Elevation of Boiling Point

Key Concepts

Boiling Point: The temperature at which a liquid's vapor pressure equals atmospheric pressure (760 mm of Hg).
Example: Water boils at 100°C. When sugar is dissolved, the boiling point increases to 102°C.

Adding a non-volatile solute lowers the solvent's vapor pressure.
According to Raoult's Law, the vapor pressure of the solution is always lower than that of the pure solvent.
The solution must be heated to a higher temperature to equal atmospheric pressure, thus increasing the boiling point.

Graph: Vapor Pressure vs. Temperature
- Vapor pressure of the solvent is always higher than that of the solution at a given temperature.
- The difference in boiling points (TB - T°B) is termed as the Elevation of Boiling Point (ΔTB).

As solute concentration increases, the solution's vapor pressure decreases further, increasing the boiling point difference.
ΔTB is directly proportional to the molarity of the solution.
- Elevation of boiling point is directly proportional to the number of moles of solute.
- Independent of the nature of the solute.

Elevation of boiling point is considered a colligative property.
Formula: ΔTB = KB × M
- Where KB is the molar elevation constant (also known as the ebullioscopic constant) and M is the molarity of the solution.

Elevation in boiling point can be used to determine the molar mass of an unknown non-volatile compound.
Method:
- Dissolve a known mass of non-volatile solute (WB) in a known mass of solvent (WA).
- Determine the elevation in boiling point (ΔTB).
- Formulas:
  - Molality (M) = (moles of solute / 2000) by weight of solvent (in grams)
  - Moles of solute = WB / MB (Molar Mass)
  - M = (WB / MB) × (2000 / WA)
  - ΔTB = KB × molality
From the above relations, the molar mass of the solute can be calculated using the known values of other quantities.