Everything about Sulfur Dioxide totally explained
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Sulfur dioxide (also
sulphur dioxide) is the
chemical compound with the formula SO
2. This important gas is the main product from the combustion of
sulfur compounds and is of significant environmental concern. SO
2 is often described as the "smell of burning sulfur" but isn't responsible for the
smell of rotten eggs.
SO
2 is produced by
volcanoes and in various industrial processes. Since
coal and
petroleum often contain sulfur compounds, their combustion generates sulfur dioxide. Further oxidation of SO
2, usually in the presence of a catalyst such as
NO2, forms
H2SO4, and thus
acid rain. It can be easily liquefied in a standard home freezer.
Preparation
Sulfur dioxide can be prepared by burning
sulfur:
» S
8 + 8 O
2 → 8 SO
2
The combustion of
hydrogen sulfide and organosulfur compounds proceeds similarly.
» 2 H
2S(g) + 3 O
2(g) → 2 H
2O(g) + 2 SO
2(g)
The roasting of sulfide ores such as iron
pyrites,
sphalerite (zinc blende) and
cinnabar (mercury sulfide) also releases SO
2:
» 4
FeS
2(s) + 11 O
2(g) → 2 Fe
2O
3(s) + 8 SO
2(g)
2
ZnS(s) + 3 O
2(g) → 2 ZnO(s) + 2 SO
2(g)
» HgS(s) + O
2(g) → Hg(g) + SO
2(g)
Sulfur dioxide is a by-product in the manufacture of
cement:
CaSiO3 and
CaSO4 is heated with
coke and sand in this process:
» 2 CaSO
4(s) + 2SiO
2(s) + C(s) → 2 CaSiO
3(s) + 2 SO
2(g) + CO
2(g)
Action of hot
sulfuric acid on copper
turnings produces sulfur dioxide.
» Cu(s) + 2H
2SO
4(aq) → CuSO
4(aq) + SO
2(g) + 2H
2O(l)
Structure and bonding
SO
2 is a bent molecule with C
2v symmetry point group. In terms of
electron-counting formalisms, the sulfur atom has an
oxidation state of +4, a
formal charge of 0, and is surrounded by 5
electron pairs and can be described as a
hypervalent molecule. From the perspective of
molecular orbital theory, most of these valence electrons are engaged in S-O bonding.
The S-O bonds are shorter in SO
2 (143.1 pm) than in
sulfur monoxide, SO (148.1 pm), whereas the O-O bonds are longer in O
3 (127.8 pm) than in
dioxygen, O
2 (120.7 pm). The mean bond energy is greater in SO
2 (548 kJ mol
−1) than in SO (524 kJ mol
−1), whereas its is less in O
3 (297 kJ mol
−1) than in O
2 (490 kJ mol
−1). These pieces of evidence lead chemists to conclude that the S-O bonds in sulfur dioxide have a
bond order of at least 2, unlike the O-O bonds in
ozone, which have a bond order of 1.5
Reactions
Treatment of basic solutions with sulfur dioxide affords
sulfite salts:
» SO
2 + 2 NaOH → Na
2SO
3 + H
2O
Featuring sulfur in the +4 oxidation state, sulfur dioxide is a
reducing agent. It is oxidized by halogens such as chlorine to give the sulfuryl halides:
» SO
2 + Cl
2 →
SO2Cl2
However, on rare occasions, it can also act as an
oxidising agent: in the
Claus process, sulfur dioxide is reduced by hydrogen sulfide to give elemental sulfur:
» SO
2 + 2 H
2S → 3 S + 2 H
2O
Sulfur dioxide can act as a metal binding
ligand, typically where the transition metal is in oxidation state 0 or +1. Up to 9 different bonding modes have been determined which include It is present even in so-called unsulphurated wine at concentrations of up to 10 milligrams per litre. It serves as an antibiotic and antioxidant, protecting wine from spoilage by bacteria and oxidation. It also helps to keep volatile acidity at desirable levels. Sulfur dioxide is responsible for the words "contains sulfites" found on wine labels. Wines with SO
2 concentrations below 10ppm don't require "contains sulfites" on the label by US and EU laws. The upper limit of SO
2 allowed in wine is 350ppm in US, in the EU is 160 ppm for red wines and 210 ppm for white and
rosé wines. In low concentrations SO
2 is mostly undetectable in wine, but at over 50ppm, SO
2 becomes evident in the nose and taste of wine.
SO
2 is also a very important element in winery sanitation. Wineries and equipment must be kept very clean, and because bleach can't be used in a winery, a mixture of SO
2, water, and citric acid is commonly used to clean hoses, tanks, and other equipment to keep it clean and free of bacteria.
