Sulfur

I

Introduction

Sulfur, symbol S, tasteless, odorless, light yellow nonmetallic element. Sulfur is in group 16 (or VIa) of the periodic table (see Periodic Law). Its atomic number is 16, and its atomic weight is 32.064.

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

II

Properties of Sulfur

All forms of sulfur are insoluble in water, but the crystalline forms are soluble in carbon disulfide. When ordinary sulfur melts, it forms a straw-colored liquid that turns darker with additional heating and then finally boils. When molten sulfur is slowly cooled, its physical properties change in accordance with the temperature, pressure, and method of crust formation. Sulfur thus exists in a variety of forms called allotropes, which consist of the liquids S_λ, and S_µ, and several solid varieties, of which the most familiar are rhombic sulfur and monoclinic sulfur (see Crystal).

The most stable variety of the element is rhombic sulfur, a yellow, crystalline solid with a density of 2.06 g/cm³ at 20°C (68°F). Rhombic sulfur is slightly soluble in alcohol and ether, moderately soluble in oils and extremely soluble in carbon disulfide. When kept at temperatures above 94.5°C (202.1°F) but below 120°C (248°F) the rhombic form changes into monoclinic sulfur consisting of elongated, transparent, needlelike structures with a density of 1.96 g/cm³ at 20°C (68°F). The temperature at which rhombic and monoclinic sulfur are in equilibrium, 94.5°C (202.1°F), is known as the transition temperature. When ordinary rhombic sulfur is melted at 115.21°C (239.38°F), it forms the mobile, pale yellow liquid S_λ, which becomes dark and viscous at 160°C (320°F) to form S_µ.

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If sulfur 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, sulfur, which consists for the most part of supercooled S_µ.

Sulfur has valences of two, four, and six, as evidenced by the compounds ferrous sulfide, FeS; sulfur dioxide, SO₂; and barium sulfate, BaSO₄, respectively. It combines with hydrogen and the metallic elements in the presence of heat to form sulfides. The most common sulfide is hydrogen sulfide, H₂S, a colorless, poisonous gas with the odor of rotten eggs. Sulfur combines also with chlorine in several proportions to produce sulfur monochloride, S₂Cl₂, and sulfur dichloride, SCl₂.

When burned in air, sulfur combines with oxygen to form sulfur dioxide, SO₂, a heavy, colorless gas with a characteristic, suffocating odor. In moist air it is slowly oxidized to sulfuric acid and is a basic constituent of other acids, such as thiosulfuric acid H₂S₂O₃, and sulfurous acid H₂SO₃. The latter has two replaceable hydrogens and forms two series of salts: normal and acid sulfites. When in solution, the acid sulfites, or bisulfites, of the alkali metals, such as sodium bisulfite, NaHSO₃, are acid in reaction. Solutions of the normal sulfites, such as sodium sulfite, Na₂SO₃, and potassium sulfite, K₂SO₃, are slightly alkaline.

Sulfur 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 sulfur 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.

III

Occurrence of Sulfur

Sulfur makes up only a small percentage of Earth’s crust in terms of its abundance among the elements. However, it is found widely distributed in both the free and combined states. In combination it occurs in many important metallic sulfides, such as lead sulfide, 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 sulfates such as barite, BaSO₄; celestite, SrSO₄; and gypsum, CaSO₄ · 2H₂O; 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 in places such as Iceland, Sicily, Mexico, and Japan, often occurring as a sublimate surrounding the volcanic apertures. Vast subterranean deposits have been found in the United States in many parts of Louisiana and Texas, as well as in Colorado, Nevada, Wyoming, and California. Free sulfur may be formed from the weathering of pyrites or may be deposited by hot sulfurous waters in which hydrogen sulfide has been oxidized by the atmosphere.

IV

Extraction of Sulfur

Several methods exist for the extraction of free sulfur from the ground. In one technique the sulfur-containing rock is placed in large piles on sloping ground and ignited. The liquid sulfur resulting from this heating is allowed to run into a series of wooden molds, in which it solidifies; in this form it is known as roll sulfur. The roll sulfur may be further purified through distillation, the vapor being passed into a large brick chamber in which it condenses on the walls as a fine powder called flowers of sulfur.

In areas where natural sulfur deposits may lie some 275 m (about 900 ft) or more below the surface of the earth, a common method used for extraction has been 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 sulfur-containing deposits. Water, heated under pressure to 170°C (338°F), is forced through the two outer pipes into the deposit, melting the sulfur. When a sufficient quantity of sulfur has been melted, hot air is forced down the inmost pipe to form a froth with the molten sulfur, and the mixture is forced up to the surface through the remaining pipe. The sulfur is run into wooden bins and solidified, yielding a product that is about 99.5 percent pure. Sulfur is also obtained from pyrites by distillation in iron or fireclay retorts, but it usually contains traces of arsenic when produced in this manner.

Sulfur can also be extracted from hydrogen sulfide gas present in petroleum and natural gas. The Claus process uses heating and chemical reactions to separate sulfur during the refinement of fossil fuels.

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