Raoult's Law Calculator (Boiling Point Elevation & Freezing Point Depression)
Calculate boiling point elevation and freezing point depression from a solute's molality and van't Hoff factor using Raoult's law. Supports six solvent constants (Kb/Kf) including water and benzene, and accounts for electrolyte dissociation.
Molal boiling/freezing point constants by solvent
Standard reference values widely cited in textbooks. Use them to check your own hand calculations.
| Solvent | Kb (°C·kg/mol) | Kf (°C·kg/mol) | Normal bp (°C) | Normal fp (°C) |
|---|---|---|---|---|
| Water | 0.512 | 1.86 | 100.0 | 0.0 |
| Benzene | 2.53 | 5.12 | 80.1 | 5.5 |
| Acetic acid | 3.07 | 3.9 | 118.1 | 16.6 |
| Cyclohexane | 2.79 | 20.2 | 80.7 | 6.5 |
| Naphthalene | 5.8 | 6.9 | 218.0 | 80.2 |
| Camphor | 5.95 | 37.7 | 204.0 | 178.4 |
What Raoult's law actually states
**The first thing Raoult's law says is that dissolving something in a solvent lowers its vapour pressure.** The rise in boiling point and the fall in freezing point are merely consequences drawn from that single fact. Since a liquid boils at the temperature where its vapour pressure matches the surrounding pressure, lowering the vapour pressure pushes that temperature higher. The elevation is written Δt = Kb × m × i and the depression Δt = Kf × m × i, where Kb and Kf are constants fixed for each solvent.
**It matters that the concentration used here is molality rather than molarity.** The volume of a solution shifts with temperature, so a volume-based measure cannot serve as a reference when the very quantity under discussion is a temperature. Counting per kilogram of solvent leaves the figure untouched however the thermometer moves. The other key quantity is the van 't Hoff factor, which records **how many particles one dissolved unit breaks into.** Colligative properties depend on the number of particles and not on their identity, so for an electrolyte the answer will not come out right unless this factor is applied.
This tool covers six solvents, including water, benzene and camphor, and accepts the concentration either directly as a molality or worked out from the mass and molar mass of the solute together with the mass of the solvent. It then returns both the change and the resulting boiling and freezing points.
How to run a calculation
- Choose the solvent Six are available, among them water, benzene and acetic acid; choosing one fills in Kb and Kf automatically.
- Decide how to supply the concentration You may enter the molality directly, or **have it worked out from the mass and molar mass of the solute and the mass of the solvent.**
- Enter the van 't Hoff factor Use 1 for a non-electrolyte, 2 for sodium chloride and 3 for calcium chloride as a first approximation.
- Read off the elevation and depression Both the size of the change and the resulting boiling and freezing points are shown.
Tips for getting more out of it
- Electrolytes such as NaCl or CaCl2 dissociate into multiple ions in water, so setting the van't Hoff factor i above 1 brings the result closer to real measurements.
- In the mass-based mode, just enter the mass of solute [g], its molar mass [g/mol], and the mass of solvent [kg] to have the molality calculated automatically.
- Both boiling point elevation and freezing point depression are colligative properties: they depend only on the number of solute particles, not on what the solute actually is.
- For a given solute amount, freezing point depression is often larger than boiling point elevation for common solvents, which is one reason calcium chloride is used as a road de-icer.
- Use the i preset buttons (non-electrolyte / NaCl-like / CaCl2-like) to quickly try out common dissociation patterns.
Where it comes in useful
Checking a chemistry assignment
Problems on colligative properties go wrong most often through a forgotten factor, which makes them worth verifying.
Estimating how well a de-icer works
**Calcium chloride outperforms sodium chloride on ice because its factor is 3 rather than 2**, so the same molality buys a larger depression.
Retracing how molar masses were measured
Freezing-point depression was the classical route to the molar mass of an unknown, and you can follow the relation in reverse.
Planning an experiment
You can work out in advance how much solute a target freezing point will require, before touching the balance.
Terms in colligative properties
- Raoult's law
- The statement that the vapour pressure of a solution is proportional to the mole fraction of the solvent. **Boiling-point elevation and freezing-point depression both follow from it.**
- Molality
- The amount of solute per kilogram of solvent. **Because no volume enters into it, the figure holds steady as the temperature changes.**
- Molal boiling-point constant
- The rise in boiling point at unit molality, fixed for each solvent. For water it is 0.515 kelvin per unit.
- Molal freezing-point constant
- The corresponding fall in freezing point. For water it is 1.853, while **camphor's value of 37.7 is so large that it was long used to measure minute quantities.**
- Van 't Hoff factor
- The number of particles one formula unit of solute yields in solution. It is 1 for a non-electrolyte and 2 in theory for sodium chloride.
- Colligative property
- A property governed by the number of dissolved particles and not by what they are. The four are vapour-pressure lowering, boiling-point elevation, freezing-point depression and osmotic pressure.
Frequently Asked Questions
Side Note — De-icers and the chemistry of colligative properties
The white granules scattered on winter roads are often calcium chloride (CaCl2) or sodium chloride (NaCl). These substances melt ice and snow not by warming them, but by lowering the freezing point of water itself once dissolved in it. This phenomenon, freezing point depression, is a classic example of the colligative properties that follow from Raoult's law.
The word "colligative" captures the idea that only the number of particles matters, not their identity. At the same molality, CaCl2 dissociates in water into Ca2+ and two Cl- ions — three particles in total — producing roughly three times the freezing point depression of a non-dissociating non-electrolyte like sucrose. That's one reason CaCl2 can act as a more powerful de-icer than NaCl in smaller quantities (though its corrosive effect on concrete is a separate consideration).
The same principle applies in the opposite direction to boiling point elevation. Adding salt to pasta water does raise its boiling point slightly, but at the amounts used in home cooking the rise is well under 1°C — nowhere near enough to make the water boil noticeably faster. The main reason cooks add salt is really to season the pasta; the boiling point elevation is just a minor chemical side effect.