Molecular solids | Intermolecular forces and properties | AP Chemistry | Khan Academy

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 negatives are attracted to the positive, the positives are attracted to the negative. 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 solids, 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. We've talked about those properties: very good at conducting electricity, malleable, etc.

Now what we're going to do is talk about what happens when you have non-metals. So the non-metals you can see in yellow right over here also include hydrogen. Now, of course, noble gases are also non-metals, but they're not reactive. So we're going to talk about the reactive non-metals. 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 non-metals.

Now, when we're talking 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 is 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.

Then that means that 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 a neighboring molecule, and then they'll be attracted to each other. We've talked about that as London dispersion forces, and at a sufficiently low temperature, that can keep them all together 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 of 103 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 do you think they're going to 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 want to 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 and 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 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—not even; it won't even melt to a liquid state—it will go straight to a gas state.

Now the last thing I want to 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 solids. 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. And in earlier videos, when we first talked about dispersion forces, we talked about temporary dipoles and induced dipoles, and they're more 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 going to 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.