Sulfur dioxide

Sulfur dioxide
General
Systematic name	sulfur dioxide
Other names	sulphur dioxide sulfur(IV) oxide sulfurous anhydride sulphurous anhydride
Molecular formula	SO2
Molar mass	64.054 g mol−1
Appearance	colorless 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 (pKa)	1.81
Structure
Molecular shape	Bent 120 [[1]
Dipole moment	1.63 D
Thermodynamic data
Standard enthalpy of formation ΔfH°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	R23, R34
S-phrases	S1/2, S9, S26 S36/37/39, S45
NFPA 704	0 3 0
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.

Preparation

Sulfur dioxide can be prepared by burning sulfur in air. This reaction, in which sulfur combines with oxygen in the air, may be written as follows:

S₈(s) + 8O₂(g) → 8SO₂(g)

The combustion of hydrogen sulfide and organosulfur compounds proceeds in a similar manner:

2H₂S(g) + 3O₂(g) → 2H₂O(g) + 2SO₂(g)

Sulfur dioxide is also produced during the roasting of sulfide ores, such as iron pyrites, sphalerite (zinc blende), and cinnabar (mercury sulfide). These reactions are:

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 calcium sulfate (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)

The 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 zero, and is surrounded by five 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) in alcoholic drinks, 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 nineteenth 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 ten ppm 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), 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 percent increase since 2000, and is roughly comparable with U.S. emissions in 1980.

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 solubility 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 solubility 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.

See also

• Oxide
• Sulfur
• Sulfur trioxide
• Sulfuric acid

References

ISBN links support NWE through referral fees

• Chang, Raymond. 2006. Chemistry, 9th ed. New York: McGraw-Hill Science/Engineering/Math. ISBN 0073221031
• Cotton, F. Albert, and Geoffrey Wilkinson. 1980. Advanced Inorganic Chemistry, 4th ed. New York: Wiley. ISBN 0471027758
• Lide, David R. 2006. CRC Handbook of Chemistry and Physics. 87th ed. Boca Raton, FL: CRC Press. ISBN 0849304873
• Wells, A. F. 1984. Structural Inorganic Chemistry. 5th ed. Oxford: Oxford University Press.

External links

All links retrieved February 26, 2023.

• NIOSH Pocket Guide to Chemical Hazards
• Sulfur Dioxide, Molecule of the Month

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