MSA University, Faculty of Pharmacy Analytical Department
NAME:MOHAMED ADELAFIFI ABDO ID:142311
FOR DR:MAHA KAMAL .
Sulfur-containing inorganic anions are commonly encountered in soil
sediments, hot springs, and lake waters. They are also frequently used in
many industrial applications. Their speciation and quantitation are very
important to the understanding of those processes. However, due to their
redox chemistry, many sulfur-containing species can readily react with each
other, decompose over time, or be oxidized in the presence of air.They are
also very sensitive to the solution pH, which would affect the distribution of the
species over time. All these make the determination of these sulfur-containing
anions a very challenging task. Currently, both wet chemistry and ion
chromatography are used for the sulfur speciation.In this poster, we will focus
on the analysis of sulfur-containing species in polysulfide matrix with
suppressed ion chromatography.The presence of large amount of polysulfide
complicates the determinationof sulfur-containing species. Various sample
pretreatment techniques were evaluated, and an optimized pretreatment
method will be shown for improving the analysis of sulfur-containing anions in
polysulfide matrix by suppressed ion chromatography.
Lime sulfur is a mixture of calcium polysuflide formed through the reaction
of calcium hydroxide and sulfur, and is an aqueous solution with
Lime sulfur is used as a spray to control fungi, bacteria, and insects. It
is also used for remediation of hexavalent chromium in chromite ore
processing residue (COPR).
• Typically lime sulfur solutions are strong alkaline
• When acified, it can produce highly toxic hydrogen sulfide
Sn2– + H+ →H2S + (n-1)/8 S8
• It can also react with sulfite to form thiosulfate
Sn2– + SO32– →(n-1) S2O32– + S2–
Table 2 lists the possible inorganic sulfur-containing anions present in
lime sulfur-based aqueous streams or solutions.
Sulfur forms polyatomic molecules with different chemical formulas, with the best-known
allotrope being octasulfur, cyclo-S8. Octasulfur is a soft, bright-yellow solid with only a faint odor,
similar to that of matches.
It melts at 115.21 °C, boils at 444.6 °C and sublimes easily.
°C, below its melting temperature, cyclo-octasulfur changes from α-octasulfur to the β-polymorph.
The structure of the S8 ring is virtually unchanged by this phase change, which affects the
intermolecular interactions. Between its melting and boiling temperatures, octasulfur changes its
allotrope again, turning from β-octasulfur to γ-sulfur, again accompanied by a lower density but
increased viscosity due to the formation of polymers.
At even higher temperatures, however, the viscosity decreases as depolymerization occurs. Molten
sulfur assumes a dark red color above 200 °C. The density of sulfur is about 2 g·cm−3
on the allotrope; all of its stable allotropes are excellent electrical insulators.
Sulfur burns with a blue flame concomitant with formation of sulfur dioxide, notable for its
peculiar suffocating odor. Sulfur is insoluble in water but soluble in carbon disulfide and, to a
lesser extent, in other nonpolar organic solvents, such as benzene and toluene. The first and the
second ionization energies of sulfur are 999.6 and 2252 kJ·mol−1
, respectively. Despite such
figures, the +2 oxidation state is rare, with +4 and +6 being more common. The fourth and sixth
ionization energies are 4556 and 8495.8 kJ·mol−1
, the magnitude of the figures caused by electron
transfer between orbitals; these states are only stable with strong oxidants as fluorine, oxygen, and
Sulfur may be found by itself and historically was usually obtained in this way, while pyrite has
been a source of sulfur via sulfuric acid.
In volcanic regions in Sicily, in ancient times, it was
found on the surface of the Earth, and the "Sicilian process" was used: sulfur deposits were piled
and stacked in brick kilns built on sloping hillsides, with airspaces between them. Then, some
sulfur was pulverized, spread over the stacked ore and ignited, causing the free sulfur to melt
down the hills. Eventually the surface-borne deposits played out, and miners excavated veins that
ultimately dotted the Sicilian landscape with labyrinthine mines. Mining was unmechanized and
labor-intensive, with pickmen freeing the ore from the rock, and mine-boys or carusi carrying
baskets of ore to the surface, often through a mile or more of tunnels. Once the ore was at the
surface, it was reduced and extracted in smelting ovens. The conditions in Sicilian sulfur mines
were horrific, prompting Booker T. Washington to write "I am not prepared just now to say to
what extent I believe in a physical hell in the next world, but a sulphur mine in Sicily is about the
nearest thing to hell that I expect to see in this life.".
Today's sulfur production is as a side product of other industrial processes such as oil refining; in
these processes, sulfur often occurs as undesired or detrimental compounds that are extracted and
converted to elemental sulfur. As a mineral, native sulfur under salt domes is thought to be a fossil
mineral resource, produced by the action of ancient bacteria on sulfate deposits. It was removed
from such salt-dome mines mainly by the Frasch process.
