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 have the general formula S_mO_n (m>2n).
• Sulphur monoxide (SO) and disulphur dioxide (S₂O₂) which are formed from the dimerisation 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 sulphide 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. It is sometimes poisonous and is 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 liquefied easily.
• 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, turning acidified K₂Cr₂O₇ solution green.

K₂Cr₂O₇+3SO₂+H₂SO₄→K₂SO₄+Cr₂(SO₄)₃+H₂O

• Decolourises KMnO4 solution and reduces ferric to ferrous salts, PbO₂ to PbSO₄ and Na₂O₂ to Na₂SO_4.
• 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

Milkiness

CaSO₃ + SO₂ + H₂O → Ca(HSO₃)₂+ H₂O

(Soluble)

• SO2, in 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 CO_2.

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 give SO₃.

2 SO₂+ O₂ → 2 SO₃

• It reacts with Cl2 in the presence of charcoal as a catalyst to form sulphuryl chloride SO₂Cl_2.

SO₂ + Cl₂ → SO₂Cl₂

Preparation

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.

Properties

• 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.

Preparation

It can be prepared by

(i) Passing a mixture of SO₂ and O₂ overheated Pt or V₂O_5.

(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

Uses

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

Comparison between Sulphur Dioxide and Sulphur Trioxide