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Lemon yellow crystals.
General properties
Name, symbol, number sulfur, S, 16
Element category nonmetal
Group, period, block 163, p
Standard atomic weight 32.065(5)g·mol−1
Electron configuration [Ne] 3s2 3p4
Electrons per shell 2, 8, 6 (Image)
Physical properties
Phase solid
Density (near r.t.) (alpha) 2.07 g·cm−3
Density (near r.t.) (beta) 1.96 g·cm−3
Density (near r.t.) (gamma) 1.92 g·cm−3
Liquid density at m.p. 1.819 g·cm−3
Melting point 388.36 K, 115.21 °C, 239.38 °F
Boiling point 717.8 K, 444.6 °C, 832.3 °F
Critical point 1314 K, 20.7 MPa
Heat of fusion (mono) 1.727 kJ·mol−1
Heat of vaporization (mono) 45 kJ·mol−1
Specific heat capacity (25 °C) 22.75 J·mol−1·K−1
Vapor pressure
P/Pa 1 10 100 1 k 10 k 100 k
at T/K 375 408 449 508 591 717
Atomic properties
Oxidation states 6, 5, 4, 3, 2, 1, -1, -2
(strongly acidic oxide)
Electronegativity 2.58 (Pauling scale)
Ionization energies
1st: 999.6 kJ·mol−1
2nd: 2252 kJ·mol−1
3rd: 3357 kJ·mol−1
Covalent radius 105±3 pm
Van der Waals radius 180 pm
Crystal structure orthorhombic
Magnetic ordering diamagnetic[1]
Electrical resistivity (20 °C) (amorphous)
Thermal conductivity (300 K) (amorphous)
0.205 W·m−1·K−1
Bulk modulus 7.7 GPa
Mohs hardness 2.0
CAS registry number 7704-34-9
Most stable isotopes
Main article: Isotopes of sulfur
iso NA half-life DM DE (MeV) DP
32S 95.02% 32S is stable with 16 neutrons
33S 0.75% 33S is stable with 17 neutrons
34S 4.21% 34S is stable with 18 neutrons
35S syn 87.32 d β 0.167 35Cl
36S 0.02% 36S is stable with 20 neutrons

Sulfur or sulphur (pronounced /ˈsʌlfər/ SUL-fər, see spelling below) is the chemical element that has the atomic number 16. It is denoted with the symbol S. It is an abundant, multivalent non-metal. Sulfur, in its native form, is a bright yellow crystalline solid. In nature, it can be found as the pure element and as sulfide and sulfate minerals. It is an essential element for life and is found in two amino acids: cysteine and methionine. Its commercial uses are primarily in fertilizers, but it is also widely used in black gunpowder, matches, insecticides and fungicides. Elemental sulfur crystals are commonly sought after by mineral collectors for their brightly colored polyhedron shapes. In nonscientific contexts, it can also be referred to as brimstone.



Rough sulfur crystal
Sulfur crystal from Agrigento, Sicily.

Sulfur (Sanskrit, गन्धक sulvari; Latin Sulphurium) was known in ancient times and is referred to in the Torah (Genesis).

English translations of the Bible commonly referred to burning sulfur as "brimstone", giving rise to the name of 'fire-and-brimstone' sermons, in which listeners are reminded of the fate of eternal damnation that await the unbelieving and unrepentant. It is from this part of the Bible that Hell is implied to "smell of sulfur" (likely due to its association with volcanic activity), although sulfur, in itself, is in fact odorless. The "smell of sulfur" usually refers to either the odor of hydrogen sulfide, e.g. from rotten egg, or of burning sulfur, which produces sulfur dioxide, the smell associated with burnt matches. The smell emanating from raw sulfur originates from a slow oxidation in the presence of air. Hydrogen sulfide is the principal odor of untreated sewage and is one of several unpleasant smelling sulfur-containing components of flatulence (along with sulfur-containing mercaptans).

According to the Ebers Papyrus, a sulfur ointment was used in ancient Egypt to treat granular eyelids. Sulfur was used for fumigation in preclassical Greece;[2] this is mentioned in the Odyssey.[3] Pliny the Elder discusses sulfur in book 35 of his Natural History, saying that its best-known source is the island of Melos. He also mentions its use for fumigation, medicine, and bleaching cloth.[4]

A natural form of sulfur known as shiliuhuang was known in China since the 6th century BC and found in Hanzhong.[5] By the 3rd century, the Chinese discovered that sulfur could be extracted from pyrite.[5] Chinese Daoists were interested in sulfur's flammability and its reactivity with certain metals, yet its earliest practical uses were found in traditional Chinese medicine.[5] A Song Dynasty military treatise of 1044 AD described different formulas for Chinese black powder, which is a mixture of potassium nitrate (KNO3), charcoal, and sulfur. Early alchemists gave sulfur its own alchemical symbol which was a triangle at the top of a cross.

In 1777, Antoine Lavoisier helped convince the scientific community that sulfur was an element and not a compound. In 1867, sulfur was discovered in underground deposits in Louisiana and Texas. The overlying layer of earth was quicksand, prohibiting ordinary mining operations; therefore, the Frasch process was developed.

Spelling and etymology

The element has traditionally been spelled sulphur in the United Kingdom (since the 14th century),[6] most of the Commonwealth including India, Malaysia, South Africa, and Hong Kong, along with the rest of the Caribbean and Ireland, but sulfur in the United States, while both spellings are used in Canada and the Philippines. IUPAC adopted the spelling “sulfur” in 1990, as did the Royal Society of Chemistry Nomenclature Committee in 1992.[7] The Qualifications and Curriculum Authority for England and Wales recommended its use in 2000.[8]

In Latin, the word is variously written sulpur, sulphur, and sulfur (the Oxford Latin Dictionary lists the spellings in this order). It is an original Latin name and not a Classical Greek loan, so the ph variant does not denote the Greek letter φ. Sulfur in Greek is thion (θείον), whence comes the prefix thio-. The simplification of the Latin words p or ph to an f appears to have taken place towards the end of the classical period.[9][10]


When burned, sulfur melts to a blood-red liquid and emits a blue flame which is best observed in the dark.

At room temperature, sulfur is a soft, bright-yellow solid. Elemental sulfur has only a faint odor, similar to that of matches. The odor associated with rotten eggs is due to hydrogen sulfide (H2S) and organic sulfur compounds rather than elemental sulfur. Sulfur burns with a blue flame that emits sulfur dioxide, notable for its peculiar suffocating odor due to dissolving in the mucosa to form dilute sulfurous acid. Sulfur itself is insoluble in water, but soluble in carbon disulfide — and to a lesser extent in other non-polar organic solvents such as benzene and toluene. Common oxidation states of sulfur include −2, +2, +4 and +6. Sulfur forms stable compounds with all elements except the noble gases. Sulfur in the solid state ordinarily exists as cyclic crown-shaped S8 molecules.

The crystallography of sulfur is complex. Depending on the specific conditions, the sulfur allotropes form several distinct crystal structures, with rhombic and monoclinic S8 best known.

A noteworthy property of sulfur is that the viscosity in its molten state, unlike most other liquids, increases above temperatures of 200 °C (392 °F) due to the formation of polymers. The molten sulfur assumes a dark red color above this temperature. At higher temperatures, however, the viscosity is decreased as depolymerization occurs.

Amorphous or "plastic" sulfur can be produced through the rapid cooling of molten sulfur. X-ray crystallography studies show that the amorphous form may have a helical structure with eight atoms per turn. This form is metastable at room temperature and gradually reverts back to crystalline form. This process happens within a matter of hours to days but can be rapidly catalyzed.


The structure of the cyclooctasulfur molecule, S8.

Sulfur forms more than 30 solid allotropes, more than any other element.[11] Besides S8, several other rings are known.[12] Removing one atom from the crown gives S7, which is more deeply yellow than S8. HPLC analysis of "elemental sulfur" reveals an equilibrium mixture of mainly S8, but also S7 and small amounts of S6.[13] Larger rings have been prepared, including S12 and S18.[14][15] By contrast, sulfur's lighter neighbor oxygen only exists in two states of allotropic significance: O2 and O3. Selenium, the heavier analogue of sulfur, can form rings but is more often found as a polymer chain.


Sulfur has 25 known isotopes, four of which are stable: 32S (95.02%), 33S (0.75%), 34S (4.21%), and 36S (0.02%). Other than 35S, the radioactive isotopes of sulfur are all short lived. 35S is formed from cosmic ray spallation of 40argon in the atmosphere. It has a half-life of 87 days.

When sulfide minerals are precipitated, isotopic equilibration among solids and liquid may cause small differences in the δS-34 values of co-genetic minerals. The differences between minerals can be used to estimate the temperature of equilibration. The δC-13 and δS-34 of coexisting carbonates and sulfides can be used to determine the pH and oxygen fugacity of the ore-bearing fluid during ore formation.

In most forest ecosystems, sulfate is derived mostly from the atmosphere; weathering of ore minerals and evaporites also contribute some sulfur. Sulfur with a distinctive isotopic composition has been used to identify pollution sources, and enriched sulfur has been added as a tracer in hydrologic studies. Differences in the natural abundances can also be used in systems where there is sufficient variation in the 34S of ecosystem components. Rocky Mountain lakes thought to be dominated by atmospheric sources of sulfate have been found to have different δS-34 values from lakes believed to be dominated by watershed sources of sulfate.


