Sulphur

Molecular chains, rubber and salt domes are prominent in the lore of sulphur


Contents

  1. Introduction
  2. Physical Properties of Sulphur
  3. Chemical Properties of Sulphur
  4. Mineralogy of Sulphur
  5. Uses of Sulphur
  6. References

Introduction

Sulphur was known in antiquity. In Latin, it was called sulpur, and in Greek, qeion. It was considered the embodiment of fire, and related to lightning. The Greek name, indeed, also means "divinity" and was derived from qeos, which referred to Zeus, who is often shown with a handful of lightning bolts. In Christian mythology, it is the fuel of Hell. A "p" in Latin was used to represent φ in words borrowed from Greek in the times when it was pronounced with a puff of air, but was not yet the "f" sound. Asparagus, in fact, should be aspharagus. Later, when the "f" sound was used, the p often changed to ph in Latin words of Greek origin. Although "sulpur" had no Greek roots, it was attracted into the form "sulphur" in late classical Latin. The spelling was altered in medieval times to "sulfur," which is the spelling that usually appears in Latin dictionaries. The English word is taken directly from Latin, traditionally in the form "sulphur." The American Chemical Society, at a time when spelling simplification was in vogue, decreed that "sulfur" was to be the accepted form in the United States. Although resisted by technical users, this form is now general in the United States, though sulphur is still occasionally seen. In the rest of the world, it is still sulphur. I shall use this spelling here in the interests of tradition and universality.

In German, sulphur is known as Schwefel, in Dutch, as zwavel. In Welsh, it is sylffur, which is simply the Welsh spelling of the sound of the English word. In French, it is soufre, in Italian zolfo, and in Spanish azufre. Sulphate is solfato in Italian, and sulfato in Spanish, where phosphorus is fósforo, since Spanish orthography tries to represent the spoken sound accurately. The old name brimstone is from "burning stone," from the Middle English brinnen, "burn." In German, however, Bernstein, with a similar etymology, means "amber," not "sulphur." "Schwefel" comes from roots expressing "slow-burning."

The chemical symbol for sulphur is S, and is the same in all languages. Dalton's symbol is shown in the heading.

Physical Properties of Sulphur

Sulphur is one of the chalcogenides, or "ore-formers," oxygen O, sulphur S, selenium Se and tellurium Te, that occupy column VIA of the periodic table. All are nonmetals, and their oxides form acidic solutions in water. The atomic number of sulphur is 16, its atomic weight 32.07, and its electron configuration is 1s22s22p63s23p4, with six valence electrons outside a neon core. The first ionization potential of the atom is 10.357 V, the second 23.405 V, so the atom holds its electrons quite strongly. It is the 15th element in order of abundance in the lithosphere, to the amount of 0.052%. Selenium and tellurium are much scarcer, while oxygen, of course, is very abundant.

The electronic structure of S has an important effect on the forms in which it exists. If a neutral atom gains two electrons, it completes the stable argonic structure. The sulphur ion S-- is stable in aqueous solution, where the polar water reduces the energy penalty of the charge. Two sulphur atoms can make a "coordinate covalent" single bond, and this S2 molecule is found in the vapor at high temperatures. Sulphur does not form S2 with a double bond, as oxygen does. If two electrons are added to this molecule, the result is the disulphide ion, S2--, which is also stable in aqueous solution. Both these ions form salts with metals that are disposed to furnish the two electrons. Lewis electron diagrams are shown for sulphur in the diagram.

Sulphur atoms can aggregate further, each one sharing one of its electron pairs with the preceding atoms, and form long chains. The end atom is always in an awkward position, but if possible can acquire two electrons for an argonic structure. In the diagram, the electron distribution makes one end polar and the other negative, so combining ends would attract one another. If one electron is transferred from one end to the other, the ends are alike but still act like free radicals. Silicon forms chains with intermediate oxygen atoms, -Si-O-Si-O-, and carbon forms chains if hydrogen satisfies the extra valences, -(CH2)n-. However, left to their own devices, each of these atoms prefers a 3-dimensional structure. Sulphur naturally forms chains if left to itself. One result of this are the polysulfide ions Sx-- that are easily made. The bond angles in these structures are expected to be 90°, since p orbitals are involved. There is free rotation about a bond. To a lesser degree, selenium also can make chains, but chain-forming is one of the outstanding peculiarities of sulphur.

It is clear that a sulphur chain could bite its own tail and become a ring. This would solve the problems of the terminal atoms, and produce a stable structure. The smallest ring that has no strains consists of 8 atoms. It forms a crown-shaped molecule with 4 atoms on top and 4 atoms on the bottom, rotated relatively by 45°. The Lewis diagram suggests this, schematically. There is a little steric interaction between nonbonded atoms, so the bond angle is stretched to 102°, and the bond length is 0.208 nm. This shows that the size of the covalent sulphur atom is close to 0.1 nm. There is also a less-stable ring of 6 atoms. In any sample of sulphur, thermal agitation is constantly breaking and reassembling chains, so they gradually assume the most stable state for the particular temperature. The yellow color of sulphur is probably due to a small number of thermally-generated unterminated chains.

