The Online Encyclopedia and Dictionary







Name, Symbol, Number sulfur, S, 16
Chemical series nonmetals
Group, Period, Block 16 (VIA), 3, p
Density, Hardness 1960 kg/m3, 2
Appearance Lemon yellow at STP
Atomic properties
Atomic weight 32.065 u [not SI]
Atomic radius (calc.) 100 pm (88 pm)
Covalent radius 102 pm
van der Waals radius 180 pm
Electron configuration [Ne]3s2 3p4
e- 's per energy level 2, 8, 6
Oxidation states (Oxide) -1,2,4,6 (strong acid)
Crystal structure orthorhombic
Physical properties
State of matter solid
Melting point 388.36 K (239.38 F)
Boiling point 717.87 K (832.50 F)
Molar volume 15.53 cm3/mol
Heat of vaporization no data
Heat of fusion 1.7175 kJ/mol
Vapor pressure 26.5 zPa at 388 K
Speed of sound __ m/s at ___ K
Electronegativity 2.58 (Pauling scale)
Specific heat capacity 710 J/(kgK)
Electrical conductivity 0.5 fS/m
Thermal conductivity 0.269 W/(mK)
1st ionization potential 999.6 kJ/mol
2nd ionization potential 2252 kJ/mol
3rd ionization potential 3357 kJ/mol
4th ionization potential 4556 kJ/mol
5th ionization potential 7004.3 kJ/mol
6th ionization potential 8495.8 kJ/mol
Most stable isotopes
iso NA half-life DM DE MeV
[not SI]
32S 95.02% S is stable with 16 neutrons
33S 0.75% S is stable with 17 neutrons
34S 4.21% S is stable with 18 neutrons
35S {syn.} 87.32 d β- 0.167 35Cl
36S 0.02% S is stable with 20 neutrons
SI units & STP are used except where noted.

Sulfur (or Sulphur; see spelling below) is the chemical element in the periodic table that has the symbol S and atomic number 16. It is an abundant, tasteless, odorless, multivalent non-metal. Sulfur, in its native form, is a yellow crystaline 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 several amino acids. Its commercial uses are primarily in fertilizers but it is also widely used in gunpowder, laxatives, matches, insecticides and fungicides.


Notable characteristics

Sulfur is a soft and light substance, bright yellow in color. Although hydrogen sulfide (H2S) has a distinct smell of rotten eggs, it should be noted that elemental sulfur is odorless. It burns with a blue flame that emits sulfur dioxide, notable for its peculiar suffocating odor. Sulfur is insoluble in water but soluble in carbon disulfide and other nonpolar solvents. Common oxidation states of sulfur include −2, +2, +4 and +6. Sulfur forms stable compounds with all elements except the noble gases.

Sulfur ordinarily exists as molecules in the solid state, mainly adopting a cyclic crown-shaped S8. Sulfur has many allotropes besides S8. Removing one atom from the crown gives S7, which is responsible for the bright yellow color associated with sulfur. Many other rings have been prepared, including S12 and S18.

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. By contrast, its lighter neighbor oxygen only exists in two states of chemical significance: O2 and O3. Selenium, the heavier analogue of sulfur, forms a few rings but is more stable as a gray-colored polymer.

A noteworthy property is that the viscosity of molten sulfur, unlike most other liquids, increases with temperature due to the formation of polymer 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, however, and gradually reverts back to crystalline form. This process happens within a matter of hours to days but can be rapidly catylized by human saliva (NOTE: never place sulfur or any of its compounds in the mouth).


Sulfur has many industrial uses. Through its major derivative, sulfuric acid (H2SO4), sulfur ranks as one of the more important elements used as an industrial raw material. 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 are used as photographic fixing agents. Magnesium sulfate, better known as Epsom salts can be used as a laxative, a bath additive, an exfoliant, or a magnesium supplement for plants.

Biological role

The amino acids cysteine and methionine contain sulfur, as do all polypeptides, proteins, and enzymes which contain these amino acids. This makes sulfur a necessary component of all living cells. Disulfide bonds between polypeptides are very important in protein assembly and structure. Homocysteine and taurine are also sulfur containing amino acids but are not coded for by DNA or part of the primary structure of proteins. Some forms of bacteria use hydrogen sulfide (H2S) in the place of water as the electron donor in a primitive photosynthesis-like process. Sulfur is absorbed by plants from soil as the sulfate ion. Inorganic sulfur forms a part of iron-sulfur clusters, and sulfur is the bridging ligand in the CuA site of cytochrome c oxidase. Sulfur is an important component of coenzyme A

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. This causes 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 archetecture. 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.