A man carrying sulfur blocks from Kawah Ijen, a volcano in East Java, Indonesia (photo 2009)

Elemental sulfur can be found near hot springs and volcanic regions in many parts of the world, especially along the Pacific Ring of Fire. Such volcanic deposits are currently mined in Indonesia, Chile, and Japan. Sicily is also famous for its sulfur mines. Sulfur deposits are polycrystalline, and the largest documented single crystal measured 22×16×11 cm3.[16][17]

Significant deposits of elemental sulfur also exist in salt domes along the coast of the Gulf of Mexico, and in evaporites in eastern Europe and western Asia. The sulfur in these deposits is believed to come from the action of anaerobic bacteria on sulfate minerals, especially gypsum, although apparently native sulfur may be produced by geological processes alone, without the aid of living organisms (see below). However, fossil-based sulfur deposits from salt domes are the basis for commercial production in the United States, Poland, Russia, Turkmenistan, and Ukraine.

Sulfur recovered from hydrocarbons in Alberta, stockpiled for shipment in North Vancouver, B.C.

Sulfur production through hydrodesulfurization of oil, gas, and the Athabasca Oil Sands has produced a surplus — huge stockpiles of sulfur now exist throughout Alberta, Canada.

Common naturally occurring sulfur compounds include the sulfide minerals, such as pyrite (iron sulfide), cinnabar (mercury sulfide), galena (lead sulfide), sphalerite (zinc sulfide) and stibnite (antimony sulfide); and the sulfates, such as gypsum (calcium sulfate), alunite (potassium aluminium sulfate), and barite (barium sulfate). It occurs naturally in volcanic emissions, such as from hydrothermal vents, and from bacterial action on decaying sulfur-containing organic matter.

The distinctive colors of Jupiter's volcanic moon, Io, are from various forms of molten, solid and gaseous sulfur. There is also a dark area near the Lunar crater Aristarchus that may be a sulfur deposit.

Sulfur is present in many types of meteorites. Ordinary chondrites contain on average 2.1% sulfur, and carbonaceous chondrites may contain as much as 6.6%. Sulfur in meteorites is normally present as troilite (FeS), but other sulfides are found in some meteorites, and carbonaceous chondrites contain free sulfur, sulfates, and possibly other sulfur compounds.[18]

Extraction and production

Extraction from natural resources

Sulfur is extracted by mainly two processes: the Sicilian process and the Frasch process.

Sicilian process

The Sicilian process, which was first used in Sicily, was used in ancient times to get sulfur from rocks present in volcanic regions. In this process, the sulfur deposits are piled and stacked in brick kilns built on sloping hillsides, and with airspaces between them. Then powdered sulfur is put on top of the sulfur deposit and ignited. As the sulfur burns, the heat melts the sulfur deposits, causing the molten sulfur to flow down the sloping hillside. The molten sulfur can then be collected in wooden buckets.

The sulfur produced by the Sicilian process must be purified by distillation.

Frasch process

In this method, three concentric pipes are used: the outermost pipe contains superheated water, which melts the sulfur, and the innermost pipe is filled with hot compressed air, which serves to create foam and pressure. The resulting sulfur foam is then expelled through the middle pipe.[19]

The Frasch process produces sulfur with a 99.5% purity content, which needs no further purification.

Production from hydrogen sulfide


The Claus process is used to extract elemental sulfur from hydrogen sulfide produced in hydrodesulfurization of petroleum or from natural gas.


In the biological route, hydrogen sulfide (H2S) from natural gas or refinery gas is absorbed with a slight alkaline solution in a wet scrubber, or the sulfide is produced by biological sulfate reduction. In the subsequent process step, the dissolved sulfide is biologically converted to elemental sulfur. This solid sulfur is removed from the reactor. This process has been built on commercial scale. The main advantages of this process are:

  1. no use of expensive chemicals,
  2. the process is safe as the H2S is directly absorbed in an alkaline solution,
  3. no production of a polluted waste stream,
  4. re-usable sulfur is produced, and
  5. the process occurs under ambient conditions.

The biosulfur product is different from other processes in which sulfur is produced because the sulfur is hydrophilic. Next to straightforward reuses as source for sulfuric acid production, it can also be applied as sulfur fertilizer.[20]


Inorganic compounds

Sulfur powder.

When dissolved in water, hydrogen sulfide is acidic and will react with metals to form a series of metal sulfides. Natural metal sulfides are common, especially those of iron. Iron sulfide is called pyrite, the so-called fool's gold. Pyrite can show semiconductor properties.[21] Galena, a naturally occurring lead sulfide, was the first semiconductor discovered and found a use as a signal rectifier in the cat's whiskers of early crystal radios.

Polymeric sulfur nitride has metallic properties even though it does not contain any metal atoms. This compound also has unusual electrical and optical properties. This polymer can be made from tetrasulfur tetranitride S4N4.

Phosphorus sulfides are useful in synthesis. For example, P4S10 and its derivatives Lawesson's reagent and naphthalen-1,8-diyl 1,3,2,4-dithiadiphosphetane 2,4-disulfide are used to replace oxygen from some organic molecules with sulfur.

The sulfate anion, SO2−4
  • Sulfides (S2−), a complex family of compounds usually derived from S2−. Cadmium sulfide (CdS) is an example.
  • Sulfites (SO2−3), the salts of sulfurous acid (H2SO3) which is generated by dissolving SO2 in water. Sulfurous acid and the corresponding sulfites are fairly strong reducing agents. Other compounds derived from SO2 include the pyrosulfite or metabisulfite ion (S2O2−5).
  • Sulfates (SO2−4), the salts of sulfuric acid. Sulfuric acid also reacts with SO3 in equimolar ratios to form pyrosulfuric acid (H2S2O7).
  • Thiosulfates (S2O2−3). Sometimes referred as thiosulfites or "hyposulfites", Thiosulfates are used in photographic fixing (HYPO) as reducing agents. Ammonium thiosulfate is being investigated as a cyanide replacement in leaching gold.[1]
  • Sodium dithionite, Na2S2O4, is the highly reducing dianion derived from hyposulfurous/dithionous acid.
  • Sodium dithionate (Na2S2O6).
  • Polythionic acids (H2SnO6), where n can range from 3 to 80.
  • Peroxymonosulfuric acid (H2SO5) and peroxydisulfuric acids (H2S2O8), made from the action of SO3 on concentrated H2O2, and H2SO4 on concentrated H2O2 respectively.
  • Sodium polysulfides (Na2Sx)
  • Sulfur hexafluoride, SF6, a dense gas at ambient conditions, is used as nonreactive and nontoxic propellant
  • Sulfur nitrides are chain and cyclic compounds containing only S and N. Tetrasulfur tetranitride S4N4 is an example.
  • Thiocyanates contain the SCN group. Oxidation of thiocyanate gives thiocyanogen, (SCN)2 with the connectivity NCS-SCN.

Organic compounds

An organic sulfur compound, dithiane.

Many of the unpleasant odors of organic matter are based on sulfur-containing compounds such as methyl mercaptan and dimethyl sulfide. Thiols and sulfides are used in the odoration of natural gas, notably, 2-methyl-2-propanethiol (t-butyl mercaptan). The odor of garlic and "skunk stink" are also caused by sulfur-containing organic compounds. Not all organic sulfur compounds smell unpleasant; for example, grapefruit mercaptan, a sulfur-containing monoterpenoid is responsible for the characteristic scent of grapefruit. It should be noted that this thiol is present in very low concentrations. In larger concentrations, the odor of this compound is that typical of all thiols, unpleasant.

Sulfur-containing organic compounds include the following (R, R', and R are organic groups such as CH3):

  • Thioethers have the form R-S-R′. These compounds are the sulfur equivalents of ethers.
  • Sulfonium ions have the formula RR'S-'R'", i.e. where three groups are attached to the cationic sulfur center. Dimethylsulfoniopropionate (DMSP; (CH3)2S+CH2CH2COO) is a sulfonium ion, which is important in the marine organic sulfur cycle.
  • Thiols (also known as mercaptans) have the form R-SH. These are the sulfur equivalents of alcohols.
  • Thiolates ions have the form R-S-. Such anions arise upon treatment of thiols with base.
  • Sulfoxides have the form R-S(=O)-R′. The simplest sulfoxide, DMSO, is a common solvent.
  • Sulfones have the form R-S(=O)2-R′. A common sulfone is sulfolane C4H8SO2.

See also Category: sulfur compounds and organosulfur chemistry


One of the direct uses of sulfur is in vulcanization of rubber, where polysulfides crosslink organic polymers. Sulfur is a component of gunpowder. It reacts directly with methane to give carbon disulfide, which is used to manufacture cellophane and rayon.[22]

Elemental sulfur is mainly used as a precursor to other chemicals. Approximately 85% (1989) is converted to sulfuric acid (H2SO4), which is of such prime importance to the world's economies that the production and consumption of sulfuric acid is an indicator of a nation's industrial development.[23] For example with 36.1 million metric tons in 2007, more sulfuric acid is produced in the United States every year than any other inorganic industrial chemical.[24] The principal use for the acid is the extraction of phosphate ores for the production of fertilizer manufacturing. Other applications of sulfuric acid include oil refining, wastewater processing, and mineral extraction.[22]

Sulfur compounds are also used in detergents, fungicides, dyestuffs, and agrichemicals. In silver-based photography sodium and ammonium thiosulfate are used as "fixing agents."

