In our previous module so far, we've been pretty well exclusively concerned with gasses. We've been measuring the properties of gasses and we've been using those to help us understand how molecules move inside of a gas, what we call the Kinetic Molecular theory. We'd now like to move past gasses and start talking about the properties of liquids. And the Kinetic molecular model to describe liquids to do that we need to begin looking at some experimental data that may tell us something about how liquids behave. And that experimental data is actually going to begin by going back and taking a look at Charles' law. You may recall that this is the apparatus that we use to study Charles' law and I actually have, such an apparatus basically here in front of me. it's just a standard syringe in which I can easily vary or, or observe variations in the volume of a trapped gas. This gas trapped in the cylinder, about 25 cc's of gas trapped in this we see, region here. And that's just like the drawing that I have drawn over, on the, the gauge here. Except that we're now going to add the capability of first, applying a constant pressure to the end of the syringe. And, as well, we will be able to keep the temperature in a constant and vary that temperature. And measure the volume, of the trapped gas, as that temperature varies. That was our Charles' law experiment before we remember over here, sort of schematically again I'm not actually lowering the temperature, as I lower the temperature, we think that the volume of the gas is going to drop. Would take us all the way through the origin. That is at zero degrees Kelvin we would predict that the volume of the gas would be zero. And so the volume of the gas proportional to the temperature of the gas in degrees Kelvin. If that is the case then we might be tempted. To attempt to collect some data down in this region here, where we continued to chill the gas. If we were to do that and again, I'm obviously just going to do this as a demonstration over here. What we would discover, is eventually we would reach a temperature where rather than see a proportional drop in the volume, we would see a very dramatic drop in the volume, all occurring at. If we were to open the syringe in the vicinity of these temperatures here, we would discover that the syringe has only liquid in it and no gas. What we could do then is observe something called a Phase transition. At the particular temperature that we were describing there, we are undergoing a change from one phase, gas To another phase, liquid. Furthermore, that phase transition always occurs for a particular liquid, at a particular pressure, at the same temperature every time. If we were, for example, instead of lowering the temperature and then observing the precipitous drop. If we were instead to start with liquid in the syringe with very low volume, what we would discover is that we maintain liquid until we get to that same temperature. And then suddenly there's a dramatic rise in the volume corresponding to the phase transition. For a fixed pressure, that phase transition always occurs at the same temperature. For any particular liquid, or for the gas corresponding to that particular liquid. We refer to that temperature, very colloquially but also technically, as the boiling point. The boiling point at each pressure is the temperature at which the liquid will turn into gas. Or, if we have the gas and we drop the temperature when it reaches the boiling point, it will fully condense back into the liquid phase. So, what does the boiling point temperature depend upon? Well one thing it depends upon is the applied pressure. If, in fact, I applied a different pressure to the end of the syringe Then what I wil discover is that the boiling point temperature is differnent for each applied pressure. So one of the things that it depends upon is the pressure. Lets look at some data, lets imagine that what I do is measure that boiling point temperature using the experiment we just described. And I'll vary the pressure that I apply to the end of the syringe down here, keeping that pressure constant for each measurement. So I'll measure a particular pressure say here and then measure the temperature at which that transition takes place here. Along these lines here. That what we've actually observed, is a decrease in the boiling point with the applied pressure. Notice, as well, that if we change to a different substance, let's try now, ethanol. That the boiling point of ethanol is higher than the boiling point of the butane. It decreases with decreasing pressure in just the same way that it does for the butane, but at every pressure, the boiling point of the buute of the ethanol is higher than the boiling point of the butane. Therefore, the boiling point clearly depends upon what the substance is, but is also clearly depends upon what the applied pressure is. This gives us some means, then, to begin to understand what's going on with the liquid. To do this we're actually going to modify our experiment a bit. Here's our same syringe again but in this particular case we're going to attach a temperature gauge to it and we're going to attach a pressure gauge to it. And one of the differences you'll notice is that in this case we have trapped a certain amount of liquid here. And we have trapped a gas above that liquid. Here's how we might make that happen. I've got a different syringe for us here, in this particular case the syringe currently contains only liquid. If I were to do this very carefully, there would be no bubbles in here at all. And now I'm going to vary, the volume above that liquid. Right now, there is no volume above that liquid. But what happens if I now pull back? It's quite difficult to do. I'm going to pull back the syringe. I actually trapped some gas above that liquid. It's hard to see in this form. So I'm going to turn my camera here, so that you can see it. So that I can hold it vertically. Here is in fact the same syringe held vertically. There's no gas above the liquid in this region. Now I'm going to pull the syringe back and you'll notice that we trap a volume above the liquid. What I want to do is measure the pressure that is in that liquid in that gas. As a function of the volume of the gas as I move this up and down. Notice, I can change that up and down and so correspondingly here what we are going to be measuring back over on this diagram here, is the pressure of the gas. But it doesn't happen when the liquid is present. The pressure, which is above the liquid, does not depend upon the volume of the gas which is above the liquid. That pressure is constant and we refer to that as the vapor pressure of the liquid. So it is the gas above the liquid, the vapor. The pressure of that vapor above the liquid and it's a constant. It doesn't vary when we change the volume. Now that's going to be mystery and we're going to learn a lot from that mystery. It's going to help us understand what's going on inside the liquid. I would observe a very different vapor pressure so it depends upon the type of the liquid. What else does the vapor pressure depend upon? Well, I did this experiment at room temperature. If I varied the temperature, I discovered that I also change the the vapor pressure of the liquid. And in fact I could change the temperature here, measure the pressure, change the temperature again, measure the pressure. And if I did that, I could construct a graph of the vapor pressure of the liquid in of the gas above the liquid as a function of the temperature of the liquid. What I discover is the graph looks something like this. The vapor pressure always increases as I increase the temperature. And it doesn't increase in a straight line, as we might have guessed from the ideal gas law. It actually varies in a very non linear kind of way here. Almost looks kind of exponential. Let's look at some real data. And notice that, at high temperatures, the vapor pressure of [INAUDIBLE] can be quite low. But even at low temperatures, the vapor pressure of, of the dimethyl either can be quite high. In fact, it's interesting to compare. Let's fix a particular vapor pressure and ask. What temperature would it take to achieve that particular vapor pressure of the liquid? And the answer is it depends a lot upon what the liquid is. That, that particular vapor pressure that I picked here. Let's see, to achieve that in the phenol I would need something that look likes, oh what have we got here? Somewhere around 170, degrees centigrade. By contrast achieve that for the dimethyl ether, is about minus 30 degrees centigrade. 200 degree difference between the two for different kinds of liquids. Interestingly though, remember What does not affect the vapor pressure, is the volume of the gas trapped. That when I vary the volume, above the liquid, I do not see a change in the vapor pressure. We're going to use this information an the variation to the vapor pressure with the temperature, to understand what's going on in the liquid and we're going to pick that up in the next lecture.
