propane on the left and acetaldehyde here on the right.

And we've already calculated their molar masses for you,

and you see that they have very close molar masses.

And so based on what you see in front of you,

which of these, you think,

would have a higher boiling point,

a sample of pure propane

or a sample of pure acetaldehyde?

Pause this video, and think about that.

All right, well, in previous videos,

when we talked about boiling points

and why they might be different,

we talked about intermolecular forces.

Because you could imagine, if you have a bunch of molecules,

let's say, in a liquid state,

the boiling point is going to be dependent

on how much energy you need to put into the system

in order for the intermolecular forces between the molecules

to be overcome so that molecules could break free

and enter into a gaseous state.

And so when we're thinking about

which might have a higher boiling point,

we really just need to think about

which one would have higher intermolecular forces.

Now, in a previous video,

we talked about London dispersion forces,

which you can view as random dipoles forming

in one molecule, and then that can induce dipoles

in a neighboring molecule.

And then the positive end, even temporarily positive end,

of one could be attracted to the temporarily

negative end of another and vice versa,

and that whole phenomenon can domino.

And we said that you're going to have more

of those London dispersion forces

the more polarizable your molecule is,

which is related to how large of an electron cloud it has,

which is related to its molar mass.

And when we look at these two molecules,

they have near identical molar masses.

So you might expect them

to have near identical boiling points,

but it turns out that that is not the case.

The boiling point of propane

is negative 42.1 degrees Celsius,

while the boiling point of acetaldehyde

is 20.1 degrees Celsius.

So what makes the difference?

Why does acetaldehyde have such a higher boiling point?

Why does it take more energy

for the molecules in liquid acetaldehyde

to be able to break free of each other

to overcome their intermolecular forces?

Well, the answer, you might imagine, is other things are

at play on top of the London dispersion forces.

And what we're going to talk about in this video is

dipole-dipole forces.

So you might already imagine where this is going.

In the video on London dispersion forces,

we talked about a temporary dipole

inducing a dipole in a neighboring molecule

and then them being attracted to each other.

Now we're going to talk about permanent dipoles.

So when you look at both of these molecules,

which one would you think has a stronger permanent dipole?

Or another way of thinking about it is

which one has a larger dipole moment?

Remember, molecular dipole moments are just the vector sum

of all of the dipole moments of the individual bonds,

and the dipole moments are all proportional

to the differences in electronegativity.

When we look at propane here on the left,

carbon is a little bit more electronegative than hydrogen

but not a lot more electronegative.

So you will have these dipole moments on each of the bonds

that might look something like this.

So you would have these things that look like that.

If that is looking unfamiliar to you,

I encourage you to review the videos on dipole moments.

But as you can see, there's a symmetry to propane as well.

So if you were to take all of these arrows that I'm drawing,

if you were to take all of these arrows

that I'm drawing and net them together,

you're not going to get much of a molecular dipole moment.

You will get a little bit of one,

but they, for the most part, cancel out.

Now what about acetaldehyde?

Well, acetaldehyde, there's a few giveaways here.

One is it's an asymmetric molecule.

So asymmetric molecules are good suspects

for having a higher dipole moment.

Another good indicator is you have some character here

that's quite electronegative.

In this case, oxygen is quite electronegative.

And even more important,

it's a good bit more electronegative than carbon.

So right over here, this carbon-oxygen double bond,

you're going to have a pretty significant dipole moment

just on this double bond.

It might look like that.

And all of the other dipole moments

for all of the other bonds aren't going

to cancel this large one out.

In fact, they might add to it a little bit

because of the molecule's asymmetry.

And so net-net, your whole molecule

is going to have a pretty significant dipole moment.

It'll look something like this,

and I'm just going to approximate it.

But we're going to point towards the more negative end,

so it might look something like this,

pointing towards the more negative end.

And I'll put this little cross here

at the more positive end.

And so you would expect a partial negative charge

at that end and a partial positive charge at this end.

And so what's going to happen

if it's next to another acetaldehyde?

Well, the partially negative end of one acetaldehyde

is going to be attracted to the partially positive end

of another acetaldehyde.

And so this is what people are talking about

when they say dipole-dipole forces.

We are talking about a permanent dipole

being attracted to another permanent dipole.

And so acetaldehyde is experiencing that

on top of the London dispersion forces,

which is why it has a higher boiling point.

Now some of you might be wondering,

hey, can a permanent dipole

induce a dipole in a neighboring molecule

and then those get attracted to each other?

And the simple answer is yes, it makes a lot of sense.

You can absolutely have a dipole

and then induced

dipole interaction.

And we might cover that in a few examples in the future,

but this can also occur.

You can have a temporary dipole

inducing a dipole in the neighbor,

and then they get attracted to each other.

And you could have a bit of a domino effect.

You can have a permanent dipole interacting

with another permanent dipole.

They get attracted to each other.

And you could have a permanent dipole inducing a dipole

in a neighboring molecule.