everything about sulfur

• 17 November 2023
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what is sulfur?

Sulfur (S), 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.

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.

History

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.

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 BCE 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. Antoine 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:

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

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 (204.8 °F). 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 (246 °F). 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 (204.8 °F), 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 (246 °F); 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 (235 °F).

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 (314.6 °F), but then rapidly increases, reaching a maximum value at about 187 °C (368.6 °F); between this temperature and the boiling point of 444.6 °C (832.3 °F), the viscosity decreases. The colour also changes, deepening from yellow through dark red, and, finally, to black at about 250 °C (482 °F). 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 (1,652 °F), S₂ is the predominant form; finally, monatomic sulfur is formed at temperatures above 1,800 °C (3,272 °F).

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. It is 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.