Sulfur or
sulphur ( ,
see spelling below) is the
chemical element that has the
atomic number 16. It is denoted with
the symbol
S. It is an abundant,
multivalent non-metal. Sulfur, in its native form, is a bright
yellow
crystalline solid. In
nature, it can be found as the pure element and as
sulfide and
sulfate
minerals. It is an
essential
element for life and is found in two
amino acids:
cysteine and
methionine. Its commercial uses are
primarily in
fertilizers, but it is also
widely used in black
gunpowder,
matches,
insecticides and
fungicides. Elemental sulfur crystals are
commonly sought after by mineral collectors for their brightly
colored
polyhedron shapes. In
nonscientific contexts, it can also be referred to as
brimstone.
History

Rough sulfur crystal
Sulfur (
Sanskrit,
sulvari;
Latin Sulphurium) was known in
ancient times and is referred to in the
Torah
(
Genesis).
English translations
of the Bible commonly referred to burning 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 await the
unbelieving and unrepentant. It is from this part of the Bible that
Hell is implied to "smell of sulfur" (likely
due to its association with volcanic activity), although sulfur, in
itself, is in fact odorless. The "smell of sulfur" usually refers
to either the odor of
hydrogen
sulfide, e.g. from rotten egg, or of burning sulfur, which
produces
sulfur dioxide, the smell
associated with burnt matches. The smell emanating from raw sulfur
originates from a slow oxidation in the presence of air. Hydrogen
sulfide is the principal odor of untreated
sewage and is one of several unpleasant smelling
sulfur-containing components of
flatulence (along with sulfur-containing
mercaptans).
According
to the Ebers Papyrus, a sulfur
ointment was used in ancient Egypt
to treat
granular eyelids. Sulfur was used for fumigation in
preclassical Greece
; this is
mentioned in the Odyssey.
Pliny the Elder discusses sulfur in book 35
of his Natural
History, saying that its best-known source is the island
of Melos
. He
also mentions its use for fumigation, medicine, and bleaching
cloth.
A natural
form of sulfur known as shiliuhuang was known in China
since the 6th century BC and found in Hanzhong
. By
the 3rd century, the Chinese discovered that sulfur could be
extracted from
pyrite. Chinese Daoists were
interested in sulfur's flammability and its reactivity with certain
metals, yet its earliest practical uses were found in
traditional Chinese medicine. A
Song Dynasty military treatise of 1044
AD described different formulas for Chinese
black powder, which is a mixture of
potassium nitrate ( ),
charcoal, and sulfur. Early
alchemists gave sulfur its own
alchemical symbol which was a triangle at
the top of a cross.
In 1777,
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 developed.
Spelling and etymology
The
element has traditionally been spelled sulphur in the
United Kingdom (since the 14th century), most of the Commonwealth including India, Malaysia
, South
Africa, and Hong Kong, along with the rest of the Caribbean
and Ireland
, but sulfur in the United States, while
both spellings are used in Canada and the Philippines
. IUPAC adopted the
spelling “sulfur” in 1990, as did the
Royal Society of Chemistry
Nomenclature Committee in 1992. The
Qualifications and
Curriculum Authority for England and Wales recommended its use
in 2000.
In Latin, the word is variously written
sulpur,
sulphur, and
sulfur (the Oxford Latin Dictionary
lists the spellings in this order). It is an original Latin name
and not a
Classical Greek loan, so
the
ph variant does not denote the Greek letter φ. Sulfur
in Greek is
thion (θείον), whence comes the prefix
thio-. The simplification of the Latin words p
or ph to an f appears to have taken place towards the end of the
classical period.
Characteristics

Sulfur melts to a blood-red
liquid.
When burned, it emits a blue flame.
At room temperature, sulfur is a soft, bright-yellow solid.
Elemental sulfur has only a faint odor, similar to that of
matches.The odor associated with rotten eggs is due
to
hydrogen sulfide ( ) and organic
sulfur compounds rather than elemental sulfur.Sulfur burns with a
blue flame that emits
sulfur dioxide,
notable for its peculiar suffocating odor due to dissolving in the
mucosa to form dilute
sulfurous acid.
Sulfur itself is insoluble in water, but
soluble in
carbon
disulfide — and to a lesser extent in other non-polar
organic solvents such as
benzene and
toluene. 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.
