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Video transcript
- [Voiceover] A liquid boils when its molecules have enough energy to break free of the attractions that exist between those molecules. And those attractions between the molecules are called the intermolecular forces. Let's compare two molecules, pentane on the left and hexane on the right. These are both hydrocarbons, which means they contain only hydrogen and carbon. Pentane has five carbons, one, two, three, four, five, so five carbons for pentane. And pentane has a boiling point of 36 degrees Celsius. Hexane has six carbons, one, two, three, four, five, and six. So six carbons, and a higher boiling point, of 69 degrees C. Let's draw in another molecule of pentane right here. So there's five carbons. Let's think about the intermolecular forces that exist between those two molecules of pentane. Pentane is a non-polar molecule. And we know the only intermolecular force that exists between two non-polar molecules, that would of course be the London dispersion forces, so London dispersion forces exist between these two molecules of pentane. London dispersion forces are the weakest of our intermolecular forces. They are attractions between molecules that only exist for a short period of time. So I could represent the London dispersion forces like this. So I'm showing the brief, the transient attractive forces between these two molecules of pentane. If I draw in another molecule of hexane, so over here, I'll draw in another one, hexane is a larger hydrocarbon, with more surface area. And more surface area means we have more opportunity for London dispersion forces. So I can show even more attraction between these two molecules of hexane. So the two molecules of hexane attract each other more than the two molecules of pentane. That increased attraction means it takes more energy for those molecules to pull apart from each other. More energy means an increased boiling point. So hexane has a higher boiling point than pentane. So as you increase the number of carbons in your carbon chain, you get an increase in the boiling point of your compound. So this is an example comparing two molecules that have straight chains. Let's compare, let's compare a straight chain to a branched hydrocarbon. So on the left down here, once again we have pentane, all right, with a boiling point of 36 degrees C. Let's write down its molecular formula. We already know there are five carbons. And if we count up our hydrogens, one, two, three, four, five, six, seven, eight, nine, 10, 11 and 12. So there are 12 hydrogens, so H12. C5 H12 is the molecular formula for pentane. What about neopentane on the right? Well, there's one, two, three, four, five carbons, so five carbons, and one, two, three, four, five, six, seven, eight, nine, 10, 11 and 12 hydrogens. So C5 H12. So these two compounds have the same molecular formula. All right? So the same molecular formula, C5 H12. The difference is, neopentane has some branching, right? So neopentane has branching, whereas pentane doesn't. It's a straight chain. All right. Let's think about the boiling points. Pentane's boiling point is 36 degrees C. Neopentane's drops down to 10 degrees C. Now, let's try to figure out why. If I draw in another molecule of pentane, all right, we just talk about the fact that London dispersion forces exist between these two molecules of pentane. So let me draw in those transient attractive forces between those two molecules. Neopentane is also a hydrocarbon. It's non-polar. So if I draw in another molecule of neopentane, all right, and I think about the attractive forces between these two molecules of neopentane, it must once again be London dispersion forces. Because of this branching, the shape of neopentane in three dimensions resembles a sphere. So it's just an approximation, but if you could imagine this molecule of neopentane on the left as being a sphere, so spherical, and just try to imagine this molecule of neopentane on the right as being roughly spherical. And if you think about the surface area, all right, for an attraction between these two molecules, it's a much smaller surface area than for the two molecules of pentane, right? We can kind of stack these two molecules of pentane on top of each other and get increased surface area and increased attractive forces. But these two neopentane molecules, because of their shape, because of this branching, right, we don't get as much surface area. And that means that there's decreased attractive forces between molecules of neopentane. And because there's decreased attractive forces, right, that lowers the boiling point. So the boiling point is down to 10 degrees C. All right. I always think of room temperature as being pretty close to 25 degrees C. So most of the time, you see it listed as being between 20 and 25. But if room temperature is pretty close to 25 degrees C, think about the state of matter of neopentane. Right? We are already higher than the boiling point of neopentane. So at room temperature and room