The S8 molecule is stable, and exists in solid, liquid and gaseous sulphur. Two crystalline forms compete, orthorhombic α-sulphur and monoclinic β-sulphur. A specimen of crystalline orthorhombic sulphur is shown at the right (© Amethyst Galleries). Below 96°, the orthorhombic form is more stable, but it is a fairly narrow thing, and conversions between the two forms are slow. α-S has a density of 2.07 and melts at 113°C. β-S has a density of 1.97 and melts at 119°C. Each melts to a straw-yellow liquid of density 1.808 called λ-sulphur. If the liquid is cooled slowly, needle-like monoclinic crystals form. When the temperature falls below 96°, these crystals slowly change to orthorhombic microcrystals, which is shown by their becoming cloudy. If the liquid is heated further, at about 200°C it darkens to a reddish color and becomes viscous. The S8 rings are thermally broken, and recombine to form long chains, called μ-sulphur. The dark color is due to the free valences at the ends of the chains. If this dark sulphur is suddenly cooled by being poured into water, it becomes a rubbery mass of tangled chains. This was called amorphous sulphur, but is really a kind of glass. On standing, it eventually reverts to S8 and crystallizes. S6 molecules make rhombohedral crystals, which is called (I think) γ-sulphur. There is apparently also an S12 molecule.

Sulphur is the best electrical insulating material known, with a resistivity of about 2 x 1023 μΩ-cm. It is not easy to measure resistivities this large. The reason for this large resistivity is probably the electron traps produced by thermal breaking of S8 rings, which means that electron-attracting terminal atoms are always present. This may also be the explanation of the yellow color of sulphur. Balls of sulphur were the preferred material for static electricity generators. They were introduced by Otto von Guericke of Magdeburg, I believe. Later, glass was used, but it is much more conductive than sulphur, and adsorbs moisture more readily. The coefficient of thermal expansion of α-S is 64 x 10-6 per °C. The thermal conductivity of α-S is 0.277 J/m-s-K, and of β-S, 0.156 J/m-s-K. These are low figures for materials of equivalent density (122 lb/cu.ft. and 129 lb/cu.ft., respectively), and especially compared with metals. The low values are consistent with the low electrical conductivity. Sulphur, however, is not generally best for thermal insulation. The specific heat of α-S is 5.40 cal/mol, β-S 5.65 cal/mol.

Flowers of sulphur is the sublimed substance from purification by distillation. It contains a little oxygen, and some sulphuric acid created by slow oxidation, but these are the only impurities. Its colloidal form makes the powder very active, and this is the only hazard involved in its use. Roll sulphur is cast rhombic sulphur, usually not as pure as flowers. Flour of sulphur is ground rhombic sulphur, and the powder size is larger than in flowers of sulphur. Flowers of sulphur should be used for chemical purposes, never as a part of any pyrotechnic mixture. It is also used as an insecticidal dust, slow oxidation producing the SO2 that is the active substance.

Selenium and tellurium also form Se8 and Te8 rings. Selenium occurs in the very nonmetallic red form, which, like sulphur, is an excellent insulator, and the gray form, which seems to be a semiconductor. Light excites charge carriers into the conduction band and increases the conductivity significantly, so gray selenium was used in light meters. Selenium also makes rectifying contacts with metals, and once was used in very important solid-state rectifiers. Silicon and other semiconductors have replaced selenium completely in these applications, since they are much easier to control and have better characteristics.

Chemical Properties of Sulphur

Sulphur exhibits oxidation numbers of -2, 0, +2, +4 and +6. Oxidation number -2 is seen in the sulphide ion, S--, and a typical compound is H2S, hydrogen sulphide or hydrosulphuric acid. Unlike H2O, this is a gas at room temperature, melting at -82.9°C and boiling at -61.8°C, because hydrogen bonding is absent. In aqueous solution, its ionization constant to H+ and HS- is 5.7 x 10-8, and further to S-- is 1.2 x 10-15. It is a stronger acid than water, but is still weak. H2S is very poisonous (more poisonous than HCN), but fresh air will overcome small concentrations. It has a strong, characteristic rotten-egg odor. The odor disappears at higher concentrations, when the gas is most dangerous. Hydrogen sulphide is used in analysis to precipitate the very insoluble sulphides of Ag, Hg and Pb.

