Sulphur and Its Inorganic Compounds

Lessons

Senior 3 · Chemistry

Sulphur and Its Inorganic Compounds

4 lessons

Lesson 2 of 4

Sulfur Dioxide

25 min

The smell of a struck match

Strike a match and, for a moment, you notice a sharp, choking smell. That is sulfur dioxide, made as the sulfur in the match head burns. The same gas is added in tiny amounts to dried fruit and some fruit juices to keep them from spoiling, is used to bleach paper and wool, and pours out of volcanoes and factory chimneys in huge amounts. In S1 you met sulfur dioxide as an air pollutant. In this lesson you prepare it, test it and find out why it is such a useful chemical, but also why it must be handled with care.

Preparing sulfur dioxide in the laboratory

Sulfur dioxide is made by adding dilute hydrochloric acid to sodium sulfite (Figure 2.1). No heating is needed:

$\mathrm{Na_2SO_3(s) + 2HCl(aq) \rightarrow 2NaCl(aq) + H_2O(l) + SO_2(g)}$

Another method is to heat copper turnings with concentrated sulfuric acid, which oxidises the copper:

$\mathrm{Cu(s) + 2H_2SO_4(l) \rightarrow CuSO_4(aq) + SO_2(g) + 2H_2O(l)}$

The gas is dried by bubbling it through concentrated sulfuric acid. Sulfur dioxide is an acidic gas, so it does not react with the acid.
It is collected by downward delivery, because it is more than twice as dense as air ( $M_{r} = 64$ , against 29) and too soluble in water to collect over it.
Everything is done in a fume cupboard.

Figure 2.1: Laboratory preparation of dry sulfur dioxide.

Warning

Safety. Sulfur dioxide is toxic. Even small amounts irritate the eyes, nose and throat and can trigger a serious asthma attack, so anyone with asthma should stay well away. Prepare and use it only in a working fume cupboard, in small amounts, and never smell it directly. Concentrated sulfuric acid causes severe burns: wear eye protection and gloves.

Physical properties

Colourless, with a sharp, choking smell.
About 2.2 times as dense as air.
Very soluble in water: about 40 volumes dissolve in one volume of water at room temperature.
Easily liquefied (it boils at −10 °C), so it is transported as a liquid in steel cylinders.

Chemical properties

1. An acidic oxide. It dissolves in water to form sulfurous acid, a weak acid that turns blue litmus red:

$\mathrm{SO_2(g) + H_2O(l) \rightleftharpoons H_2SO_3(aq)}$

With alkalis it forms sulfites: $\mathrm{SO_2(g) + 2NaOH(aq) \rightarrow Na_2SO_3(aq) + H_2O(l)}$ .

2. A reducing agent. In the presence of water, sulfur dioxide removes oxygen from, or gives electrons to, oxidising agents, and is itself oxidised to sulfate. This gives the two tests for the gas:

It turns acidified potassium dichromate(VI) from orange to green: $\mathrm{K_2Cr_2O_7(aq) + H_2SO_4(aq) + 3SO_2(g) \rightarrow K_2SO_4(aq) + Cr_2(SO_4)_3(aq) + H_2O(l)}$
It turns acidified potassium manganate(VII) from purple to colourless: $\mathrm{5SO_2(g) + 2KMnO_4(aq) + 2H_2O(l) \rightarrow K_2SO_4(aq) + 2MnSO_4(aq) + 2H_2SO_4(aq)}$

3. A bleaching agent. Moist sulfur dioxide bleaches coloured flowers, wool, silk and straw. It works by reduction: it removes oxygen from the coloured substance, turning it into a colourless one. The bleaching is not permanent: over weeks, oxygen in the air re-oxidises the substance and the colour slowly returns.

4. An oxidising agent, with stronger reducing agents. Mixed with hydrogen sulfide, it forms a yellow deposit of sulfur: $\mathrm{SO_2(g) + 2H_2S(g) \rightarrow 3S(s) + 2H_2O(l)}$ . Burning magnesium keeps burning in it: $\mathrm{2Mg(s) + SO_2(g) \rightarrow 2MgO(s) + S(s)}$ .

5. With oxygen. With a catalyst, it combines with more oxygen to form sulfur trioxide, the key step in making sulfuric acid (Lesson 3): $\mathrm{2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g)}$ .

Definition

Bleaching: removing the colour from a substance by a chemical reaction. Sulfur dioxide bleaches by reduction, which can slowly be undone by air; chlorine (Unit 4) bleaches by oxidation, which is permanent. Example: straw hats and wool are bleached with sulfur dioxide because chlorine would damage the fibres.

Table 2.1: Sulfur dioxide compared with carbon dioxide

Property	Sulfur dioxide, $\mathrm{SO_2}$	Carbon dioxide, $\mathrm{CO_2}$
Smell	sharp, choking	none
$M_{r}$ ; collected by	64; downward delivery	44; downward delivery
Solubility in water	very soluble	fairly soluble
Acidified potassium dichromate(VI)	orange → green	no change
Acidified potassium manganate(VII)	purple → colourless	no change
Limewater	turns milky (calcium sulfite forms)	turns milky
Bleaches damp coloured petals	yes	no

Both gases turn limewater milky, so limewater cannot tell them apart. Use acidified potassium dichromate(VI): only sulfur dioxide turns it green.