In this method, superheated water
was pumped into a native sulfur deposit to melt the sulfur, and
then compressed air returned the 99.5% pure melted product to the surface. Throughout the 20th
century this procedure produced elemental sulfur that required no further purification. However,
due to a limited number of such sulfur deposits and the high cost of working them, this process for
mining sulfur has not been employed in a major way anywhere in the world since 2002.
Today, sulfur is produced from petroleum, natural gas, and related fossil resources, from which it
is obtained mainly as hydrogen sulfide.Organosulfur compounds, undesirable impurities in
petroleum, may be upgraded by subjecting them to hydrodesulfurization, which cleaves the C–S
R-S-R + 2 H2 → 2 RH + H2S
The resulting hydrogen sulfide from this process, and also as it occurs in natural gas, is converted
into elemental sulfur by the Claus process. This process entails oxidation of some hydrogen
sulfide to sulfur dioxide and then the comproportionation of the two:
3 O2 + 2 H2S → 2 SO2 + 2 H2O
SO2 + 2 H2S → 3 S + 2 H2O
Owing to the high sulfur content of the Athabasca Oil Sands, stockpiles of elemental sulfur from
this process now exist throughout Alberta,
Another way of storing sulfur is as a binder for concrete, the resulting product having
many desirable properties (see sulfur concrete).
The world production of sulfur in 2011
amounted to 69 million tonnes (Mt), with more than 15 countries contributing more than 1 Mt
each. Countries producing more than 5 Mt are China (9.6), US (8.8), Canada (7.1) and Russia
While the production has been slowly increasing from 1900 to 2010, the price was much
less stable, especially in the 1980s and around 20.
Common oxidation states of sulfur range from −2 to +6. Sulfur forms stable compounds with all
elements except the noble gases.
Treatment of sulfur with hydrogen gives hydrogen sulfide. When dissolved in water, hydrogen
sulfide is mildly acidic:
H2S ـــــــ> HS–
Hydrogen sulfide gas and the hydrosulfide anion are extremely toxic to mammals, due to their
inhibition of the oxygen-carrying capacity of hemoglobin and certain cytochromes in a manner
analogous to cyanide and azide (see below, under precautions).Reduction of elemental sulfur gives
which consist of chains of sulfur atoms terminated with S–
2 Na + S8 → Na2S8
This reaction highlights arguably the single most distinctive property of sulfur: its ability to
catenate (bind to itself by formation of chains). Protonation of these polysulfide anions gives the
polysulfanes, H2Sx where x = 2, 3, and 4.
Ultimately reduction of sulfur gives sulfide salts:
16 Na + S8 → 8 Na2S
The interconversion of these species is exploited in the sodium-sulfur battery. The radical anion
gives the blue color of the mineral lapis lazuli With very strong oxidants, S8 can be oxidized,
for example, to give bicyclic S8
Elemental sulfur is mainly used as a precursor to other chemical Approximately 85%
(1989) is converted to sulfuric acid (H2SO4): 2 S + 3 O2 + 2 H2O → 2 H2SO4
With sulfuric acid being of central importance to the world's economies,its production and
consumption is an indicator of a nation's industrial development.
For example with 32.5 million
tonnes in 2010,
the United States produces more sulfuric acid every year than any other inorganic industrial
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 is increasingly used as a component of fertilizers. The most important form of sulfur for
fertilizer is the mineral calcium sulfate. Elemental sulfur is hydrophobic (that is, it is not soluble in
water) and, therefore, cannot be directly utilized by plants. Over time, soil bacteria can convert it
to soluble derivatives, which can then be utilized by plants. Sulfur improves the use efficiency of
other essential plant nutrients, particularly nitrogen and phosphorus.
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.
Plant requirements for sulfur are equal to or exceed those for phosphorus. It is 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 continue to decrease, the deficit in the sulfur input/output is
likely to increase, unless sulfur fertilizers are used.
Oxides and oxyanions
The principal sulfur oxides are obtained by burning sulfur:
S + O2 → SO2
2 SO2 + O2 → 2 SO3
Other oxides are known, e.g. sulfur monoxide and disulfur mono- and dioxides, but they are
The sulfur oxides form numerous oxyanions with the formula SOn
. Sulfur dioxide and sulfites
(SO2−3) are related to the unstable sulfurous acid (H2SO3). Sulfur trioxide and sulfates (SO2−4)
are related to sulfuric acid. Sulfuric acid and SO3combine to give oleum, a solution of pyrosulfuric
acid (H2S2O7) in sulfuric acid.
