Freezing Point Depression - Chemistry LibreTexts
solute concentration, the greater the freezing point depression of the solution. The freezing point The equation describing the change in freezing point from pure solvent m = molality = moles of solute per kilogram of solvent. Experiment. Objectives of the Data Analysis. •. Use the mathematical relationship between freezing point depression and solution molality. The freezing points of solutions are all lower than that of the pure solvent. The freezing point depression is directly proportional to the molality of.
So at the surface, we said if I have a bunch of water molecules in the liquid state, we knew that although the average temperature might not be high enough for the water molecules to evaporate, that there's a distribution of kinetic energies. And some of these water molecules on the surface because the surface ones might be going fast enough to escape. And when they escape into vapor, then they create a vapor pressure above here.
And if that vapor pressure is high enough, you can almost view them as linemen blocking the way for more molecules to kind of run behind them as they block all of the other ambient air pressure above them. So if there's enough of them and they have enough energy, they can start to push back or to push outward is the way I think about it, so that more guys can come in behind them. So I hope that lineman analogy doesn't completely lose you.
Now, what happens if you were to introduce solute into it? Some of the solute particle might be down here. It probably doesn't have much of an effect down here, but some of it's going to be bouncing on the surface, so they're going to be taking up some of the surface area.
And because, and this is at least how I think of it, since they're going to be taking up some of the surface area, you're going to have less surface area exposed to the solvent particle or to the solution or the stuff that'll actually vaporize. You're going to have a lower vapor pressure. And remember, your boiling point is when the vapor pressure, when you have enough particles with enough kinetic energy out here to start pushing against the atmospheric pressure, when the vapor pressure is equal to the atmospheric pressure, you start boiling.
But because of these guys, I have a lower vapor pressure. So I'm going to have to add even more kinetic energy, more heat to the system in order to get enough vapor pressure up here to start pushing back the atmospheric pressure.
So solute also raises the boiling point. So the way that you can think about it is solute, when you add something to a solution, it's going to make it want to be in the liquid state more. Whether you lower the temperature, it's going to want to stay in liquid as opposed to ice, and if you raise the temperature, it's going to want to stay in liquid as opposed to gas.
I found this neat-- hopefully, it shows up well on this video. I have to give due credit, this is from chem. This is just the surface of water molecules, and it gives you a sense of just how things vaporize as well. There's some things on the surface that just bounce off. And here's an example where they visualized sodium chloride at the surface.
Freezing Point Depression
And because the sodium chloride is kind of bouncing around on the surface with the water molecules, fewer of those water molecules kind of have the room to escape, so the boiling point gets elevated. Now, the question is by how much does it get elevated? And this is one of the neat things in life is that the answer is actually quite simple.
The change in boiling or freezing point, so the change in temperature of vaporization, is equal to some constant times the number of moles, or at least the mole concentration, the molality, times the molality of the solute that you're putting into your solution. So, for example, let's say I have 1 kilogram of-- so let's say my solvent is water.
And I have 1 kilogram of water, and let's say we're just at atmospheric pressure. And let's say I have some sodium chloride, NaCl. And let's say I have 2 moles of NaCl. I'll have 2 moles. The question is how much will this raise the boiling point of this water? So first of all, you just have to figure out the molality, which is just equal to the number of moles of solute, this 2 moles, divided by the number of kilograms of solvent.
So let's say we have 1 kilogram of solvent. This was, of course, moles. So our molality is 2 moles per kilogram. So we just have to figure out what this constant is, and then we'll know the temperature elevation.
And actually, that same Purdue site, they gave a list of tables. I haven't run the experiments myself. They have some neat charts here. But they say, OK water, normal boiling point is degrees Celsius at standard atmospheric pressure.
And then they say that the constant is 0. So let's just say 0.
How does molality affect the freezing point?
So it equals 0. So k is equal to 0. And I want to be very clear here because this is a very-- I won't say a subtle point, but it's an interesting point. So I said that there's the molality of-- I just realized I made a mistake. I said the molality of sodium chloride is 2. But that would be if sodium chloride stayed in this molecular state, if it stayed together, right?
But what happens is that the sodium chloride actually disassociates, and we learned all about it in that previous video. Each molecule or each sodium chloride pair disassociates into two molecules, into a sodium ion and a chlorine anion. And because of that, because this disassociates into two, the molality is actually going to be two times the number of moles of sodium chloride I have.
So it's going to be two times this. So my molality will actually be 4. And this is an interesting point. If I was dealing with-- and I wrote it here. So this right here is glucose, and this is sodium chloride, or at least sodium chloride in its crystal form. One molecule, I guess you can view it, or one salt of it.
I guess you could just view it as one of these little pairs right here. But the interesting thing is is you could have the same number of moles of sodium chloride when you view it as a compound and glucose. But glucose, when it goes into water, it just stays as one molecule of glucose. In this case, i must be known, and the procedure is primarily useful for organic compounds using a nonpolar solvent. Cryoscopy is no longer as common a measurement method as it once was, but it was included in textbooks at the turn of the 20th century.
Boiling point elevation and freezing point depression
As an example, it was still taught as a useful analytic procedure in Cohen's Practical Organic Chemistry of in which the molar mass of naphthalene is determined using a Beckmann freezing apparatus.
Freezing-point depression can also be used as a purity analysis tool when analysed by differential scanning calorimetry. This is also the same principle acting in the melting-point depression observed when the melting point of an impure solid mixture is measured with a melting-point apparatussince melting and freezing points both refer to the liquid—solid phase transition albeit in different directions. In principle, the boiling-point elevation and the freezing-point depression could be used interchangeably for this purpose.
However, the cryoscopic constant is larger than the ebullioscopic constantand the freezing point is often easier to measure with precision, which means measurements using the freezing-point depression are more precise. Also this phenomenon is applicable in preparing a freezing mixture for use in an ice-cream machine.
Boiling point elevation and freezing point depression (video) | Khan Academy
For this purpose, NaCl or another salt is used to lower the melting point of ice. Last but not the least, FPD measurements are used in the dairy industry to ensure that milk has not had extra water added.
Milk with a FPD of over 0. This typically occurs simply because the solute molecules do not fit well in the crystal, i. In this case, for low solute concentrations, the freezing point depression depends solely on the concentration of solute particles, not on their individual properties.
The freezing point depression thus is called a colligative property. The resulting reduced entropy of the solute particles thus is independent of their properties. This approximation ceases to hold when the concentration becomes large enough for solute-solute interactions to become important. In that case, the freezing point depression depends on particular properties of the solute other than its concentration.
KF, the cryoscopic constantwhich is dependent on the properties of the solvent, not the solute. When conducting experiments, a higher KF value makes it easier to observe larger drops in the freezing point.