Sulfur Compounds

• Type of physical science: Chemistry
• Field of study: Chemical compounds

Sulfur, the thirteenth-most-abundant element in the Earth’s crust (before carbon and chlorine), forms a wide variety of compounds of importance in industrial chemistry, in inorganic and organic laboratory chemistry, and in the chemistry of living organisms.

Overview

Sulfur is a Group 16 element, which means that it has six electrons in its valence shell.

Thus, in its inorganic compounds, it can either pick up two electrons to form a -2 ion (sulfide, S^2-) or give up electrons, usually in twos, to form various oxides and oxyanions, among them sulfur dioxide (SO₂), sulfur trioxide (SO₃), sulfite (SO₃^2-), thiosulfate (S₂O₃^2-), and sulfate (SO₄^2-) ions, the last frequently in the form of sulfuric acid (H₂SO₄). Except for the sulfides, these inorganic compounds are covalently bonded about the sulfur atom, with the entire oxyanion acting as a building block in an ionic crystal. The organic compounds of sulfur are covalent, with sulfur typically bonded to carbon or hydrogen atoms by single bonds, or to oxygen by double bonds. Examples include mercaptans (R-SH), sulfides (R-S-R), disulfides (R-S-S-R), sulfonic acids (R-SO₃H), and sulfoxides (R-S(=O) -R) (in all cases, R- stands for an organic molecular structure of almost any kind).

Many of these compounds are found in plant and animal chemistry. Sulfur can also enter into polymeric structures, as part of the primary polymer chain or as a cross-linking agent. Finally, sulfur is an integral part of many molecules in physiological chemistry, where it forms a cross-linking bond between protein chains, as well as part of the ring structure of some pharmaceuticals.

Among inorganic sulfur compounds, sulfide minerals include galena (PbS), a major ore of lead; cinnabar (HgS), the chief ore of mercury; and pyrite (FeS2), an important sulfur-bearing mineral. The metal is recovered by roasting the ore in air, which oxidizes the sulfide to sulfur dioxide. This can be recovered and used to make sulfuric acid. Mercuric sulfide is also used as a pigment, as are zinc and cadmium sulfides (ZnS and CdS). The last two compounds give light when struck by electrons, and are used as phosphors in cathode-ray tubes and some laboratory instruments. Hydrogen sulfide (H₂S) is a highly toxic gas with the odor of rotten eggs.

Burning elemental sulfur in air produces sulfur dioxide; sulfur trioxide is produced mainly by catalytic oxidation of sulfur dioxide. Sulfur dioxide has the sharp, acrid odor observed in fumes from burning matches and fireworks. It can be dissolved in water and neutralized with soda to make sulfites (Na₂SO₃) and bisulfites (NaHSO₃). Sulfites, bisulfites, and the parent sulfur dioxide are used as disinfectants, preservatives, and bleaches, particularly for foodstuffs and natural products. Textile fibers, wicker, gelatin, and beet sugar are bleached in this way. Sulfur dioxide has been used as a disinfectant in breweries and as a preservative in table wines.

When sulfur trioxide reacts with water, the product is sulfuric acid:

SO₃ + H₂O → H₂SO₄. Sulfuric acid is by a considerable margin the largest-volume material produced by chemical industry, and has been for at least two centuries. Production in the United States in the 1970s and 1980s averaged nearly 40 million tons per year. Some 70 percent of sulfuric acid production goes into fertilizers, as ammonium sulfate or calcium acid phosphates. The rest is used in other manufacturing processes.

Some metal sulfates are of historical interest or industrial importance. Glauber’s salt, sodium sulfate decahydrate (Na₂SO₄·10H₂O), the “sal mirabile” of alchemy, is still used in dyeing and printing, and as an ionic bulking agent in some detergents. Epsom salt, magnesium sulfate heptahydrate (MgSO₄·7H₂O), has a place in the pharmacopeia. Gypsum, calcium sulfate dihydrate (CaSO₄·2H₂O), can be dehydrated to the hemihydrate, CaSO₄·1/2H₂O, which is plaster of Paris, or to anhydrite, CaSO₄, which is a useful drying agent. Plaster of Paris hardens with water by returning to the gypsum structure. Barium sulfate (BaSO₄) is the opacifying agent used in X-rays of the gastrointestinal system and is used industrially as a pigment base.

Some other inorganic sulfur compounds are important. The ions thiosulfate (S₂O₃^2-) and dithionite (S₂O₄^2-) are valuable reducing agents in the laboratory, usually as the sodium salts. Sodium thiosulfate pentahydrate (Na₂SO₃·5H₂O) is the “hypo” used in photographic development. Its purpose there is as a complexing agent, to remove unreacted silver ions, which, if left in the negative, would cause overall blackening.

