Statement I: The boiling point of a 0.1 m urea solution is lower than…

The correct answer is Option (1) → Both Statement I and Statement II are correct and Statement II is the correct explanation of Statement I.

Elevation of boiling point is a colligative property: it depends on the number of solute particles, not on their nature. With the van’t Hoff factor i counting the particles each formula unit gives:

    \( \displaystyle \Delta T_b=i\,K_b\,m \)

Statement I: The boiling point of a 0.1 m urea solution is lower than that of a 0.1 m KCl solution. — True. Urea is a non-electrolyte, so i = 1. KCl dissociates, \( \ce{KCl -> K^+ + Cl^-} \), so i is close to 2. At the same molality the KCl solution has about twice as many solute particles, a larger ΔT_b and so a higher boiling point.

Statement II: The elevation of boiling point of a solution is directly proportional to the number of solute particles present in a given amount of the solvent. — True. This is what makes ΔT_b a colligative property.

Why Statement II explains Statement I? Both solutions have the same molality, so they differ only in the number of particles each gives in the solvent. Since the elevation is proportional to that number, the urea solution (fewer particles) shows the smaller elevation and boils at the lower temperature.

NCERT (2026-27 edition):
1.6.2 Elevation of Boiling Point – “… Similar to lowering of vapour pressure, the elevation of boiling point also depends on the number of solute molecules rather than their nature.”

1.7 Abnormal Molar Masses – “… For example, if we dissolve one mole of KCl (74.5 g) in water, we expect one mole each of K⁺ and Cl^– ions to be released in the solution. If this happens, there would be two moles of particles in the solution. If we ignore interionic attractions, one mole of KCl in one kg of water would be expected to increase the boiling point by 2 × 0.52 K = 1.04 K.”