Name: Sulfur
Symbol: S
Atomic Number: 16
Atomic Weight: 32.066000
Family: Non Metals
CAS RN: 7704-34-9
Description: A yellow solid.
State (25C): Solid
Oxidation states: +2, +4, +6, -2

Molar Volume: 15.5 cm3/mole
Valence Electrons: 3p4

Boiling Point:  717.9K, 444.75C, 832.55F
Melting Point:
388.51K, 115.36C, 239.65F
Electrons Energy Level: 2, 8, 6
Isotopes: 18 + 4 Stable
Heat of Vaporization: unknown
Heat of Fusion: 1.7175 kJ/mol
Density: 2.07 g/cm3 @ 300K
Specific Heat: 0.71 J/gK
Atomic Radius: 1.09
Ionic Radius: 0.37
Electronegativity: 2.58 (Pauling); 2.44 (Allrod Rochow)
Vapor Pressure: 2.65E-20 Pa @ 115.36C
Sulfur (Sanskrit, sulvari; Latin, sulfur or sulpur) was known in ancient times, and is referred to in the Biblical Pentateuch (Genesis).  The word itself probably is from the Arabic sufra meaning yellow, from the bright color of the naturally occurring form, although the Sanskrit name for sulfur, sulvari could also be interpreted as meaning "enemy of copper".

English translations of the Bible commonly refer to 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 awaits the nonbelieving and unrepented.  It is from this part of the Bible that Hell is implied to "smell of sulfur", although sulfur is in fact odorless. Other than Sulfuric Acid, H2SO4, perhaps the most familiar compound of sulfur in the chemistry lab is the foul-smelling Hydrogen Sulfide gas, H2S, e.g. from rotten eggs also associated with burnt matches.

Homer mentioned "pest-averting sulfur" in the 8th century BC and in 424 BC, the tribe of Boeotia destroyed the walls of a city by burning a mixture of coal, sulfur, and tar under them.  Sometime in the 12th century, the Chinese invented gun powder which is a mixture of Potassium Nitrate (KNO3), Carbon, and Sulfur.  Early alchemists gave sulfur its own alchemical symbol which was a triangle at the top of a cross.  In the late 1770s, 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 utilized.

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Alchemical Symbol, Fire & Brimstone

1s2 2s2p6 3s2p4

Known from ancient times (mentioned in the Hebrew scriptures as brimstone) sulfur was classified as an element in 1777 by Lavoisier.  Pure sulfur is tasteless and odorless with a light yellow color.  Sulfur is the tenth most abundant element in the known universe.

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Additional Representations of Alchemical Symbols for Sulfur

There are three allotropic forms of sulfur.  Two are crystalline and one is amorphous (sometimes called "plastic" sulfur).  This last form gradually reverts to one of the more stable crystalline forms.  Most sulfur is recovered directly as the element from underground deposits by injecting super-heated water and piping out molten sulfur (sulfur melts at 112oC).

Perhaps the most significant compound of sulfur used in modern industrialized societies is Sulfuric Acid (H2SO4).  Sulfur Dioxide, SO2, finds practical applications in bleaching and refrigeration but it is also a nuisance gas resulting from the burning of sulfurous coals.  Sulfur Dioxide gas then reacts with the water vapor in the air to produce a weak acid, Sulfurous Acid (H2SO3) which contributes to the acid rain problem.

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Sulfur or sulphur is an abundant, tasteless, odorless, multivalent non-metal.  Sulfur, in its native form, is a yellow crystalline solid.  In nature, it can be found as the pure element or 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 gunpowder, matches, insecticides and fungicides.


Sulfur melts to a blood-red liquid. When burned, it emits a blue flame.  Sulfur is a pale yellow, odorless and brittle material. It displays three allotropic forms: orthorhombic, monoclinic and amorphous.  The orthorhombic form is the most stable form of sulfur.  Monoclinic sulfur exists between the temperatures of 96C and 119C and reverts back to the orthorhombic form when cooled.  Amorphous sulfur is formed when molten sulfur is quickly cooled.  Amorphous sulfur is soft and elastic and eventually reverts back to the orthorhombic form.

2s2 2p6
3s2 3p4

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 from Hydrogen Sulfide, H2S, and organic sulfur compounds. Sulfur burns with a blue flame that emits Sulfur Dioxide, SO2, notable for its peculiar suffocating odor.  Sulfur is insoluble in water bu soluble in Carbon Disulfide, CS2, and to a lesser extent in other organic solvents such as Benzene, C6H6.   Common oxidations 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.  Sulfur has many allotropes besides S8. Removing one atom from the crown gives S7, which is responsible for sulfur's distinctive yellow color.   Many other rings have been prepared, including S12 and S18.   By contrast, its lighter neighbor oxygen only exists in two states of allotropic significance: O2 and O3.  Selenium, Se, the heavier analogue of sulfur can form rings but is more often found as a polymer chain.

