Elevation of boiling point

A liquid boils when its vapour pressure equals the external pressure: water boils at 373.15 K because there its vapour pressure is 1.013 bar (1 atm).

A non-volatile solute lowers the vapour pressure, so the solution has to be heated past the solvent's boiling point before its vapour pressure reaches 1.013 bar. Such a solution therefore always boils higher than its solvent: a solution of 1 mol of sucrose per 1000 g of water, at 1 atm, boils at 373.52 K.

The elevation ΔTb = Tb − Tb° is, for dilute solutions, directly proportional to molality: ΔTb = Kb m.

Kb, the molal elevation (ebullioscopic) constant, has units K kg mol⁻¹ and depends only on the solvent: 0.52 for water, 2.53 for benzene.

With w₂ g of solute of molar mass M₂ in w₁ g of solvent, m = (w₂ × 1000)/(M₂ × w₁), so M₂ = (Kb × w₂ × 1000)/(ΔTb × w₁).

Kb = R × M₁ × Tb² / (1000 × ΔvapH), where M₁ is the solvent's molar mass, Tb its boiling point in kelvin and ΔvapH its enthalpy of vaporisation.