Boiling-Point Elevation
Also known as colligative boiling
Worked example: Kb = 0.512, b = 1 mol/kg → dT = 0.512 K — press Try an example to run it live, then adjust anything.
Enter your known values, leave one input blank, and solves for the missing one. Tap a variable’s symbol to see what it means, with a typical value. Try different units for next level excitement!
Colligative counting →
Grade 12Grade 12 Chemistry
Test your skills in the Exam Room: new numbers every attempt — free lessons for students, no sign-up, just pure learning.
share your results
Boiling-Point Elevation explained
A non-volatile solute lowers the solvent's vapor pressure, so the solution must be heated a little hotter before that pressure reaches the atmosphere's — the boiling point rises. The rise depends only on how many particles are dissolved, not what they are, scaled by the solvent's ebullioscopic constant Kb. Water's Kb is a modest 0.512 K·kg/mol: a hearty 1.0 mol/kg sugar solution boils at just 100.5 °C, which is why salting pasta water changes its flavor far more than its physics.
Other solvents respond much more strongly — benzene's Kb is 2.53 and camphor's freezing-side cousin is famously huge — and remember that ionic solutes count each ion: 1 mol/kg NaCl acts as nearly 2 mol/kg of particles. Solved backwards, a measured ΔTb divided by Kb gives the molality, the classical route to a dissolved compound's molar mass before mass spectrometers existed.
Boiling-Point Elevation formula
- = Boiling-point elevation (C°)
- = Ebullioscopic constant (K·kg/mol)
- = Molality of solution (mol/kg)
Missing one of these? Work it out first, then come back
- Boiling-point elevation — Freezing-Point Depression, Freezing-Point Depression with the van 't Hoff Factor
- Ebullioscopic constant — Freezing-Point Depression, Freezing-Point Depression with the van 't Hoff Factor
- Molality of solution — Freezing-Point Depression, Freezing-Point Depression with the van 't Hoff Factor