- [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.