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The Element
Sulfur
Sulfur
| General |
| 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
|
| Atomic properties |
| Atomic weight |
32.065 amu |
| 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) |
±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 (832.5 °F) |
| Molar volume |
15.53 ×10-6
m3/mol |
| Heat of vaporization
|
no data |
| Heat of fusion |
1.7175 kJ/mol |
| Vapor pressure |
2.65 E-20 Pa at 388 K |
| Speed of sound |
__ m/s at 293.15 K |
| Miscellaneous |
| Electronegativity
|
2.58 (Pauling scale) |
| Specific heat
capacity |
710 J/(kg*K) |
| Electrical conductivity
|
5.0 E-22 106/m ohm |
| Thermal conductivity
|
0.269 W/(m*K) |
| 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 |
|
|
| SI units & STP
are used except where noted. |
Sulfur (sulphur) is a chemical element in the
periodic table that has
the symburol S and atomic number 16. An abundant
tasteless odorless multivalent non-metal, sulfur is best known
as yellow crystals and occurs in many sulfide and sulfate minerals
and even in its native form (especially in volcanic regions). It is an essential
element in all living organisms and is needed in several amino acids and hence in many
proteins. It is primarily used
in fertilizers but is also widely
used in gunpowder, laxatives, matches
and insecticides.
This non-metal is pale yellow in appearance, soft, light, with a distinct
odor when allied with hydrogen (rotten egg smell).
It burns with a blue flame that emits a peculiar suffocating odor
(sulfur dioxide, SO2).
Sulfur is insoluble in water but soluble in
carbon disulfide. Common
oxidation states of sulfur
include -2, +2, +4 and +6. In all states, solid, liquid, and gaseous,
sulfur has allotropic forms, whose relationships
are not completely understood. Crystalline sulfur can be shown to
form an 8 membered sulfur ring, S8.
Sulfur can be obtained in two crystalline modifications, in orthorhombic
octahedra, or in monoclinic prisms, the former of which is the more
stable at ordinary temperatures.
It is used for many industrial processes such as the production of
sulfuric acid (H2SO4) for
batteries, the production
of gunpowder, and the vulcanization of rubber. Sulfur is used as a fungicide, and in the manufacture
of phosphate fertilizers.
Sulfites are used to bleach papers and dried fruits.
Sulfur also finds use in matches and fireworks. Sodium
or ammonium thiosulfate are used as photographic fixing agents. Epsom salts, magnesium sulfate, can be used
as a laxative, as a bath additive,
as an exfoliant, or a magnesium supplement in plant
nutrition.
The amino acids
cysteine, methionine,
homocysteine, and taurine contain sulfur, as do
some common enzymes, making sulfur a necessary
component of all living cells. Disulfide bonds between
polypeptides are very important
in protein assembly and structure.
Some forms of bacteria use hydrogen sulfide (H2S)
in the place of water as the electron doner in a primitive
photosynthesis-like process.
Sulfur is absorbed by plants from soil as 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 (Sanskrit, sulvere; Latin sulpur)
was known in ancient times and was called brimstone in the Biblical story of Pentateuch
(Genesis). Homer mentioned "pest-averting sulfur"
in the 9th century BC and in 424 BC,
the tribe of Bootier 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. Through
experimentation, alchemists knew that the element mercury can be combined
with sulfur. In the late 1770s, Antoine Lavoisier helped
convince the scientific community that sulfur was an element and
not a compound.
Occurrence
Sulfur occurs naturally in large quantities compounded to other elements
in sulfides (example: pyrites) and sulfates
(example: gypsum). It is found in its free
form near hot springs and volcanic regions and in ores like
cinnabar, galena, sphalerite and stibnite. This element is also
found in small amounts in coal and petroleum, which
produce sulfur dioxide when burned.
Fuel standards increasingly require sulfur to be extracted from fossil fuels because sulfur
dioxide combines with water droplets to produce acid rain.
This extracted sulfur is then refined and represents a large portion
of sulfur production. It is also mined along the US Gulf coast, by
pumping hot water into sulfur containing deposits (such as salt domes)
which melts the sulfur. The molten sulfur is then pumped to the surface.
Through its major derivative, sulfuric acid, 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 industrial
and fertilizer complexes. Sulfuric
acid production is the major end use for sulfur, and consumption of
sulfuric acid has been regarded as one of the best indexes of a nation's
industrial development. More sulfuric acid is produced in the United
States every year than any other chemical.
The distinctive colors of Jupiter's volcanic moon Io, are from
various forms of multen, 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.
Many of the unpleasant odors of organic matter are based on sulfur-containing
compounds such as hydrogen sulfide, which
has the characteristic smell of rotten eggs. Dissolved in water, hydrogen
sulfide is acidic (pKa1 = 7.00, pKa2 = 12.92)
and will react with metals to form a series of metal sulfides. Natural
metal sulfides are found, especially those of iron. Iron sulfides
are called iron pyrites, the so called
fool's gold. Interestingly, pyrites can show semiconductor
properties.[1] Galena, a naturally
occurring lead sulfide (as the detector in a "cat's hair" rectifier)
was of course the original semiconductor discovered.
Polymeric sulfur nitride has metallic properties even though it
doesn't contain any metal
atoms. This compound also has unusual electrical and optical properties.
Amorphous or "plastic" sulfur is produced through fast cooling crystalline
sulfur. X-ray studies show that the amorphous
form may have an eight atom per spiral helical structure
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. Epsom salts are magnesium
sulfate.
- sulfuric acid reacting with SO3 in equimolar ratios
forms pyrosulfuric acid.
- peroxymonosulfuric acid and peroxydisulfuric acids, made from
the action of SO3 on concentrated H2O2,
and H2SO4
on concentrated H2O2 respectively.
- thiocyanogen, (SCN)2.
- tetrasulfur tetranitride S4N4.
- A thiol 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 sulfide is a molecule with the
form R-S-R', where R and R' are organic groups. These are the
sulfur equivalents of ethers.
Sulfur has 18 isotopes, of which four stable isotopes: S-32 (95.02%), S-33
(0.75%), S-34 (4.21%), and S-36 (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 contributes 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 hydologic 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. In addition to being quite
toxic (more toxic than cyanide), sulfur dioxide reacts
with atmospheric water to produce acid rain. In high atmospheric
concentration, it reacts with water in the lungs to form sulfuric
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. Although very
smelly in low concentrations, in higher concentrations sulfur quickly
deadens the sense of smell, so potential victims may be unaware of
its presence until they experience its possibly deadly effects.
The element is traditionally spelled sulfur in the US and Canada, but sulphur
in Britain, New Zealand,
and Australia. The IUPAC
has adopted the spelling "sulfur", as has the Royal Society of Chemistry
Nomenclature Committee.
Reference
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