|
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 (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 ×10-3 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 in
British English) 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 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, S
8.
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
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 (
H2S
O4) 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 (H
2S) 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 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.
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 (pK
a1 = 7.00, pK
a2 = 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] (
http://home.earthlink.net/~lenyr/iposc.htm)
Galena, a naturally occurring lead sulfide (as the detector in a "cat's hair" rectifier) was of course the original
semiconductor discovered.
Some 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] (http://doccopper.tripod.com/gold/AltLixiv.html)
- 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.
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 concentration this element can kill quickly by preventing
respiration. Sulfur quickly deadens the sense of smell so potential victims may be unaware of its presence.
Sulfur is traditionally spelled "sulphur" in
British English, but
IUPAC has adopted the spelling "sulfur", as has the
Royal Society of Chemistry Nomenclature Committee[?]. Increasingly "sulfur" is being used in British English instead.
See also: sulfur cycle[?], disulfide bond