pressure, neopentane is a gas, right? The molecules have enough energy already to break free of each other. And so neopentane is a gas at room temperature and pressure. Whereas, if you look at pentane, pentane has a boiling point of 36 degrees C, which is higher than room temperature. So we haven't reached the boiling point of pentane, which means at room temperature and pressure, pentane is still a liquid. So pentane is a liquid. And let's think about the trend for branching here. So we have the same number of carbons, right? Same number of carbons, same number of hydrogens, but we have different boiling points. Neopentane has more branching and a decreased boiling point. So we can say for our trend here, as you increase the branching, right? So not talk about number of carbons here. We're just talking about branching. As you increase the branching, you decrease the boiling points because you decrease the surface area for the attractive forces. Let's compare three more molecules here, to finish this off. Let's look at these three molecules. Let's see if we can explain these different boiling points. So once again, we've talked about hexane already, with a boiling point of 69 degrees C. If we draw in another molecule of hexane, our only intermolecular force, our only internal molecular force is, of course, the London dispersion forces. So I'll just write "London" here. So London dispersion forces, which exist between these two non-polar hexane molecules. Next, let's look at 3-hexanone, right? Hexane has six carbons, and so does 3-hexanone. One, two, three, four, five and six. So don't worry about the names of these molecules at this point if you're just getting started with organic chemistry. Just try to think about what intermolecular forces are present in this video. So 3-hexanone also has six carbons. And let me draw another molecule of 3-hexanone. So there's our other molecule. Let's think about electronegativity, and we'll compare this oxygen to this carbon right here. Oxygen is more electronegative than carbon, so oxygen withdraws some electron density and oxygen becomes partially negative. This carbon here, this carbon would therefore become partially positive. And so this is a dipole, right? So we have a dipole for this molecule, and we have the same dipole for this molecule of 3-hexanone down here. Partially negative oxygen, partially positive carbon. And since opposites attract, the partially negative oxygen is attracted to the partially positive carbon on the other molecule of 3-hexanone. And so, what intermolecular force is that? We have dipoles interacting with dipoles. So this would be a dipole-dipole interaction. So let me write that down here. So we're talk about a dipole-dipole interaction. Obviously, London dispersion forces would also be present, right? So if we think about this area over here, you could think about London dispersion forces. But dipole-dipole is a stronger intermolecular force compared to London dispersion forces. And therefore, the two molecules here of 3-hexanone are attracted to each other more than the two molecules of hexane. And so therefore, it would take more energy for these molecules to pull apart from each other. And that's why you see the higher temperature for the boiling point. 3-hexanone has a much higher boiling point than hexane. And that's because dipole-dipole interactions, right, are a stronger intermolecular force compared to London dispersion forces. And finally, we have 3-hexanol over here on the right, which also has six carbons. One, two, three, four, five, six. So we're still dealing with six carbons. If I draw in another molecule of 3-hexanol, let me do that up here. So we sketch in the six carbons, and then have our oxygen here, and then the hydrogen, like that. We know that there's opportunity for hydrogen bonding. Oxygen is more electronegative than hydrogen, so the oxygen is partially negative and the hydrogen is partially positive. The same setup over here on this other molecule of 3-hexanol. So partially negative oxygen, partially positive hydrogen. And so hydrogen bonding is possible. Let me draw that in. So we have a hydrogen bond right here. So there's opportunities for hydrogen bonding between two molecules of 3-hexanol. So let me use, let me use deep blue for that. So now we're talking about hydrogen bonding. And we know that hydrogen bonding, we know the hydrogen bonding is really just a stronger dipole-dipole interaction. So hydrogen bonding is our strongest intermolecular force. And so we have an increased attractive force holding these two molecules of 3-hexanol together. And so therefore, it takes even more energy for these molecules to pull apart from each other. And that's reflected in the higher boiling point for 3-hexanol, right? 3-hexanol has a higher boiling point than 3-hexanone and also more than hexane. So when you're trying to figure out boiling points, think about the intermolecular forces that are present between two molecules. And that will allow you to figure out which compound has the higher boiling point.