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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 |
| SI units
& STP are used except where noted. |
Sulfur (sulphur) is a chemical element
in the periodic table that has the symbol 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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