Sulfur dioxide

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Sulfur dioxide
IUPAC name
Sulfur dioxide
Other names
Sulfur(IV) oxide
Sulfurous anhydride
Molar mass 64.066 g/mol
Appearance Colorless gas
Odor Sulfurous
Density 2.6288 kg/m3
Melting point −72 °C (−98 °F; 201 K)
Boiling point −10 °C (14 °F; 263 K)
9.4 g/100 ml
Solubility Soluble in acetone, chloroform, diethyl ether, ethanol, formic acid
Vapor pressure 237.2 kPa
Acidity (pKa) 1.81
248.223 J·K−1·mol−1
−296.81 kJ/mol
Safety data sheet Sigma-Aldrich
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
3,000 ppm (mouse, 30 min)
2,520 ppm (rat, 1 hr)
Related compounds
Related compounds
Sulfur trioxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Infobox references

Sulfur dioxide is a gas with an irritating smell, industrially used mainly to make sulfuric acid. It has the chemical formula SO2.



Sulfur dioxide will react with alkali to form sulfites:

SO2 + 2 NaOH → Na2SO3 + H2O

Sulfur dioxide will reduce hydrogen sulfide to elemental sulfur:

SO2 + 2 H2S → 3 S + 2 H2O

Sulfur dioxide can be further oxidized by halogens, to sulfuryl halides, such as sulfuryl chloride:

SO2 + Cl2 → SO2Cl2

Sulfur dioxide can be further oxidized by oxygen to sulfur trioxide:

2 SO2 + O2 → 2 SO3

Sulfur dioxide dissolves in water, forming sulfurous acid. This acid is very similar to carbonic acid in that it only exists in very small quantities as an ionic species in such solutions, and can only be observed as an anhydrous substance in very extreme conditions.


Sulfur dioxide is a toxic white gas, with a poignant smell of burned matches. It is soluble in water at low temperature. SO2 boils at −10 °C, and melts at −72 °C.


Liquified sulfur dioxide is sold as compressed gas in chemical industry.


Sulfur is usually cheap, so burning sulfur in air produces plenty of sulfur dioxide. However, set up limitations sometimes prevent the use of this method, usually due to having a closed system where the sulfur cannot be open to the air. Mixing an oxidizer such as potassium chlorate or potassium nitrate can solve this issue but a wet method may be preferred. In that case, the action of a strong acid on a sulfite or bisulfite will create sulfur dioxide gas.

When burning elemental sulfur, a small amount will aerosolize and may contaminate the installation and the purity of the product. There are a few ways to remove the impurities: one requires to dissolve the gas in cold water, filter the resulting sulfurous acid solution and then boil the solution to desorb the sulfur dioxide from solution. Another way is to wash the resulting gas in a wash column, usually with warm water to reduce the adsorption of SO2 or aprotic solvents that do not react with it. Electrostatic precipitators are even better, as they do not require adsorption-desorption processes and are more efficient.

Sulfur dioxide can also be produced by roasting metal sulfides, a common process applied in the sulfuric acid industry. It can also be produced by mixing calcium sulfate with sand (silicon dioxide) and carbon (coke), though this process requires very high temperatures.

"A lovely generator can be made by mixing H2SO4 with an equal volume of water, and then adding sulfite or (meta)bisulfite to this. Slight heating gives a smooth and not too fast generation of lots of SO2. You don't need strong heating, and that is a pleasant thing" - Thread on SO2 generation




Sulfur dioxide is corrosive and toxic. Inhalation of a high quantity can irritate and damage human tissues.


Storing liquified sulfur dioxide should only be done in corrosion-resistant cylinders.


Sulfur dioxide can be neutralized by bubbling it in alkaline solutions or mixing it with hydrogen sulfide or ammonia. Bubbling it through hydrogen peroxide yields sulfuric acid.


Relevant Sciencemadness threads