# Boiling Point Elevation by a Non-Volatile Solute

The boiling point of a liquid is the temperature at which its vapour pressure becomes equal to 760 mm (i.e. 1 atmospheric pressure).
Since the addition of a non-volatile solute lowers the vapour pressure of the solvent, the vapour pressure of a solution is always lower than that of the pure solvent, and hence it must be heated to a higher temperature to make its vapour pressure equal to atmospheric pressure. Thus the solution boils at a higher temperature than the pure solvent.

If T°b is the boiling point of the solvent and Tb is the boiling point of the solution, the difference in boiling points (ΔTb) is called the elevation of boiling point.

Thus, Tb − T°b = ΔTb

ΔTb ∝ molality where ΔTb = elevation of boiling point

n = no. of moles of non-volatile solute

N = Total no. of moles in the solution

This implies that the boiling point elevation in a dilute solution is directly proportional to the number of moles of the solute dissolved in a given amount of the solvent and is quite independent of the nature of the solute. Hence, boiling point elevation is a colligative property.

ΔTb = Kb × m

Kb : molal elevation constant or Ebullioscopic constant

m : molality of the solution

Molal boiling point elevation constant or ebullioscopic constant of the solvent, is defined as the elevation in boiling point which may theoretically be produced by dissolving one mole of any solute in 1000 g of the solvent.

or , $\large \Delta T_b = \frac{1000 \times K_b \times w}{m_1 \times W}$

where m1 = molecular weight of solute and w and W are weights of solute and solvent