The structure of the cyclooctasulfur molecule, S8

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

A noteworthy property of sulfur is that its viscosity in its molten state, unlike most other liquids, increases above temperatures of 200C due to the formation of polymer chains.  The molten sulfur also becomes dark red in colour above this temperature due to the presence of free valences on terminal atoms of the polymer chains.  However, after a specific temperature is reached, the viscosity is reduced because there is enough energy to break the chains.

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.


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 exploited in Indonesia, Chile, and Japan.

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 (Calcium Sulfate, CaSO4), 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 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, FeS2), Cinnabar (Mercury Sulfide, HgS), Galena (Lead Sulfide, PbS), Sphalerite (Zinc Sulfide, ZnS) and Stibnite (Antimony Sulfide, Sb2S3); and the sulfates, such as Gypsum (Calcium Sulfate, CaSO4), Alunite (Potassium Aluminium Sulfate, KAl(SO4)2), and Barite (Barium Sulfate, BaSO4).  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 also present in many types of meteroites.


Sulfur has many industrial uses. Through its major derivative, Sulfuric Acid, H2SO4, sulfur ranks as one of the more important industrial raw materials.  It is of prime importance to every sector of the world's economies.

Sulfuric Acid production is the major end use for sulfur, and consumption of Sulfuric Acid has been regarded as one of the best indices of a nation's industrial development.   More Sulfuric Acid is produced in the United States every year than any other industrial chemical.

Sulfur is also used in batteries, detergents, the vulcanization of rubber, fungicides, and in the manufacture of phosphate fertilizers.  Sulfites are used to bleach paper and as a preservative in wine and dried fruit.  Because of its flammable nature, sulfur also finds use in matches, gunpowder, and fireworks. Sodium or Ammonium Thiosulfate [Na2S2O3 and (NH4)2S2O3 respectively] are used as photographic fixing agents.  Magnesium Sulfate, MgSO4, better known as Epsom salts, can be used as a laxative, a bath additive, an exfoliant, or a Magnesium supplement for plants.  Sulfur is used as the light-generating medium in the rare lighting fixtures known as sulfur lamps.

In the late 1700s, furniture makers used molten sulfur to produce decorative inlays in their craft.  Because of the Sulfur Dioxide, SO2, produced produced during the process of melting sulfur, the craft of sulfur inlays was soon abandoned.


Sulfur is extracted by mainly two processes: the Sicilian process and the Frasch process. The Sicilian process, which was first used in Sicily, was utilized 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 then put on top of the sulfur deposit and lit on fire.  As the sulfur burns it produces heat, which melts the sulfur deposits, causing the molten sulfur to flow down the sloping hillsides.  The molten sulfur can then be collected in wooden buckets.

The second process used to obtain sulfur is the 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 out through the middle pipe.

The Frasch process produces sulfur with a 99.5% purity content, and which needs no further purification.  On the other hand, the sulfur produced by the Sicilian process must be purified by distillation.


Sulfur is a component of many common minerals, such as Galena (PbS), Gypsum (CaSO42H2O), Pyrite (FeS2), Sphalerite (ZnS or FeS), Cinnabar (HgS), Stibnite (Sb2S3), Epsomite (MgSO47H2O), Celestite (SrSO4) and Barite (BaSO4).  Nearly 25% of the sulfur produced today is recovered from petroleum refining operations and as a byproduct of extracting other materials from sulfur containing ores. The majority of the sulfur produced today is obtained from underground deposits, usually found in conjunction with salt deposits, with a process known as the Frasch process.

Hydrogen Sulfide, H2S, has the characteristic smell of rotten eggs.   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, FeS2, is called pyrite, the so called fool's gold.   Interestingly, pyrite can show semiconductor properties.  Galena, PbS, 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.

Many of the unpleasant odors of organic matter are based on sulfur-containing compounds such as Methyl and Ethyl Mercaptan used to scent natural gas so that leaks are easily detectable.  The odor of garlic and "skunk stink" are also caused by sulfur-containing organic compounds.  However, not all organic sulfur compounds smell unpleasant; for example, grapefruit mercaptan, a sulfur-containing monoterpenoid is responsible for the characteristic scent of grapefruit.

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 important 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.

