Colligative Properties of Solutions

Yes, the word is a mouthful, but it also sounds a lot like another word you are probably more familiar with. No, I don’t mean “Colgate”. “Collective” is the word I had in mind. Better yet, the two words (colligative and collective) are very similar in meaning. “Collective” refers to a group as a whole. “Colligative properties” are the collective properties of chemicals in solution.

Let’s make sure you’re up on your solution chemistry jargon before we go on. A solution is a homogenous (uniform) mixture of one (or more) chemical(s) dissolved in another. A solute is the chemical that is dissolved, and the solvent is the chemical the solute got dissolved in. Most commonly, people think of solutions as a solid (like salt) dissolved in a liquid (like water). In such a case, salt is the solute, water is the solvent. Solutes are not always solids, and solvents are not always liquids, but that’s a discussion for another day.

When a solute is dissolved in a solvent, the resulting solution has different properties than the solvent does. Moreover, changing the amount of the solute that is dissolved will affect the degree to which certain properties change. In general, a colligative property is a property that depends on the number of solute particles in solution. Note that this definition says “particles” and not “molecules”. Some solutes break into pieces (dissociate) when they dissolve (ionic compounds, like salt), and they change the properties of the solution more rapidly than a solute that remains intact (like sugar) as a result.

There are four commonly cited colligative properties of solutions. We’ll deal with each of these in turn. They are: Boiling Point Elevation, Freezing Point Depression, Vapor Pressure Lowering, and Osmotic Pressure. Each of these is specific to liquid solvents.

Boiling Point Elevation

Boiling Point Elevation (BPE) is exactly what it says. As more solute is added to a solution, the boiling point of that solution increases beyond the original boiling point of the solvent alone. You can calculate this change in boiling point very easily, so long as you have two pieces of information. You need to know the total molality (m) of the solute particles in your solution (molality is a measure of concentration – if you don’t know how to calculate it, you can read another article of mine –, and you need to know the “molal boiling point elevation constant” (Kb) for the solvent. You can look up Kb for your solvent in tables in chemistry texts and other reference materials. If you’re doing chemistry homework, your book does have the values you need. The equation for finding the change in boiling point is an easy one:

Increase in Boiling Point = molality times the molal boiling point elevation constant
dT = m * Kb

If you want to know the new boiling point, then you just add the change in temperature to the solvent’s original boiling point.

A brief list of Kb’s and boiling points can be found at the following site.
Better yet, it also has the values for freezing point too.

Freezing Point Depression

Freezing Point Depression (FPD) is almost identical to boiling point elevation. The only real differences are that we’re dealing with the solvent’s freezing point, the freezing point gets lower as more solute is dissolved, and Kf (the molal freezing point constant) is a different number than Kb was.

Decrease in Freezing Point = molality times the molal freezing point depression constant
dT = m * Kf

Vapor Pressure Lowering

To understand Vapor Pressure Lowering (VPL), it’s important that you know something about solvation. (Not salvation – I’m only trying to bring you science, not religion.) Solvation is the process by which a solute gets dissolved, and it is really neat. To make it short and sweet, when a solute is added to a solvent, the solvent molecules surround the solute particles, effectively forming cages of solvent molecules around each “invading” solute molecule or ion. You might imagine that these “cages” are a bit more bulky than just a single molecule, and you’d be right, especially since these cages can be several layers thick.

Now, going back to vapor pressure – vapor pressure refers to the amount of liquid solvent that “escapes” into the air (or vacuum) above the solvent or solution. At any given temperature and pressure, a solvent has a specific vapor pressure. (This is why a glass of water will slowly evaporate away, even though it is never heated to boiling.) This happens because the solvent molecules are constantly in motion, and when they reach the surface, their motion can carry them off into the air/vacuum if they have sufficient speed. This happens at a constant rate, but adding a solute lowers that rate. You can easily visualize why this is. Imagine that you are in the deep end of a swimming pool, underwater, and ready to come up for air. Then imagine that some thoughtless person has dumped hundreds of beach balls into the pool. All those beach balls will obstruct your path to the surface. They don’t make it impossible, fortunately for you, but they sure do slow you down. The same thing happens with solvent cages – they function as the beach balls, blocking access to the surface, so fewer solvent molecules make their way to the surface with enough speed to escape, and vapor pressure is lowered.

Osmotic Pressure

Osmotic pressure is often the most familiar of the colligative properties. People are already familiar with “osmosis” from biology, so they recognize the word. Indeed the two are closely related. Osmotic pressure is the pressure exerted against a semi-permeable membrane as a result of osmosis. Remember that the solvent will migrate towards areas of higher concentration, so in the end, there will be more solvent on the high concentration side of the barrier than the low concentration side. In a closed container, this means that the liquid level rises of the high concentration side, and gravity works to restore things to normal, creating pressure.

Osmotic pressure is calculated with an equation that is startlingly similar to the ideal gas law (P=nRT/V). We replace the P with a capital Greek letter pi ( ), and n/V is the same as Molarity (M). This gives the normal form of the osmotic pressure equation:


is the osmotic pressure
M is molarity (which is also in my article on finding concentration)
R is the gas constant
T is the temperature (in Kelvin)

There you have it, a quick overview of colligative properties. Perhaps you’ll keep them in mind, and recognize what’s happening when they dump salt on icy winter roads, or your lemonade doesn’t seem to evaporate quite so fast in the sun as your friend’s water. Or perhaps you’ll just do better on a chemistry test one day. Either way, it’ll make me proud.