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Sulfur dioxide
Sulfur dioxide Sulfur dioxide
Other names	sulphur dioxide sulfur(IV) oxide sulfurous anhydride sulphurous anhydride
Molecular formula	SO2
Molar mass	64.054 g mol−1
Appearance	colourless gas
CAS number	[7446-09-5]
EINECS number	231-195-2
Properties
Density and phase	2.551 g/L, gas
Solubility in water	9.4 g/100 mL (25 °C)
Melting point	−72.4 °C (200.75 K)
Boiling point	−10 °C (263 K)
Critical Point	157.2°C at 7.87 MPa
Acidity (p''K''a)	1.81
Structure
Molecular shape	Bent 120 [1]
Dipole moment	1.63 D
Thermodynamic data
Standard enthalpy offormation Δf''H''°gas	−296.84 kJ mol−1
Standard molar entropy ''S''°gas	248.21 J K−1 mol−1
'Safety data'
EU classification	Toxic
R-phrases	,
S-phrases	S1/2, S9, S26 S36/37/39, S45
NFPA 704
PEL-TWA (OSHA)	5 ppm (13 mg m−3)
IDLH (NIOSH)	100 ppm
Flash point	non-flammable
RTECS number	WS4550000
Supplementary data page
Structure and properties	''n'', ''εr'', etc.
Thermodynamic data	Phase behaviour Solid, liquid, gas
Spectral data	UV, IR, NMR, MS
Related compounds
Other cations	Selenium dioxide Tellurium dioxide
Related compounds	Sulfur trioxide Sulfuric acid
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)

'Sulfur dioxide' (also 'sulphur dioxide') is the chemical compound with the formula SO₂. This important gas is the main product from the combustion of sulfur compounds and is of significant environmental concern. SO₂ is often described as the "smell of burning sulfur" but is ''not'' responsible for the smell of rotten eggs.
SO₂ is produced by volcanoes and in various industrial processes. Since coal and petroleum contain various amounts of sulfur compounds, their combustion generates sulfur dioxide. Further oxidation of SO₂, usually in the presence of a catalyst such as NO₂, forms H₂SO₄, and thus acid rain.^[1]

Contents

Preparation

Structure and bonding

Uses

Emissions

Temperature dependence of aqueous solubility

References

See also

External links

Preparation

Sulfur dioxide can be prepared by burning sulfur:
:S₈(s) + 8O₂(g) → 8SO₂(g)
The combustion of hydrogen sulfide and organosulfur compounds proceeds similarly.
:2H₂S(g) + 3O₂(g) → 2H₂O(g) + 2SO₂(g)
The roasting of sulfide ores such as iron pyrites, sphalerite (zinc blende) and cinnabar (mercury sulfide) also emits SO₂:
:4FeS₂(s) + 11O₂(g) → 2Fe₂O₃(s) + 8SO₂(g)
:2ZnS(s) + 3O₂(g) → 2ZnO(s) + 2SO₂(g)
:HgS(s) + O₂(g) → Hg(g) + SO₂(g)
When anhydrous CaSO₄ is heated with coke and sand in the manufacture of cement, CaSiO₃, sulfur dioxide is a by-product.
:2CaSO₄(s) + 2SiO₂(s) + C(s) → 2CaSiO₃(s) + 2SO₂(g) + CO₂(g)
Action of hot concentrated sulfuric acid on copper turnings will produce sulfur dioxide.
:Cu(s) + 2H₂SO₄(aq) → CuSO₄(aq) + SO₂(g) + 2H₂O(l)

Structure and bonding

SO₂ 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. From the perspective of molecular orbital theory, most of these electron pairs are non-bonding in character, as is typical for hypervalent molecules.
One conventional covalent bond is present between each oxygen and the central sulfur atom, with two further electrons delocalised between the oxygens and the sulfur atom.

Uses

Sulfur dioxide is sometimes used as a preservative (E number: E220^[2]) in alcoholic drinks^[3], or dried apricots and other dried fruits due to its antimicrobial properties. The preservative is used to maintain the appearance of the fruit rather than prevent rotting. This can give fruit a distinctive chemical taste.
Sulfur dioxide is also a good reductant. In the presence of water, sulfur dioxide is able to decolorize substances that can be reduced by it; thus making it a useful reducing bleach for papers and delicate materials such as clothes.
This bleaching effect normally does not last very long. Oxygen in the atmosphere reoxidizes the reduced dyes, restoring the color.
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. This is called the contact process.
According to Claude Ribbe in ''The Crime of Napoleon,'' sulfur dioxide gas was used as an execution poison by the French emperor to suppress a slave revolt in Haiti early in the 19th century.
Sulfur dioxide blocks nerve signals from the pulmonary stretch receptors (PSR's) and abolishes the Hering-Breuer inflation reflex.
Prior to the development of freons, sulfur dioxide was used as a refrigerant in home refrigerators.
Sulfur dioxide is the anhydride of sulfurous acid, H₂SO₃.
Sulfur dioxide is a very important element in winemaking, and is designated as parts per million in wine. It acts as an antibiotic and antioxidant, protecting wine from spoilage organisms, bacteria, and oxidation, and also helps to keep volatile acidity at desirable levels. Sulfur dioxide is responsible for the words "contains sulfites" found on wine labels. Wines with SO₂ concentrations below 10ppm do not require "contains sulfites" on the label by US and EU laws. The upper limit of SO₂ 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 SO₂ is mostly undetected in wine, but at over 50ppm, SO₂ becomes evident in the nose and taste of wine.
SO₂ is also a very important element in winery sanitation. Wineries and equipment must be kept very clean, and because bleach cannot be used in a winery, a mixture of SO₂, water, and citric acid is commonly used to clean hoses, tanks, and other equipment to keep it clean and free of bacteria.

Emissions

According to the U.S. EPA (as presented by the ''2002 World Almanac'' or in chart form ^[4]), 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 SO₂ 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 + SO₂ → CaSO₃
Aerobic oxidation converts this CaSO₃ into CaSO₄, 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^[5].
Al-Mishraq, an Iraqi sulfur plant, was the site of a 2004 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", i.e., 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.

References

1. Dr. Mike Thompson,
Winchester College, UK http://www.chm.bris.ac.uk/motm/so2/so2h.htm
2. Current EU approved additives and their E Numbers, The Food Standards Agency website.
3. "All wines contain sulphur dioxide in various forms, collectively known as sulphites. Even in completely unsulphured wine it is present at concentrations of up to 10 milligrams per litre", Sulphites in wine, MoreThanOrganic.com.
4. National Trends in Sulfur Dioxide Levels, United States Environmental Protection Agency.
5. China has its worst spell of acid rain, United Press International.

See also

★ Sulfur

★ Sulfur trioxide

★ Sulfur-iodine cycle

★ National Ambient Air Quality Standards

★ Homer City Generating Station

★

External links

★ United States Environmental Protection Agency Sulfur Dioxide page

★ International Chemical Safety Card 0074

★ IARC Monograph "Sulfur Dioxide and some Sulfites, Bisulfites and Metabisulfites"

★ NIOSH Pocket Guide to Chemical Hazards

★ Food Intolerance Network - Sulfite factsheet

★ Sulfur Dioxide, Molecule of the Month