{{#if:64.066 g mol−1Colorless gasSO22.6288 kg m−394 g dm−3[1]237.2 kPa1.8112.190.403 cP (at 0 °C)201|! style="background: #F8EABA; text-align: center;" colspan="2" | Properties
Sulfur dioxide
CAS number 7446-09-5 YesY
PubChem 1119
ChemSpider 1087 YesY
UNII 0UZA3422Q4 YesY
EC number 231-195-2
UN number 1079, 2037
KEGG D05961 YesY
MeSH Sulfur+dioxide
ChEBI CHEBI:18422 YesY
RTECS number WS4550000
Beilstein Reference 3535237
Gmelin Reference 1443
Jmol-3D images Image 1
Molecular formula SO2
Molar mass 64.066 g mol−1
Appearance Colorless gas
Density 2.6288 kg m−3
Melting point

-72 °C, 201 K, -98 °F

Boiling point

−10 °C, 263 K, 14 °F

Solubility in water 94 g dm−3[1]
Vapor pressure 237.2 kPa
Acidity (pKa) 1.81 Basicity (pKb) 12.19
Viscosity 0.403 cP (at 0 °C)
Space group C2v
Molecular shape Dihedral
Dipole moment 1.62 D
Std enthalpy of
-296.81 kJ mol−1
Standard molar
248.223 J K−1 mol−1
EU Index 016-011-00-9
EU classification Toxic T
R-phrases R23, R34, R50
S-phrases (S1/2), S9, S26, S36/37/39, S45
NFPA 704
NFPA 704.svg
LD50 3000 ppm (30 min inhaled, mouse)
Related compounds
Related sulfur oxides Sulfur monoxide
Sulfur trioxide
Related compounds Ozone

Selenium dioxide
Sulfurous acid
Tellurium dioxide

 N (verify) (what is: YesY/N?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Sulfur dioxide (also sulphur dioxide) is the chemical compound with the formula SO2. It is a toxic gas with a pungent, irritating smell, that is released by volcanoes and in various industrial processes. Since coal and petroleum often contain sulfur compounds, their combustion generates sulfur dioxide unless the sulfur compounds are removed before burning the fuel. Further oxidation of SO2, usually in the presence of a catalyst such as NO2, forms H2SO4, and thus acid rain.[2] Sulfur dioxide emissions are also a precursor to particulates in the atmosphere. Both of these impacts are cause for concern over the environmental impact of these fuels.

Structure and bonding

SO2 is a bent molecule with C2v symmetry point group. In terms of electron-counting formalism, the sulfur atom has an oxidation state of +4 and a formal charge of +1.

Two resonance structures of sulfur dioxide

The Lewis structure of sulfur dioxide consists of an S=O double bond and an S–O dative bond without utilizing d-orbitals,[3] resulting in a bond order of 1.5.


Combustion routes

Sulfur dioxide is the product of the burning of sulfur or of burning materials that contain sulfur:

S8 + 8 O2 → 8 SO2

The combustion of hydrogen sulfide and organosulfur compounds proceeds similarly.

2 H2S + 3 O2 → 2 H2O + 2 SO2

The roasting of sulfide ores such as pyrite, sphalerite, and cinnabar (mercury sulfide) also releases SO2:[4]

4 FeS2 + 11 O2 → 2 Fe2O3 + 8 SO2
2 ZnS + 3 O2 → 2 ZnO + 2 SO2
HgS + O2 → Hg + SO2
4 FeS + 7O2 → 2 Fe2O3 + 4 SO2

A combination of these reactions is responsible for the largest source of sulfur dioxide, volcanic eruptions. These events can release millions of tonnes of SO2.

Reduction of higher oxides

Sulfur dioxide is a by-product in the manufacture of calcium silicate cement: CaSO4 is heated with coke and sand in this process:

2 CaSO4 + 2 SiO2 + C → 2 CaSiO3 + 2 SO2 + CO2

The action of hot sulfuric acid on copper turnings produces sulfur dioxide.

Cu + 2 H2SO4 → CuSO4 + SO2 + 2 H2O

From sulfite

Sulfite results from the reaction of aqueous base and sulfur dioxide. The reverse reaction involves acidification of sodium metabisulfite:

H2SO4 + Na2S2O5 → 2 SO2 + Na2SO4 + H2O


Industrial reactions

Treatment of basic solutions with sulfur dioxide affords sulfite salts:

SO2 + 2 NaOH → Na2SO3 + H2O

Featuring sulfur in the +4 oxidation state, sulfur dioxide is a reducing agent. It is oxidized by halogens to give the sulfuryl halides, such as sulfuryl chloride:

SO2 + Cl2 → SO2Cl2

Sulfur dioxide is the oxidising agent in the Claus process, which is conducted on a large scale in oil refineries. Here sulfur dioxide is reduced by hydrogen sulfide to give elemental sulfur:

SO2 + 2 H2S → 3 S + 2 H2O

The sequential oxidation of sulfur dioxide followed by its hydration is used in the production of sulfuric acid.