Peroxides convert sulfur into unstable oxides such as S8O, a sulfoxide. Peroxymonosulfuric acid
(H2SO5) and peroxydisulfuric acids (H2S2O8), made from the action of SO3 on concentrated H2O2,
and H2SO4 on concentrated H2O2 respectively.Thiosulfate salts (S2O2−3), sometimes referred as
used in photographic fixing (HYPO) and as reducing agents, feature sulfur in two oxidation states.
Sodium dithionite (Na2S2O4), contains the more highly reducing dithionite anion (S
2O2−4).Sodium dithionate (Na2S2O6) contains the dithionate anion (S2O6
) and is the first member
of the polythionic acids(H2SnO6), where n can range from 3 to many. Thiosulfurous acid (HS-
S(=O)-OH) is formed in trace amounts when hydrogen sulfide and sulfur dioxide gases are mixed
at room temperature, but its salts (thiosulfites) are unknown.
Halides and oxyhalides
The two main sulfur fluorides are sulfur hexafluoride, a dense gas used as nonreactive and
nontoxic propellant, and sulfur tetrafluoride, a rarely used organic reagent that is highly toxic.
Their chlorinated analogs are sulfur dichloride and sulfur monochloride. Sulfuryl chloride and
chlorosulfuric acid are derivatives of sulfuric acid; thionyl chloride (SOCl2) is a common reagent
in organic synthesis.
The principal ores of copper, zinc, nickel, cobalt, molybdenum, and other metals are sulfides.
These materials tend to be dark-colored semiconductors that are not readily attacked by water or
even many acids. They are formed, both geochemically and in the laboratory, by the reaction of
hydrogen sulfide with metal salts. The mineral galena (PbS) was the first demonstrated
semiconductor and found a use as a signal rectifier in the cat's whiskers of early crystal radios. The
iron sulfide called pyrite, the so-called "fool's gold," has the formula FeS2.
The upgrading of
these ores, usually by roasting, is costly and environmentally hazardous. Sulfur corrodes many
metals via the process called tarnishing
Some of the main classes of sulfur-containing organic compounds include the following:Thiols or
mercaptans (as they are mercury capturers as chelators) are the sulfur analogs of alcohols;
treatment of thiols with base gives thiolate ions.Thioethers are the sulfur analogs of
ethers.Sulfonium ions have three groups attached to a cationic sulfur center.
Dimethylsulfoniopropionate (DMSP) is one such compound, important in the marine
organicsulfur cycle.Sulfoxides and sulfones are thioethers with one and two oxygen atoms
attached to the sulfur atom, respectively. The simplest sulfoxide, dimethyl sulfoxide, is a common
solvent; a common sulfone is sulfolane.Sulfonic acids are used in many detergents.Compounds
with carbon–sulfur bonds are uncommon with the notable exception of carbon disulfide, a volatile
colorless liquid that is structurally similar to carbon dioxide.It is used as a reagent to make the
polymer rayon and many organosulfur compounds. Unlike carbon monoxide, carbon monosulfide
is only stable as a dilute gas,
as in the interstellar medium.Organosulfur compounds are responsible for some of the unpleasant
odors of decaying organic matter. They are used in the odoration of natural gas and cause the odor
of garlic and skunk spray. Not all organic sulfur compounds smell unpleasant at all concentrations:
the sulfur-containing monoterpenoid grapefruit mercaptan in small concentrations is responsible
for the characteristic scent of grapefruit, but has a generic thiol odor at larger concentrations.
Sulfur mustard, a potent vesicant, was used in World War I as a disabling agent.Sulfur-sulfur
bonds are a structural component to stiffen rubber, in a way similar to the biological role of
disulfide bridges to rigidify proteins (see biological below). In the most common type of industrial
"curing" or hardening and strengthening of natural rubber, elemental sulfur is heated with the
rubber to the point that chemical reactions form disulfide bridges between isoprene units of the
polymer. This process, patented in 1843, allowed rubber to become a major industrial product,
especially automobile tires. Because of the heat and sulfur, the process was named vulcanization,
after the Roman god of the forge and volcanism.
Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and
Jump up^ "Sulfur History". Georgiagulfsulfur.com. Retrieved 2008-09-12.
^ Jump up to:a b c
Greenwood, N. N.; & Earnshaw, A. (1997).Chemistry of the Elements (2nd
ed.), Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.
Jump up^ Online Etymology Dictionary entry for sulfur. Retrieved 2011-08-18.
^ Jump up to:a b
Michie, C. A.; Langslow, D. R. (1988). "Sulphur or sulfur? A tale of two
spellings".British Medical Journal 297 (6664): 1697–1699. doi:10.1136/bmj.297.6664.1697.
Jump up^ McNaught, Alan (1991). "Journal style update". The Analyst 116 (11):
Jump up^ Sulphur, Worldwidewords
Jump up^ "General Certificate of Secondary Education (Science A Chemistry". 2011-02-27.