If you're seeing this message, it means we're having trouble loading external resources on our website.

If you're behind a web filter, please make sure that the domains *.kastatic.org and *.kasandbox.org are unblocked.

Main content
SAP‑5 (EU)
SAP‑5.B (LO)
SAP‑5.B.1 (EK)
SAP‑5.B.2 (EK)
SAP‑5.B.3 (EK)

Video transcript

- [Instructor] 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. So, for example, we could look at group one elements here, especially things like lithium, or sodium, or potassium. And 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. And it'd be pretty easy for them to loose 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. While if you go on the other side of the periodic table, if you look at the halogens right over here, they're one electron away from having a noble gas electron configuration, from having a full outer shell. So, they really like to grab electrons. And, if you look at elements like oxygen and sulfur, they really like to grab two electrons, if they can. So, what do you think happens if you have some metals on the left end here mixed with some nonmetals on the right end here? Well, you might imagine there would be a reaction. So, 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. And now if you have a bunch of positive ions hanging around a bunch of negative ions, what do you think is going to happen? They're going to get attracted to each other. And 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'd taste salty. But, there's many other ionic salt solids, many of them would actually be categorized as salt, generally. You could have a potassium chloride. You could have a sodium chloride. You could have, for example, a magnesium oxide. What's going on there? Well, in that situation, the magnesium, each magnesium might lose two electrons, so they become a 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. And 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. So first of all, let's think about the melting points. So, these solids, the electrostatic attraction between these ions is strong. And so they tend to have high melting points. Now what if we were to compare melting points between ionic solids? So 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. Well, as you can imagine the electrostatic attraction, it's going to be dependent on two things. The magnitude of the charge and the radius of the atoms that ae forming this lattice structure. And 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 and so you're going to have a higher melting point right over here. The melting point of magnesium oxide? 2,825 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 and each of them gain an electron, then the fluoride anion is still going to be a reasonable bit smaller than the chloride anion. Or 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 gonna 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. But 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? And to understand that, let me draw a two dimensional representation of this. So 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. And 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 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, or I were 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 had enough that this side begins to budge? So it begins to budge. Would it just bend, or what do you think's gonna happen when I get right about there? Well, when I get right about there, all of a sudden I've, not only have I broken the lattice, but the negatives are next to the negatives and the positives are next to the positives and 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. And when it's just in its solid form like this, the, even though you do have these ions they're not going to move about. So ionic solids in their solid form, they 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 this 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. Sorry, I know, I couldn't help it. But hopefully you know a little bit more about ionic solids now.
AP® is a registered trademark of the College Board, which has not reviewed this resource.