Boiling point elevation is ΔTb = i × Kb × m, where i is the van’t Hoff factor, Kb is the solvent’s ebullioscopic (boiling-point) constant, and m is the molality of the solute in mol/kg. The new boiling point is the pure solvent’s boiling point plus ΔTb. For example, 1 mol/kg of sugar in water raises the boiling point by 1 × 0.512 × 1 = 0.512 °C, to 100.5 °C.

The van’t Hoff factor (i) is the number of dissolved particles produced per formula unit of solute. Non-electrolytes like sugar or urea stay as single molecules, so i = 1. Ionic compounds dissociate: NaCl gives Na⁺ and Cl⁻ (i ≈ 2), and CaCl₂ gives one Ca²⁺ and two Cl⁻ (i ≈ 3). Colligative effects depend on particle count, so a higher i means a larger boiling-point rise.

The ebullioscopic constant Kb is a property of the solvent that sets how strongly its boiling point rises per unit molality. For water Kb = 0.512 °C·kg/mol, meaning a 1 mol/kg ideal solution boils 0.512 °C higher. Other solvents have larger constants — benzene is 2.53 and chloroform 3.63 °C·kg/mol — so the same molality raises their boiling points much more.

Dissolved particles lower the vapor pressure of the solvent, so a higher temperature is needed before the vapor pressure reaches atmospheric pressure and the liquid boils. This is a colligative property — it depends on the number of dissolved particles, not their identity. Salting pasta water raises its boiling point only slightly (a teaspoon per litre adds well under 0.1 °C), so the effect on cooking time is negligible.