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.