Sulfur is an ingredient in some acne treatments.[25][26]

An increasing application is as fertilizer. Standard sulfur is hydrophobic and therefore has to be covered with a surfactant by bacteria in the ground before it can be oxidized to sulfate. This makes it a slow release fertilizer, which cannot be taken up by the plants instantly, but has to be oxidized to sulfate over the growth season. Sulfur also improves the use efficiency of other essential plant nutrients, particularly nitrogen and phosphorus.[27] Biologically produced sulfur particles are naturally hydrophilic due to a biopolymer coating. This sulfur is therefore easier to disperse over the land (via spraying as a diluted slurry), and results in a faster release.

Sulfites, derived from burning sulfur, are heavily used to bleach paper. They are also used as preservatives in dried fruit.

Magnesium sulfate, better known as Epsom salts, can be used as a laxative, a bath additive, an exfoliant, a magnesium supplement for plants, or a desiccant.

Specialized applications

Sulfur is used as a light-generating medium in the rare lighting fixtures known as sulfur lamps.

Historical applications

In the late 18th century, furniture makers used molten sulfur to produce decorative inlays in their craft. Because of the sulfur dioxide produced during the process of melting sulfur, the craft of sulfur inlays was soon abandoned. Molten sulfur is sometimes still used for setting steel bolts into drilled concrete holes where high shock resistance is desired for floor-mounted equipment attachment points. Pure powdered sulfur was also used as a medicinal tonic and laxative. Sulfur was also used in baths for people who had seizures.

Fungicide and pesticide

Sulfur is one of the oldest fungicides and pesticides. Dusting sulfur, elemental sulfur in powdered form, is a common fungicide for grapes, strawberry, many vegetables and several other crops. It has a good efficacy against a wide range of powdery mildew diseases as well as black spot. In organic production, sulfur is the most important fungicide. It is the only fungicide used in organically farmed apple production against the main disease apple scab under colder conditions. Biosulfur (biologically produced elemental sulfur with hydrophilic characteristics) can be used well for these applications.

Standard-formulation dusting sulfur is applied to crops with a sulfur duster or from a dusting plane. Wettable sulfur is the commercial name for dusting sulfur formulated with additional ingredients to make it water soluble. It has similar applications, and is used as a fungicide against mildew and other mold-related problems with plants and soil.

Sulfur is also used as an "organic" (i.e. "green") insecticide (actually an acaricide) against ticks and mites. A common method of use is to dust clothing or limbs with sulfur powder. Some livestock owners set out a sulfur salt block as a salt lick.

Biological role

See sulfur cycle for more on the inorganic and organic natural transformations of sulfur.

Sulfur is an essential component of all living cells.

Inorganic sulfur forms a part of iron-sulfur clusters, and sulfur is the bridging ligand in the CuA site of cytochrome c oxidase, a basic substance involved in utilization of oxygen by all aerobic life.

Sulfur may also serve as chemical food source for some primitive organisms: some forms of bacteria use hydrogen sulfide (H2S) in the place of water as the electron donor in a primitive photosynthesis-like process in which oxygen is the electron receptor. The photosynthetic green and purple sulfur bacteria and some chemolithotrophs use elemental oxygen to carry out such oxidization of hydrogen sulfide to produce elemental sulfur (So), oxidation state = 0. Primitive bacteria which live around deep ocean volcanic vents oxidize hydrogen sulfide in this way with oxygen: see giant tube worm for an example of large organisms (via bacteria) making metabolic use of hydrogen sulfide as food to be oxidized.

The so-called sulfur bacteria, by contrast, "breathe sulfate" instead of oxygen. They use sulfur as the electron acceptor, and reduce various oxidized sulfur compounds back into sulfide, often into hydrogen sulfide. They also can grow on a number of other partially oxidized sulfur compounds (e. g. thiosulfates, thionates, polysulfides, sulfites). The hydrogen sulfide produced by these bacteria is responsible for the smell of some intestinal gases and decomposition products.

Sulfur is a part of many bacterial defense molecules. For example, though sulfur is not a part of the lactam ring, it is a part of most beta lactam antibiotics, including the penicillins, cephalosporins, and monobactams.

Sulfur is absorbed by plants via the roots from soil as the sulfate ion and reduced to sulfide before it is incorporated into cysteine and other organic sulfur compounds (see sulfur assimilation for details of this process).

Sulfur is regarded as secondary nutrient although plant requirements for sulfur are equal to and sometimes exceed those for phosphorus. However sulfur is recognized as one of the major nutrients essential for plant growth, root nodule formation of legumes and plants protection mechanisms. Sulfur deficiency has become widespread in many countries in Europe.[28][29][30] Because atmospheric inputs of sulfur will continue to decrease, the deficit in the sulfur input/output is likely to increase, unless sulfur fertilizers are used.

In plants and animals the amino acids cysteine and methionine contain sulfur, as do all polypeptides, proteins, and enzymes which contain these amino acids. Homocysteine and taurine are other sulfur-containing acids which are similar in structure, but which are not coded for by DNA, and are not part of the primary structure of proteins. Glutathione is an important sulfur-containing tripeptide which plays a role in cells as a source of chemical reduction potential in the cell, through its sulfhydryl (-SH) moiety. Many important cellular enzymes use prosthetic groups ending with -SH moieties to handle reactions involving acyl-containing biochemicals: two common examples from basic metabolism are coenzyme A and alpha-lipoic acid.

Disulfide bonds (S-S bonds) formed between cysteine residues in peptide chains are very important in protein assembly and structure. These strong covalent bonds between peptide chains give proteins a great deal of extra toughness and resiliency. For example, the high strength of feathers and hair is in part due to their high content of S-S bonds and their high content of cysteine and sulfur (eggs are high in sulfur because large amounts of the element are necessary for feather formation). The high disulfide content of hair and feathers contributes to their indigestibility, and also their odor when burned.

Traditional medical role for elemental sulfur

In traditional medical skin treatment which predates modern era of scientific medicine, elemental sulfur has been used mainly as part of creams to alleviate various conditions such as psoriasis, eczema and acne. The mechanism of action is not known, although elemental sulfur does oxidize slowly to sulfurous acid, which in turn (through the action of sulfite) acts as a mild reducing and antibacterial agent.


Elemental sulfur is non-toxic, but it can burn as an oxidizer or a reducing agent, producing combustion products that are toxic, such as carbon disulfide, carbon oxysulfide, hydrogen sulfide, and sulfur dioxide.

Although sulfur dioxide is sufficiently safe to be used as a food additive in small amounts, at high concentrations it reacts with moisture to form sulfurous acid which in sufficient quantities may harm the lungs, eyes or other tissues. In organisms without lungs such as insects or plants, it otherwise prevents respiration.

Hydrogen sulfide is toxic. Although very pungent at first, it quickly deadens the sense of smell, so potential victims may be unaware of its presence until death or other symptoms occur.

Sulfur trioxide, a volatile liquid at standard temperature and pressure, is extremely dangerous, especially in contact with water, which reacts with it to form sulfuric acid with the generation of much heat. Sulfuric acid poses extreme hazards to many objects and substances.

Environmental impact

The burning of coal and/or petroleum by industry and power plants generates sulfur dioxide (SO2), which reacts with atmospheric water and oxygen to produce sulfuric acid (H2SO4). This sulfuric acid is a component of acid rain, which lowers the pH of soil and freshwater bodies, sometimes resulting in substantial damage to the environment and chemical weathering of statues and structures. Fuel standards increasingly require sulfur to be extracted from fossil fuels to prevent the formation of acid rain. This extracted sulfur is then refined and represents a large portion of sulfur production. In coal fired power plants, the flue gases are sometimes purified. In more modern power plants that use syngas the sulfur is extracted before the gas is burned.