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 of sulfur is that the
viscosity in its molten state, unlike most other
liquids, increases above temperatures of due to the formation of
polymers. The molten sulfur assumes a dark
red color above this temperature. At higher temperatures, however,
the viscosity is decreased as depolymerization occurs.
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.
Allotropes
Sulfur forms more than 30 solid
allotrope,
more than any other element. Besides S
8, several other
rings are known. Removing one atom from the crown gives
S
7, which is more deeply yellow than S
8.
HPLC analysis of "elemental sulfur" reveals an
equilibrium mixture of mainly S
8, but also S
7
and small amounts of S
6. Larger rings have been
prepared, including S
12 and S
18. By contrast,
sulfur's lighter neighbor
oxygen only exists
in two states of allotropic significance: O
2 and
O
3.
Selenium, the heavier
analogue of sulfur, can form rings but is more often found as a
polymer chain.
Isotopes
Sulfur has 25 known
isotopes, four of which
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
40argon 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.
Occurrence
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 mined in Indonesia
, Chile
, and
Japan. Sicily is also famous for its
sulfur mines. Sulfur deposits are polycrystalline, and the largest
documented single crystal measured
22×16×11 cm
3.
Significant deposits 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, although apparently native sulfur may be
produced by geological processes alone, without the aid of living
organisms (see below).
However, fossil-based sulfur deposits from
salt domes are the basis for commercial production in the United
States, Poland, Russia, Turkmenistan
, and Ukraine
.
Sulfur
production through hydrodesulfurization of oil, gas, and
the Athabasca
Oil Sands
has produced a surplus — huge stockpiles of
sulfur now exist throughout Alberta, Canada.
Common naturally occurring sulfur compounds include the
sulfide minerals, such as
pyrite (iron sulfide),
cinnabar (mercury sulfide),
galena (lead sulfide),
sphalerite (zinc sulfide) and
stibnite (antimony sulfide); and the sulfates, such
as gypsum (calcium sulfate),
alunite
(potassium aluminium sulfate), and
barite
(barium sulfate). 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 present in many types of
meteorites. Ordinary chondrites contain on average
2.1% sulfur, and carbonaceous chondrites may contain as much as
6.6%. Sulfur in meteorites is normally present as
troilite (FeS), but other sulfides are found in
some meteorites, and carbonaceous chondrites contain free sulfur,
sulfates, and possibly other sulfur compounds.
Extraction and production
Extraction from natural resources
Sulfur is extracted by mainly two processes: the Sicilian process
and the
Frasch process. The Sicilian
process, which was first used in
Sicily, was
used in ancient times to get sulfur from rocks present in volcanic
regions. In this process, the sulfur deposits are piled and stacked
in brick kilns built on sloping hillsides, and with airspaces
between them. Then powdered sulfur is put on top of the sulfur
deposit and ignited. As the sulfur burns, the heat melts the sulfur
deposits, causing the molten sulfur to flow down the sloping
hillside. The molten sulfur can then be collected in wooden
buckets.
The second process used to obtain sulfur is the Frasch process. In
this method, three concentric pipes are used: the outermost pipe
contains superheated water, which melts the sulfur, and the
innermost pipe is filled with hot compressed air, which serves to
create foam and pressure. The resulting sulfur foam is then
expelled through the middle pipe.
The Frasch process produces sulfur with a 99.5% purity content, and
which needs no further purification. The sulfur produced by the
Sicilian process must be purified by distillation.
Production from hydrogen sulfide
Chemically
The
Claus process is used to extract
elemental sulfur from
hydrogen
sulfide produced in
hydrodesulfurization of petroleum or
from
natural gas.
Biologically
In the biological route, hydrogen sulfide (H
2S) from
natural gas or refinery gas is absorbed with a slight alkaline
solution in a
wet scrubber, or the
sulfide is produced by biological sulfate reduction. In the
subsequent process step, the dissolved sulfide is biologically
converted to elemental sulfur. This solid sulfur is removed from
the reactor. This process has been built on commercial scale. The
main advantages of this process are:
- no use of expensive chemicals,
- the process is safe as the H2S is directly absorbed
in an alkaline solution,
- no production of a polluted waste stream,
- re-usable sulfur is produced, and
- the process occurs under ambient conditions.