2 SO2 + 2 H2O + O2 → 2 H2SO4

Laboratory reactions

Sulfur dioxide can react with certain 1,3-dienes in a cheletropic reaction to give organosulfur compounds.

Sulfur dioxide can bind to metal ions as a ligand to form metal sulfur dioxide complexes, typically where the transition metal is in oxidation state 0 or +1. Many different bonding modes (geometries) are recognized, but in most cases the ligand is monodentate, attached to the metal through sulfur, which can be either planar and pyramidal η1.[5]


Precursor to sulfuric acid

Sulfur dioxide is an intermediate in the production of 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. Several billion kilograms are produced annually for this purpose.

As a preservative

Sulfur dioxide is sometimes used as a preservative for dried apricots, dried figs, and other dried fruits owing to its antimicrobial properties, and it is sometimes called E220 when used in this way. As a preservative, it maintains the colorful appearance of the fruit and prevents rotting. It is also added to sulfured molasses.

In winemaking

Sulfur dioxide is an important compound in winemaking, and is designated as parts per million in wine, E number: E220.[6] It is present even in so-called unsulfurated wine at concentrations of up to 10 mg/L.[7] It serves as an antibiotic and antioxidant, protecting wine from spoilage by bacteria and oxidation. Its antimicrobial action also helps to minimize volatile acidity. Sulfur dioxide is responsible for the words "contains sulfites" found on wine labels.

Sulfur dioxide exists in wine in free and bound forms, and the combinations are referred to as total SO2. Binding, for instance to the carbonyl group of acetaldehyde, varies with the wine in question. The free form exists in equilibrium between molecular SO2 (as a dissolved gas) and bisulfite ion, which is in turn in equilibrium with sulfite ion. These equilibria depend on the pH of the wine. Lower pH shifts the equilibrium towards molecular (gaseous) SO2, which is the active form, while at higher pH more SO2 is found in the inactive sulfite and bisulfite forms. It is the molecular SO2 which is active as an antimicrobial and antioxidant, and this is also the form which may be perceived as a pungent odour at high levels. Wines with total SO2 concentrations below 10 parts per millon (ppm) do not require "contains sulfites" on the label by US and EU laws. The upper limit of total SO2 allowed in wine in the US is 350 ppm; in the EU it is 160 ppm for red wines and 210 ppm for white and rosé wines. In low concentrations, SO2 is mostly undetectable in wine, but at free SO2 concentrations over 50 ppm, SO2 becomes evident in the nose and taste of wine.[citation needed]

SO2 is also a very important compound in winery sanitation. Wineries and equipment must be kept clean, and because bleach cannot be used in a winery due the risk of cork taint,[8] a mixture of SO2, water, and citric acid is commonly used to clean and sanitize equipment. Compounds of ozone (O3) are now used extensively as cleaning products in wineries[citation needed] due to their efficiency, and because these compounds do not affect the wine or equipment.

As a reducing agent

Sulfur dioxide is also a good reductant. In the presence of water, sulfur dioxide is able to decolorize substances. Specifically it is 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. In municipal wastewater treatment, sulfur dioxide is used to treat chlorinated wastewater prior to release. Sulfur dioxide reduces free and combined chlorine to chloride.[9]

Sulfur dioxide is fairly soluble in water, and by both IR and Raman spectroscopy, it is known that the hypothetical sulfurous acid, H2SO3, is not present to any extent. However, such solutions do show spectra of the hydrogen sulfite ion, HSO3, by reaction with water, and it is in fact the actual reducing agent present:

SO2 + H2O HSO3 + H+

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. The role of sulfur dioxide in mammalian biology is not yet well understood.[10] Sulfur dioxide blocks nerve signals from the pulmonary stretch receptors (PSRs) and abolishes the Hering–Breuer inflation reflex.

As a refrigerant

Being easily condensed and possessing a high heat of evaporation, sulfur dioxide is a candidate material for refrigerants. Prior to the development of CFCs, sulfur dioxide was used as a refrigerant in home refrigerators.

