Ozone, sulphur and sulphur dioxide

The smell after a thunderstorm

After a thunderstorm there is a clean, sharp smell in the air. That smell is ozone, $\text{O}_{3}$, an allotrope of oxygen. Allotropes are different forms of the same element in the same state. Ozone has three oxygen atoms in each molecule, and the oxygen of the air has two.

Ozone

Making it. A slow stream of dry oxygen is passed through a silent electric discharge. This is a discharge without a spark. A part of the oxygen changes to ozone: $3\text{O}_{2}\ \rightleftharpoons \ 2\text{O}_{3}$. The gas that leaves is "ozonised oxygen" with up to about 10 percent ozone. The change takes in heat. The standard enthalpy of formation of ozone is about $+142.7$ kJ/mol, so ozone is less stable than oxygen. It is a gas of the stratosphere, where sunlight makes it.

Structure. Ozone is bent, with an angle of 117°. Its two O-O bonds are equal. Each has a length of 128 pm, between a single and a double bond. We write two resonance forms, and the real molecule is a mix of them. The bond order is 1.5.

Properties. Ozone is a pale blue gas with a pungent smell. It is a strong oxidising agent, because it gives up an oxygen atom easily: $\text{O}_3 \rightarrow \text{O}_2 + [\text{O}]$.

• It turns lead(II) sulphide into lead(II) sulphate: $\text{PbS}\ +\ 4\text{O}_{3}\ \rightarrow \ \text{PbSO}_{4}\ +\ 4\text{O}_{2}$. This is how old paintings with blackened lead are restored.
• It frees iodine from potassium iodide: $2\text{KI}\ +\ \text{H}_{2}\text{O}\ +\ \text{O}_{3}\ \rightarrow \ 2\text{KOH}\ +\ \text{I}_{2}\ +\ \text{O}_{2}$. The iodine turns starch blue, and the same reaction is used to measure ozone.
• It makes mercury lose its shiny curved surface. This is the "tailing" of mercury.
• It bleaches many dyes and kills germs, so it is used to purify drinking water.

The ozone layer. About 15 to 35 km above the ground, the ozone absorbs most of the harmful ultraviolet light of the sun. Some man-made compounds destroy it. The chlorofluorocarbons, or CFCs, such as $\text{CF}_{2}\text{Cl}_{2}$, drift up and are broken by ultraviolet light to give chlorine atoms. A chlorine atom has an unpaired electron, and A-Level and IB books write it as a radical, $\text{Cl}^{\bullet}$. NCERT writes it as Cl. The chain is:

• $\text{Cl}^{\bullet} + \text{O}_3 \rightarrow \text{ClO}^{\bullet} + \text{O}_2$
• $\text{ClO}^{\bullet} + \text{O} \rightarrow \text{Cl}^{\bullet} + \text{O}_2$

The chlorine atom comes back in the second step. It is a catalyst, so one atom can destroy many thousands of ozone molecules. Nitrogen oxide, $\text{NO}$, from aircraft works in the same way. Most countries have stopped using CFCs.

Sulphur

Sulphur is a yellow solid found as the free element near volcanoes and in the ores of metals, as sulphides and sulphates. It has several allotropes. Two are important.

• Rhombic sulphur (alpha) is stable at room temperature. It forms yellow crystals.
• Monoclinic sulphur (beta) is stable above 369 K. It forms needle-shaped crystals.

The temperature of 369 K, where the two forms are in equilibrium, is the transition temperature. Both forms are made of puckered rings of eight atoms, $\text{S}_{8}$. The ring is shaped like a crown. They differ in how the rings are packed in the crystal. Both dissolve in carbon disulphide, and neither dissolves in water. In the vapour near 1000 K, the main species is $\text{S}_{2}$. It has two unpaired electrons, like $\text{O}_2$, and is paramagnetic. Sulphur differs from oxygen in forming rings, because the S-S single bond is strong. The O-O single bond is weak.

Start with 600 kg of an ore that is 80 percent pyrites. You get 512 kg of sulphur dioxide. Raise the purity to 100 percent. The blue line moves up to meet the grey one.

Practise with new numbers

Sulphur dioxide

Making it. Burning sulphur in air gives it: $\text{S}\ +\ \text{O}_{2}\ \rightarrow \ \text{SO}_{2}$. In industry, it comes from roasting metal sulphides, such as iron pyrites: $4\text{FeS}_{2}\ +\ 11\text{O}_{2}\ \rightarrow \ 2\text{Fe}_{2}\text{O}_{3}\ +\ 8\text{SO}_{2}$. In the laboratory, a sulphite and a dilute acid give it: $\text{Na}_{2}\text{SO}_{3}\ +\ 2\text{HCl}\ \rightarrow \ 2\text{NaCl}\ +\ \text{H}_{2}\text{O}\ +\ \text{SO}_{2}$.

Structure. The molecule is bent, with an angle of about 119°. The two S-O bonds are equal, 143 pm long, and the bond order is 1.5, as in ozone.

Properties. Sulphur dioxide is a colourless gas with a choking smell. It is easy to liquefy and very soluble in water.