The biosulfur product is different from other processes in which
sulfur is produced because the sulfur is hydrophilic. Next to
straightforward reuses as source for sulfuric acid production, it
can also be applied as sulfur fertilizer.
Chemistry
Inorganic compounds
When 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. Pyrite can show
semiconductor properties.
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.
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 useful 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.

The
sulfate
anion,
- Sulfides (S2−), a complex
family of compounds usually derived from S2−. Cadmium sulfide (CdS) is an example.
- Sulfites ( ), the salts of sulfurous acid (H2SO3)
which is generated by dissolving SO2 in water. Sulfurous
acid and the corresponding sulfites are fairly strong reducing
agents. Other compounds derived from SO2 include the
pyrosulfite or metabisulfite ion ( ).
- Sulfates ( ), the salts of sulfuric acid. Sulfuric acid also reacts with
SO3 in equimolar ratios to form pyrosulfuric acid
(H2S2O7).
- Thiosulfates ( ). Sometimes
referred as thiosulfites or "hyposulfites", Thiosulfates are used
in photographic fixing (HYPO) as reducing agents. Ammonium
thiosulfate is being investigated as a cyanide replacement in leaching gold.[4476]
- Sodium dithionite, , is the
highly reducing dianion derived from hyposulfurous/dithionous
acid.
- Sodium dithionate
(Na2S2O6).
- Polythionic acids
(H2SnO6), where
n can range from 3 to 80.
-
Peroxymonosulfuric acid
(H2SO5) and peroxydisulfuric acids
(H2S2O8), made from the action of
SO3 on concentrated H2O2, and H2SO4 on concentrated
H2O2 respectively.
- Sodium polysulfides
(Na2Sx)
- Sulfur hexafluoride,
SF6, 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 S4N4 is an example.
- Thiocyanates contain the
SCN− group. Oxidation of thiocyanate gives thiocyanogen, (SCN)2 with the
connectivity NCS-SCN.
Organic compounds
Many of the unpleasant odors of organic matter are based on
sulfur-containing compounds such as
methyl
mercaptan and dimethyl sulfide. Thiols and sulfides are used in
the odoration of natural gas, notably, 2-methyl-2-propanethiol
(t-butyl mercaptan). The odor of
garlic and
"
skunk stink" are also caused by
sulfur-containing organic compounds. Not all organic sulfur
compounds smell unpleasant; for example,
grapefruit mercaptan, a
sulfur-containing
monoterpenoid is
responsible for the characteristic scent of grapefruit. It should
be noted that this thiol is present in very low concentrations. In
larger concentrations, the odor of this compound is that typical of
all thiols, unpleasant.
Sulfur-containing organic compounds include the following (R, R',
and R
are organic groups such as CH3):
- 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;
(CH3)2S+CH2CH2COO−)
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 have the form
R-S-. Such anions arise upon treatment of thiols with base.
- Sulfoxides have the form
R-S(=O)-R′. The simplest sulfoxide, DMSO, is a common solvent.
- Sulfones have the form
R-S(=O)2-R′. A common sulfone is
sulfolane C4H8SO2.
See also Category: sulfur
compounds and organosulfur
chemistry
Applications
One of the direct uses of sulfur is in
vulcanization of rubber, where
polysulfides crosslink organic polymers. Sulfur
is a component of
gunpowder. It reacts
directly with methane to give
carbon
disulfide, which is used to manufacture
cellophane and
rayon.
Elemental sulfur is mainly used as a precursor to other chemicals.
Approximately 85% (1989) is converted to
sulfuric acid (
H2S
O4),
which is of such prime importance to the
world's economies that the production and
consumption of sulfuric acid is an indicator of a nation's
industrial development. For example with 36.1 million metric tons
in 2007, more sulfuric acid is produced in the United States every
year than any other inorganic industrial chemical. The principal
use for the acid is the extraction of phosphate ores for the
production of fertilizer manufacturing. Other applications of
sulfuric acid include oil refining, wastewater processing, and
mineral extraction.
Sulfur compounds are also used in
detergents,
fungicides,
dyestuffs, and agrichemicals. In
silver-based
photography sodium and
ammonium
thiosulfate are used as
"fixing agents."