Hydrogen sulphide is prepared for laboratory use by the action of hydrochloric acid on ferrous sulphate, FeS + 2HCl → H2S + FeCl2. Another method, for smaller quantities, is to heat a mixture of sulphur, asbestos and paraffin. The asbestos catalyzes the reduction of the sulphur with hydrogen from the paraffin. The mix was available under the trade name of Aitch-Two-Ess, but asbestophobia has probably suppressed it. A third method is to heat an 8% solution of thionacetamide to 80°C: CH3(CSH)=NH + 2H2O → CH3COOH + NH3 + H2S. The thioacetamide solution is added to the solution in which sulphide ion is desired, which is then heated.

Silver is quite resistant to corrosion. After a long time in a dry atmosphere, it may acquire a transparent film of oxide that appears bluish or violet, often seen in the "toning" of silver coins, that is considered attractive and increases their value. Hydrogen sulphide, however, produces a black tarnish of Ag2S that is not attractive. Eggs, mustard, garlic and rubber bands are all sources of sulphide that will tarnish silver. Sodium bicarbonate will remove the black tarnish gently. Commercial "silver cleaners" may include abrasives and chemicals that will attack the silver as well as the sulphide. An alloy of silver with a small amount of beryllium is said not to tarnish.

Sodium sulphide, Na2S, is an ionic compound of sodium with (-2) sulphur. The compounds Na2S4 and Na2S5, and many other similar sulphides, are easily recognized as short bits of sulphur chains that have picked up a sodium ion at each end, or more probably just two electrons, becoming a polysulphide ion with charge -2.

Carbon disulphide, CS2, is the analogue of carbon dioxide. It is a volatile clear liquid of density 1.26 melting at -111.5°C and boiling at 46.3°C. The difference from carbon dioxide probably indicates that the bonds are not of the same character. Like hydrogen sulphide, it has a distinctive odor, but is not nearly as poisonous. Sulphur dissolves readily in CS2, and can be crystallized from the solution. This is also a compound associated with the oxidation number -2. Sulphur also dissolves in CCl4.

The radical -SH is called sulphydryl or mercaptan. It appears in many organic compounds, like its more familiar oxygen analogue -OH. Methyl mercaptan, CH3SH, is a gas added to natural gas to give it an odor so that leaks will be noticed before an explosive mixture occurs. This is the simplest example of a thioalcohol. Thioethers are also derivatives of H2S. All these compounds tend to have unpleasant odors.

Sulphur with oxidation number -2 occurs in two amino acid side chains, as -SH in cysteine, and as -S- (thioether) in methionine. In proteins, these amino acids can form disulphide bonds, -S-S-, that cross-link protein chains. Disulphide bonds are frequently seen in extracellular proteins. They can be formed in reactions like 2GSH + RO-OH ↔ GS-SG + ROH + H2O, where GSH is glutathione and R is an alkyl radical. The important Coenzyme A has an -SH group as well. Sulphur is essential to metabolism and life. The green at the surface of the yolk in a boiled egg is FeS, formed from the Fe from the yolk and the S from the albumin of the white, which is a protein.

Oxidation number 0 occurs in S8, and in the polysulphides as well. However, it does not correspond to a distinct group of compounds. +2 is represented by the unusual compound SO, which is not stable like the oxygen analogue, the very stable CO, because S does not like to make double bonds of this type. When sulphur is burned in air, the dioxide SO2 is formed, in which the oxidation number appears to be +4. The structure of this molecule is surprising, since the S kernel is surrounded by 12 electrons, not the usual 8 of the argonic structure. Pauling calls this a transargonic structure. It is formed by promoting one electron to a d level, so that five bonding molecular orbitals can be based on the configuration sp3d. This gives a lower energy than the alternatives. In fact, SO2 is quite stable. The bond angle is 119.5°, and the bond length is 0.143 nm. Carbon dioxide, for comparison, is linear.

SO2 is made by burning sulphur in air, S + O2 → SO2, or by roasting sulphide ores, such as pyrite: 4FeS2 + 11O2 → 2Fe2O3 + 8SO2. It is a very irritating gas, and is used for fumigation. The structures of several sulphur-oxygen compounds and radicals are shown in the diagram.

Sulphur dioxide is a cheap and excellent refrigerant, and was once widely used in home refrigerators before it was replaced by the freons. The only problem with SO2 is its irritating nature. The freons are nontoxic and noninflammable, but are now forbidden because of the damage that the chlorine does to the ozone layer. Perhaps sulphur dioxide will come back as a common refrigerant. Its critical temperature is 143.12°C, and its critical pressure is 77.65 atm, so it is easy to liquefy at room temperature by simple compression. The liquid, of density 1.4601 g/cc, boils at -10°C at one atmosphere, with a latent heat of 172.3 Btu/lb. SO2 melts at -75.2°C. Its ratio of specific heats is 1.256, and the molecular weight is 64.06, from which the speed of sound can be calculated.

