Boiling Point Elevation

Boiling point elevation occurs when the boiling point of a solution becomes higher than the boiling point of a pure solvent. The temperature at which the solvent boils is increased by adding any non-volatile solute. A common example of boiling point elevation can be observed by adding salt to water. The boiling point of the water is increased (although in this case, not enough to affect the cooking rate of food).

Boiling point elevation, like freezing point depression, is a colligative property of matter. This means it depends on the number of particles present in a solution and not on the type of particles or their mass. In other words, increasing the concentration of the particles increases the temperature at which the solution boils.

How Boiling Point Elevation Works

In a nutshell, boiling point increases because most of the solute particles remain in the liquid phase rather than enter the gas phase. In order for a liquid to boil, its vapor pressure needs to exceed ambient pressure, which is harder to achieve once you add a nonvolatile component. If you like, you could think of adding a solute as diluting the solvent. It doesn't matter whether the solute is an electrolyte or not. For example, boiling point elevation of water occurs whether you add salt (an electrolyte) or sugar (not an electrolyte).

Boiling Point Elevation Equation

The amount of boiling point elevation can be calculated using the Clausius-Clapeyron equation and Raoult's law. For an ideal dilute solution:

Boiling Point_total = Boiling Point_solvent + ΔT_b

where ΔT_b = molality * K_b * i

with K_b = ebullioscopic constant (0.52°C kg/mol for water) and i = Van't Hoff factor

The equation is also commonly written as:

ΔT = K_bm

The boiling point elevation constant depends on the solvent. For example, here are constants for some common solvents: