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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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