SO2 dissolves somewhat in water to form sulphurous acid, H2SO3. This is a reasonably strong acid, with a first dissociation constant of 1.72 x 10-2 and a second dissociation constant of 6.24 x 10-8. It is the basis for bisulphite or acid sulphite salts, and sulphite salts. Sodium bisulphite, NaHSO3, is used to prevent unwanted bacterial action in winemaking, and calcium bisulphite, Ca(HSO3)2, is used to dissolve lignin in wood pulp to release cellulose fibres for making paper.

SO2 can be further oxidized by oxygen to SO3, where the oxidation number is +6. At low temperatures, the yield of the reaction is good, but it proceeds with extreme slowness. At high temperatures, the reaction speed is satisfactory, but now the yield is very poor. This dilemma is solved by catalysis. Nitrogen oxides are a suitable catalyst, but the reaction must be carried out in the gaseous phase, which necessitates large reaction volumes. This process was carried out in large, lead-lined rooms, and so was called the chamber process. It was found that platinum was a good catalyst. Since the reaction occurred on the surface of the platinum, much less space is required, and the plant is more economical. This is called the contact process, which produces most of the SO3 today. Vanadium pentoxide, V2O5, is now the catalyst used, since it is not poisoned by the arsenic impurities frequently found in the sulphides that are roasted to obtain SO2. SO3 is planar and triangular, with bond angles 120° and bond lengths 0.143 nm.

SO3 dissolves readily with water to form sulphuric acid, H2SO4. It is actually much quicker to dissolve the SO3 in concentrated sulphuric acid, and then to dilute the result to the desired concentration. Sulphuric acid can be obtained with 100% concentration, density 1.838 g/cc. More SO3 makes fuming sulphuric acid or "oleum," which contains the molecule H2S3O7, pyrosulphuric acid. A definite chemical reaction occurs when SO3 is dissolved in water, which is exothermic (130 kJ/mole). Pauling says that this is due to the hydration of the H+ ion, which cannot be the case or it would occur with all strong acids. It is the hydration of the SO4-- that must be mainly responsible. At any rate, pour concentrated acid slowly into water with stirring, so the heat can be dissipated. The other way around will simply put all the heat into a little water, and make steam and spatter. Sulphuric acid is produced in the largest tonnage of any industrial chemical (second is ammonia).

Sulphuric acid boils at 330°C, which is higher than the boiling points of most acids. To make any acid from one of its salts, simply treat the salt with sulphuric acid and distill. Nitric and hydrochloric acids are easily prepared in this way. One of the primary uses of sulphuric acid is the treatment of phosphate rock to produce phosphoric acid fertilizer.

Hydrogen sulphide burns in air to form sulphur dioxide: 2H2S + 3O2 → 2H2O + 3SO2. In this reaction, the sulphur is oxidized from -2 to +4. In the reaction H2S + 2O2 → H2SO4 + 790 kJ/mol, the oxidation is from -2 to +6. This reaction is very exothermic, but does not happen when H2S is burned. Hydrogen sulphide is said to be a reducing agent in these reactions. On the other hand, +6 sulphur, as in sulphuric acid, is a strong oxidizing agent. In the reaction Cu + 2H2SO4 → CuSO4 + 2H2O + SO2, the copper is oxidized from 0 to +2, while one sulphur atom is reduced from +6 to +4. The hot concentrated acid soaks up the water, while the sulphur dioxide is evolved, so the reaction goes to completion. Sulphur furnishes very good examples of oxidation and reduction.

The sulphate radical can be written -OSO3 to show that one of the oxygens is different from the remaining three. H2SO4 is actually HO-SO2-OH. This oxygen can be replaced by a hydrogen to get -SSO3, the thiosulphate ion, whose structure is shown in the diagram. One sulphur is +6, the other is -2. The sodium salt is NaS2O3, sodium thiosulphate. This salt was once believed to be a hyposulphate (-2O4), so it came to be called "hypo" in photography. It dissolves silver halides, so it can be used to remove unexposed silver halide from a developed photographic image. If you think that the +6 sulphur could be reduced to +4 at the same time as the -2 sulphur is oxidized to 0, you are quite right. If a solution of hypo is acidified, this occurs: Na2S2O3 + 2HCl → 2NaCl + H2O + SO2 + S. Thiosulphuric acid, H2S2O3 undergoes this autooxidation-reduction or disproportionation spontaneously when a solution of it is heated.