As a reagent and solvent in the laboratory

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 and cuprous chloride affords the corresponding aryl sulfonyl chloride, for example:[11]

Preparation of m-trifluoromethylbenzenesulfonyl chloride.svg

As an air pollutant

A sulfur dioxide plume from the Halemaʻumaʻu vent, glows at night

Sulfur dioxide is a noticeable component in the atmosphere, especially following volcanic eruptions.[12] According to the United States Environmental Protection Agency (as presented by the 2002 World Almanac or in chart form[13]), the following amount of sulfur dioxide was released in the U.S. per year:

Year SO2 (thousands of short tons)
1970 31,161
1980 25,905
1990 23,678
1996 18,859
1997 19,363
1998 19,491
1999 18,867

Sulfur dioxide is a major air pollutant and has significant impacts upon human health.[14] In addition the concentration of sulfur dioxide in the atmosphere can influence the habitat suitability for plant communities as well as animal life.[15] Sulfur dioxide emissions are a precursor to acid rain and atmospheric particulates. Due largely to the US EPA’s Acid Rain Program, the U.S. has witnessed a 33% decrease in emissions between 1983 and 2002. This improvement resulted in part 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 of the CaSO3 gives CaSO4, anhydrite. Most gypsum sold in Europe comes from flue-gas desulfurization.

Sulfur can be removed from coal during the burning process by using limestone as a bed material in Fluidized bed combustion.[16]

Sulfur can also be removed from fuels prior to burning the fuel, preventing the formation of SO2 because there is no sulfur in the fuel from which SO2 can be formed. The Claus process is used in refineries to produce sulfur as a byproduct. The Stretford process has also been used to remove sulfur from fuel. Redox processes using iron oxides can also be used, for example, Lo-Cat[17] or Sulferox.[18]

Fuel additives, such as calcium additives and magnesium oxide, are being used in gasoline and diesel engines in order to lower the emission of sulfur dioxide gases into the atmosphere.[19]

As of 2006, China was 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.[20]



Inhaling sulfur dioxide is associated with increased respiratory symptoms and disease, difficulty in breathing, and premature death.[21] In 2008, the American Conference of Governmental Industrial Hygienists reduced the short-term exposure limit from 5 ppm to 0.25 ppm. The OSHA PEL is currently set at 5 ppm (13 mg/m3) time weighted average. NIOSH has set the IDLH at 100 ppm.[22]

A 2011 systematic review concluded that exposure to sulfur dioxide is associated with preterm birth.[23]


In the United States, the Center for Science in the Public Interest lists the two food preservatives, sulfur dioxide and sodium bisulfite, as being safe for human consumption except for certain individuals who may be sensitive to them, especially in large amounts.[24]


  1. Template:RubberBible87th
  2. Holleman, A. F.; Wiberg, E. (2001), Inorganic Chemistry, San Diego: Academic Press, ISBN 0-12-352651-5
  3. Script error
  4. Shriver, Atkins. Inorganic Chemistry, Fifth Edition. W. H. Freeman and Company; New York, 2010; p. 414.
  5. Script error
  6. Current EU approved additives and their E Numbers, The Food Standards Agency website.
  7. Sulphites in wine, MoreThanOrganic.com.
  8. http://www.extension.purdue.edu/extmedia/FS/FS-50-W.pdf
  9. Script error
  10. Script error
  11. R. V. Hoffman (1990), "m-Trifluoromethylbenzenesulfonyl Chloride", Org. Synth., http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=CV7P0508; Coll. Vol. 7: 508
  12. Volcanic Gases and Their Effects. Volcanoes.usgs.gov. Retrieved on 2011-10-31.
  13. National Trends in Sulfur Dioxide Levels, United States Environmental Protection Agency.
  14. Sulfur Dioxide. United States Environmental Protection Agency
  15. C.Michael Hogan. 2010. Abiotic factor. Encyclopedia of Earth. eds Emily Monosson and C. Cleveland. National Council for Science and the Environment. Washington DC
  16. Script error
  17. Lo-Cat Process[dead link]
  18. Process screening analysis of alternative gas treating and sulfur removal for gasification. (December 2002) Report by SFA Pacific, Inc. prepared for U.S. Department of Energy (PDF) . Retrieved on 2011-10-31.
  19. Walter R. May Marine Emissions Abatement. SFA International, Inc., p. 6.
  20. China has its worst spell of acid rain, United Press International (2006-09-22).
  21. Sulfur Dioxide U.S. Environmental Protection Agency
  22. "NIOSH Pocket Guide to Chemical Hazards". http://www.cdc.gov/niosh/npg/npgd0575.html.
  23. Script error
  24. "Center for Science in the Public Interest – Chemical Cuisine". http://www.cspinet.org/reports/chemcuisine.htm. Retrieved March 17, 2010.

External links