Elevation Boiling Point

Bookmark added to your notes.
View Notes

What is Boiling Point Elevation?

The elevation of the boiling point refers to the rise of a solvent's boiling point upon the addition of a solute. The resulting solution has a higher boiling point when a non-volatile solute is applied to a solvent than that of the pure solvent. The boiling point of sodium chloride (salt) solution and water, for example, is higher than that of pure water.

The elevation of boiling points is a colligative property of matter, i.e., it depends on the solvent-to-solvent ratio but not on the identity of the solute. This means that the elevation of a solution's boiling point depends on the amount of solution applied to it. The higher the solute concentration in the solution, the greater the elevation of the boiling point.

Boiling Point Elevation

The vapor pressure of a solvent can decrease when a solution is applied. This occurs because of the solute displacement of solvent molecules. This means that some of the solvent molecules on the liquid's surface are replaced by the solvent; both electrolytic and non-electrolytic solutions will occur. The lower number of solvent molecules on the surface means that less can evaporate, thus reducing the vapor pressure. A higher temperature is needed for the vapor pressure to equal the ambient pressure, and a higher boiling point is observed.

(image will be uploaded soon)

A graph describing the elevation at the boiling point of water when sucrose is added is given above. At a pressure of 1atm, pure water boils at 100℃. However, in water, a 10-molal solution of sucrose boils at around 105℃.

Why Does Boiling Point Elevation Occur?

The temperature at which its vapor pressure is equal to the pressure of its surrounding atmosphere is the boiling point of a liquid. Non-volatile liquids do not evaporate quickly and have very low vapor pressures (assumed to be zero). The vapor pressure of the resulting solution is lower than that of the pure solvent when a non-volatile solute is applied to the solvent.

Therefore the solution must be supplied with a larger amount of heat for it to boil. The boiling point elevation is this rise in the solution's boiling point. A rise in the concentration of the added solution is followed by a further decrease in the solution's vapor pressure and a further increase in the solution's boiling point.

A temperature graph of pressure v/s detailing the boiling point elevation of a solution is given below.

(image will be uploaded soon)

Here, ΔTb represents the elevation of the solution's boiling point. It can be observed from the graph that:

  • The solution's freezing point is lower than that of the pure solvent (freezing point depression).

  • The solution's boiling point is higher than that of the pure solvent.

Note: The liquid's boiling point often depends on the pressure of its surroundings (which is why water boils at temperatures lower than 100℃ at high altitudes, where the surrounding pressure is low).

Boiling Point Elevation Formula

The boiling point of a non-volatile solute containing solution can be expressed as follows:

Boiling point of solution = pure solvent boiling point + elevation of the boiling point.

The boiling point elevation (ΔTb) is proportional to the solute concentration in the solution. The following equation allows it to be measured.

ΔTb = i*Kb*m


  • It is the Van’t Hoff factor.

  • Kb is the ebullioscopic constant.

  • m is the molality of the solute.

It is important to remember that when the solute concentration is very high, this formula becomes less accurate. Also, this formula for volatile solvents does not hold true.

In terms of ℃/molal, or ℃.kg.mol-1, the ebullioscopic constant (Kb) is also expressed. Below the Kb values for some common solvents are tabulated.

Kb Values for Some Common Solvents


Kb Value ( in ℃.kg.mol-1 )





Acetic Acid






With the support of the boiling point elevation formula, the degree of dissociation of the solute and the molar mass of the solute can be measured.

The Relationship Between Boiling Point Elevation and Vapor Pressure

In terms of vapor pressure, boiling point elevation can be clarified. Vapor pressure is defined as the pressure exerted at a given temperature by a vapor in thermodynamic equilibrium with its condensed phases. It is simply a measure of the ability of the solvent molecules, in layman's words, to escape by entering the gas phase. When the vapor pressure is equal to the air pressure, a liquid boils.

Boiling Point - The boiling point of a liquid in its purest form. The liquid can boil when the vapor pressure of the liquid equals the ambient pressure.

FAQ (Frequently Asked Questions)

Question 1. Explain Boiling Point Elevation with Examples?

Ans. Salted water's boiling point is higher than pure water's boiling point. Salt is an electrolyte that dissociates into solution ions, thereby having a relatively significant effect on the point of boiling. Remember that non-electrolytes, including sugar, often raise the boiling point. However, since a nonelectrolyte does not dissociate to form several particles, the effect per mass is less than that of a soluble electrolyte.

Boiling Point Elevation Equation -

A combination of the Clausius-Clapeyron equation and Raoult's law is the formula used for measuring boiling point elevation. It is believed that the solvent is non-volatile.

ΔTb = Kb · bB


  • ΔTb is the boiling point elevation.

  • Kb is the ebullioscopic constant, which depends on the solvent.

  • bB is the molality of the solution (typically found in a table).

Thus, boiling point elevation is directly proportional to a chemical solution's molal concentration.

Question 2. 10 Grams of a Non-Volatile and Non-Dissociating Solute is Dissolved in 200 Grams of Benzene. The Resulting Solution Boils at a Temperature of 81.2℃. Find the Molar Mass of the Solute.

Ans. Let x = number of moles of solute. The boiling point of pure benzene is 80.1℃ and its ebullioscopic constant is 2.53℃/molal. From the boiling point elevation formula, the following relationship can be obtained:

(81.2℃ – 80.1℃) = (1)*(2.53℃.kg.mol-1)(x/0.2 kg).

x = (1.1℃ * 0.2kg)/(2.53℃.kg.mol-1).

x = 0.0869 moles.

Since 0.0869 moles of the solute has a mass of 10 grams, 1 mole of the solute will have a mass of 10/0.0869 grams, which is equal to 115.07 grams. Therefore, the molar mass of the solute is 115.07 grams per mole.