If anyhdrous sulphur trioxide reacts with sulphur, blue-green crystals of sulphur sesquioxide are produced. "Sesquioxide" means that there are 1-1/2 oxygens for each sulphur, and the formula is usually written S2O3, which might suggest that the oxidation number of sulphur is +3 in this compound. Actually, the molecule is twice this size, and the structure is as shown in the diagram at the left. The molecular formula is properly S4O6. There is also a potassium salt called potassium tetrathionate, K2S4O6, in which sulphur appears to have an oxidation number +5/2! The grouping S4O6 is neutral in one case; in the other, it has charge -2. The structure of this ion is shown at the left; it has an interesting Lewis electron diagram in which every kernel is surrounded by 8 electrons. The solution of the paradox is that the sulphur atoms have different formal charges even in the same ion, and the sesquioxide is an oxidized version of the tetrathioniate, in which two sulphurs have gone from +4 to +6, and the structure has changed from argononic to transargononic. The "double bonds" are an imperfect way to show the difference in structures. Both these substances are unstable when heated, or when dissolved in water, and disproportionate. These compounds contain a -S-S- linkage, analogous to the -O-O- linkage of a peroxide.

To test for sulphide ion, dissolve the solid in HCl, and hold a piece of filter paper moistened with lead acetate solution over it. If nothing happens, heat the solution. If still nothing happens, add a little granular zinc and the hydrogen evolved will reduce the sulphides to H2S. A brownish-black or silvery stain of PbS on the filter paper will show that sulphide is present. If the solid is not soluble, heat the material strongly and see if SO2 is evolved by odor or by turning a piece of moist blue litmus paper red. SO2 is also produced when the solid is heated on charcoal with the blowpipe flame. To test for sulphate, acidify the solution with HCl, and add a few drops of BaCl2. A white precipitate of BaSO4 indicates the presence of sulphates. To test for sulphites, proceed as for sulphates but filter and discard any precipitate of BaSO4 (sulphites almost always contain some sulphate, since they are oxidized by the air). Then add a drop of 3% H2O2, which oxidizes sulphites to sulphates, and test again with BaCl2. A white precipitate shows that sulphites were present.

Sulphur, and most of its compounds, are non-toxic and completely safe. The notable exceptions to this rule have been mentioned above, especially H2S, SO2 and CS2. Of course, sulphuric acid is very deleterious to tissue, but it is not a poison, and dilution eliminates any hazard. Selenium and tellurium compounds are strong poisons, and should be handled with care. Tellurium compounds give a remarkably foul breath to anyone that has inhaled or ingested them. The elements themselves do not seem to be hazardous.

Mineralogy of Sulphur

An approximation of present-day sea water can be made by dissolving 26.9g of NaCl, 3.2g of MgCl2, 2.2g of MgSO4 and 1.4g of CaSO4 in a litre of pure water. Saturated solutions of these salts would contain 358g, 535g, 309g and 1.9g, respectively, of the salts. Since there are roughly 2.65g of SO4-- per litre, the seas probably contain most of the crust's sulphur as sulphate. Concentration by evaporation of sea water would quickly result in the precipitation of CaSO4, probably as the dihydrate, gypsum. Salt would not precipitate until most of the water had evaporated, and the magnesium would come down last of all. In geological history, the drying up of water trapped in basins under arid conditions has made bedded deposits of gypsum, and of CaSO4 without water of hydration as anhydrite, which is even less soluble than gypsum. Where the evaporation has gone to completion, salt beds are produced, as at Stassfurt in Germany and many other places. The Great Salt Lake is an example of inland waters concentrated in a closed basin since the Pleistocene.

In volcanic areas, the gases H2S and SO2 are often evolved in quantity. They may react according to 2H2S + SO2 → 3S + 2H2O and deposit elemental sulphur, sometimes in well-formed crystals. The hydrogen sulphide may be oxidized by atmospheric oxygen with the same effect. The bacterium Rhodospirillum rubrum obtains energy by oxidizing H2S to elemental sulphur. Volcanic deposits of elementary sulphur exist in the Andes of South America, in Japan, and in Sicily, but they are not a major source of sulphur because of difficult access.

Most of the sources of elemental sulphur are biogenic. Deposits of gypsum and anyhdrite are processed by anaerobic bacteria, especially where hydrocarbons are available as food for the bacteria, which reduce the sulphate to sulphur with the emission of CO2 and H2S. The CO2 reacts with the calcium present to make the mineral calcite, and limestone. The H2S may escape, or, especially if trapped, may be oxidized by oxygen-containing surface waters that have percolated into the area to elemental sulphur, or by biological action. In this way, bedded sulphur is created, associated with limestone and gypsum or anhydrite. Such deposits were mined in south-central Sicily by underground methods, and long were the main source of industrial sulphur. The sulphur-bearing rocks are of Miocene age.