Organic sulfur compounds can be grouped as reduced or oxidized sulfur compounds, the structural difference being that the reduced compounds contain only single bonds to the sulfur atom(s) while the oxidized have double bonds, usually to oxygen. Many of the reduced compounds have pungent odors, some quite unpleasant, or are components of flavor in vegetables such as onions and garlic. This is particularly true of the thiols (R-SH), or of compounds that contain the -SH group. We may take the mercaptans in ascending order of complexity of the carbon-containing group: Methyl mercaptan (CH₃SH) has been identified in radish root. Ethanethiol (CH₃CH₂SH) is not found in nature but is used as an odorant in natural (piped) and cylinder gases; it can be detected at levels as low as 0.02 part per billion. The compound n-propyl mercaptan (CH₃CH₂CH₂SH) is found in onions. Isoamyl mercaptan, (CH₃)₂CHCH₂CH₂SH, and trans-2-butenethiol, CH₃CH=CHCH₂SH, are components of the spray of skunks. Dimercaptans are found in asparagus and cabbage: The compound found in asparagus gives rise to pungent sulfur metabolites that appear in the urine after one eats asparagus: CH₃SCH₂CH₂COSH₃ and CH₂ = CHCOSCH₃.

Sulfides (R-S-R) and disulfides (R-S-S-R) include dimethyl sulfide, (CH₃)₂S, isolated from paper-pulp-processing liquors.

Various sulfides are found in plants of the Asteraceae family, such as marigolds and chrysanthemums. The principal component of garlic oil is diallyl disulfide, (CH₂=CHCH₂)₂S₂. Pineapple juice contains sulfur-containing esters such as methyl 3-(methylthio) propanoate (CH₃SCH₂CH₂COOCH₃). The seeds and oil of mustard plants contain complex sulfur compounds whose molecules also incorporate a molecule of the sugar glucose.

The vesicant (blister-producing) mustard gas, used in war, is bis(2-chloroethyl) sulfide, ClCH₂CH₂-S-CH₂CH₂Cl.

Various biochemical compounds are sulfides, many of them heterocycles (ring structures containing carbon with other elements) that also contain nitrogen. Coenzyme A, abbreviated CoA or CoA-SH, is a large molecule with a mercapto group that can be converted into a thioester. Acetyl CoA has the structure CoA-S-COCH₃, and can transfer the acetyl group to other biological molecules. The natural penicillins are sulfur-nitrogen heterocycles. The compound luciferin, which on biological oxidation provides the light of glowworms and fireflies, is also a sulfur-nitrogen heterocycle.

Sulfide and disulfide cross-linkages play a part in both biochemistry and industrial chemistry. Disulfide links are found in protein molecules; mercapto groups in the amino acid cysteine form a disulfide bond that links one part of the long protein chain to another, to hold it in a particular shape required for its job as an enzyme or other biologically active compound.

Presence of many disulfide cross-links produces the rigid structural proteins found in hair, horn, hoof, and nail materials, which cannot be broken down by most animal digestive systems. This is why owls cough up pellets containing feathers, claws, and so forth of creatures they have eaten whole. It is, conversely, why moth larvae can eat wool fibers; they have the enzymes to handle the disulfides.

Sulfide linkages are of immense importance in rubber chemistry. The isoprene polymer that is rubber (as found, for example, in rubber cement) has poor structural properties. It can be firmed up by mixing with elemental sulfur and heating, a process called vulcanization. The sulfur reacts with residual double bonds in the polymer chains, attaching them to one another by sulfide cross-links. Depending on the number of cross-links, the rubber can be made elastic, or firm, or even totally rigid. The chemicals used in the rubber industry to accelerate vulcanization are sulfur compounds; mercaptobenzothiazole is only one of many such compounds.

Oxidized sulfur compounds include sulfoxides, RS(=O) R; sulfonic acids, RSO₃H; and sulfate esters, R-OSO₃H. The most important of the sulfoxides is dimethyl sulfoxide, or DMSO, (CH₃)₂SO. It is a powerful solvent in both laboratory and industrial chemistry. In the latter, it is used to dissolve the synthetic polymer Orlon; the solution is extruded through fine jets into a water bath, where the DMSO disperses into the water and leaves a polymer strand. DMSO has the interesting property of crossing the skin barrier of the body, carrying certain dissolved drugs with it.

Although sulfonic acids and sulfate esters have important laboratory uses, their most common application is as detergents. Detergent molecules have the form R-SO₃^-Na⁺ or R-OSO₃^-Na⁺, where R is a long hydrocarbon chain, typically C₁₂H₂₅. They act by combining the hydrocarbon end of the molecule with greasy dirt, while the ionic sulfonate end combines with water, forming a bridge that allows the dirt to be rinsed away in the water. The virtue of sulfonate detergents over soaps, which act similarly but have a carboxylate ion as the water-soluble end, is that the sulfonate ion does not precipitate with the cations of hard water, remaining effective even in seawater.

