six, as evidenced by the compounds ferrous sulphide, FeS; sulphur dioxide, SO2; and barium sulphate, BaSO4, respectively. It combines with hydrogen and the metallic elements in the presence of heat to form sulphides. The most common sulphide is hydrogen sulphide, H2S, a colourless, poisonous gas with the odour of rotten eggs. Sulphur combines also with chlorine in several proportions to produce sulphur monochloride, S2Cl2, and sulphur dichloride, SCl2. When burned in air, sulphur combines with oxygen to form sulphur dioxide, SO2, a heavy, colourless gas with a characteristic, suffocating odour. In moist air it is slowly oxidized to sulphuric acid and is a basic constituent of other acids, such as thiosulphuric acid H2S2O3, and sulphurous acid H2SO3. The latter has two replaceable hydrogens and forms two series of salts: normal and acid sulphites. When in solution, the acid sulphites, or bisulphites, of the alkali metals, such as sodium bisulphite, NaHSO 3, are acid in reaction. Solutions of the normal sulphites, such as sodium sulphite, Na2SO3, and potassium sulphite, K2SO3, are slightly alkaline.

Sulphur dioxide is released into the atmosphere in the combustion of fossil fuels, such as gas, petroleum, and coal, and constitutes one of the most troublesome air pollutants. The concentration of sulphur dioxide in air may range from 0.01 to several parts per million, and it may be responsible for the decay of buildings and monuments, for acid rain, and for human discomfort and disability. See Air Pollution.

Occurrence

Sulphur ranks 16th in abundance among the elements in the earth's crust and is found widely distributed in both the free and combined states. In combination it occurs in many important metallic sulphides, such as lead sulphide, or galena, PbS; zinc blende, ZnS; copper pyrite, (Cu,Fe)S2; cinnabar, HgS; stibnite, Sb2S3; and iron pyrite FeS2. It is also combined with other elements in the form of sulphates such as barite, BaSO4; celestite, SrSO4; and gypsum, CaSO42H2O; and it is present in the molecules of many organic substances such as mustard, eggs, hair, proteins, and oil of garlic. In the free state it is found mixed with gypsum and pumice stone in volcanic regions throughout Iceland, Sicily, Mexico, and Japan, often occurring as a sublimate surrounding the volcanic apertures. Vast subterranean deposits are found in the United States in many parts of Louisiana and Texas. Free sulphur may be formed from the weathering of pyrites or may be deposited by hot sulphurous waters in which hydrogen sulphide has been oxidized by the atmosphere. Annual world production of elemental sulphur in the early 1990s amounted to about 52.7 million metric tons.

Extraction

Several methods exist for the extraction of free sulphur from the earth. In Sicily the sulphur-containing rock is placed in large piles on sloping ground and ignited. The liquid sulphur resulting from this heating is allowed to run into a series of wooden moulds, in which it solidifies; in this form it is known as roll sulphur. The roll sulphur may be further purified through distillation, the vapour being passed into a large brick chamber in which it condenses on the walls as a fine powder called flowers of sulphur. In areas where natural sulphur deposits may lie some 275 m (about 900 ft) or more below the surface of the earth, as in Louisiana and Texas, the method most commonly used for extraction is the Frasch process, invented in 1891 by the American chemist Herman Frasch. In this method four concentric pipes, the largest being 20 cm (8 in) in diameter, are driven down into the sulphur-containing deposits. Water, heated under pressure to 170° C (338° F), is forced through the two outer pipes into the deposit, melting the sulphur. When a sufficient quantity of sulphur has been melted, hot air is forced down the inmost pipe to form a froth with the molten sulphur, and the mixture is forced up to the surface through the remaining pipe. The sulphur is run into wooden bins and solidified, yielding a product that is about 99.5 per cent pure. Sulphur is also obtained from pyrites by distillation in iron or fireclay retorts, but it usually contains traces of arsenic when produced in this manner.

Uses

The most important use of sulphur is in the manufacture of sulphur compounds, such as sulphuric acid, sulphites, sulphates, and sulphur dioxide, all mentioned above. Medicinally, it has assumed importance because of its widespread use in sulpha drugs and in many skin ointments. Sulphur is also employed in the production of matches, vulcanized rubber, dyes, and gunpowder. In a finely divided state and, frequently, mixed with lime, sulphur is used as a fungicide on plants. The salt, sodium thiosulphate, Na2S2O35H2O, commonly called hypo, is used in photography for fixing negatives and prints. When combined with various inert mineral fillers, sulphur forms a special cement used to anchor metal objects, such as railings and chains, in stone. Sulphuric acid is one of the most important of all industrial chemicals because it is employed not only in the manufacture of sulphur-containing molecules but also in the manufacture of numerous other materials that do not themselves contain sulphur, such as phosphoric acid.