In the late Palaeozoic, the Gulf of Mexico was much smaller and further north, extending into what is now Texas and Louisana, and was an arm of the Pacific, formed when South America pulled away after its collision with North America. Some say storms blew water into the shallow basin. In the Mesozoic, terranes riding in on subducting plates to the west, which now are part of Mexico, blocked the entrance to this proto-Gulf, making it an inland sea. This sea dried up in the arid climate, depositing thick beds of salt, gypsum and anyhdrite. As South America further receded in the Tertiary, muddy sediments from the rising North America accumulated in a deepening trough in huge amounts, until the salt beds were buried under 50,000 feet of stratified soft sediments of various kinds. The proto-Gulf was not connected to the Atlantic until late Jurassic time, and the salt was divided into northern and southern parts. There are no fossils or palaeomagnetic evidence to show ages or spreading history, so study of this region has been difficult.

Salt is a less dense (2.16 g/cc) substance than the sediments that were piled on it (2.5 g/cc and more), and the pressure on it was immense. It is also a plastic substance that flows by recrystallization on an extended time scale. Regions that happened to be thicker bowed upwards from buoyant forces, while the salt around them was squeezed into the middle. The salt rose in thin columns that pierced and shoved aside the soft sediments, called diapirs, or "through-piercers," that rose toward the surface, fed from below with salt that flowed in from the sides. These diapirs developed a caprock of insoluble anhydrite derived from the gypsum beds which they pierced. When the salt came within a few thousand feet of the surface, dark shales that contained abundant organic matter were eaten by the anaerobic bacteria that make crude oil and evolve CH4. This petroleum accumulated in traps formed by sandy beds forced upwards by the rising salt and the impermeable salt. The principal salt bed of the Texas-Lousiana area is the middle Jurassic Louann salt. The Werner and Metate formations are anhydrite. The Challenger formation is the analogue south of the Gulf.

The huge amount of salt represented in part by the Louann salt was deposited when the Atlantic was first opening, and was still narrow. The new ocean basin was poorly connected with the rest of the ocean, and the climate was arid, so a large amount of evaporites was deposited. The thickness should be no surprise, since the basin was deepening to accommodate it, and evaporation is geologically a very rapid process. The Gulf of Mexico had not yet widened to its current width, splitting the Louann salt into a northern and a southern part, and a connection between the Atlantic and Pacific did not yet exist. When such a connection was finally established as South America moved away, the salt was buried beneath new sediments, and the Atlantic was no longer an evaporating basin. The isthmus of Panama was later formed by volcanic activity, and the oceans were separate again. Salt corresponding to the Louann is found on the African west coast.

Everything was now perfect for the sulphate-reducing bacteria to make elemental sulphur, and to create brecciated (broken) limestone layers above the anhydrite. The sulphur filled fractures and voids in this limestone, so that about 25% of the rock was sulphur. If the salt rose to the surface, erosion destroyed the caprock and all the sulphur, leaving only the salt core. However, if the salt was again buried by sediments and gypsum, the process could begin again with the same result. There are over 200 of these salt domes, some inland as far as Dallas, others in Gulf waters, and more around the periphery of the Gulf in Mexico. The deep trough was split in two by the later widening of the Gulf, and conditions are similar all around its current periphery. Studies with isotopic tracers have confirmed the biological origin of salt-dome sulphur.

Oil was discovered near Beaumont, Texas in 1901 on the flanks of a salt dome that became famous as Spindletop (a few years earlier, a prominent geologist said he would personally drink all the oil found in Texas). This was made possible by a new method of drilling wells, the rotary method, imported from Russia, where there were similar conditions. The search for oil prompted the prospecting for salt domes, so their locations became well-known. Usually, there are not many indications on the surface. Drilling for oil revealed the abundance of sulphur in some (not all) salt domes.

We have noted that if a salt dome reached the surface, the sulphur was removed by erosion. The sulphur in subsurface salt domes was at depths of 1000 ft. and more, beyond the reach of normal mining techniques. Herman Frasch (1851-1914) quickly devised an ingenious method for reaching the sulphur. Three concentric pipes were lowered in the casing of a well (only an 8" diameter hole was required). Superheated water, at 160°-170°C was forced down the outer annulus. The sulphur melted at 110°C and puddled around the bottom of the hole. Hot compressed air was then blown down in innermost tube, which made a froth of the molten sulphur, which rose willingly in the remaining annulus, to be produced at the surface and led into large ponds to solidify. This sulphur was 99.5% pure, and did not contain annoying impurities like arsenic which accompanied sulphur from pyrite. This process was introduced in 1904, and soon broke the Sicilian monopoly.

The Frasch process uses the same drilling technology as the petroleum industry, and the same exploratory techniques. One dome can produce 30,000,000 tons of sulphur. About 25 of the 200 domes in the Texas-Lousiana region contain economic amounts of sulphur. A great deal of water is also necessary, as well as fuel for the boilers, which is usually the abundant natural gas of the region.

