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Sulfur dioxide can be prepared by burning sulfur:

S8 + 8 O2 → 8 SO2

The combustion of hydrogen sulfide and organosulfur compounds proceeds similarly.

2 H2S (g) + 3 O2 (g) → 2 H2O (g) + 2 SO2 (g)

The roasting of sulfide ores such as pyrites, sphalerite (zinc blende), and cinnabar (mercury sulfide) also releases SO2:

4 FeS2 (s) + 11 O2 (g) → 2 Fe2O3 (s) + 8 SO2 (g)

2 ZnS (s) + 3 O2 (g) → 2 ZnO (s) + 2 SO2 (g)

HgS (s) + O2 (g) → Hg (g) + SO2 (g)

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 (s) + 2 SiO2 (s) + C (s) → 2 CaSiO3 (s) + 2 SO2 (g) + CO2 (g)

Action of hot sulfuric acid on copper turnings produces sulfur dioxide.

Cu (s) + 2 H2SO4 (aq) → CuSO4 (aq) + SO2 (g) + 2 H2O (l)

It can also be prepared with sodium metabisulfite:

H2SO4 (aq) + Na2S2O5 (aq) → 2 SO2 (g) + Na2SO4 (s) + H2O (l)

This is an exothermic reaction.

Structure and bonding

SO2 is a bent molecule with C2v 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.

three resonance structures of sulfur dioxide

The S–O bonds are shorter in SO2 (143.1 pm) than in sulfur monoxide, SO (148.1 pm), whereas the O–O bonds are longer in ozone (127.8 pm) than in dioxygen, O2 (120.7 pm). The mean bond energy is greater in SO2 (548 kJ/mol) than in SO (524 kJ/mol), whereas it is less in O3 (297 kJ/mol) than in O2 (490 kJ/mol). 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.


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

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:
SO2 + 2 H2S → 3 S + 2 H2O

Sulfur dioxide can react with 1,3-dienes in a cheletropic reaction.

Sulfur dioxide can bind to metal ions as a ligand, 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.


Precursor to sulfuric acid


As a preservative

Sulfur dioxide is sometimes used as a preservative for dried apricots 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 appearance of the fruit and prevents rotting.

In winemaking

Sulfur dioxide is an important compound in winemaking, and is designated as parts per million in wine, E number: E220. 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 SO2 concentrations below 10 ppm do not require "contains sulfites" on the label by US and EU laws. The upper limit of SO2 allowed in wine is 350 ppm in the US, in the EU 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 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 clean, and because bleach cannot be used in a winery, 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 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.

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 possessing 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 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:


According to the United States Environmental Protection Agency (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 USmarker 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 reacts with sulfur dioxide to form calcium sulfite:
CaO + SO2 → CaSO3
Aerobic oxidation of the CaSO3 gives 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, Chinamarker 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.

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.




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