Boiling Point Elevation
Definition and meaning of Boiling Point Elevation in chemistry.
Boiling point elevation is the increase in a solvent's boiling point that occurs when a non-volatile solute is dissolved in it, relative to the boiling point of the pure solvent.
In more detail
Dissolved solute particles lower the solvent's vapor pressure (a consequence of Raoult's law), so the solution must be heated to a higher temperature before its vapor pressure reaches atmospheric pressure and boiling begins. As a colligative property, the magnitude of the effect depends only on the number of dissolved particles, not their chemical identity. It is calculated as ΔTb = i·Kb·m, where i is the van't Hoff factor (particles produced per formula unit), Kb is the solvent's molal boiling point elevation constant, and m is the solution's molality.
Key facts
| Formula | ΔTb = i·Kb·m |
|---|---|
| Field | Physical Chemistry |
| Water's Kb | 0.512 °C·kg/mol |
| Property type | Colligative (depends on particle count, not identity) |
Dissolving 1 mole of sucrose in 1 kg of water gives a 1 m solution. Since sucrose does not dissociate (i = 1) and water's Kb is 0.512 °C·kg/mol, ΔTb = 1 × 0.512 × 1 = 0.512°C, raising the boiling point from 100.000°C to about 100.512°C at 1 atm.
Frequently asked questions
Does the identity of the solute affect boiling point elevation?
No. As a colligative property, ΔTb depends only on the number of dissolved particles per unit of solvent, not on what the solute is chemically, though ionic solutes that dissociate (like NaCl, with i ≈ 2) produce a larger effect than nonionic solutes (like sucrose, with i = 1) at the same molality.
Why must molality, not molarity, be used in the equation?
Molality (mol solute per kg solvent) is temperature-independent, since it is based on mass rather than solution volume, which changes slightly with temperature.