A great deal of sulphur comes from the "sweetening" of crude oil and natural gas. Much petroleum is "sour," containing significant amounts of H2S, which must be removed before the refined products are sold, or the gas supplied to the pipeline. Since petroleum is liquid or gaseous, this purification is much easier than it is for the solid coal. Other sulphur comes from the roasting of sulphide ores of lead, zinc and copper, and pyrite. The recovery of the SO2 in the roasting process, and using it to make sulphuric acid that can either be sold or employed in the ore processing, is not only economical, but very beneficial to the surrounding countryside.

Besides elemental sulphur, sulphur minerals are chiefly metal sulphides and sulphates. The sulphides are produced by hot, active gases and fluids containing the sulphide ion. Dana lists 26 sulphides, which include most of the common ores, 12 sulphates, and 8 sulphosalts, which are thioantimonates and thioarsenates.

Pyrite, FeS2, is a very common mineral, often well-crystallized in characteristic forms (pyritohedrons). It is found, exceptionally, as flat radial crystal discs near Sparta, Illinois, called "Sparta dollars." The sulphur in bituminous coal is often in the form of pyrite. It is hard, heavy (5.02) and has a metallic lustre and light-brass color. It often occurs in flakes associated with placer gold, and is called "fool's gold." Pyrite is a compound of ferrous iron and the bisulphide ion, S2--. Marcasite has the same composition, but a different crystal structure. It is whiter, and has different crystal habits. The regular sulphides, FeS and Fe2S3, are not found in nature in pure form. Pyrrhotite is Fe1-xS, and pentlandite is a mixed sulphide of nickel and iron, the principal ore of nickel at Sudbury, Ontario. Chalcopyrite and bornite have copper with the iron, and arsenopyrite has As replacing one S. The sulphur content ruins pyrite as an ore of iron. SO2 produced from roasting pyrite often contains arsenic as an impurity, which poisons Pt catalysts in making sulphuric acid.

Pyrite is more properly called pyrites, a three-syllable word from Greek (pye-rye-tees), but this was thought to be a plural, so it was singularized to pyrite by American geologists. Cherry is a similar case (originally cherries from the French cerise). The reference is to the striking of sparks from pyrite to start a fire.

Galena, PbS, is another familiar sulphide mineral. Not only does it crystallize in the cubic system, its crystals are very often shiny metallic-looking cubes. The metallic luster is due to free electrons in the conduction band, since PbS is a semiconductor. It can make a PN-junction with a metal whisker, as in "crystal" set radios. Galena is lead-gray and heavy (7.4-7.6), but not very hard, and so scratches easily. It has excellent cubic cleavage, and so breaks up into smaller cubes when shattered. It is very often associated with silver, whose sulphide ore is argentite, Ag2S. Argentite can be cut with a knife like lead, which distinguishes it from galena, which it otherwise resembles, being heavy and soft. These are the most important ores of lead and silver, respectively.

An ore is a mineral from which a metal can be extracted profitably, and this means that the metal has to be concentrated in the ore by some process. Iron, lead and silver may be widely disseminated in amounts too small to be useful. Fluids may leach out these metals as ions, and carry them to a point where hydrogen sulphide causes them to precipitate in a small region, as in a vein that follows a crack in the rocks. This may happen with fluids coming from above, called supergene enrichment, or more commonly from below, in hypogene enrichment, usually by active volcanic gases. Mining districts are those in which such fluids have been active in the past. The great insolubility of sulphides explains their occurrence as minerals. Copper, lead, zinc, silver, mercury, nickel, cobalt, tin, antimony, bismuth and molybdenum are won from sulphide ores. Of the common metals, only iron, aluminium and magnesium are obtained from non-sulphide ores. For further discussion of the sulphide ores, see the pages on the individual metals.

Uses of Sulphur

Several important uses of sulphur and its compounds have aready been discussed. The uses of sulphur can be classified under the oxidation number, as sulphides (-2), elemental sulphur (0), SO2 and sulphites (+4), and SO3 and sulphates (+6). Sulphuric acid (+6) has already been mentioned as the most important product of the chemical industry, the "King of Chemicals."

A very important use of sulphur, if not of large amounts, is the vulcanization of rubber. Understanding where this fits in is an excuse for reviewing the nature of the material. Rubber was found in use by the South American natives when the Spanish arrived. They played with solid rubber balls, wore clothing made impermeable by rubber, and used rubber bottles. The Spanish could not make any use of this new material, and it remained a curiosity for over 200 years. In 1770, Joseph Priestley named the substance "rubber" from its ability to rub out pencil marks. If it was dissolved in turpentine, fabrics could be impregnated with it and made waterproof (the "MacIntosh"). It also made waterproof boots and similar objects, and by 1830 was widely used. Rubber is soluble in benzene, carbon disulphide, turpentine, and similar solvents. Rubber cement (originally rubber in benzene) was popular for art paste-ups, because any excess was easily removed by rubbing with a finger.