See also


  1. ^ Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics. CRC press. 2000. ISBN 0849304814. 
  2. ^ p. 242, Archaeomineralogy, George Rapp, 2nd ed., Springer: 2009, ISBN 978-3-540-78593-4.
  3. ^ Odyssey, book 22, lines 480–495.
  4. ^ pp. 247–249, Pliny the Elder on science and technology, John F. Healy, Oxford University Press, 1999, ISBN 0198146876.
  5. ^ a b c Zhang Yunming (1986). "The History of Science Society: Ancient Chinese Sulfur Manufacturing Processes". Isis 77: 487. doi:10.1086/354207. 
  6. ^, retrieved 2nd April 2009 18:29 GMT.
  7. ^ Spelling of Sulfur (PDF)
  8. ^ Worldwidewords, 9 December 2000.
  9. ^
  10. ^ Kelly DP (1995) Sulfur and its Doppelgänger. Arch. Microbiol. 163: 157-158.
  11. ^ Ralf Steudel, Bodo Eckert (2003). "Solid Sulfur Allotropes Sulfur Allotropes". Topics in Current Chemistry 230: 1–80. doi:10.1007/b12110. 
  12. ^ Steudel, R. (1982). "Homocyclic Sulfur Molecules". Topics Curr. Chem. 102: 149. 
  13. ^ Tebbe, F. N.; Wasserman, E.; Peet, W. G.; Vatvars, A. and Hayman, A. C. (1982). "Composition of Elemental Sulfur in Solution: Equilibrium of S6, S7, and S8 at Ambient Temperatures". J. Am. Chem. Soc. 104: 4971. doi:10.1021/ja00382a050. 
  14. ^ Beat Meyer (1964). "Solid Allotropes of Sulfur". Chem. Rev. 64 (4): 429–451. doi:10.1021/cr60230a004. 
  15. ^ Beat Meyer (1976). "Elemental sulfur". Chem. Rev. 76: 367–388. doi:10.1021/cr60301a003. 
  16. ^ P. C. Rickwood (1981). "The largest crystals". American Mineralogist 66: 885–907. 
  17. ^ "The giant crystal project site". Retrieved 2009-06-06. 
  18. ^ B. Mason (1962). Meteorites. New York: John Wiley & Sons. p. 160. 
  19. ^ Botsch, Walter (2001). "Chemiker, Techniker, Unternehmer: Zum 150. Geburtstag von Hermann Frasch" (in German). Chemie in unserer Zeit 35 (5): 324–331. doi:10.1002/1521-3781(200110)35:5<324::AID-CIUZ324>3.0.CO;2-9. 
  20. ^ Zessen, E. van, et al. (2004). "Application of THIOPAQ(TM) biosulphur in agriculture". Proceedings of Sulphur 2004, Barcelona (Spain), 24 - 27 Oct. 57 - 68. 
  21. ^ Nyle Steiner (22 February 1). "Iron Pyrites Negative Resistance Oscillator". Retrieved 2007-08-15. 
  22. ^ a b Nehb, Wolfgang; Vydra, Karel (2006). "Sulfur". Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH Verlag. doi:10.1002/14356007.a25_507.pub2. 
  23. ^ Sulfuric Acid Growth
  24. ^ Ober, Joyce A.. "Mineral Yearbook 2007: Sulfur". United States Geological Survey. 
  25. ^ Lin, A; Reimer, R; Carter, D (1988). "Sulfur revisited†". Journal of the American Academy of Dermatology 18 (3): 553. doi:10.1016/S0190-9622(88)70079-1. PMID 2450900. 
  26. ^ Kaminsky, Ana (2003). "Less Common Methods to Treat Acne". Dermatology 206 (1): 68. doi:10.1159/000067824. PMID 12566807. 
  27. ^ Sulfur as a fertilizer
  28. ^ Zhao, F (1999). "Sulphur Assimilation and Effects on Yield and Quality of Wheat". Journal of Cereal Science 30: 1. doi:10.1006/jcrs.1998.0241. 
  29. ^ Blake-Kalff, M.M.A. (2000). Plant and Soil 225: 95. doi:10.1023/A:1026503812267. 
  30. ^ Ceccotti, S. P. (1996). "Plant nutrient sulphur-a review of nutrient balance, environmental impact and fertilizers". Fertilizer Research 43: 117. doi:10.1007/BF00747690. 

External links

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SULPHUR [[[symbol]] S, atomic weight 32.07 (0 = 16)], a non-metallic chemical element, known from very remote times and regarded by the alchemists, on account of its inflammable nature, as the principle of combustion; it is also known as brimstone. The element occurs widely and abundantly distributed in nature both in the free state and in combination. Free or native sulphur, known also as "virgin sulphur," occurs in connexion with volcanoes and in certain stratified rocks in several modes, viz. as crystals, and as stalactitic, encrusting, reniform, massive, earthy and occasionally pulverulent forms as "sulphur meal." It seems rather doubtful whether the unstable monoclinic modification of sulphur (0 - sulphur) is ever found in a native state.

The crystals belong to the orthorhombic system, and have usually a pyramidal habit (fig.), but may be sphenoidal or tabular. Twins are rare. The cleavage is imperfect, but there is a well-marked conchoidal fracture. The hardness ranges from about I to 2, and the from I. 9 to 2.1. Crystals of sulphur are transparent or translucent and highly refractive with strong birefringence; they have a resinous or slightly adamantine lustre, and present the characteristic sulphur-yellow colour. Impurities render the mineral grey, greenish or reddish, bituminous matter being often present in the massive varieties. Sulphur containing selenium, such as occurs in the isle of Vulcano in the Lipari Isles, may be orange-red; and a similar colour is seen in sulphur which contains arsenic sulphide, such as that from La Solfatara near Naples. The presence of tellurium in native sulphur is rare, but is known in certain specimens from Japan.

Volcanic sulphur usually occurs as a sublimate around or on the walls of the vents, and has probably been formed in many cases by the interaction of sulphur dioxide and hydrogen sulphide. Sublimed sulphur also results from the spontaneous combustion of coal seams containing pyrites. Deposits of sulphur are frequently formed by the decomposition of hydrogen sulphide, on exposure to the atmosphere: hence natural sulphureous waters, especially hot springs, readily deposit sulphur. The reduction of sulphates to sulphides by means of organic matter, probably through the agency of sulphur-bacteria, may also indirectly furnish sulphur, and hence it is frequently found in deposits of gypsum. Free sulphur may also result from the decomposition of pyrites, as in pyritic shales and lignites, or from the alteration of galena: thus crystals of sulphur occur, with anglesite, in cavities in galena at Monteponi near Iglesias in Sardinia; whilst the pyrites of Rio Tinto in Spain sometimes yield sulphur on weathering. It should be noted that the oxidation of sulphur itself by atmospheric influence may give rise to sulphuric acid, which in the presence of limestone will form gypsum: thus the sulphur-deposits of Sicily suffer alteration of this kind, and have their outcrop marked by a pale earthy gypseous rock called briscale. Some of the most important deposits of sulphur in the world are worked in Sicily, chiefly in the provinces of Caltanisetta and Girgenti, as at Racalmuto and Cattolica; and to a less extent in the provinces of Catania, Palermo (Lercara) and Trapani (Gibellina). The sulphur occurs in Miocene marls and limestone, associated with. gypsum, celestine, aragonite and calcite. It was formerly believed that the sulphur had a volcanic origin, but it is now generally held that it has either been reduced from gypsum by organic agencies, or more probably deposited from sulphur-bearing waters. Liquid occasionally enclosed in the sulphur and gypsum has been found by 0. Silvestri and by C. A. H. Sjdgren to contain salts like those of sulphur-springs. An important zone of sulphur-bearing Miocene rocks occurs on the east side of the Apennines, constituting a great part of the province of Forli and part of Pesaro. Cesena and Perticara are well-known localities in this district, the latter yielding crystals coated with asphalt. Sulphur is occasionally found crystallized in Carrara marble; and the mineral occurs also in Calabria. Fine crystals occur at Conil near Cadiz; whilst in the province of Teruel in Aragon, sulphur in a compact form replaces fresh-water shells and plant-remains, suggesting its origin from sulphur-springs. Nodular forms of sulphur occur in Miocene marls near Radoboj in Croatia, and near Swoszowic, south of Cracow. Russia possesses large deposits of sulphur in Daghestan in Transcaucasia, and in the Transcaspian steppes. Important deposits of sulphur are worked at several localities in Japan, especially at the Kosaka mine in the province of Rikuchiu, and at Yatsukoda-yama, in the province of Mutsu. Sulphur is worked in Chile and Peru. A complete list of localities for sulphur would include all the volcanic regions of the world. In the United States, sulphur occurs in the following states, in many of which the mineral has been worked: Louisiana (q.v.),Utah,Colorado, California, Nevada, Alaska, Idaho, Texas and Wyoming. The Rabbit Hole sulphur-mines are in Nevada, and a great deposit in Utah occurs at Cove Creek, Beaver (disambiguation)|Beaver county. In the British Islands native sulphur is only a mineralogical rarity, but it occurs in the Carboniferous Limestone of Oughterard in Co. Galway, Ireland.' In combination the element chiefly occurs as metallic sulphides and sulphates. The former are of great commercial importance, being, in most cases, valuable ores, e.g. copper pyrites (copper), galena (lead), blende (zinc), cinnabar (mercury), &c. Of the sulphates we notice gypsum and anhydrite (calcium), barytes (barium) and kieserite (magnesium). Gaseous compounds, e.g. sulphur dioxide and sulphuretted hydrogen, are present in volcanic exhalations (see Volcano) and in many mineral waters. The element also occurs in the animal and vegetable kingdoms. It is present in hair and wool, and in albuminous bodies; and is also a constituent of certain vegetable oils, such as the oils of garlic and mustard. There is, in addition, a series of bacteria which decompose sulphureous compounds and utilize the element thus liberated in their protoplasm (see Bacteriology).


As quarried or mined free sulphur is always contaminated with limestone, gypsum, clay, &c.; the principle underlying its extraction from these impurities is one of simple liquation, i.e. the element is melted, either by the heat of its own combustion or other means, and runs off from the earthy residue.

In the simplest and crudest method, as practised in Sicily, a mass of the ore is placed in a hole in the ground and fired; after a time the heat melts a part of the sulphur which runs down to the bottom of the hole and is then ladled out. This exceptionally wasteful process, in which only one-third of the sulphur is recovered, has been improved by conducting the fusion in a sort of kiln. A semicircular or semi-elliptical pit (calcarone) about 33 ft. in diameter and 8 ft. deep is dug into the slope of a hill, and the sides are coated with a wall of stone. The sole consists of two halves slanting against each other,. the line of intersection forming a descending gutter which runs to the outlet. This outlet having been closed by small stones and sulphate of lime cement, the pit is filled with sulphur ore, which is heaped up considerably beyond the edge of the pit and covered with a layer of burnt-out ore. In building up the heap a number of narrow vertical passages are left to afford a draught for the fire. The ore is kindled from above and the fire so regulated (by making or unmaking air-holes in the covering) that, by the heat produced References. - A very full article ("Zolfo") by G. Aichino, of the Geological Survey of Italy, will be found in the Enciclopedia delle arte e industrie (Turin, 1898). This includes a full bibliography. See also J. F. Kemp in Rothwell's Mineral Industry (1893), vol. ii.; Jules Brunfaut, De l'Exploitation des soufres (2nd ed., 1874); Georgio Spezia, Sull' origine del solfo nei giacementi solfiferi della Sicilia (Turin, 1892). For Japanese sulphur see T. Wada, Minerals of Japan (Tokyo, 1904).

