SULFUR

AMOS B. SMITH III, UNIVERSITY OF PENNSYLVANIA

To many people, the element sulfur quickly brings to mind unpleasant memories of rotten eggs and unwelcome encounters with skunks. While true that many sulfur-containing compounds, such as mercaptans, thioethers, and disulfides, possess remarkably foul odors, elemental sulfur (S) in pure form is completely odorless, tasteless, and by and large nontoxic.

Referred to in Genesis as brimstone, meaning "a stone that burns," elemental sulfur, which does readily burn in air, has been known since ancient times. The word derives from sulvere and sulphur (Sanskrit and Latin, respectively). Major sulfur deposits are found in the salt domes ubiquitous to the Gulf Coast of the U.S.; sulfur is the 16th most abundant element in nature. Recovery is by melting with superheated water, a commercial process developed by Herman Frasch in the 1890s and still in use today.

Sulfur is also a major by-product of and can be extracted from coal, ores, and minerals such as gypsum, cinnabar, barite, and pyrite (fool's gold). From an economic perspective, most of the world's sulfur production is used to make sulfuric acid (H₂SO₄), approximately 40 million tons every year just in the U.S. Fertilizers and lead-storage automobile batteries consume a large portion of this supply, with smaller amounts used as insecticides, as dyeing agents, in the manufacture of gunpowder, and to vulcanize natural and synthetic rubbers to impart desirable mechanical properties.

Since sulfur is positioned directly under oxygen in the periodic table, one might suspect relatively simple behavior at the atomic level. This is certainly not the case! Indeed, Jerry Donohue, in his authoritative treatise, "The Structures of the Elements," argues, "Of all of the elements, sulfur presents the most confusion and complexity in this respect." Unlike sulfur's near neighbors, such as N, O, Cl, and Br, or P, As, Se, and Te, which in elemental form exist respectively as diatomic molecules or take on polymorphic networks, the four isotopes of sulfur [³²S (95.1%), ³³S (0.74%), ³⁴S (4.2%), and ³⁶S (0.016%)] combine to form a uniquely large number of allotropic forms (compare S_n, where n = 1–12, 18, and infinity). Of these, only eight have been characterized crystallographically. The most common form, stable at room temperature and atmospheric pressure, is orthorhombic sulfur (S₈). Here, eight sulfur atoms bond covalently in crownlike rings. This allotrope, also known as rhombic sulfur, Muthmann's sulfur, and a-S, was among the first substances to be examined crystallographically by William Bragg in 1914.

	CLUMPED Sulfur produces a monoclinic shape when crystallized out of solution. ANDREW LAMBERT PHOTOGRAPHY/SCIENCE PHOTO LIBRARY

The chemistry of sulfur is equally complex; it combines with many elements to yield a bewildering array of both organic and inorganic compounds. Common inorganic compounds include sulfur hydrides, carbon disulfide (CS₂), sulfur selenides, and sulfur halides (SX₆). Oxides of sulfur are particularly important, possessing both beneficial and deleterious properties. Sulfur dioxide (SO₂), for example, finds beneficial use in preserving fruits and vegetables and in the brewing and wine-making industry as both an antioxidant and an antibiotic. Sulfur dioxide and its close relative, sulfur trioxide (SO₃), represent serious hazards, arising in the environment principally by burning of sulfur-rich fuels such as coal and oil or by smelting ores. Released into the atmosphere and combined with water, these pollutants form sulfuric acid and in turn acid rain, a cause of huge economic damage.

In combination with nitrogen, sulfur forms sulfur nitride (S₄N₄), precursor to the sulfur-nitrogen inorganic polymer (SN)_x, which was shown in 1975 by my colleagues at the University of Pennsylvania, Alan G. MacDiarmid and Alan J. Heeger, to be a metal at room temperature with a conductivity similar to iron. This pioneering work eventually led to the fundamental discovery of organic conducting polymers, and in turn to the award of the Nobel Prize in Chemistry to MacDiarmid, Heeger, and Hideki Shirakawa in 2000.

Organosulfur compounds, equally widespread in nature, are essential for all living organisms. Particularly important are the amino acids cysteine and methionine. Cysteine possesses the unique thiol (–SH) functionality, which plays a critical role in the folding and three-dimensional structures of proteins through formation of cross-linking disulfide bonds. Methionine, in the form of S-adenosylmethione, is nature's methylating agent. Other important sulfur-containing biomolecules include keratin, biotin, thiamine, coenzyme A, glutathione, and lipoic acid.

In synthetic chemistry, sulfur plays a central role, especially in its oxidized forms. Dimethylsulfoxide (DMSO), a commonly used organic solvent with invaluable solubility properties, is the key reagent in the widely used Swern oxidation of alcohols. Enantiomerically pure sulfoxides, sulfoximines, and derivatives make superb chiral auxiliaries and reagents for many asymmetric reactions (for example, Davis oxidation and Johnson resolution), while sulfones and dithianes are utilized extensively to form - and -carbon-carbon bonds (for example, Julia olefination and Corey-Seebach umpolung chemistry). In our and others' laboratories, dithianes have found extensive use in mono- and multicomponent fragment unions for the construction of complex natural products possessing important bioregulatory properties.

Amos B. Smith III is the Rhodes-Thompson Professor of Chemistry and a member of the Monell Chemical Senses Center at the University of Pennsylvania.

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