 So let's talk about a different feature of solutions. Solutions that are actually more common than the ones I've been talking about previously. So far I've been careful when I describe a solution to always describe a solution with a molecular solute. If we talk about a sugar water solution, sucrose dissolves in water and I've got as many molecules of sucrose in the solution as I had in the solid that I had to begin with. But for electrolytic solvents, solvents that dissociate into ions when you dissolve them. For example, sodium chloride is an electrolytic solvent, meaning when I dissolve it in water, what I end up with is aqueous sodium ions and aqueous chloride ions. And of course water in the solution as well. Point being that my sodium chloride as a solid has become two different species. Sodium ions and chloride ions in solution. The fact that it dissociates into ions in solution makes it an electrolytic solute. There's two important things to recognize about electrolytic solutes that are very different from non-electrolytes. Those are number one, what we've just been talking about, the fact that the solute dissociates in solution. That changes the concentrations. For example, if I make a .1 molal solution of sodium chloride, if I dissolve .1 moles of sodium chloride in a kilogram of water, I'll end up with a solution that's .1 molal in sodium ions and .1 molal in chloride ions. Altogether that concentration will be .2 molal. So if we think back to the colligative properties and ask ourselves how much the freezing point will depress or how much the boiling point will be elevated in the solution, that may only be proportional to the amount of solute I have, but I have twice as much solute as I thought I had if I was not considering that the fact that the solute dissociates. So certainly important to consider the fact that the solute dissociates to calculate the concentration of these solutions. It's also very important to recognize that in electrolytic solutions, the solutions are almost never ideal. We can almost never get away with treating the solutions as an ideal solution, so we can't actually use many of the expressions that we've developed for colligative properties or for vapor pressure and so on. So to give you an idea of what I mean by non-ideal and how non-ideal these solutions may or may not be, let's consider a few examples that I can give you some data for. So for example, we'll start with a molecular solvent, a non-electrolyte. So again, sucrose, when sugar dissolves in water, we just have sucrose molecules in solution. They don't dissociate into ions. If I prepare a .1 molal solution of sucrose in water, if I, we won't go through the calculations, but if we calculate the mole fraction of sucrose in that solution, sucrose is quite a bit heavier than water, the mole fraction of sucrose in that solution will be much less than .1 moles of sucrose is a pretty large mass of sucrose. So the molar, I'm sorry, the mole fraction works out to be .00180. So about two-tenths of a percent mole-wise, molecule-wise, is sucrose in that solution. So that's the mole fraction. If I were to measure the activity of the sucrose in that solution, turns out the activity is pretty close to that value. So this is not that unideal a solution. The activity of sucrose in that solution is .00179. And that's not something we can compute or predict at least at this point. That's just something that would have to be measured by, for example, measuring a colligative property and seeing, or some property of the sucrose and seeing what activity that gives us. So if we measure the activity of sucrose, that turns out to be pretty close to the mole fraction, which means that the activity coefficient, because the activity is a little less than the mole fraction, the activity coefficient is a little less than 1. So this is a fairly ideal solution. Sucrose is fairly ideal. It has an activity coefficient pretty close to 1. So far so good. But now let's see what happens when I prepare a solution that's let's say .1 molal in an ionic or an electrolytic solute, sodium chloride, for example. If I prepare a solution that's .1 molal and sodium chloride, and I'll skip the intermediate steps, I'll just tell you what the activity coefficient is for this .1 molal sodium chloride solution, so we can judge how ideal it is. The activity coefficient in a .1 molal sodium chloride solution, .778. So if 1 is 100% ideal, fully ideal, this is an activity that's 22 or 23% lower than we'd expect based on an ideal solution. So it's deviating quite a bit from ideality. Even worse, if I go to a solution that's .1 molal in calcium chloride, as we'll see shortly, the fact that this is a plus 2 ion rather than a plus 1 ion makes it much less ideal. The activity coefficient for this solution, or at least for the calcium chloride in the solution, is down to .518. So the activity is barely half what we would expect it to be if we considered the solution to be ideal. So you can see that for these electrolytic solutions, the activity coefficient is very far away from 1. That's very typical, that's very common. So the solutions are very non-ideal. So we're going to have to consider almost always the activity of the solutions when we deal with electrolytes. That raises a couple of questions, which we'll address next. And those questions are, number one, why would this be the case? Why is it the case that electrolytic solutions are so much more non-ideal than non- electrolytic solutions? And number two, if you think about it carefully, what do I even mean when I tell you the activity coefficient of the solution is .778. Remember, I've told you to remember that the electrolytic solutes dissolve, dissociate into two different ions. So we could talk about the activity or the activity coefficient of sodium ions, the activity coefficient of chloride ions. I've just given you one activity coefficient number for the entire solution. This is essentially the activity coefficient of the NACL solute. But what justification do I have for talking about NACL as a single solute when we know it's going to dissociate into two species? So those are two mysteries I'll leave you with right now and we'll explore those a little bit later coming up next.