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by the combustion of the least sufficient quantity of sulphur, the rest is liquefied. The molten sulphur accumulates on the sole, whence it is from time to time run out into a square stone receptacle, from which it is ladled into damp poplar-wood moulds and so brought into the shape of truncated cones weighing 110 to 130 lb each. These cakes are sent out into commerce. A calcarone with a capacity of 28,256 cub. ft. burns for about two months, and yields about 200 tons of sulphur. The yield is about 50%. The immense volumes of sulphurous acid evolved give rise to many complaints; all the minor pits suspend work during the summer to avoid destruction of the crops. A calcarone that is to be used all the year round must be at least 220 yds. from any inhabited place and 110 yds. from any field under cultivation.

More efficient is the Gill kiln which uses coke as a fuel. The kiln consists of two (or more) connected cells which are both charged with the ore. The first cell is heated and the products of combustion are led into the second cell where they give up part of their heat to the contained ore, so that by the time the first cell is exhausted the mass in the second cell is at a sufficiently high temperature to ignite spontaneously when air is admitted. Other methods have been employed, but with varying commercial success. For example, in the Gritti and Orlando processes the ore is charged into retorts and the fusion effected by superheated steam, the sulphur being run off as usual; or as was suggested by R. E. Bollman in 1867 the ore may be extracted by carbon bisulphide.

Crude sulphur, as obtained from kilns, contains about 3% of earthy impurities, and consequently needs refining. The following apparatus (invented originally by Michel of Marseilles and improved subsequently by others) enables the manufacturer to produce either of two forms of "refined" sulphur which commerce demands. It consists of a large stone chamber which communicates directly with two slightly slanting tubular retorts of iron. The retorts are charged with molten sulphur from an upper reservoir, which is kept at the requisite temperature by means of the lost heat of the retort fires. The chamber has a safety value at the top of its vault, which is so balanced that the least surplus pressure from within sends it up. The first puff of sulphur vapour which enters the chamber takes fire and converts the air of the chamber into a mixture of nitrogen and sulphur dioxide. The next following instalments of vapour, getting diffused throughout a large mass of relatively cold gas, condense into a kind of "snow," known in commerce and valued as "flowers of sulphur" (fibres sulphuris). By conducting the distillation slowly, so that the temperature within the chamber remains at a sufficiently low degree, it is possible to obtain the whole of the product in the form of "flowers." If compact ("roll") sulphur is wanted the distillation is made to go on at the quickest admissible rate. The temperature of the interior of the chamber soon rises to more than the fusing-point of sulphur (113° C.), and the distillate accumulates at the bottom as a liquid, which is tapped off from time to time to be cast into the customary form of rods.

The Louisiana deposits are worked by a process devised by Herman Frasch in 1891. It consists in sinking a bore-hole, after the manner of a petroleum well, and letting in four pipes centrally arranged, the outer pipe being 10 in. in diameter, the next 6 in., the next 3 in. and the innermost I in. The operation consists in forcing down the 3-in. pipe superheated steam at 330° F. to melt the sulphur. Compressed air is now driven down the 1-in. pipe and bubbles into the melted sulphur and water; the specific gravity of which is greatly diminished, so that it rises to the surface through the outer pipes; it is then run off to settling tanks. The sulphur so obtained is 98% pure.

in some places sulphur is extracted from iron pyrites by one of two methods. The pyrites is subjected to dry distillation from out of iron or fire-clay tubular retorts at a bright red heat. Onethird of the sulphur is volatilized-3FeS 2 = Fe3S4 -12S-and obtained as a distillate. The second method is analogous to the calcarone method of liquation: the ore is placed in a limekiln-like furnace over a mass of kindled fuel to start a partial combustion of the mineral, and the process is so regulated that, by the heat generated, the unburnt part is decomposed with elimination of sulphur, which collects in the molten state on an inverted roof-shaped sole below the furnace and is thence conducted into a cistern. Such pyrites sulphur is usually contaminated with arsenic, and conse- quently is of less value than Sicilian sulphur, which is characteristically free from this impurity.

Large quantities are also recovered from alkali waste (see Alkali Manufacture); another source is the spent oxide of gas manufacture (see GAs).

The substance known as "milk of sulphur" (lac sulphuris) is very finely divided sulphur produced by the following, or some analogous, chemical process. One part of quicklime is slaked with 6 parts of water, and the paste produced diluted with 24 parts of water; 2.3 parts of flowers of sulphur are added; and the whole is boiled for about an hour or longer, when the sulphur dissolves. The mixed solution of poiysulphides and thiosulphate of calcium thus produced is clarified, diluted largely, and then mixed with enough of pure dilute hydrochloric acid to produce a feebly alkaline mixture when sulphur is precipitated. The addition of more acid would produce an additional supply of sulphur (by the action of the H2S203 on the dissolved H 2 S); but this thiosulphate sulphur is yellow and compact, while the polysulphide part has the desired qualities, forming an extremely fine, almost white, powder. The precipitate is washed, collected, and dried at a very moderate heat.

Properties.-Sulphur exists in several allotropic modifications, but before considering these systematically we will deal with the properties of ordinary (or rhombic) sulphur. Commercial sulphur forms yellow crystals which melt at 113° and boil at 444'53° C. under ordinary pressure (H. L. Callendar, Chem. News, 1891, 63, p. 1); just above the boiling point the vapour is orange-yellow, but on continued heating it darkens, being deep red at 50o; at higher temperatures it lightens, becoming straw-yellow at 650°. These colour changes are connected with a dissociation of the molecules. At 524° Dumas deduced the structure S6 from vapour-density determinations, whilst for the range 860 0 to 1040 0, Sainte-Claire Deville and Troost deduced the formula S2. Biltz (Ber., 1888, 21, p. 2013; 1901, 34, p. 2490) showed that the vapour density decreased with the temperature, and also depended on the pressure. G. Preuner and W. Schupp (Zeit. phys. Chem., 1909, 6 9, p. 1 57), in a study of the dissociation isotherms over 300°-850°, detected molecules of Ss, S6 and S2, whilst S i appears to exist below pressures of 30 mm. Boiling and freezing-point determinations of the molecular weight in solution indicate the formula S8. The density of solid sulphur is 2 062 to 2'070, and the specific heat 0.1712; it is a bad conductor of electricity and becomes negatively electrified on friction. It ignites in air at 363° and in oxygen at 275-280 (H. Moissan, Compt. rend., 1903, 1 37, p. 547), burning with a characteristic blue flame and forming much sulphur dioxide, recognized by its pungent odour. At the same time a little trioxide is formed, and, according to Hempel (Ber., 1890, 2 3, p. 1 455), half the sulphur is converted into this oxide if the combustion be carried out in oxygen at a pressure of 40 to 50 atmospheres. Sulphur also combines directly with most of the elements to form sulphides. The atomic weight was determined by Berzelius, Erdmann and Marchand, Dumas and Stas. Thomsen (Zeit. phys. Chem., 1894, 13, p. 726) obtained the value 32.0606.

Allotropic Modifications.-Sulphur assumes crystalline, amorphous and (possibly) colloidal forms. Historically the most important are the rhombic (Sa) and monoclinic (So) forms, discussed by E. Mitscherlich in 1822 (see Ann. chim. phys., 1823, 24, p. 264). The transformations of these two forms are discussed in Chemistry: Physical. Rhombic sulphur may be obtained artificially by slowly crystallizing a solution of sulphur in carbon bisulphide, or, better, by exposing pyridine saturated with sulphuretted hydrogen to atmospheric oxidation (Ahrens, Ber., 1890, 23, p. 2708). It is insoluble in water,' but readily soluble in carbon bisulphide, sulphur chloride and oil of turpentine. The common monoclinic variety is obtained by allowing a crust to form over molten sulphur by partially cooling it, and then breaking the crust and pouring off the still liquid portion, whereupon the interior of the vessel will be found coated with long needles of this variety. Like S. it is soluble in carbon bisulphide. Three other monoclinic forms have been described. By acting upon a solution of sodium hyposulphite with potassium bisulphate, Gernez (Compt. rend., 188 4, 9 8, p. 1 44) obtained a form which he termed nacre (or pearly) sulphur; the same modification was obtained by Sabatier (ibid., 1885, 100, p. 1346) on shaking hydrogen persulphide with alcohol or ether. It is readily transformed into rhombic sulphur. Another form, mixed with the variety just described, is obtained by adding 3 to 4 volumes of alcohol to a solution of ammonium sulphide saturated with sulphur and exposing the mixture to air at about 5°. Engel's monoclinic form (Compt. rend., 1891, 112, p. 866) is obtained by mixing a solution of sodium hyposulphite with double its volume of hydrochloric acid, filtering and extracting with chloroform; the extract yielding the variety on evaporation. A triclinic form is claimed to be obtained by Friedel (Bull. soc. chim., 1879, 32, p. 14) on subliming ordinary sulphur.

1 It is a common practice of keepers of dogs to place a piece of roll sulphur in the animal's water but this serves no useful purpose owing to this property.

Amorphous sulphur or Sy exists in two forms, one soluble in carbon bisulphide, the other insoluble. Milk of sulphur (see above), obtained by decomposing a polysulphide with an acid, contains both forms. The insoluble variety may also be obtained by decomposing sulphur chloride with water and by other reactions. It gradually transforms itself into rhombic sulphur.