Uses of sulfur dioxide

Making sulfuric acid by the Contact process (Lesson 3).
A food preservative (food additive E220) in dried fruit, fruit juices and wine: it kills bacteria and moulds and stops fruit going brown.
Bleaching wood pulp for paper, and wool, silk and straw, which chlorine would damage.
Fumigating (killing pests in) stores and containers.

Worked example

Problem. 6.3 g of sodium sulfite reacts completely with excess dilute hydrochloric acid. What volume of sulfur dioxide, at room temperature and pressure, is produced? (Na = 23, S = 32, O = 16; molar volume $24\,\mathrm{dm^3\,mol^{-1}}$ )

Solution. $M(\mathrm{Na_2SO_3}) = 2 \times 23 + 32 + 3 \times 16 = 126\,\mathrm{g\,mol^{-1}}$ , so $n = \frac{6.3}{126} = 0.050\,\mathrm{mol}$ .

1 mol $\mathrm{Na_2SO_3}$ gives 1 mol $\mathrm{SO_2}$ , so volume $= 0.050 \times 24 = 1.2\,\mathrm{dm^3}$ .

Worked example

Problem. A gas X turns acidified potassium manganate(VII) from purple to colourless. Damp red rose petals placed in a jar of X turn white. After a few days in the open air, the petals turn pink again. Identify X and explain the colour changes of the petals.

Solution. X decolourises acidified potassium manganate(VII), so it is a reducing agent, and it bleaches, so X is sulfur dioxide. It bleaches the petals by reduction, removing oxygen from the red dye to form a colourless compound. In air, oxygen slowly re-oxidises the compound back to the coloured dye, so the colour returns. (If the bleach had been chlorine, which bleaches by oxidation, the petals would have stayed white.)

Tip

SO₂ is the green-maker. Orange dichromate turns green, purple manganate(VII) turns colourless: both because sulfur dioxide is a reducing agent. Limewater cannot tell it from carbon dioxide.

Sergei's Chemistry: Sulfur burns in oxygen; properties of sulfur dioxide (8 min). Watch the blue flame of burning sulfur, the indicator turning red, and the dichromate paper turning from orange to green.

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Practice

Write the equation for the preparation of sulfur dioxide from sodium sulfite and dilute hydrochloric acid.
Explain why sulfur dioxide is (a) dried with concentrated sulfuric acid, (b) collected by downward delivery.
Describe the test for sulfur dioxide and the result.
Why can limewater not be used to tell sulfur dioxide from carbon dioxide?
Write equations to show that sulfur dioxide is an acidic oxide.
Explain the difference between bleaching by sulfur dioxide and bleaching by chlorine.
Give two uses of sulfur dioxide and the property each depends on.
What mass of copper reacts with excess hot concentrated sulfuric acid to give $480\,\mathrm{cm^3}$ of sulfur dioxide at room temperature and pressure? (Cu = 64; molar volume $24\,\mathrm{dm^3\,mol^{-1}}$ )

Show answersHide answers

$\mathrm{Na_2SO_3(s) + 2HCl(aq) \rightarrow 2NaCl(aq) + H_2O(l) + SO_2(g)}$ .
(a) It is an acidic gas, so it does not react with the acid, which removes the water. (b) It is much denser than air ( $M_{r}$ 64 against 29) and very soluble in water.
Pass the gas over (or through) acidified potassium dichromate(VI): it turns from orange to green.
Both turn limewater milky (sulfur dioxide forms insoluble calcium sulfite).
$\mathrm{SO_2(g) + H_2O(l) \rightleftharpoons H_2SO_3(aq)}$ and $\mathrm{SO_2(g) + 2NaOH(aq) \rightarrow Na_2SO_3(aq) + H_2O(l)}$ .
Sulfur dioxide bleaches by reduction, and the colour can slowly return as air re-oxidises the substance. Chlorine bleaches by oxidation, which is permanent.
Any two: food preservative (kills microbes, a reducing agent); bleaching wool or paper pulp (reduces coloured substances); making sulfuric acid (combines with oxygen to form $\mathrm{SO_3}$ ); fumigant (toxic to pests).
$n(\mathrm{SO_2}) = \frac{0.480}{24} = 0.0200\,\mathrm{mol}$ ; 1 mol Cu gives 1 mol $\mathrm{SO_2}$ , so mass of Cu $= 0.0200 \times 64 = 1.28\,\mathrm{g}$ .

Sources and credits

Chemistry 2e, 18.10: Occurrence, Preparation, and Properties of Sulfur (sulfur dioxide and its reactions), Paul Flowers, Klaus Theopold, Richard Langley and William R. Robinson, OpenStax, via LibreTexts, licensed under CC BY 4.0. Adapted by Ganzaa.
Video: Sulfur burns in oxygen: properties of sulfur dioxide; test with acidified potassium dichromate paper, Sergei's Chemistry, embedded from YouTube.
The diagram, table, worked examples and Practice items were written and drawn by Ganzaa.

1The smell of a struck match

2Preparing sulfur dioxide in the laboratory

3Physical properties

4Chemical properties

5Uses of sulfur dioxide

The smell of a struck match

Preparing sulfur dioxide in the laboratory

Physical properties

Chemical properties

Uses of sulfur dioxide