 I figured while I had the time, I would throw up a quick video on these intermolecular forces since they do represent about 12% of what your quarter exam grade is going to be. And we need to get all the points we can get on this thing. So let's talk a little bit about it. Intermolecular forces are the forces that hold solids and liquids together. If we were to look inside a solid, we would see nice, neatly arranged particles that are close together. They have a shape of their own. They don't assume the shape of their container. All stuff that I covered back when we did the kinetic molecular theory during the first semester. In liquids, the particles are further apart. One more space between them. And they're kind of random. We don't have the rows and columns anymore. In a gas, we have to put a container with a lid on top so they don't get out. They're the furthest apart. Again, all random scattered throughout the container at this time, not just down in the liquid. Intermolecular forces help us understand the behaviors here and what's going on in these substances. First, I'll start by saying that both solids and liquids have intermolecular forces. And I'll abbreviate intermolecular force IMF just to make it quicker and simpler. Solids and liquids both have intermolecular forces, whereas gases do not. Of the two, solids have the stronger ones, while liquids are weaker in terms of intermolecular forces. It works out a lot like magnets do. If you take two magnets and put them close together, you can fill a strong force of attraction or repulsion. But then as you move those magnets further and further apart, you get less and less force. The distance makes the force weaker. In solids, the particles are very close together, so the intermolecular forces are stronger. In liquids, they are further apart, thus the intermolecular forces are weaker. So again, solids and liquids have intermolecular forces, gases do not. Between the solids and liquids, solids generally have stronger intermolecular forces, while liquids have weaker ones. We talked a little bit in class today about the heating graph and how that relates to intermolecular forces. We have temperature on our y-axis, usually measured in degrees Celsius, and we either have time or heat on the x-axis, increasing as we go to the right. Here's what happens. We take some solid, like ice, and we put it in a closed container and we heat it up over a burner on a hot plate. We measure its temperature every 10 seconds or so. What we would see is the solid would heat up in a nice, linear fashion, until that solid reached the melting point. It's water at zero degrees Celsius, and then the temperature would remain constant while all the ice melted. Then once the ice was gone, we had liquid water in our container, we'd see the temperature rise in a nice, linear fashion again. Until for water, we reach 100 degrees Celsius, the boiling point, in which case the temperature would be constant again. Until all that water had boiled, and we had nothing but a gas, and then we could start heating the gas up. This is all solid down here. This is all liquid through here. This is all gas through there. Heating up a solid, heating up a liquid, heating up a gas. Again, this first plateau is where melting is taking place. The second plateau is where boiling happens. And again, while it melts, while it boils, the temperature remains constant. This is why, as long as there are ice cubes in your drink, the temperature remains the same, even if you keep your drink in a hot car. While that ice melts, the temperature will remain constant. This is why in a survival situation, you could boil water in a plastic bottle. If you hang a plastic bottle full of water over a campfire, you can boil the water in it without melting the plastic bottle. And that's because the water inside can't get any hotter than 100 degrees Celsius, the boiling point. Let's talk about why these regions are flat. In this part, this first part where it's melting, the heat of fusion is being absorbed. And what the heat of fusion does is it gets us from here to here. It is weakening these intermolecular forces. Instead of being used to raise the average kinetic energy of the substance, it's being used to weaken the intermolecular forces. Up here, we have something called heat of vaporization being absorbed. And what that heat of vaporization has to do is get us from here to here, from where we have these weakened intermolecular forces to none at all. So we can say this heat of vaporization is removing them. Generally speaking, that top plateau line is longer than the bottom one. It takes less energy to weaken those intermolecular forces than it does to remove those intermolecular forces. But again, while those two things are happening, while we are either weakening those intermolecular forces or