- [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.