Sulfur is an ingredient in some
acne
treatments.
An increasing application is as fertilizer. Standard sulfur is
hydrophobic and therefore has to be covered with a surfactant by
bacteria in the ground before it can be oxidized to sulfate. This
makes it a slow release fertilizer, which cannot be taken up by the
plants instantly, but has to be oxidized to sulfate over the growth
season. Sulfur also improves the use efficiency of other essential
plant nutrients, particularly nitrogen and phosphorus. Biologically
produced sulfur particles are naturally hydrophilic due to a
biopolymer coating. This sulfur is therefore easier to disperse
over the land (via spraying as a diluted slurry), and results in a
faster release.
Sulfites, derived from burning sulfur, are
heavily used to
bleach paper. They are also used as preservatives in dried
fruit.
Magnesium sulfate, better known as
Epsom salts, can be used as a
laxative, a bath additive, an
exfoliant, a
magnesium
supplement for plants, or a
desiccant.
Specialized applications
Sulfur is used as a light-generating medium in the rare lighting
fixtures known as
sulfur lamps.
Historical applications
In the late 18th century,
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.
Molten sulfur is sometimes still used for setting steel bolts into
drilled concrete holes where high shock resistance is desired for
floor-mounted equipment attachment points. Pure powdered sulfur was
also used as a medicinal tonic and laxative. Sulfur was also used
in baths for people who had seizures.
Fungicide and pesticide
Sulfur is one of the oldest fungicides and pesticides. Dusting
sulfur, elemental sulfur in powdered form, is a common fungicide
for grapes, strawberry, many vegetables and several other crops. It
has a good efficacy against a wide range of powdery mildew diseases
as well as black spot. In organic production, sulfur is the most
important fungicide. It is the only fungicide used in
organic farmed apple production against
the main disease
apple scab under colder
conditions. Biosulfur (biologically produced elemental sulfur with
hydrophilic characteristics) can be used well for these
applications.
Standard-formulation dusting sulfur is applied to crops with a
sulfur duster or from a dusting plane. Wettable sulfur is the
commercial name for dusting sulfur formulated with additional
ingredients to make it water soluble. It has similar applications,
and is used as a
fungicide against
mildew and other mold-related problems with
plants and soil.
Sulfur is also used as an "
organic"
(i.e. "green")
insecticide (actually an
acaricide) against
ticks and
mites. A common method
of use is to dust clothing or limbs with sulfur powder. Some
livestock owners set out a sulfur salt
block as a
salt lick.
Biological role
See
sulfur cycle for more on the
inorganic and organic natural transformations of sulfur.
Sulfur is an essential component of all living
cells.
Inorganic sulfur forms a part of
iron-sulfur clusters, and sulfur is the
bridging ligand in the
CuA site of
cytochrome c oxidase, a basic
substance involved in utilization of oxygen by all aerobic
life.
Sulfur may also serve as chemical food source for some primitive
organisms: some forms of
bacteria use
hydrogen sulfide (H
2S)
in the place of water as the
electron donor
in a primitive
photosynthesis-like
process in which oxygen is the electron receptor. The
photosynthetic green and purple sulfur
bacteria and some
chemolithotrophs use elemental oxygen to
carry out such oxidization of hydrogen sulfide to produce elemental
sulfur (S
o), oxidation state = 0. Primitive bacteria
which live around deep ocean volcanic vents oxidize hydrogen
sulfide in this way with oxygen: see
giant tube worm for an example of large
organisms (via bacteria) making metabolic use of hydrogen sulfide
as food to be oxidized.
The so-called
sulfur bacteria, by
contrast, "breathe sulfate" instead of oxygen. They use sulfur as
the electron acceptor, and reduce various oxidized sulfur compounds
back into sulfide, often into hydrogen sulfide. They also can grow
on a number of other partially oxidized sulfur compounds
(e. g. thiosulfates, thionates, polysulfides, sulfites). The
hydrogen sulfide produced by these bacteria is responsible for the
smell of some intestinal gases and decomposition products.
Sulfur is a part of many bacterial defense molecules. For example,
though sulfur is not a part of the
lactam
ring, it is a part of most
beta lactam
antibiotics, including the
penicillins,
cephalosporins, and
monobactams.
