Sulfur
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 two
amino acids. Its commercial uses are primarily in
fertilizers but it is also widely used in
gunpowder,
matches,
insecticides and
fungicides.
|
A piece of sulfur melts to a blood-red liquid. When burned, it emits a blue flame. |
At room temperature, sulfur is a soft bright yellow solid. Although sulfur is infamous for its smell—frequently compared to rotten eggs—the odor is actually characteristic of
hydrogen sulfide (H
2S); 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 to a lesser extent in other organic solvents such as
benzene. 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 S
8 molecules. Sulfur has many
allotropes besides S
8. Removing one atom from the crown gives S
7, which is responsible for sulfur's distinctive yellow color. Many other rings have been prepared, including S
12 and S
18. By contrast, its lighter neighbor
oxygen only exists in two states of chemical significance: O
2 and O
3.
Selenium, the heavier analogue of sulfur can form rings but is more often found as a polymer chain.
|
The structure of the S8 molecule |
The
crystallography of sulfur is complex. Depending on the specific conditions, the sulfur
allotropes form several distinct
crystal structures, with
rhombic and
monoclinic S
8 best known.
A noteworthy property is that the
viscosity of molten sulfur, unlike most other liquids, increases with temperature due to the formation of
polymer chains. However, after a certain 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.
Sulfur has many industrial uses. Through its major derivative,
sulfuric acid (
H2S
O4), 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. 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 produced during the process of melting sulfur, the craft of sulfur inlays was soon abandoned.
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 nor are they part of the
primary structure of proteins. Some forms of
bacteria use
hydrogen sulfide (H
2S) in the place of water as the
electron donor in a primitive
photosynthesis-like process. 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 (
sulfur assimilation). 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 .
The burning of
coal and
petroleum by industry and power plants liberates huge amounts of
sulfur dioxide (S
O2) which reacts with atmospheric water and oxygen to produce
sulfuric acid. 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.
Sulfur (
Sanskrit,
sulvere;
Latin sulpur) was known in ancient times, and is referred to in the
Biblical Pentateuch (
Genesis). The word itself is almost certainly from the
Arabic sufra meaning yellow, from the bright color of the naturally occurring form.
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 as mentioned above sulfur is in fact odorless. The "smell of sulfur" usually refers to the odor of
hydrogen sulfide, e.g. from rotten eggs. Burning sulfur produces
sulfur dioxide, the smell 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.
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 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
Ukraine.
Sulfur extracted from oil, gas and the
Athabasca Oil Sands has become a glut on the market, with huge stockpiles of sulfur in existence throughout Alberta.
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.
See also Sulfide minerals, Sulfate minerals.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
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 S
4N
4.
Phosphorus sulfides are important in synthesis. For example, P
4S
10 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:*
Sulfides (S
2âˆ'), a complex family of compounds usually derived from S
2âˆ'.
Cadmium sulfide (CdS) is an example.
*
Sulfites (SO
32âˆ'), the salts of
sulfurous acid (H
2SO
3) which is generated by dissolving SO
2 in water. Sulfurous acid and the corresponding sulfites are fairly strong reducing agents. Other compounds derived from SO
2 include the pyrosulfite or metabisulfite ion (S
2O
52âˆ').
*
Sulfates (SO
42âˆ'), the salts of
sulfuric acid. Sulfuric acid also reacts with SO
3 in equimolar ratios to form
pyrosulfuric acid (H
2S
2O
7).
*
Thiosulfates (sometimes referred to as thiosulfites or "hyposulfites") (S
2O
32âˆ'). Thiosulfates are used in photographic fixing (HYPO) as reducing agents. Ammonium thiosulfate is being investigated as a
cyanide replacement in leaching
gold.[
2]
*
Sodium dithionite, Na
2S
2O
4, is the highly reducing dianion derived from hyposulfurous/dithionous acid.
*
Sodium dithionate (Na
2S
2O
6).
*
Polythionic acids (H
2S
nO
6), where
n can range from 3 to 80.
*
Peroxymonosulfuric acid (H
2SO
5) and
peroxydisulfuric acids (H
2S
2O
8), made from the action of SO
3 on concentrated
H2O2, and
H2SO4 on concentrated H
2O
2 respectively.
*
Sodium polysulfides (Na
2S
x)
*
Sulfur hexafluoride, SF
6, 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 S
4N
4 is an example.
*
Thiocyanates contain the SCN
âˆ' group. Oxidation of thiocyanoate gives
thiocyanogen, (SCN)
2 with the connectivity NCS-SCN.
Organic sulfur compounds (where R, R', and R'' are organic groups such as CH
3):
*
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; (CH
3)
2S
+CH
2CH
2COO
âˆ') 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 s have the form R-S
-. Such anions arise upon treatment of
thiols with base.
*
Sulfoxides have the form
R-S(=O)-
R′. A common sulfoxide is
DMSO.
*
Sulfones have the form
R-S(=O)
2-
R′. A common sulfone is sulfolane C
4H
8SO
2.
See also Category: sulfur compounds and organosulfur chemistrySulfur 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.
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 δS-34 values of co-genetic minerals. The differences between minerals can be used to estimate the temperature of equilibration. The δ
C-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.
Carbon disulfide, carbon oxysulfide, 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, at high concentrations it reacts with moisture to form
sulfurous acid 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.
Hydrogen sulfide is quite
toxic (more toxic than
cyanide). Although very pungent 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,
Ireland,
Hong Kong and
India, but
sulfur in the
United States, while both spellings are used in
Australia,
Canada and
New Zealand. The
IUPAC adopted the spelling "sulfur" in 1990, as did the
Royal Society of Chemistry Nomenclature Committee in 1992. This spelling has begun to replace its variant in official use, unlike
aluminum, which is not commonly used outside
North America, and which
IUPAC rejected in 1990 in favor of
aluminium.
*
Sulfur cycle*
Disulfide bond*
Sulfonium S
+, S
+R
3*
Los Alamos National Laboratory – Sulfur* R. Steudel (ed.): Elemental Sulfur and Sulfur-Rich Compounds (part I & II), Topics in Current Chemistry Vol. 230 & 231, Springer, Berlin 2003.
*
Sulfur phase diagram.*
WebElements.com – Sulfur*
chemicalelements.com/sulfur
