Q.Give reasons for the following :

(a) Dioxygen is a gas but sulphur a solid

Concept: Molecular structure and intermolecular forces.

• $O_2$ is a small, diatomic molecule held by weak London dispersion forces — hence a gas at room temperature. …

(a) Dioxygen is a gas but sulphur a solid.

Concept & Intuition

The physical state of an element depends on the strength of intermolecular forces, which in turn depends on molecular size and shape. Oxygen exists as small, diatomic O₂ molecules. Sulphur, however, forms larger, ring-shaped S₈ molecules. Bigger molecules have more electrons and larger surface areas, leading to stronger London dispersion forces. These forces hold S₈ molecules together as a solid at room temperature, while the weak forces between O₂ molecules keep it a gas.

Step-by-step reasoning

1. Molecular form of oxygen

   Oxygen exists as O₂ — a linear, diatomic molecule with only 16 electrons. The van der Waals (London) forces between such small molecules are very weak. At room temperature, these forces are easily overcome by thermal energy, so O₂ remains a gas.

2. Molecular form of sulphur

   Sulphur, in its most stable allotrope (rhombic sulphur), forms S₈ rings. Each ring has 8 atoms, giving a much larger electron cloud and molecular mass (256 g/mol vs 32 g/mol for O₂). The larger surface area and more electrons produce significantly stronger London dispersion forces.

3. Comparison of melting/boiling points

   • O₂: m.p. –218°C, b.p. –183°C (gas at room temp)
   • S₈: m.p. 115°C, b.p. 444°C (solid at room temp)

   The stronger intermolecular forces in sulphur require much more energy to overcome, so it is a solid.

(b) NO (g) released by jet aeroplanes is slowly depleting the ozone layer.

Concept & Intuition

Ozone (O₃) in the stratosphere absorbs harmful UV radiation. Nitric oxide (NO) acts as a catalyst in a chain reaction that converts ozone into oxygen. A catalyst is not consumed in the net reaction, so a single NO molecule can destroy thousands of ozone molecules before being removed.

Step-by-step reasoning

1. Source of NO Jet aeroplanes fly in the upper troposphere and lower stratosphere. At high combustion temperatures, nitrogen and oxygen from air combine:

This NO is released directly into the stratosphere.

2. The catalytic cycle

   NO reacts with ozone in two steps:

   • Step 1: $\text{NO} + \text{O}_3 \rightarrow \text{NO}_2 + \text{O}_2$
   • Step 2: $\text{NO}_2 + \text{O} \rightarrow \text{NO} + \text{O}_2$ (The oxygen atom comes from photolysis of O₃ or O₂.)

3. Net reaction

   Adding the two steps:

Notice that NO is regenerated — it is not used up. It can go on to destroy more ozone.

4. Why it’s “slowly depleting” Each NO molecule can cycle through this reaction many times before being removed (e.g., by reacting with OH or being transported away). Even small concentrations of NO cause significant ozone loss over time. This is a chain reaction, and NO is a free radical catalyst.

(c) Interhalogens are more reactive than pure halogens.

Concept & Intuition

Interhalogens are compounds of two different halogens, e.g., ClF, BrF₃, ICl. Their reactivity is higher because the bond between two different halogens is polar — one atom pulls electron density more than the other. This polarity makes the bond weaker and more easily broken than the nonpolar bond in a pure halogen (e.g., Cl–Cl). Also, the more electronegative halogen carries a partial negative charge, making the molecule susceptible to attack.

Step-by-step reasoning

1. Bond polarity …