Two other sulfonates should be mentioned; both appear as amides. The first is an example of the group of antibacterial agents called sulfa drugs, all of which have the p-aminobenzenesulfonamide structure of sulfanilamide, as shown. The differences lie in the nitrogen-containing structures (some quite complex) that replace the -NH₂ group at the right of the structure. The second is saccharin, also called o-sulfobenzimide; its current systematic name is 1,2-benzisothiazolin-3-one 1,1-dioxide.

Applications

Sulfur compounds are ubiquitous in chemistry. Both inorganic and organic sulfur compounds are found everywhere in laboratory chemistry, industrial chemistry, and physiological chemistry. For all that, however, they are little in evidence in ordinary applications.

Most people know, for example, about chlorine compounds as hazards to the environment; about heavy metals as toxins; about acids in general, even in household chemistry. Yet, except for sulfuric acid (which most people never have seen as such) and possibly sulfites, people are not likely to know which sulfur compounds affect them and their surroundings, and how.

This is in part because the compounds do not advertise themselves as containing sulfur.

One must look closely at the label of lawn and garden fertilizer to find that it contains ammonium sulfate; one must know some chemistry to realize that the hydrogen ion in the acid phosphates came from sulfuric acid. The alum used to crisp up homemade pickles is typically potassium aluminum sulfate (potash alum), not ammonium aluminum sulfate. Similarly, small print identifies California wine as containing sulfites. As for pharmaceuticals, the system adopted for condensing their long chemical names more often than not conceals all information about the chemical elements they contain. Certainly, onions and garlic (or the friendly skunk at the back door) carry no labels at all to identify their flavor- and odor-producing materials. It is easy to overlook the sulfur compounds that affect everyday life.

It is even easier to overlook, or know nothing of, the larger volumes of sulfur compounds that are consumed internally by the chemical industry or in other manufacturing processes. As one example, the important industrial base soda ash (sodium carbonate, Na₂CO₃) was at one time manufactured by the Leblanc process, using sulfuric acid as a starting material. Salt cake (sodium sulfate) was reacted with limestone and coal to produce sodium carbonate and calcium sulfide (not calcium sulfate). Thus, the presence of sulfuric acid in soda ash manufacture was concealed, and as the soda ash was generally used in further manufacture, its presence was not found in the final products, either.

The fertilizer and rubber industries have already been mentioned as internal consumers of sulfur products. The paper industry consumes sulfite and sulfate pulp. In cloth dyeing and finishing, enormous quantities of alums—sulfates of iron (III), chromium (III), or aluminum (III)—are used as mordants in the dyeing process. Aluminum alum is used in the manufacture of dyes, to make “lakes.” It is also used to clear water of clay and other colloidal material in public water systems; in sugar manufacture, to clarify the product; in tanning; in marble and porcelain cement; and in a dozen other applications that are seldom seen because the products used by consumers show no traces of the chemicals used in manufacture. Sulfur compounds are truly ubiquitous.

Context

Sulfur is one of the handful of elements known as elements since ancient times. This fact is reflected in the name itself, which comes to us from Sanskrit via Latin, and refers in meaning only to the element itself, not to its origin or chemical properties (as do “francium” and “germanium,” for example, and “silicon”—from “flint,” silex). Another clue to its age is that each language has its own name for sulfur, just as for iron or gold; this is in contrast to elements like ruthenium or argon, discovered in the nineteenth century, whose names are identical in all tongues.

Compounds of sulfur have long been known, as well. Some of the ancient names are still familiar. “Oil of vitriol,” meaning (more or less) concentrated sulfuric acid, comes from alchemical times, as do “blue vitriol” (copper sulfate) and “green vitriol” (ferrous sulfate), and the various “alums” mentioned above. Other examples of sulfates as vitriols can be given, as well as of sulfates named for their discoverer or place of origin (Glauber’s salt, Epsom salt).

“Orpiment” (arsenic trisulfide, As₂S₃), “liver of sulfur” (mixed polysulfide and thiosulfate of potassium), and “flowers of sulfur” (elemental sulfur as a finely divided powder) are other old names, still current, that show our long familiarity with sulfur compounds.

Sulfur is found in elemental form at the rims of volcanoes (the “brimstone” of the Bible) and, more important, in underground deposits. In the nineteenth century, the most important source was Sicily, where the sulfur-bearing earth from the mines was mounded up and burned. The heat from burning part of the sulfur melted the rest, which flowed out and solidified as it cooled. In 1894, the extensive deposits in Texas and Louisiana were first tapped by the Frasch process, which melts the sulfur in situ by pumping superheated steam down to the underground bed, then forces the molten sulfur to the surface with compressed air. However, most sulfur is recovered as a by-product of petroleum refining and natural-gas processing rather than mined by the Frasch process.

