Microscale Gas Chemistry:

 Sulfur Dioxide Information
A. Appearance

B. Physical Properties of SO2

Sulfur Dioxide, O2
Atomic mass: 64.06 g/mol
melting point -72.7 oC
boiling point -10 oC
C. History
     Sulfur dioxide has a number of important uses.  It has been used as a fumigant since ancient times.  In the Odyssey, Homer describes the burning of sulfur (formation of SO2) as a way that homes were 'fumigated'.

D. Natural Abundance
     Volcanic activity is the primary source of sulfur dioxide in nature.  Human activity, specifically the combustion of sulfur-rich coal and petroleum, accounts for much more of the SO2 in nature.  Sulfur dioxide is the main culprit in acid rain.

E. Industrial Production
     Sulfur burns in air with a blue flame.  The product of this combustion is SO2(g).  Sulfur dioxide is produced by the combustion of sulfur or hydrogen sulfide (Chapter 8.)  It is also produced by the roasting of metal sulfides such as FeS2.

F. Industrial Uses
     The vast majority of the sulfur produced in the world is converted to SO2 as the first step in the production of sulfuric acid.

     Sulfur dioxide is used as a fungicide.  It also has widespread toxic effects on microorganisms and is thus used to disinfect and preserve food and wine.  Dried fruits are preserved with SO2 which prevents discoloration (browning) as well as preserves the fruit.  Traces of SO2 can be noticed when smelling or eating dried fruit.  Sulfur dioxide is also extensively used in the brewing industry.  Numerous foodstuffs, including wines, use it as a preservative.  It is used as a bleaching agent for paper, textiles, oils, etc.

G. Gas Density of SO2
      Sulfur dioxide is over twice (2.21 X) as dense as air.  The density of SO2 is 2.618 g/L at 25 oC and 1 atm.

H. Gas Solubility of SO2
     Sulfur dioxide is highly soluble in water.  At room temperature, the dissolving process is exothermic:


At room temperature, the solubility of SO2 is approximately 200 g SO2/L.  Thus, 1 mL of water could dissolve 76 mL SO2!  The solubility of SO2 in water is highly temperature dependent and is about 400 times more soluble at 0 oC (228 g/L) than it is at 90 oC (5.8 g/L).

    SO2 is often thought of as the 'anhydride of sulfurous acid', H2SO3.  However, it is questionable if sulfurous acid exists at all.  If it does, the equilibrium constant is so small that fewer than 1 molecule per billion is in the form of sulfurous acid:

SO2(aq) + H2O  <====>  H2SO3(aq)  K<< 1 x 10-9

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