As a reducing bleach
Sulfur dioxide is also a good
reductant. In the presence of water, sulfur dioxide is able to decolorize substances. Specifically it's a useful reducing
bleach for
papers and delicate materials such as clothes. This bleaching effect normally doesn't last very long.
Oxygen in the atmosphere reoxidizes the reduced dyes, restoring the color.
Precursor to sulfuric acid
Sulfur dioxide is also used to make sulfuric acid, being converted to
sulfur trioxide, and then to
oleum, which is made into
sulfuric acid. Sulfur dioxide for this purpose is made when sulfur combines with oxygen. The method of converting sulfur dioxide to sulfuric acid is called the
contact process.
Biochemical and biomedical roles
Sulfur dioxide is toxic in large amounts. It or its conjugate base bisulfite is produced biologically as an intermediate in both sulfate-reducing organisms and in sulfur oxidizing bacteria as well. Sulfur dioxide has no role in mammalian biology. Sulfur dioxide blocks nerve signals from the
pulmonary stretch receptors (PSR's) and abolishes the
Hering-Breuer inflation reflex.
As a refrigerant
Being easily condensed and with a high heat of evaporation, sulfur dioxide is a candidate material for refrigerants. Prior to the development of
freons, sulfur dioxide was used as a
refrigerant in home refrigerators.
As a reagent and solvent
Sulfur dioxide is a versatile inert solvent that has been widely used for dissolving highly oxidizing salts. It is also used occasionally as a source of the sulfonyl group in
organic synthesis. Treatment of aryl
diazonium salts with sulfur dioxide affords the corresponding aryl sulfonyl chloride.
Dechlorination
In municipal wastewater treatment sulfur dioxide is used to treat chlorinated wastewater prior to release. Sulfur dioxide reacts with free and combined chlorine to form negatively charged chloride ions.
Emissions
According to the
U.S. EPA (as presented by the
2002 World Almanac or in chart form), the following amount of sulfur dioxide was released in the
U.S. per year, measured in thousands of
short tons:
| *1999 |
18,867 |
| *1998 |
19,491 |
| *1997 |
19,363 |
| *1996 |
18,859 |
| *1990 |
23,678 |
| *1980 |
25,905 |
| *1970 |
31,161 |
Due largely to the
US EPA’s
Acid Rain Program, the U.S. has witnessed a 33 percent decrease in emissions between 1983 and 2002. This improvement resulted from
flue gas desulfurization, a technology that enables SO
2 to be chemically bound in
power plants burning sulfur-containing
coal or
oil. In particular,
calcium oxide (lime) reacts with sulfur dioxide to form
calcium sulfite:
» CaO + SO
2 → CaSO
3
Aerobic oxidation converts this CaSO
3 into CaSO
4,
gypsum. Most gypsum sold in Europe comes from flue gas desulfurization.
New fuel additive catalysts, such as
ferox, are being used in gasoline and diesel engines in order to lower the emission of sulfur oxide gases into the atmosphere. This is also done by forcing the sulfur into stable mineral salts and mixed mineral sulfates as opposed to sulfuric acid and sulfur oxides.
As of 2006,
China is the world's largest sulfur dioxide polluter, with 2005 emissions estimated to be 25.49 million tons. This amount represents a 27% increase since 2000, and is roughly comparable with U.S. emissions in 1980.
Al-Mishraq, an Iraqi sulfur plant, was the site of a 2003 disaster resulting in the release of massive amounts of sulfur dioxide into the atmosphere.
Temperature dependence of aqueous solubility
| 22 g/100ml (0 °C) |
15 g/100ml (10 °C) |
| 11 g/100ml (20 °C) |
9.4 g/100 ml (25 °C) |
| 8 g/100ml (30 °C) |
6.5 g/100ml (40 °C) |
| 5 g/100ml (50 °C) |
4 g/100ml (60 °C) |
| 3.5 g/100ml (70 °C) |
3.4 g/100ml (80 °C) |
| 3.5 g/100ml (90 °C) |
3.7 g/100ml (100 °C) |
The values are tabulated for 101.3 kPa partial pressure of SO2. Solubility of gas in a liquid depends on the gas partial pressure according to Henry's law.
The solublity is given for "pure water", for example, water that contains only SO2 in the amount at equilibrium with the gas phase. This "pure water" is going to be acidic. The solublity of SO2 in neutral (or alkaline) water is generally going to be higher because of the pH-dependent speciation of SO2 in the solution with the production of bisulfite and some sulfite ions.
Threats to Health
Sulfur dioxide acts as an acid. Inhalation results in labored breathing, coughing, and/or a sore throat and may cause permanent pulmonary damage. When mixed with water and contacted by skin, frostbite may occur. When it makes contact with eyes, redness and pain will occur.
Further Information
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