Rubber is elastic, water repellent and a good electrical insulator. However, it became brittle at low temperatures, sticky at high temperatures, had little mechanical durability and oxidized rapidly. It was found that sulphur improved the properties somewhat, but it was reserved for Charles Goodyear (1800-1860) to find in 1839 that if rubber was not only mixed with sulphur, but cooked at 120°-160°C for a sufficient time, it became practically a new substance. It was elastic over a wide range of temperature, becoming neither brittle nor tacky, mechanically more durable, and quite impermeable. If about 5% sulphur is used, the rubber is soft and flexible. When 32% sulphur is used, the product, called ebonite or hard rubber, is an excellent, durable thermosetting plastic, once popular as a structural material in all kinds of electrical apparatus. It has now been replaced by cheaper plastics.

Rubber was obtained from the latex of Hevea brasiliensis, a tree that grew wild in South America. The latex is a colloidal suspension of rubber that comes from just under the bark of the tree (it is not the sap). In this tree it is particularly pure and free of the resins that such liquids (which are fairly common--as in the dandelion and milkweed) usually contained. The suspension is coagulated with formic acid, and the liquid pressed out between rollers, forming sheets about 1/4" thick that are often smoked to sterilize them, and packed in bales for shipment. Seeds of H. brasiliensis were smuggled out of Brazil in 1876 by Sir Henry Wickham (1846-1928), and used to establish plantations in southeast Asia, the source of most of the natural rubber since that time.

Rubber is a polymer with the empirical formula (C5H8)x, with one C=C double bond to each unit. Sulphur cross-links the polymer chains, taking advantage of these double bonds. The density of rubber is about 0.915 g/cc, and it is a white elastic solid. It is essentially a (1,4) polymer of isoprene, the molecule shown in the diagram, but efforts to reproduce natural rubber exactly have been in vain. Natural latex contains about 35% rubber substance. However, many very good rubber-like materials have been produced as artificial polymers. Natural rubber will retain its market only so long as it is cost-effective. The spur to artificial rubber production was the disruption of plantation rubber during World War II. Germany had produced artificial rubbers even during World War I.

About 80% of industrial rubber is used to make pneumatic tires and tubes. Much rubber is reclaimed from used rubber products, but not enough is done along these lines even now. The pneumatic tire was invented in 1877, and the cord tire, which replaced fabric with parallel cords and made the tire much more durable, in 1910. The mix for making a tire includes, besides rubber and sulphur, accelerators (CaO, MgO, organic compounds), softeners (stearic acid, paraffin, petroleum jelly), reinforcing agents (carbon black, zinc oxide), and antioxidants. The green tire is assembled with all its fabric, cord, and wire reinforcement, and put into the vulcanizing mold. Heat and pressure then make the finished tire.

Sulphur is an important fuel in pyrotechnic mixtures, because it is cheap and stable. It occurs in match heads, the most common pyrotechnical device, and was an ingredient of black powder. Black powder is a special mixture of 75% potassium nitrate, 15% charcoal, and 10% sulphur, more or less. The sulphur and charcoal are ground together dry so that the thixotropic sulphur thoroughly coats the active surface of the charcoal. Then the nitrate is added, and the mix is wet ground until homogenized. The dried mix is formed into grains of the desired size, and the powder is ready to use. More details can be found in Bang!.

I have found an intriguing reference to philosophers' eggs, which are made from a mixture of NaCl, S and Hg. I know nothing more about them. Perhaps they are made up into a small ball, and do something strange when set on fire. They are not the Pharoah's Serpent eggs made from mercuric thiocyanate, Hg(SCN)2 (now banned due to mercurophobia). Perhaps they are a medical treatment. Dictionaries do not stoop to define them.

References

W. N. Jones, Inorganic Chemistry (Philadelphia: Blakiston, 1949). Chapter 21. A traditional text with considerable descriptive material.

L. Pauling, General Chemistry (New York: Dover, 1988). An excellent modern text, but with material on sulphur scattered through the book.

C. S. Hurlbut, Jr., Dana's Manual of Mineralogy, 16th ed. (New York: John Wiley, 1952).

Any good encyclopedia will have an article on rubber and the rubber industry, as well as articles on sulphur and sulphuric acid.

R. L. Bates, Geology of the Industrial Rocks and Minerals (New York: Dover, 1969). pp. 406-414. Good bibliography.

M. Kay and E. H. Colbert, Stratigraphy and Life History (New York: John Wiley & Sons, 1965). pp. 559-566. There has been important more recent work on the development of the Gulf of Mexico. For example, see A. Salvador, Bull. A.A.P.G., April 1987.

Picture of rhombic sulphur kindly furnished by Amethyst Galleries, Inc.. This is an excellent website and specimens can be purchased online. This is probably the best mineral website, and the company should be supported for making it available.


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Composed by J. B. Calvert
Created 10 January 2003
Last revised 23 February 2003