The colloidal sulphur, Ss, described by Debus as a product of the interaction of sulphuretted hydrogen and sulphur dioxide in aqueous solution, is regarded by Spring (Rec. tra y. chim., 1906, 2 5, p. 2 53) as a hydrate of the formula S 8 H 2 O. The "blue sulphur," described by Orloff, has been investigated by Paterno and Mazzucchelli (Abs. Journ. Chem. Soc., 1907, ii. 451).

Molten Sulphur. - Several interesting phenomena are witnessed when sulphur is heated above its melting point. The solid melts to a pale yellow liquid which on continued heating gradually darkens and becomes more viscous, the maximum viscosity occurring at 180°, the product being dark red in colour. This change is associated with a change in the spectrum (N. Lockyer). On continuing the heating, the viscosity diminishes while the colour remains the same. If the viscous variety be rapidly cooled, or the more highly heated mass be poured into water, an elastic substance is obtained, termed plastic sulphur. This substance, however, on standing becomes brittle. The character of molten sulphur has been mainly elucidated by the researches of A. Smith and his collaborators. Smith (Abs. Journ. Chem. Soc., 1907, H. 20, 451, 757) regards molten sulphur as a mixture of two isomers SA and Sµ in dynamic equilibrium, SA being light in colour and mobile, and S, dark and viscous. At low temperatures SA predominates, but as the temperature is raised S, increases; the transformation, however, is retarded by some gases, e.g. sulphur dioxide and hydrochloric acid, and accelerated by others, e.g. ammonia. The solid derived from SA is crystalline and soluble in carbon bisulphide, that from S, is amorphous and insoluble. As to the formation of precipitated sulphur, Smith considers that the element first separates in the liquid S,,, condition, which is transformed into SA and finally into Sa; the insoluble (in carbon bisulphide) forms arise when little of the Sw has been transformed; whilst the soluble consist mainly of Sa. Similar views are adopted_ by H. Erdmann (Ann., 1908, 362, p. 133), but. he regards Sµ as the polymer S,, analogous to ozone 03; Smith, however, regards S, as S8.

Compounds. Sulphuretted hydrogen, H 2 S, a compound first examined by C. Scheele, may be obtained by heating sulphur in a current of hydrogen, combination taking place between 200° C. and 358° C., and being complete at the latter temperature, dissociation taking place above this temperature (M. Bodenstein, Zeit. phys. Chem., 18 99, 2 9, p. 315); by heating some metallic sulphides in a current of hydrogen; by the action of acids on various metallic sulphides (ferrous sulphide and dilute sulphuric acid being most generally employed); by the action of sulphur on heated paraffin wax or vaseline, or by heating a solution of magnesium sulphydrate. It is also produced during the putrefaction of organic substances containing sulphur and is found among the products obtained in the destructive distillation of coal. To obtain pure sulphuretted hydrogen the method generally adopted consists in decomposing precipitated antimony sulphide with concentrated hydrochloric acid. As an alternative, H. Moissan (Comp. rend., 1903, 1 37, p. 363) condenses the gas by means of liquid air and fractionates the product.

Sulphuretted hydrogen is a colourless gas possessing an extremely offensive odour. It acts as a strong poison. It burns with a pale blue flame, forming sulphur dioxide and water. It is moderately soluble in water, the solution possessing a faintly acid reaction. This solution is not very stable, since on exposure to air it slowly oxidizes and becomes turbid owing to the gradual precipitation of sulphur. The gas is much more soluble in alcohol. It forms a hydrate of composition H 2 S 7H 2 0. (De Forcrand, Compt. rend., 1888, 106, p. 1357.) The gas may be liquefied by a pressure of about 17 atmospheres, the liquid so obtained boiling at - 61.8° C.; and by further cooling it yields a solid, the melting point of which is given by various observers as - 82° to - 86° C. (see Ladenburg, Ber., 1900, 33, p. 6 37). It is decomposed by the halogens, with liberation of sulphur. Concentrated sulphuric acid also decomposes it: H 2 SO 4 +H 2 S = 2H 2 0 +S02+S. It combines with many metals to form sulphides, and also decomposes many metallic salts with consequent production of sulphides, a property which renders it extremely useful in chemical analysis. It is frequently used as a reducing agent: in acid solutions it reduces ferric to ferrous salts, arsenates to arsenites, permanganates to manganous salts, &c., whilst in alkaline solution it converts many organic nitro compounds into the corresponding amino derivatives. Oxidizing agents rapidly attack sulphuretted hydrogen, the primary products of the reaction being water and sulphur.

By the action of dilute hydrochloric acid on metallic polysulphides, an oily product is obtained which C. L. Berthollet considered to be H3S5. L. Thenard, on the other hand, favoured the formula H2S2. It was also examined by W. Ramsay (Journ. Chem. Soc., 1874, 12, p. 857). Hofmann, who obtained it by saturating an alcoholic solution of ammonium sulphide with sulphur and mixing the product with an alcoholic solution of strychnine, considered the resulting product to be H2S3; while P. Sabatier by fractionating the crude product in vacuo obtained an oi l which boiled between 60° and 85° C. and possessed the composition H4S5.

Several halogen compounds of sulphur are known, the most stable of which is sulphur fluoride, SF 6, which was first prepared by H. Moissan and Lebeau (Compt. rend., 1900, 130, p. 865) by fractionally distilling the product formed in the direct action of fluorine on sulphur. It is tasteless, colourless and odourless gas, which is exceedingly stable and inert. It may be condensed and yields a solid which melts at - 55° C. Sulphuretted hydrogen decomposes it with formation of hydrofluoric acid and liberation of sulphur. Sulphur chloride, S2C12, is obtained as a by-product in the manufacture of carbon tetrachloride from carbon bisulphide and chlorine, and may also be prepared on the small scale by distilling sulphur in a chlorine gas, or by the action of sulphur on sulphuryl chloride in the presence of aluminium chloride (0. Ruff). It is an ambercoloured, fuming liquid possessing a very unpleasant irritating smell. It boils at 139° C. and is solid at - 80° C. It is soluble in carbon bisulphide and in benzene. It is gradually decomposed by water: 2S 2 C1 2 + 3H 2 0 = 4HC1 + 2S + H2S203, the thiosulphuric acid produced in the primary reaction gradually decomposing into water, sulphur and sulphur dioxide. Sulphur chloride dissolves sulphur with great readiness and is consequently used largely for vulcanizing rubber; it also dissolves chlorine. The chloride SC1 2 according to the investigations of 0. Ruff and Fischer (Ber., 1903, 36, p. 418) did not appear to exist, but E. Beckmann (Zeit. phys. Chem., 1909, 42, p. 1839) obtained it by distilling the product of the interaction of chlorine and S 2 C1 2 at low pressures. The tetrachloride, SC14, is formed by saturating S 2 C1 2 with chlorine at - 22° C. (Michaelis, Ann., 1873, 170, p. 1). It is a yellowish-brown liquid which dissociates rapidly with rise of temperature. On cooling it solidifies to a crystalline mass which fuses at - 80° C. (Ruff, ibid.). Water decomposes it violently with formation of hydrochloric and sulphurous acids. Sulphur bromide, S 2 Br 2, is a dark red liquid which boils with decomposition at about 200° C. The products obtained by the action of iodine on sulphur are probably mixtures, although E. Mclvor (Chem. News, 1902, 86, p. 5) obtained a substance of composition S312 (which in all probability is a chemical individual) as a reddish-coloured powder by the action of sulphuretted hydrogen on a solution of iodine trichloride.

Four oxides of sulphur known, namely sulphur dioxide, S02, sulphur trioxide, S03, sulphur sesquioxide, S203, and persulphuric anhydride, S 2 0 7. The dioxide has been known since the earliest times and is found as a naturally occurring product in the gaseous exhalations of volcanoes and in solution in some volcanic springs. It was first collected in the pure condition by J. Priestley in 1775 and its composition determined somewhat later by A. L. Lavoisier. It is formed when sulphur is burned in air or in oxygen, or when many metallic sulphides are roasted. It may also be obtained by heating carbon, sulphur and many metals with concentrated sulphuric acid: C + 2H 2 SO 4 = 2SO 2 }- CO 2 + 2H 2 O; S + 2H 2 SO 4 = 3S0 2 + 2H 2 0; Cu + 2H 2 SO 4 = SO 2 -fCuSO 4 + 2H 2 0; and by decomposing a sulphite, a thiosulphate or a thionic acid with a dilute mineral acid. It is a colourless gas which possesses a characteristic suffocating odour. It does not burn, neither does it support combustion. It is readily soluble in alcohol and in water, the solution. in water possessing a strongly acid reaction. It is easily liquefied, the liquid boiling at - 8° C., and it becomes crystalline at - 72.7° C. (Walden, Zeit. phys. Chem., 1902, 43, p. 43 2). Walden (ibid.) has shown that certain salts dissolve in liquid sulphur dioxide forming additive compounds, two of which have been prepared in the case of potassium iodide: a yellow crystalline solid of composition, KI 14 S0 2, and a red solid of composition, KI 4S0 2. It is decomposed by the influence of strong light or when strongly heated. It combines directly with chlorine to form sulphuryl chloride and also with many metallic peroxides, converting them into sulphates. In the presence of water it frequently acts as a bleaching agent, the bleaching process in this case being one of reduction. It is frequently used as an "antichlor," since in presence of water it has the power of converting chlorine into hydrochloric acid: SO 2 + C12 + 2H 2 0 = 2HC1 + H 2 SO 4. In many cases it acts as a reducing agent (when used in the presence of acids); thus, permanganates are reduced to manganous salts, iodates are reduced with liberation of iodine, &c., 2KMnO 4 + 550 2 + 2H 2 0 = K 2 SO 4 + 2MnSO 4 + 2H 2 SO 4; 2K103+ 550 2 + 4H 2 O =1 3 + 2KHSO 4 + 3H2S04.