removing them, the temperature doesn't change. All the energy is going into changing the intermolecular forces in our substance. Last thing I talked about today was understanding the three types of intermolecular forces. And in covalent substances, we have hydrogen bonding, we have dipole-dipole attraction, and we have dispersion forces. These forces are listed in order from the strongest to the weakest. So hydrogen bonding is the strongest of the intermolecular forces in covalent substances simply because it has the most polar bonds in it, the most polar molecules. That creates a relatively strong force of attraction between the positive end of one molecule and the negative end of the other. Dispersion forces have the weakest polarities, temporary polarities in either non-polar molecules or individual atoms of elements produced under extreme conditions. Very high pressure, very low temperatures. Because of that, these are the weakest. Because these are the strongest intermolecular forces of hydrogen bonding, we see certain properties that we can associate with hydrogen bonding, like high melting points and high boiling points. Again, due to the strong force of attraction between the particles and the substance, if the particles are being held together by strong forces, then it's going to take a lot of heat to either weaken those forces or remove those forces. Down here, these are weak, so these have the lowest melting points and boiling points. Now, something related to boiling point is vapor pressure. And we say substances with hydrogen bonding have low vapor pressure. It seems kind of contradictory when you think about what vapor pressure is, but hear me out. Vapor pressure deals with boiling. Substances boil when the vapor pressure is equal to the atmospheric pressure on top of the liquid. So this is atmospheric pressure. Substance with hydrogen bonding have a low vapor pressure, so the vapor pressure is way down here. We have to push this all the way up here to get the substance to boil, and that takes a lot of heat. It takes a lot of heat to push that vapor pressure up to where it's the same as the atmospheric pressure in the substance and start to boil. Over here on this side of the spectrum, these substances have high vapor pressure, and they very easily move into the gaseous state, so there's generally a lot of particles above the liquid, so they already have a high vapor pressure. So our atmospheric pressure is in the same spot. Our atmospheric pressure is up here, but now our vapor pressure is right here. It's already high. It doesn't have far to go to hit that atmospheric pressure and make the thing boil, so it doesn't take a whole lot of heat to get it there. So it seems kind of contradictory at first. We have a high boiling point, but a low vapor pressure? Yeah, we got a long way to go because the vapor pressure is low. That's why the boiling point is so high. Over here again, we have a low boiling point, but that's because we have a high vapor pressure. We have a lot of particles that are already in the gaseous state, which means we don't have to bump it up that much further to get it into the boiling area. Hydrogen bonding tends to have strong surface tension. We have strong attractive forces, so the particles at the surface hold together very well. They make kind of a network of attraction across the surface, which gives it a strong surface tension. This is why if you jump into the water from too high an elevation, it hurts. That surface tension is hard to break because it's relatively strong, so it feels like you're getting slapped when you hit the water. Down here, this is weak. These tend to be liquids. Well, these tend to be gases. Dipole dipole is just everywhere in between. It's that middle sibling that's kind of ignored because it's kind of the middle of the road. This is high. That's low. This is medium. That's high. That's low. This is medium. This is low. This is high. This is medium. Everything about it's kind of just in the middle there. Last thing I'll do is talk a little bit about how you decide what kind of intermolecular force you have. And to do that, I want to go to the practice sheet that I gave you on these. And again, this one here, I meant to type silicon dioxide, SiO2, not SO2. Now here's how it works. All of these are binary. All of those are two element compounds, which is the way it should be in CP chemistry. Which means there are only three combinations that will produce hydrogen bonds. It has to be a hydrogen-oxygen bond, so that's going to be water. It can be a hydrogen-nitrogen bond, which means it can be ammonia. Or it can be a hydrogen-fluorine bond, which we don't actually have on our list here. HF would be the other simple binary compound with hydrogen bonding in it. Water, H2O, ammonia, NH3. And finally, hydrogen fluoride. Those are