As already noted, sulfuric acid is the leading heavy-industrial chemical. (Other inorganic contenders are nitrogen and oxygen, from liquid air, lime, ammonia, and caustic soda.)

It can be made from elemental sulfur or from the sulfur dioxide obtained when sulfur-bearing ores are roasted. Environmental regulations have led to strict control of sulfur dioxide emissions and the widespread use of sulfur recovery technologies in industry. Without sulfuric acid, much of the chemical industry would come to a halt. This is the most obvious sense in which sulfur is vital in our contemporary world. A large share of sulfur is used to produce fertilizers, reflecting its essential role in plant growth and global food production. Sulfur is also increasingly important in battery production and metal processing for the energy sector, including nickel refining. The countless other appearances of the element in everyday chemistry and physiology, however, of which the examples given above are only a sampling, show that sulfuric acid is merely center stage in a cast of thousands, every member of which is important to humans.

Principal terms

AMIDE: an organic compound derived from a carboxylic or sulfonic acid in which the acidic hydroxy group is replaced by an amino or substituted amino group

CROSS-LINKS: connections between two polymer chains, or between different parts of a single polymer chain; the connections may be direct bonds, or bonds through one or more other atoms

HETEROCYCLE: an organic ring compound (a compound in which a group of atoms forms a closed, cyclic structure) in which the ring is composed of atoms of more than one element

OXIDATION: in inorganic chemistry, removal of electrons from an atom or ion; in organic chemistry, removal of hydrogen atoms from a compound, or addition of oxygen atom(s)

POLYMER: a long chain compound whose molecules are hundreds or thousands of atoms in length and are composed of small groups of atoms that repeat like links in a chain

PROTEIN: a biological polymer in which the links are amino acids

REDUCTION: in inorganic chemistry, addition of electrons to an atom or ion; in organic chemistry, addition of hydrogen atoms to a compound, or removal of oxygen atom(s)

THIO-: a prefix in chemical nomenclature that indicates that one or more atoms in a compound, usually oxygen atoms, have been replaced by sulfur atoms (for example, a thioester, RCOSR, from the regular ester, RCOOR)

Bibliography

Baum, Stuart J. Introduction to Organic and Biological Chemistry. 4th ed., Macmillan, 1987.

Baynes, Katie. “Sulfur Compounds.” NASA, Feb. 2025, www.earthdata.nasa.gov/topics/atmosphere/sulfur-compounds. Accessed 23 Apr. 2026.

Embree, Harland D. Organic Chemistry. Scott, Foresman, 1983.

Fessenden, Ralph J., and Joan S. Fessenden. Organic Chemistry. 4th ed., Brooks/Cole, 1990.

Francioso, Antonio, et al. “Chemistry and Biochemistry of Sulfur Natural Compounds: Key Intermediates of Metabolism and Redox Biology.” Oxidative Medicine and Cellular Longevity, vol. 2020, 2020, p. 8294158, doi:10.1155/2020/8294158. Accessed 23 Apr. 2026.

Greenwood, Norman N., and Alan Earnshaw. Chemistry of the Elements. Pergamon Press, 1984.

Haynes, Williams. The Stone That Burns: The Story of the American Sulphur Industry. Van Nostrand, 1942.

Hill, Caroline R., et al. “Sulfur Compounds: From Plants to Humans and Their Role in Chronic Disease Prevention.” Critical Reviews in Food Science and Nutrition, vol. 63, no. 27, 2023, pp. 8616–38, doi:10.1080/10408398.2022.2057915. Accessed 23 Apr. 2026.

Kutney, Gerald W., and Kenneth Turnbull. “The Sulfur Chemist: The Bearer of Ill Wind?” Journal of Chemical Education, vol. 61, no. 4, 1984, p. 372, doi:10.1021/ed061p372. Accessed 23 Apr. 2026.

Linstromberg, Walter W., and Henry E. Baumgarten. Organic Chemistry: A Brief Course. 6th ed., D. C. Heath, 1987.

Pratt, Christopher J. “Sulfur.” Scientific American, vol. 222, May 1970, pp. 63–72, www.scientificamerican.com/article/sulfur. Accessed 23 Apr. 2026.

Radel, Stanley R., and Marjorie H. Navidi. Chemistry. West, 1990.

Research and Markets. “Sulfur Market Report 2026.” ResearchAndMarkets.com, 2026, www.researchandmarkets.com/reports/5769469/sulfur-market-report. Accessed 23 Apr. 2026.

West, James R., editor. New Uses of Sulfur. American Chemical Society, 1975.