It is prepared on the industrial scale for the manufacture of sulphuric acid, for the preparation of sodium sulphate by the Hargreaves process, and for use as a bleaching-disinfecting agent and as a preservative. When compressed it is also used largely as a refrigerating agent, and in virtue of its property of neither burning nor supporting combustion it is also used as a fire extinctor. The solution of the gas in water is used under the name of sulphurous acid. The free acid has not been isolated, since on evaporation the solution gradually loses sulphur dioxide. This solution possesses reducing properties,and gradually oxidizes to sulphuric acid on exposure. When heated in a sealed tube to 180° C. it is transformed into sulphuric acid, with liberation of sulphur. Numerous salts, termed sulphites, are known. Since the free acid would be dibasic, two series of salts exist, namely, the neutral and acid salts. The neutral alkaline salts are soluble in water and show an alkaline reaction, the other neutral salts being either insoluble or difficultly soluble in water. The acid salts have a neutral or slightly acid reaction. The sulphites are prepared by the action of sulphur dioxide on the oxides, hydroxides or carbonates of the metals, or by processes of precipitation. Sulphurous acid may have either of the constitutions <OH O)S ‹ OH 0: S or or be an equilibrium mixture of these OH O H two substances. Although the correct formula for the acid is not known, sulphites are known of both types. Sodium sulphite is almost certainly of the second and unsymmetrical type. Two ethyl sulphites are known, the first or symmetrical form being derived from sulphuryl chloride and alcohol, and the second and unsymmetrical from sodium sulphite and ethyl iodide; the junction of 'one ethyl group with a sulphur atom in the second salt follows because it yields ethyl sulphonic acid, also obtainable from ethyl mercaptan, C 2 H 5 SH. Two isomeric sodium potassium sulphites are known, and may be obtained by neutralizing acid sodium sulphite with potassium carbonate, and acid potassium sulphite with sodium carbonate; their formulae are: O 2 SK(ONa) and 02SNa(OK).

There are various haloid derivatives of sulphurous acid. Thionyl fluoride, SOF 21 has been obtained as a fuming, gas by decomposing arsenic fluoride with thionyl chloride (Moissan and Lebeau, Corn pt. rend., 1900, 130, p. 1436). It is decomposed by water into hydrofluoric and sulphurous acids. Thionyl chloride, SOC1 21 may be obtained by the action of phosphorus pentachloride on sodium sulphite; by the action of sulphur trioxide on sulphur dichloride at 75 -80° C. (Journ. Chem. Soc., 1903, p. 420); and by the action of chlorine monoxide on sulphur at low temperature. It is a colourless, highly refracting liquid, boiling at 78°; it fumes on exposure to moist air. Water decomposes it into hydrochloric and sulphurous acids. On treatment with potassium bromide it yields thionyl bromide, SOBr2, an orange-yellow liquid which boils at 68° C. (40 mm.) (Hartoz and Sims, Chem. News, 1893, 67, p. 82).

Sulphur trioxide, SO 3, mentioned by Basil Valentine in the 15th century, was obtained by N. Lemery in 1675 by distilling green vitriol. It may be prepared by distilling fuming sulphuric acid, or concentrated sulphuric acid over phosphorus pentoxide, or by the direct union of sulphur dioxide with oxygen in the presence of a catalyst, such as platinized asbestos (see Sulphuric Acid). This oxide exists in two forms. The aform is readily fusible and melts at 14.8° C. It corresponds to the simple molecular complex S03. The avariety is infusible, but on heating to 50° C. is transformed into the aform. It corresponds to the molecular complex (S03)2. When perfectly dry this oxide has no caustic properties; it combines rapidly, however, with water to form sulphuric acid, with the development of much heat. It combines directly with concentrated sulphuric acid to form pyrosulphuric acid, H 2 S 2 0 7. It reacts most energetically with many organic compounds, removing the elements of water in many cases and leaving a carbonized mass. It combines directly with many elements and compounds and frequently acts as energetic oxidizing agent. It finds considerable application in the colour industry.

Sulphuryl fluoride, SO 2 F 2, formed by the action of fluorine on sulphur dioxide (H. Moissan, Cornpt. rend. 1 3 2, p. 374), is an exceedingly stable colourless gas at ordinary temperatures, becoming solid at about -120° C. Sulphuryl chloride, SO 2 C1 2, first obtained in 1838 by Regnault (Ann. chim. phys., 1838, (2), 69, p. 170), by the action of chlorine on a mixture of ethylene and sulphur dioxide, may also be obtained by the direct union of sulphur dioxide and chlorine (especially in the presence of a little camphor); and by heating chlorsulphonic acid in the presence of a catalyst, such as mercuric sulphate (Pawlewski, Ber., 18 97, 3 0, p. 765): 2S0 2 C1. OH =S02C12+ H 2 SO It is a colourless fuming liquid which boils at 69° C. and which is readily decomposed by water into sulphuric and hydrochloric acids. Fluorsulphonic acid, SO 2 F OH, is a mobile liquid obtained by the action of an excess of hydrofluoric acid on well-cooled sulphur trioxide. It boils at 162.6° and is decomposed violently by water. Chlorsulphonic acid, SO 2 CI. OH, first prepared by A. Williamson (Proc. Roy. Soc., 1856, 7, p. 11) by the direct union of sulphur trioxide with hydrochloric acid gas, may also be obtained by distilling concentrated sulphuric acid with phosphorus oxychloride: 2H 2 SO 4 +POC1 3 =2SO 2 C1. OH+HC1+HP0 3. It is a colourless fuming liquid which boils at 152-153° C. When heated under pressure it decomposes, forming sulphuric acid, sulphuryl chloride, &c. (Ruff, Ber., 1901, 34, p. 35 0 9). It is decomposed by water with explosive violence. Disulphuryl chloride, S 2 O 5 C1 2, corresponding to pyrosulphuric acid, is obtained by the action of sulphur trioxide on sulphur dichloride, phosphorus oxychloride, sulphuryl chloride or dry sodium chloride: 650 3- + 2POC1 3 = P 2 O 5 + 3S 2 O 5 C1 2; S2C12+ 5503 = S 2 0 5 C1 2 + 550 2; SO 3 + SO 2 C1 2 = S 2 0 5 C1 2; 2NaC1 + 3SO 3 = S 2 0 5 C1 2 -1 Na 2 SO 4. It may also be obtained by distilling chlorsulphonic acid with phosphorus pentachloride: 2S0 2 C1. OH +PC1 5 = S 2 0 5 C1 2 + POC1 3 + 2HC1. It is a colourless, oily, fuming liquid which is decomposed by water into sulphuric and hydrochloric acids. An oxychloride of composition S 2 0 3 C1 4 has been described.

Sulphur sesquioxide, S203, is formed by adding well-dried flowers of sulphur to melted sulphur trioxide at about 12-15° C. The sulphur dissolves in the form of blue drops which sink in the liquid and finally solidify in blue-green crystalline crusts. It is unstable at ordinary temperatures and rapidly decomposes into its generators on warming. It is readily decomposed by water with formation of sulphurous, sulphuric and thiosulphuric acids, with simultaneous liberation of sulphur. Hyposulphurous acid, H 2 S 2 0 4, was first really obtained by Berthollet in 1789 when he showed that iron left in contact with an aqueous solution of sulphur dioxide dissolved without any evolution of gas, whilst C. F. Schonbein subsequently showed the solution possessed reducing properties. P. Schutzenberger (Compt. rend., 1869, 69, p. 169) obtained the sodium salt by the action of zinc on a concentrated solution of sodium bisulphite: Zn + 4NaHSO 3 = Na 2 S 2 O 4 + ZnSO 3 + Na 2 SO 3 + 2H 2 O, the salt being separated from the sulphites formed by fractional precipitation. A solution of the free acid may be prepared by adding oxalic acid to the solution of the sodium salt. This solution is yellow in colour, and is very unstable decomposing at ordinary temperature into sulphur and sulphur dioxide. A pure zinc salt has been prepared by Nabl (Monats., 1899, 20, p. 679) by acting with zinc on a solution of sulphur dioxide in absolute alcohol, whilst H. Moissan (Compt. rend., 1902, 135, p. 647) has also obtained salts by the action of dry sulphur dioxide on various metallic hydrides. Considerable controversy arose as to the constitution of the salts of this acid, the formula of sodium salt, for example, being written as NaHSO 2 and Na 2 S 2 O 4; but the investigations of C. Bernthsen (Ann., 1881, 208, p. 142; 1882, 211, p. 285; Ber., 1900, 33, p. 126) seem to decide definitely in favour of the latter (see also T. S. Price, Journ. Chem. Soc.; also Bucherer and Schwalbe, Zeit. angew. Chem., 1904, 17, p. 1 447). Although this acid appears to be derived from an oxide S203, it is not certain that the known sesquioxide is its anhydride.

