Oxides of Sulphur

The oxides of sulphur are inorganic compounds made up entirely of sulphur and oxygen atoms. In the Earth’s lower atmosphere, the most commonly found oxides of sulphur are sulphur dioxide (SO2) and sulphur trioxide (SO3). Some other notable classes of sulphur oxides are listed below.

  • The lower sulphur oxides, which have the general formula SmOn (m>2n).
  • Sulphur monoxide (SO) and disulphur dioxide (S2O2), which is formed from the dimerization of sulphur monoxide.
  • Disulphur monoxide (S2O)
  • The higher sulphur oxides, in which sulphur exhibits an oxidation state of +6.

Normally, oxides of sulphur are formed when substances containing sulphur are burnt in air containing plenty of oxygen. It can be found during the roasting of sulfide ores, burning of fossil fuels, coals, etc. One of the most common sources of sulphur oxide that we can relate to is the emission from vehicles. Sulphur dioxide can be formed naturally due to volcanic activity and also as a byproduct during the metallurgy of copper. Sulphur trioxide, on the other hand, is prepared industrially as a precursor to sulphuric acid and is, therefore, referred to as sulphuric anhydride. The lower oxides of sulphur are formed as intermediates during the combustion of elemental sulphur and are relatively less stable when compared to SO2 and SO3.

Important Oxides of Sulphur

While there are many types of oxides of sulphur, the two most important ones are;

  1. Sulphur dioxide (SO2)
  2. Sulphur trioxide (SO3)

Sulphur dioxide

Sulphur dioxide is one of the most common sulphur oxides that is found on the earth and even in space. It is a colourless gas and sometimes poisonous also soluble in water. Exposure to this gas in high concentration can be harmful to living beings. It can cause adverse health effects in humans.

Properties of Sulphur Dioxide

  • Sulphur dioxide is a colourless, acidic gas with a pungent and suffocating smell.
  • It can be easily liquefied.
  • It is highly soluble in water and its aqueous solution (H2SO3), is acidic in nature.
  • It acts as a strong reducing agent and as such reduces halogens to halogen acids, turns acidified K2Cr2O7 solution green,


  • Decolourises KMnO4 solution and reduces ferric to ferrous salts, PbO2 to PbSO4 and Na2O2 to Na2SO4
  • Being acidic, it reacts with NaOH solution to give sodium sulphite (Na2SO3) which then reacts with more SO2 to form sodium hydrogen sulphite (NaHSO3).

2NaOH + SO2→ Na2SO3+ H2O

Na2SO3+ SO2+H2O→ NaHSO3

  • It also decomposes carbonates and bicarbonates evolving CO2 gas.

Na2CO3 + 2SO2+ H2O → NaHSO3 + CO2

NAHCO3 + SO2 → NaHSO3 + CO2

  • With lime water, it forms milkiness due to the formation of insoluble calcium sulphite (CaSO3) which disappears on further passing SO2 for a long time as a result of the formation of soluble sodium bisulphite (NAHSO3).

Ca(OH)2 + SO2 → CaSO3 + H2O


CaSO3 + SO2 + H2O → Ca(HSO3)2+ H2O


  • SO2 on reaction with PCL5 gives thionyl chloride (SOCI) which fumes in moist air and is used in organic chemistry.

PCI5 + SO2 → SOCl2 + POCI3

  • It also acts as an oxidising agent and a Lewis base. For example, it oxidises H2S to S

2H2S + SO2 → 2H2O +3S.

Lighted magnesium ribbon and heated potassium metal keep on burning in SO2 and are oxidised.

2Mg + SO2 → 2MgO + S

3Mg + SO2 → 2MgO + MgS

4K+3SO2 → K2SO3 + K2S2O3

3Fe + SO2 → 2FeO + FeS

CO is oxidised to CO2

2CO+ SO2 → 2CO2 +S

In the presence of HCl, stannous and mercurous salts are oxidised to stannic and mercuric salts.

2SnCl2+ SO2+ 4HCl→ 2SnCl4 + 2H2O + S

2 Hg2Cl2+ SO2 +4HCl → 4HgCl2+ 2H20 + S

It combines with O2 in the presence of platinised asbestos at 723 K or in the presence of V2O5 at 773 K to gives SO3.

2 SO2+ O2 → 2 SO3

  • It reacts with Cl2 in the presence of charcoal as a catalyst to form sulphuryl chloride SO2Cl2

SO2 + Cl2 → SO2Cl2


It can be prepared by;

● Burning of sulphur in the air

S + O2 → SO2

● Heating of iron pyrites and by the action of dilute acids on sulphites and bisulphites

4FeS2+ 11O2→Fe2O3+ 8SO2

In the laboratory, it is prepared by the action of concentrated H2SO4 on copper turnings

Cu + 2H2SO4 → CuSO4 + SO2 + 2H2O

Uses of Sulphur Dioxide

  • Sulphur dioxide is used in the manufacture of H2SO4.
  • It is used in the refining of cane juice in the sugar industry.
  • For fumigation, as a germicide and for preserving fruits.
  • Liquid SO2 is used as a non-aqueous solvent and as a refrigerant.

Sulphur trioxide (SO3)

Sulphur trioxide is often formed when sulphur dioxide is oxidised. This chemical compound can occur in different forms more widely in a white crystalline solid. When it is in liquid form it is colourless. It is a highly reactive substance and tends to react violently with water. When sulphur trioxide is in vapour form it is considered to be a major pollutant and is one of the components of acid rain. It also tends to fume to a great degree in the atmosphere and its vapour is also highly corrosive. This compound should be handled with extreme caution.


  • It is an acidic oxide and dissolves in H2O to form H2SO4. As such it reacts with CaO to form CaSO4 and decomposes carbonates to evolve CO.
  • It reacts with HCl to form chlorosulphonic acid (HOSO2CI).
  • It acts as a strong oxidising agent.


It can be prepared by;

(i) Passing a mixture of SO2 and O2 overheated Pt or V2O5

(ii) by dehydration of H2SO4 with P2O5 or by heating ferric sulphate

2 SO2+ O2 ⇔ 2 SO3

H2SO4+P2O5 → SO3+ 2 HPO3

Fe2(SO4)3 → Fe2O3 +3 SO3


  • Used in the preparation of sulphuric acid and other chemicals.
  • It is an important reagent in sulfonation reactions.

Comparison between Sulphur Dioxide and Sulphur Trioxide

Properties Sulphur Dioxides (SO2) Sulphur Trioxides (SO3)
Molecular mass 64 80
Melting point  (oC) -75 16.8
Boiling points (oC) -10 43.7
Density 0 oC, 101.3kPa 1.250 g/dm3 2.052 g/dm3
20 oC 101.3kPa 2.93 g/dm3 1.916 g/dm3
Solubility in water 80 dm3/dm3 (97.7 ppmm) Decay


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