- [Instructor] So let's talk a little bit
about molecular solids.
So just as a little bit of review,
we've talked about ionic solids
where ions form these lattices.
So those might be the positive ions right over there,
and then you have your negative ions.
And the negative is attracted to the positive.
The positive is attracted to the negative.
And I'm just showing a two-dimensional version of it,
but it forms a three-dimensional lattice.
So that's an ionic solid.
We have also seen metallic solid
where you have metals that all contribute
some valence electrons to the sea of electrons.
So what you end up having
is essentially these positive cations
that are in this sea of electrons.
And we've talked about those properties,
very good at conducting electricity, malleable, et cetera.
Now, what we're gonna do is talk about what happens
when you have nonmetals.
So the nonmetals, you can see in yellow right over here,
also includes hydrogen.
Now, of course, noble gases are also nonmetals,
but they're not reactives.
So we're gonna talk about the reactive nonmetals.
They can form molecules with each other.
For example, one iodine can bond
to another iodine with covalent bonds.
So you could have a molecule like I2.
You have things like carbon dioxide.
Each carbon can bond to two oxygens.
These are each molecules formed
due to covalent bonds between nonmetals.
Now, when we talk about molecular solids,
we're talking about putting a bunch of these together.
So let's say putting a bunch of iodine molecules together,
and the intermolecular forces
at a sufficiently low temperature
are sufficient to hold together those molecules as a solid.
So what do I mean by that?
Let's look at a few examples.
This right over here is a picture of solid iodine,
and the way it's made up is you have these iodine molecules.
Now, each of these iodine molecules are formed
by a covalent bond between two iodine atoms.
Now, the reason why it's a solid
is there's enough dispersion forces.
We talked about these London dispersion forces
that are formed by temporary dipoles
inducing dipoles in neighboring molecules.
For example, just by random chance,
for a moment, you might have more electrons
on this end of this iodine molecule,
creating a partially negative charge.
And then that means some of the electrons
on this end of this neighboring iodine molecule
might be repulsed by that negative charge,
so it forms a partially positive charge.
And so you have a temporary dipole
inducing a dipole in the neighboring molecule,
and then they'll be attracted to each other,
and we've talked about that as London dispersion forces.
And at a sufficiently low temperature,
that can keep them altogether as a solid.
Now, it's important to point out,
I keep saying sufficiently low temperature
because these molecular solids,
because they are only held together
not by the covalent bonds,
the covalent bonds hold together each of the molecules,
but the molecules are held together
by these fairly weak dispersion forces.
They tend to have relatively low melting points.
For example, solid iodine right over here
has a melting point,
has a melting point
of 113.7
degrees Celsius.
And I know what you're saying.
That's not that low.
That's higher than the temperature at which water boils.
It would be quite uncomfortable for any of us
to be experiencing 113.7 degrees Celsius.
But this is relatively low when you talk about solids.
Think about the temperatures
it requires to melt, say, table salt.
We've talked about that.
Think about the temperatures it takes to melt iron.
There, you're talking about hundreds of degrees,
in certain solids, thousands of degrees Celsius.
This is much lower.
And so as a general principle,
molecular solids tend to have relatively low melting points.
Now, how good you think they're gonna be
as conductors of electricity?
Pause the video and think about that.
Well, in order to be conductors of electricity,
somehow charge needs to move through the solid.
And unlike metallic solids,
you don't have the sea of electrons
that can just move around,
so these tend to be bad conductors of electricity.
If you wanna see another example of a molecular solid,
this right over here is solid carbon dioxide,
often known as dry ice.
What you see here is each of these molecules, each carbon,
is bonded to two oxygens.
It has a double-bond with each of those oxygens.
These are covalent bonds
that form each of these molecules.
But what keeps all of the molecules attracted to each other
is, once again, those dispersion forces.
And these forces between the molecules are so weak
that solid carbon dioxide doesn't even really melt.
It doesn't even go to a liquid state.
If you heat it up enough
to overcome these intermolecular forces,
these dispersion forces, it will sublime,
which means it goes directly from a solid to a gas state,
and it does that at a very low temperature.
It sublimes
at negative 78.5
degrees Celsius.
And if you've ever handled a dry ice,
which I don't recommend you doing without gloves
because it will hurt your skin if you do touch it,
I actually did that recently at my son's birthday party,
we were playing around with dry ice,
you don't mess around with this thing
because it is so incredibly cold.
And at that temperature, it will go from a solid.
It won't even melt to a liquid state.
It will go straight to a gas state.
Now, the last thing I wanna do
is think about why different molecular solids
will have different melting points.
So let's compare, for example, molecular iodine
to molecular chlorine.
Each of these can form molecular solid.
We looked at iodine a few minutes ago.
Which of these would you think would form molecular solids
with higher melting points?
Pause the video and think about that.
Well, as we talked about it, each of these molecules,
they're formed by covalent bonds between two atoms,
and what keeps the solid together
are these dispersion forces.
In an earlier videos,
when we first talked about dispersion forces,
we talked about temporary dipoles and induced dipoles,
and they were likely to form between heavier atoms
and molecules because they have larger electron clouds
and are more polarizable.
So if you compare molecular iodine to molecular chlorine,
you can see that iodine is clearly made up of larger atoms
and is therefore a larger molecule,
which is more polarizable.
So it's larger,
which means it's more polarizable, generally speaking,
polarizable,
which means it has stronger,
generally speaking, dispersion forces,
stronger dispersion forces.
Now, just as a reminder, these dispersion forces
are between molecules.
Each molecule has a covalent bond between two iodines,
and then the dispersion forces are between the molecules.
But because it has stronger dispersion forces,
we would expect that a molecular solid formed by iodine
is gonna have a higher melting point
than a molecular solid formed by chlorine.
And I actually do have the numbers here.
So the melting point
of a molecular solid formed by iodine,
we've already talked about that,
that's 113.7 degrees Celsius,
while the melting point of a molecular solid
formed by molecular chlorine
has a melting point of negative 101.5 degrees Celsius,
which is very cold,
and so iodine has a higher melting point
because of the stronger dispersion forces.
Now, as I said, those dispersion forces
are still not that strong.
This is still not that high of a temperature
compared to melting points of other types of solids
we have looked at in the past.