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 (see
sulfur
assimilation for details of this process).
Sulfur is regarded as secondary nutrient although plant
requirements for sulfur are equal to and sometimes exceed those for
phosphorus. However sulfur is recognized as one of the major
nutrients essential for plant growth, root nodule formation of
legumes and plants protection mechanisms. Sulfur deficiency has
become widespread in many countries in Europe. Because atmospheric
inputs of sulfur will continue to decrease, the deficit in the
sulfur input/output is likely to increase, unless sulfur
fertilizers are used.
In
plants and
animals
the
amino acids cysteine and
methionine
contain sulfur, as do all
polypeptides,
proteins, and
enzymes
which contain these amino acids.
Homocysteine and
taurine
are other sulfur-containing acids which are similar in structure,
but which are not coded for by
DNA, and are not
part of the
primary structure of
proteins.
Glutathione is an important
sulfur-containing tripeptide which plays a role in cells as a
source of chemical reduction potential in the cell, through its
sulfhydryl (-SH) moiety. Many important cellular enzymes use
prosthetic groups ending with -SH moieties to handle reactions
involving acyl-containing biochemicals: two common examples from
basic metabolism are
coenzyme A and
alpha-lipoic acid.
Disulfide bonds (S-S bonds) formed
between cysteine residues in peptide chains are very important in
protein assembly and structure. These strong covalent bonds between
peptide chains give proteins a great deal of extra toughness and
resiliency. For example, the high strength of feathers and hair is
in part due to their high content of S-S bonds and their high
content of cysteine and sulfur (eggs are high in sulfur because
large amounts of the element are necessary for feather formation).
The high disulfide content of hair and feathers contributes to
their indigestibility, and also their odor when burned.
Traditional medical role for elemental sulfur
In traditional medical skin treatment which predates modern era of
scientific medicine, elemental sulfur has been used mainly as part
of creams to alleviate various conditions such as psoriasis, eczema
and acne. The mechanism of action is not known, although elemental
sulfur does oxidize slowly to sulfurous acid, which in turn
(through the action of
sulfite) acts as a
mild reducing and antibacterial agent.
Precautions
Elemental sulfur is non-toxic, but it can burn as an oxidizer or a
reducing agent, producing combustion products that are toxic, such
as
carbon disulfide,
carbon oxysulfide,
hydrogen sulfide, and
sulfur dioxide.
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 organisms without lungs such
as insects or plants, it otherwise prevents
respiration.
Hydrogen sulfide is
toxic. Although very pungent at first, it quickly
deadens the sense of smell, so potential victims may be unaware of
its presence until death or other symptoms occur.
Sulfur trioxide, a volatile liquid
at standard temperature and pressure, is extremely dangerous,
especially in contact with water, which reacts with it to form
sulfuric acid with the generation of
much heat. Sulfuric acid poses extreme hazards to many objects and
substances.
Environmental impact
The burning of
coal and/or
petroleum by industry and
power plants generates
sulfur dioxide (S
O2), which reacts with atmospheric water
and oxygen to produce
sulfuric acid
(H
2SO
4). This sulfuric acid is a component of
acid rain, which lowers the
pH of
soil and freshwater bodies,
sometimes resulting in substantial damage to the
environment and
chemical weathering of statues and
structures. 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. In coal fired
power plants, the flue gases are sometimes purified. In more modern
power plants that use
syngas the
sulfur is extracted before the gas is burned.
See also
References
- p. 242, Archaeomineralogy, George Rapp, 2nd ed.,
Springer: 2009, ISBN 978-3-540-78593-4.
- Odyssey, book 22, lines 480–495.
- pp. 247–249, Pliny the Elder on science and
technology, John F. Healy, Oxford University Press, 1999, ISBN
0198146876.
- http://www.rod.beavon.clara.net/sulphur.htm, retrieved 2nd
April 2009 18:29 GMT.
- Spelling of Sulfur (PDF)
- Worldwidewords, 9 December 2000.
- Vanderkrogt.net.
- Kelly DP (1995) Sulfur and its Doppelgänger. Arch.
Microbiol. 163: 157-158.
- Sulfuric Acid Growth
- Sulfur as a fertilizer
External links