Sulfur (Sanskrit, sulvere; Latin sulpur) was known in ancient times, and is referred to in the Biblical story of Pentateuch (Genesis). English translations of this commonly refer to sulfur as "brimstone", giving rise to the name of 'Fire and brimstone' sermons, which are sermons where hell and eternal damnation for sinners is stressed. It is from this part of the bible that hell is thought to smell of sulfur. The word itself is almost certainly from the Arabic "sufra" meaning yellow, from the bright color of the naturally-occurring form.

Homer mentioned "pest-averting sulfur" in the 9th 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.


Elemental sulfur can be found near hot springs and volcanic regions in many parts of the world, especially along the Pacific Ring of Fire. These occurrences are the basis for the traditional name brimstone, since sulfur could be found near the brims of volcanic craters. Such volcanic deposits are currently exploited in Indonesia, Chile, and Japan.

Significant desposits 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. Such deposits are the basis for commercial production in the United States, Poland, Russia, Turkmenistan, and the Ukraine.

Common naturally-occurring sulfur compounds include the metal sulfides, such as pyrite (iron sulfide), cinnabar (mercury sulfide), galena (lead sulfide), sphalerite (zinc sulfide) and stibnite (antimony sulfide); and the metal sulfates, such as gypsum (calcium sulfate), alunite (potassium aluminium sulfate), and barite (barium sulfate). Hydrogen sulfide is the gas responsible for the odor of rotten eggs. 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 meteorites.


Hydrogen sulfide 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 is called pyrite, the so called fool's gold. Interestingly, pyrite can show semiconductor properties.[1] 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.

Many of the unpleasant odors of organic matter are based on sulfur-containing compounds such as ethyl and methyl 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.

Polymeric sulfur nitride has metallic properties even though it does not contain any metal atoms. This compound also has unusual electrical and optical properties.

Other important compounds of sulfur include:

  • sodium dithionite , Na2S2O4, a powerful reducing agent.
  • sulfurous acid, H2SO3, created by dissolving SO2 in water. Sulfurous acid and the corresponding sulfites are fairly strong reducing agents. Other compounds derived from SO2 include the pyrosulfite ion (S2O52−).
  • The thiosulfates (S2O32−). Thiosulfates are used in photographic fixing, are oxidizing agents, and ammonium thiosulfate is being investigated as a cyanide replacement in leaching gold.[2]
  • Compounds of dithionic acid (H2S2O6)
  • The polythionic acids , (H2SnO6), where n can range from 3 to 80.
  • The sulfates, the salts of sulfuric acid.
  • Sulfides are simple compounds of sulfur with some other chemical element.
  • Sulfuric acid reacts with SO3 in equimolar ratios to form pyrosulfuric acid .
  • peroxymonosulfuric acid and peroxydisulfuric acids , made from the action of SO3 on concentrated H2O2, and H2SO4 on concentrated H2O2 respectively.
  • tetrasulfur tetranitride S4N4.
  • Thiocyanates are compounds containing the thiocyanate ion, SCN-
  • thiocyanogen , (SCN)2.
  • A thioether is a molecule with the form R-S-R', where R and R' are organic groups. These are the sulfur equivalents of ethers.
  • A thiol (also known as a mercaptan) is a molecule with an -SH functional group. These are the sulfur equivalents of alcohols.
  • A thiolate ion has an -S- functional group attached. These are the sulfur equivalent of alkoxide ions.
  • A sulfone is a molecule with an R-S(=O)-R' functional group where R and R' are organic groups.
  • A sulfoxide is a molecule with an R-S(=O)(=O)-R' functional group where R and R' are organic groups. A common example of a sulfoxide is DMSO.


Sulfur has 18 isotopes, of which four 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. Sulfur-35 is formed from cosmic ray spallation of argon-40 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 co-existing 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 S-34 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.


Carbon disulfide, hydrogen sulfide, and sulfur dioxide should all be handled with care.

Although sulfur dioxide is sufficiently safe to be used as a food additive in small amounts, in higher atmospheric concentration it reacts with water in the lungs to form sulfurous acid there; this causes immediate bleeding, the lungs fill up with blood and suffocation results. In creatures without lungs such as insects or plants, it otherwise prevents respiration.

Hydrogen sulfide is quite toxic (more toxic than cyanide). Although very smelly at first, it quickly deadens the sense of smell, so potential victims may be unaware of its presence until it is too late.


The element has traditionally been spelled sulphur in the United Kingdom, India, and New Zealand, but sulfur in the United States, while both spellings are used in Australia and Canada. The IUPAC adopted the spelling "sulfur" in 1990, as did the Royal Society of Chemistry Nomenclature Committee in 1992.

See also


External links

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