Persulphuric anhydride, S207, is a thick viscous liquid obtained by the action of the silent discharge upon a mixture of sulphur trioxide and oxygen. It solidifies at about 0° C, to a mass of long needles, and is very volatile. It is decomposed readily into sulphur trioxide and oxygen when heated. Water decomposes it with formation of sulphuric acid and oxygen: 25207 + 4H 2 0 = 4H 2 SO 4 + 02. Persulphuric acid, HS04, the acid corresponding to S207, has not been obtained in the free state, but its salts were first prepared in 1891 by H. Marshall (Journ. Chem. Soc., 1891, p. 771) by electrolysing solutions of the alkaline bisulphates. The potassium salt, after recrystallization from warm water, separates in large tabular crystals. Its aqueous solution gradually decomposes with evolution of oxygen, behaves as a strong oxidant, and liberates iodine from potassium iodide. Solutions of persulphates in the cold give no precipitate with barium chloride, but when warmed barium sulphate is precipitated with simultaneous liberation of chlorine: K 2 S 2 0 8 + BaC1 2 = BaSO 4 + K 2 SO 4 + C1 2. The conductivity measurements of G. Bredig point to the salt possessing the double formula.

Thiosulphuric acid, formerly called hyposulphurous acid, H2S203, cannot be preserved in the free state, since it gradually decomposes with evolution of sulphur dioxide and liberation of sulphur: H 2 S 2 O 3 = S+S0 2 +H 2 O. The salts of the acid, however, are stable, the sodium salt in particular being largely used for photographic purposes under the name of "hypo." This salt may be prepared by digesting flowers of sulphur with sodium sulphite solution or by boiling sulphur with milk of lime. In this latter reaction the deep yellow solution obtained is exposed to air when the calcium polysulphide formed is gradually converted into thiosulphate by oxidation, and the calcium salt thus formed is converted into the sodium salt by sodium carbonate or sulphate. The thiosulphates are readily decomposed by mineral acids with liberation of sulphur dioxide and precipitation of sulphur: Na 2 S 2 0 3 + 2HC1 = 2NaC1 + S + SO 2 + H 2 O. They form many double salts and give a dark violet coloration with ferric chloride solution, this colour, however, gradually disappearing on standing, sulphur being precipitated. The acid is considered to possess the structure 0 2 S(SH) (OH), since sodium thiosulphate reacts with ethyl bromide to give sodium ethyl thiosulphate, which on treatment with barium chloride gives presumably barium ethyl thiosulphate. This salt, on standing, decomposes into barium dithionate, BaS206, and diethyl disulphide, (C2H5)2S2, which points to the presence of the SH group in the molecule.

The thionic acids are a group of sulphur-containing acids of general formula H2S006, where n=2, 3, 4, possibly 6. Dithionic acid, H2S206, prepared by J. Gay-Lussac in 181q, is usually obtained in the form of its barium salt by suspending freshly precipitated hydrated manganese dioxide in water and passing sulphur dioxide into the mixture until all is dissolved; the barium salt is then precipitated by the careful addition of barium hydroxide. Much manganese sulphate is formed during the reaction, and H. C. Carpenter (Journ. Chem. Soc., 1902, 81, p. I) showed that this can be almost entirely avoided by replacing the manganese oxide by hydrated ferric oxide, the reaction proceeding according to the equation: 2Fe(OH) 3 3S0 2 = FeS 2 0 6 FeS0 3 3H 2 0. He points out that the available oxygen in the oxides may react either as SO 2 + H 2 O ?-- O = H 2 SO 4 or as 2S0 2 -IH20 + 0 = H 2 S 2 0 6; and that in the case of ferric oxide 96% of the theoretical yield of dithionate is obtained, whilst manganese oxide only gives about 75%. A solution of the free acid may be obtained by decomposing the barium salt with dilute sulphuric acid and concentrating the solution in vacuo until it attains a density of about 1.35 (approximately), further concentration leading to its decomposition into sulphur dioxide and sulphuric acid. The dithionates are all soluble in water and when boiled with hydrochloric acid decompose with evolution of sulphur dioxide and formation of a sulphate. Trithionic acid, H2S306, is obtained in the form of its potassium salt by the action of sulphur dioxide on a solution of potassium thiosulphate: 2K 2 S 2 0 3 -f3S0 2 = 2K 2 S 3 0 6 -{- S; or by warming a solution of silver potassium thiosulphate KAgS 2 0 3 = Ag 2 S K 2 S 3 0 6; whilst the sodium salt may be prepared by adding iodine to a mixture of sodium thiosulphate and sulphite: Na 2 S0 3 -fNa 2 S 2 0 3 -f12 = Na 2 S 3062NaI. The salts are unstable; and a solution of the free acid (obtained by the addition of hydrofluosilicic acid to the potassium salt) on concentration in vacuo decomposes rapidly: H 2 S 3 0 6 = H 2 SO 4 -{- S S02. Tetrathionic acid, H 2 S 4 0 6, is obtained in the form of its barium salt by digesting barium thiosulphate with iodine: 2Ba 2 S 2 0 3 -f12 = BaS406 -F 2BaI, the barium iodide formed being removed by alcohol; or in the form of sodium salt by the action of iodine on sodium thiosulphate. The free acid is obtained (in dilute aqueous solution) by the addition of dilute sulphuric acid to an aqueous solution of the barium salt. It is only stable in dilute aqueous solution, for on concentration the acid decomposes with formation of sulphuric acid, sulphur dioxide and sulphur.

Wackenroder's solution (Debus, Journ. Chem. Soc., 1888, 53, p. 278) is prepared by passing sulphuretted hydrogen gas into a nearly saturated aqueous solution of sulphur dioxide at about o° C. The solution is then allowed to stand for 48 hours and the process repeated many times until the sulphur dioxide is all decomposed. The reactions taking place are complicated, and the solution contains ultimately small drops of sulphur in suspension, a colloidal sulphur (which Spring (Rec. tra y. chim., 1906, 2 5, p. 2 53) considers to be a hydrate of sulphur of composition S $ H 2 0), sulphuric acid, traces of trithionic acid, tetraand pentathionic acids and probably hexathionic acid. The solution obtained may be evaporated in vacuo until it attains a density of 1.46 when, if partially saturated with potassium hydroxide and filtered, it yields crystals of potassium pentathionate, K 2 S 5 0 6.3H 2 0. The formation of the pentathionic acid may be represented most simply as follows: 5S0 2 -15H 2 S = H 2560 6 + 5S -{ - 4H 2 0. The aqueous solution of the acid is fairly stable at ordinary temperatures. The pentathionates give a brown colour on the addition of ammoniacal solutions of silver nitrate and ultimately a black precipitate. Hexathionic acid, H 2 S 6 0 6, is probably present in the mother liquors from which potassium pentathionate is prepared. The solution on the addition of ammoniacal silver nitrate behaves similarly to that of potassium pentathionate, but differs from it in giving an immediate precipitate of sulphur with ammonia, whereas the solution of the pentathionate only gradually becomes turbid on standing.

The per-acids of sulphur were first obtained in 1898 by Caro (Zeit. angew. Chem., 1898, p. 845) who prepared monopersulphuric acid by the action of sulphuric acid on a persulphate. This acid may also be prepared by the electrolysis of concentrated sulphuric acid, and it is distinguishable from persulphuric acid by the fact that it immediately liberates iodine from potassium iodide. It behaves as a strong oxidant and in aqueous solution is slowly hydrolysed. It most probably corresponds to the formula H2S05.

See H. E. Armstrong and Lowry, Chem. News (1902), 8 5, p. 193; Lowry and West, Journ. Chem. Soc. (1900), 77, p. 95 o; H. E. Armstrong and Robertson, Proc. Roy. Soc., 50, p. 105; T. S. Price, Ber., 1902, 35, p. 291; Journ. Chem. Soc. (1906), p. 53; A. v. Baeyer and V. Villiger, Ber., passim. Pharmacology. - The sources of all sulphur preparations used in medicine (except calx sulphurata) are native virgin sulphur and the sulphides of metals. Those contained in the British Pharmacopoeia are the following: (1) Sulphur sublimatum, flowers of sulphur (U.S.P.), which is insoluble in water. From it are made (a) confectio sulphuris; (b) unguentum sulphuris; (c) sulphur praecipitatum, milk of sulphur (U.S.P.) which has a sub-preparation trochiscus sulphuris each lozenge containing 5 grs. of precipitated sulphur and 1 gr. of potassium acid tartrate; (d) potassa sulphurata (liver of sulphur), a mixture of salts of which the chief are sulphides of potassium; (e) sulphuris iodidum (U.S.P.), which has a preparation unguentum sulphuris iodidi, strength 1 in 25. From the heating of native calcium sulphate and carbon is obtained calx sulphurata (U.S. and B.P.), or sulphurated lime, a greyish-white powder.

XXVI. 3 Therapeutics. - Externally, sulphur is of use in skin affections. Powdered, it has little effect upon the skin, but in ointment or used by fumigation it has local therapeutic properties. In scabies (itch) it is the best remedy, killing the male parasite, which remains on the surface of the skin. To get at the female and the ova prolonged soaking in soap and water is necessary, the epiderm being rubbed away and the ointment then applied. Precipitated sulphur is also useful in the treatment of acne, but sulphurated lime is more powerful in acne pustulosa and in the appearance of crops of boils. Internally, sulphur is a mild laxative, being converted in the intestine into sulphides. Milk of sulphur, the confection and the lozenge, is used for this purpose. Sulphur and sulphur waters such as those of Harrogate, Aix-la-Chapelle and Aix-les-Bains, have a powerful effect in congested conditions of the liver and intestines, haemorrhoids, gout and gravel. Sulphur is of use in chronic bronchial affections, ridding the lungs of mucus and relieving cough. In chronic rheumatism sulphur waters taken internally and used as baths are effectual. Sulphur in some part escapes unchanged in the faeces.

When sulphur is burned in air or oxygen, sulphur dioxide is produced, which is a powerful disinfectant, used to fumigate rooms which have been occupied by persons suffering from some infectious disease.

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