- [Instructor] So I have
these two molecules here,
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.