Inorganic Sulfur Compounds:

Organic Sulfur Compounds: (where R, R', and R are organic groups such as CH3):

Common Minerals
Galena (PbS) Gypsum (CaSO42H2O)
Pyrite (FeS2) Sphalerite (ZnS or FeS)
Cinnabar (HgS) Stibnite (Sb2S3)
Epsomite (MgSO47H2O) Celestite (SrSO4)
Barite (BaSO4)


Sulfur has 18 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 40Ar 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 dS-34 values of co-genetic minerals. The differences between minerals can be used to estimate the temperature of equilibration. The dC-13 and dS-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 dS-34 values from lakes believed to be dominated by watershed sources of sulfate.

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Isotopes Atomic Mass Half-Life
S26 26.028  
S27 27.019 21 ms
S28 28.004 125 ms
S29 28.9966 187 ms
S30 29.9849 1.178 seconds
S31 30.9796 2.572 seconds
S32 31.9721 Stable
S33 32.9715 Stable
S34 33.9679 Stable
S35 34.969 87.32 days
S36 35.9671 Stable
S37 36.9711 5.05 minutes
S38 37.9712 170.3 minutes
S39 38.9751 11.5 seconds
S40 39.975 8.8 seconds
S41 40.98 >1 us
S42 41.981 0.56 seconds
S43 42.987 220 ms
S44 43.988 123 ms
S45 44.995 82 ms
S46 45.999 >200 ns
S47 47.008 >200 ns
S48 48.013 >200 ns
S49 49.02  


Carbon Disulfide, CS2, Carbon Oxysulfide, COS, Hydrogen Sulfide, H2S, and Sulfur Dioxide, SO2, should all be handled with care.

KCNS + H2SO4 + H2O rarrow.gif (63 bytes) KHSO4 + NH4HSO4 + COS (Carbon Oxysulfide)

Although Sulfur Dioxide, SO2, is sufficiently safe to be used as a food additive in small amounts, at high concentrations it reacts with moisture to form Sulfurous Acid, H2SO3, which in sufficient quantities may harm the lungs, eyes or other tissues.  In creatures without lungs such as insects or plants, it otherwise prevents respiration.

40px-Skull_and_crossbones.svg.jpg (1420 bytes) Hydrogen Sulfide, H2S, is quite toxic (more toxic than cyanide).  although very pungent at first, it quickly deadens the senses of smell, so potential victims may be unaware of its presence until it is too late.

Biological Role

Sulfur is an essential component of all living cells.

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 elctron donor in a primitive photosynthesis-like process.  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 is absorbed by plants via the roots from soil as the Sulfate Ion, SO4=,   and reduced to sulfide before it is incorporated into Cysteine and other organic sulfur compounds (sulfur assimilation).

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 resiliancy. 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 also contributes to their indigestability, and also their disagreeable odor when burned.

Environmental Impact

The burning of coal and petroleum by industry and power plants liberates huge amounts of 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, resulting in substantial damage to the natural environment and chemical weathering of statues and architecture.  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.


The element has traditionally been spelled sulphur in the United Kingdom, Ireland, Hong Kong, the Commonwealth Caribbean and India, but sulfur in the United States, while both spellings are used in Australia, Canada and New Zealand.  IUPAC adopted the spelling "sulfur" in 1990, as did the Royal Society of Chemistry Nomenclature Committee in 1992.  The spelling of the term in non-official texts is gradually becoming uniform as sulfur.

The Latin name of the element is sulfur with an F. Since it is an original Latin name and not a Classical Greek loan, the fricative phoneme is indeed denoted with f rather than ph (which would denote the Greek letter f). Sulfur in Greek is theion.


Christian countries often associate sulfur (in English usually under its ancient name, brimstone) with Hell and divine wrath, mostly due to the phrase "fire and brimstone", which occurs in the Bible in descriptions of Hell and divine punishment.   "Fire and brimstone" sermons are those used by some preachers to compel belief by depictions of the horrors of Hell and its punishments.  A joke among scientists has used those descriptions of Hell to conclude that, whereas Heaven's temperature would be a scorching 525oC (because it is bathed in a light of the sun... sevenfold as the light of seven days) Hell can be no hotter than the boiling point of brimstone (a mere 444.6oC), and thus cannot be as hot as Heaven.