the simple, binary, easy to draw Lewis structures, easy to use Vespert molecules that would be hydrogen bonding. There are others. We could talk hydrogen peroxide in here. That's a possible hydrogen bond, one H2O2. We could talk ammonium ions, NH4. That would have hydrogen bonding potentially. But this, again, just not much else that's going to do it. Look for water, look for ammonia, look for hydrogen fluoride for hydrogen bonders. For dispersion forces, the weakest of all of them, you can look for elements. If you have any elements in there like neon, any or helium, HE, those would all have to be dispersion forces. And there's none of those in there either. So everything else in here, we have to determine the polarity of the molecule to figure out if it's dipole-dipole attraction, or if it's the dispersion forces. Because again, hydrogen bonding, these are hydrogen oxygen, hydrogen nitrogen, or hydrogen fluorine bonds in polar molecules. So again, that's going to be water, ammonia, and hydrogen fluoride, if we keep it simple. And I really hope that the people who wrote the quarter test keep it simple. They start throwing stuff on there that we didn't do structures for, then you might as well just guess anyways because you're not going to be able to do the structures. But that's the simple ones right there. For dipole-dipole, it's just got to be any other polar molecule. Again, any other polar molecule besides those. And then for dispersion forces, it's going to be non-polar or element. So let's get into it. This is H2S. H2S looks like that. That's a bent molecule like water. Remember oxygen, sulfur, or selenium, if those are your metal atoms, middle atoms, not metal atoms. If those are your metal atoms, then it's going to be a bent molecule because of those unpaired electrons. If you need to review that, go watch the video I have on shapes and vesper. The shortcut video in particular where I talk about those metal atoms and what they mean. So sulfur's electronegativity is 2.5 and hydrogens is 2.1 for an electronegativity difference of 2, of 0.4. We have a 0.4 electronegativity difference here, which means the bonds are not polar. If it's not polar, then this is going to be dispersion forces. Moving on to SiO2. SiO2 is like carbon dioxide. And again, silicon and carbon make identical molecules because they both have four bonding sites or four lone electrons in their Lewis structures. Again, watch my shortcut video on covalent bonding if you still have in trouble with that, where I go through what it means to switch out atoms, what it means when they have the same Lewis structure and all that. Well, this is a 3-atom linear molecule, and it doesn't matter what the electronegativity difference is. It does come out to be polar. Silicone's 1.8 and oxygen's 3.5, so we're definitely in the 0.5 to 2.0 range. This would be positive and these would be negative, but there's no separation of charge there. So it is non-polar, which means it's dispersion force again because if it's non-polar or an element, it's got to be dispersion force. And again, silicon dioxide and carbon dioxide are the same. We have polar bonds here as well. Electronegativity difference of 0.5, so it's at the very bottom end of our range, but even though we have polar bonds, we don't have separation of charge, so it is non-polar. So it'd be dispersion force again. CH4, I don't even have to work anything out on that. That's a 5-atom molecule. And a 5-atom molecule is a tetrahedron, and when they're binary, they're never polar. So that would be dispersion force. PCL3, that's pure middle because phosphorus has a pair of electrons there. Again, watch my shortcut video if you haven't watched it. Phosphorus is 2.1 and chlorine is 3.0, so those are polar bonds. Phosphorus is the lower one, so it's positive. Chlorine is the higher one. We got separation of charge. That's polarity, dipole, dipole. And again, I'm just rushing to the polarity on here. There is a whole separate video about polarity. If you still don't know how to do polarity or shapes or anything like that, then make sure you watch those videos. Hydrogen iodide, HI, 2.1 and 2.5, that is a difference of 0.4. That is not polar, so those are dispersion forces again. And that's how that works out. Now down here where it asks, which of the above subs would you expect to have a high melting point? Those are the ones with the strongest bonds. That's the hydrogen bonding, so that's the water and the ammonia. Which of the above subs would you expect to be low? Those are the weak ones, the dispersion force ones. The H2S, the SiO2, the CO2, there were a lot of them. The CH4 and the hydrogen iodide, HI. Those would all be dispersion forces, so those would be the lowest of them. I think that's enough. It's already coming up on a 20-minute video. Good luck on that quarter exam.