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

Let's talk a little bit about ionic solids, which you can imagine are solids formed by ions. So let's think a little bit about these ions. For example, we could look at group one elements here, especially things like lithium, sodium, or potassium. In many other videos, we have talked about these elements wanting maybe not so much to keep their outermost electron because they only have one electron in their outermost shell. It'd be pretty easy for them to lose that electron to get to a noble gas configuration to have a full outer shell.

So these characters like to lose one electron. The group two elements like to lose two electrons. Well, if you go on the other side of the periodic table, if you look at the halogens right over here, they are one electron away from having a noble gas electron configuration, from having a full outer shell. So they really like to grab electrons. If you look at elements like oxygen and sulfur, they really like to grab two electrons if they can.

What do you think happens if you have some metals on the left end here mixed with some non-metals on the right end here? Well, you might imagine there would be a reaction. For example, if you mixed sodium with chlorine, the sodiums might lose an electron to the chlorines, in which case you're going to have sodium cations—positively charged ions. And if the chlorines are now taking those electrons, they then become chloride anions.

Now, if you have a bunch of positive ions hanging about, hanging around a bunch of negative ions, what do you think is going to happen? They're going to get attracted to each other. They’re going to get attracted to each other and form a lattice structure like this. I like to use sodium chloride as an example because this is probably the one that we see most in our life. This is table salt. If you were to lick it, it tastes salty.

But there's many other ionic solids. Many of them would actually be categorized as salts. Generally, you could have a potassium chloride, you could have a sodium fluoride, you could have, for example, a magnesium oxide. What's going on there? Well, in that situation, each magnesium might lose two electrons so they become an ion with a positive two charge, and each of the oxygens would gain two electrons so then they are anions with a negative two charge.

These characters once again are going to be attracted to each other and form an ionic solid in a regular lattice structure like this. So let's think a little bit about their properties. First of all, let's think about the melting points. These solids, the electrostatic attraction between these ions is strong, and so they tend to have high melting points.

What if we were to compare melting points between ionic solids? For example, if you wanted to compare the melting point of sodium chloride to the melting point of magnesium oxide, which one do you think has a higher melting point? Pause this video and think about it.

As you can imagine, the electrostatic attraction is going to be dependent on two things: the magnitude of the charge and the radius of the atoms that are forming this lattice structure. The magnitude of the charge here is clear. Here you have a plus two charge being attracted to a negative two charge. So this has a stronger electrostatic attraction.

You're going to have a higher melting point right over here: the melting point of magnesium oxide is 2825 degrees Celsius, while the melting point of table salt or sodium chloride is 801 degrees Celsius. You could also try to compare sodium chloride to something like sodium fluoride. Which one do you think is going to have a higher melting point, sodium chloride or sodium fluoride?

Well, fluorines are smaller than chlorines. If each of them gains an electron, then the fluoride anion is still going to be a reasonable bit smaller than the chloride anion. When you have smaller constituent ions, the electrostatic attraction is actually stronger. Remember, we've seen in Coulomb's law that the closer two charges are to each other, the stronger the attractive or the repulsive force; and if they're opposite charges, it's going to be an attractive force.

So sodium fluoride is actually going to have a higher melting point than sodium chloride by a little bit. It actually turns out that the melting point of sodium fluoride is 996 degrees Celsius. If you're comparing these three, the highest melting point is magnesium oxide, followed by sodium fluoride, followed by sodium chloride. So charge is what's really dominating over here.

Now, the next question you might be wondering is, all right, I can imagine these solids are really hard, but what would happen if I were to try to break it? Would it bend like a lot of the metals we know, and we'll study that in other videos, or would something else happen? To understand that, let me draw a two-dimensional representation of this.

Let me draw the chlorine, or I should say the chloride anions, and this is just a two-dimensional version of that lattice. Obviously not drawing it to scale. Then let me draw the sodiums, sodium cations. As you can see, the positives are attracted to the negative; that's why they're next to each other. The negatives aren't next to each other because they repel each other. The positives aren't next to each other.

But what would happen if I were to try to press down really hard on this side? And if I were to press really hard up on this side? So what would happen if I press hard enough that this side begins to budge? So it begins to budge. Would it just bend, or what do you think's going to happen when I get right about there?

Well, when I get right about there, all of a sudden, I've broken—not only have I broken the lattice, but the negatives are next to the negatives, and the positives are next to the positives. So it's not just going to bend and be malleable like a lot of the metals we've seen; it's just going to break.

So this is going to be, even though it's going to be hard, it is going to be brittle. Now, the last question we'll address in this video is how good do you think ionic solids conduct electricity? Pause this video and think about that.

Well, in order to conduct electricity, either electrons or charge generally has to be able to move about. When it's just in its solid form like this, even though you do have these ions, they're not going to move about. So ionic solids in their solid form aren't good at conducting electricity. They can be good at conducting electricity if you were to dissolve it in a solution. For example, if you were to dissolve this salt in water, now the ions can move around, and then they're good at conducting electricity.

Or if you were to heat the sodium chloride up beyond 801 degrees Celsius and it turns into a liquid, then once again, the ions can move around, and you can actually conduct electricity. Take everything I say with a grain of salt. I'm sorry, I know I couldn't help it, but hopefully, you know a little bit more about ionic solids now.