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Sulfur Data

Atomic Radius (): 1.09
Atomic Volume cm3/mol : 15.5cm3/mol
Covalent Radius: 1.02
Crystal Structure: Orthorhombic
Ionic Radius: 0.37

Chemical Properties

Electrochemical Equivalents: 0.299 g/amp-hr
Electron Work Function: unknown
Electronegativity: 2.58 (Pauling); 2.44 (Allrod Rochow)
Heat of Fusion: 1.7175 kJ/mol
Incompatibilities: unknown
First Ionization Potential: 10.36
Second Ionization Potential: 23.33
Third Ionization Potential: 34.83
Valence Electron Potential: 160
Ionization Energy (eV): 10.360 eV

Physical Properties

Atomic Mass Average: 32.066
Boiling Point: 717.9K, 444.75C, 832.55F
Melting Point: 388.51K, 115.36C, 239.65F
Heat of Vaporization: unknown
Coefficient of Lineal Thermal Expansion/K-1: 74.33E-6
Electrical Conductivity: 5.0E-24 106/cm
Thermal Conductivity: 0.00269 W/cmK
Density: 2.07 g/cm3 @ 300K
Enthalpy of Atomization: 278.7 kJ/mole @ 25C
Enthalpy of Fusion: 1.72 kJ/mole
Enthalpy of Vaporization: 9.62 kJ/mole
Flammability Class: unknown
Molar Volume: 15.5 cm3/mole
Optical Refractive Index: 1.001111
Relative Gas Density (Air=1): unknown
Specific Heat: 0.71 J/gK
Vapor Pressure: 2.65E-20 Pa @ 115.36C
Estimated Crustal Abundance: 3.50102 milligrams per kilogram
Estimated Oceanic Abundance:
9.05102 milligrams per liter

(Sanskrit, sulvere; L. sulpur) Known to the ancients; referred to in Genesis as brimstone. Sulfur is found in meteorites. A dark area near the crater Aristarchus on the moon has been studied by R.W. Wood with ultraviolet light. This study suggests strongly that it is a sulfur deposit. Sulfur occurs native in the vicinity of volcanoes and hot springs. It is widely distributed in nature as iron pyrites, galena, sphalerite, cinnabar, stibnite, gypsum, epsom salts, celestite, barite, etc. Sulfur is commercially recovered from wells sunk into the salt domes along the Gulf Coast of the U.S. It is obtained from these wells by the Frasch process, which forces heated water into the wells to melt the sulfur, which is then brought to the surface. Sulfur also occurs in natural gas and petroleum crudes and must be removed from these products. Formerly this was done chemically, which wasted the sulfur. New processes now permit recovery, and these sources promise to be very important. Large amounts of sulfur are being recovered from Alberta gas fields. Sulfur is pale yellow, odorless, brittle solid, which is insoluble in water but soluble in carbon disulfide. In every state, whether gas, liquid or solid, elemental sulfur occurs in more than one allotropic form or modification; these present a confusing multitude of forms whose relations are not yet fully understood. Amorphous or "plastic" sulfur is obtained by fast cooling of the crystalline form. X-ray studies indicate that amorphous sulfur may have a helical structure with eight atoms per spiral. Crystalline sulfur seems to be made of rings, each containing eight sulfur atoms, which fit together to give a normal X-ray pattern. eleven isotopes of sulfur exist. Four occur in natural sulfur, none of which is radioactive. A finely divided form of sulfur, known as flowers of sulfur, is obtained by sublimation. Sulfur readily forms sulfides with many elements. sulfur is a component of black gunpowder, and is used in the vulcanization of natural rubber and a fungicide. It is also used extensively in making phosphatic fertilizers. A tremendous tonnage is used to produce sulfuric acid, the most important manufactured chemical. It is used in making sulfite paper and other papers, as a fumigant, and in the bleaching of dried fruits. The element is a good insulator. Organic compounds containing sulfur are very important. Calcium sulfur, ammonium sulfate, carbon disulfide, sulfur dioxide, and hydrogen sulfide are but a few of the many other important compounds of sulfur. Sulfur is essential to life. It is a minor constituent of fats, body fluids, and skeletal minerals. Carbon disulfide, hydrogen sulfide, and sulfur dioxide should be handled carefully. Hydrogen sulfide in small concentrations can be metabolized, but in higher concentrations it quickly can cause death by respiratory paralysis. It is insidious in that it quickly deadens the sense of smell. Sulfur dioxide is a dangerous component in atmospheric air pollution. In 1975, University of Pennsylvania scientists reported synthesis of polymeric sulfur nitride, which has the properties of a metal, although it contains no metal atoms. The material has unusual optical and electrical properties. High-purity sulfur is commercially available in purities of 99.999+%.

Source: CRC Handbook of Chemistry and Physics, 1913-1995. David R. Lide, Editor in Chief. Author: C.R. Hammond