• It is an acidic oxide. It makes sulphurous acid, $\text{SO}_{2}\ +\ \text{H}_{2}\text{O}\ \rightleftharpoons \ \text{H}_{2}\text{SO}_{3}$, which turns litmus red.
• With sodium hydroxide: $\text{SO}_{2}\ +\ 2\text{NaOH}\ \rightarrow \ \text{Na}_{2}\text{SO}_{3}\ +\ \text{H}_{2}\text{O}$. With more gas: $\text{Na}_{2}\text{SO}_{3}\ +\ \text{H}_{2}\text{O}\ +\ \text{SO}_{2}\ \rightarrow \ 2\text{NaHSO}_{3}$.
• With chlorine, in the presence of charcoal: $\text{SO}_{2}\ +\ \text{Cl}_{2}\ \rightarrow \ \text{SO}_{2}\text{Cl}_{2}$. The product is sulphuryl chloride.
• With oxygen, using a catalyst of vanadium(V) oxide: $2\text{SO}_{2}\ +\ \text{O}_{2}\ \rightleftharpoons \ 2\text{SO}_{3}$. This is the second step of the contact process, in the next lesson.

Sulphur is at $+4$ in $\text{SO}_2$. It can go up to $+6$ or down to 0, so the gas is both a reducing agent and a weaker oxidising agent.

• It reduces iron(III) to iron(II): $2\text{Fe}^{3+}\ +\ \text{SO}_{2}\ +\ 2\text{H}_{2}\text{O}\ \rightarrow \ 2\text{Fe}^{2+}\ +\ \text{SO}_{4}^{2-}\ +\ 4\text{H}^{+}$.
• It decolourises purple acidified permanganate: $5\text{SO}_{2}\ +\ 2\text{MnO}_{4}^{-}\ +\ 2\text{H}_{2}\text{O}\ \rightarrow \ 5\text{SO}_{4}^{2-}\ +\ 2\text{Mn}^{2+}\ +\ 4\text{H}^{+}$.
• It turns orange acidified dichromate green: $\text{Cr}_{2}\text{O}_{7}^{2-}\ +\ 3\text{SO}_{2}\ +\ 2\text{H}^{+}\ \rightarrow \ 2\text{Cr}^{3+}\ +\ 3\text{SO}_{4}^{2-}\ +\ \text{H}_{2}\text{O}$. The green is from chromium(III).
• It oxidises hydrogen sulphide, $\text{SO}_{2}\ +\ 2\text{H}_{2}\text{S}\ \rightarrow \ 3\text{S}\ +\ 2\text{H}_{2}\text{O}$. Here it is the oxidising agent.

Bleaching. Moist sulphur dioxide bleaches coloured materials by reduction. The colour often comes back in air, because oxygen oxidises the product. This is a temporary bleaching. The bleaching by chlorine is by oxidation and is permanent.

Uses and harm. The gas is used to make sulphuric acid, to bleach wool and silk, and as a preservative and a disinfectant. It is also the main cause of acid rain. Normal rain has a pH of about 5.6 because of dissolved carbon dioxide. Rain with sulphur dioxide and its oxidation products has a lower pH. It corrodes marble and metals and harms plants and fish.

Try it in Python

The program finds the amount of sulphur dioxide in a water sample from a titration with potassium permanganate. It uses the ratio 5 : 2 from the equation.

Common mistakes

• Saying that sulphur dioxide bleaches by oxidation. It bleaches by reduction, and the colour can return.
• Writing the ozone molecule as straight. It is bent, at 117°.
• Thinking that the bond in ozone is a single bond on one side and a double bond on the other. The two bonds are equal, because of resonance.
• Forgetting the 5 : 2 ratio of sulphur dioxide to permanganate.
• Calling the chlorine atom a reactant used up in the ozone chain. It is regenerated, so it is a catalyst.

Remember these

• Ozone: $3\text{O}_{2}\ \rightleftharpoons \ 2\text{O}_{3}$ in a silent discharge, bent at 117°, bond order 1.5, strong oxidising agent, absorbs ultraviolet light.
• CFCs give chlorine atoms, which catalyse the destruction of ozone.
• Sulphur: rhombic and monoclinic forms, both $\text{S}_{8}$ crowns, transition temperature 369 K.
• $\text{SO}_{2}$: bent, acidic, reducing, a temporary bleach, and a cause of acid rain.
• Test: it turns acidified dichromate green and decolourises permanganate.

Check yourself

Try each one on your own first. If you are stuck, open a hint. Then check with Show the answer.

1. Why is ozone a stronger oxidising agent than oxygen?

2. What is the oxidation state of sulphur in $\text{SO}_{2}$, and why can the gas act as both an oxidising and a reducing agent?

3. How can you tell rhombic sulphur from monoclinic sulphur?

4. Name the colour change when sulphur dioxide is passed into acidified potassium dichromate solution. Name the species that causes the new colour.

References

• NCERT, Chemistry Part I for Class XII, the chapter The p-Block Elements, on ozone, sulphur and sulphur dioxide.
• OpenStax, Chemistry 2e, Chapter 18, on sulphur and its compounds.