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AP Chem: SAP‑3 (EU), SAP‑3.C (LO), SAP‑3.C.1 (EK)

Video transcript

- [Instructor] In this video, we're gonna think about how ions will arrange themselves when they form solid crystals. When they form these lattice structures. So just in very broad brush terms, let's say that we have a bunch of this white cation, and we have a bunch of this green, or this blue-green anion. So let's say they're in a one-to-one ratio. How will that look? How will the solid look, if you were to take a two-dimensional slice of it? To imagine that, we can draw what we could call particulate models. We're just imagining a two-dimensional slice of the solid, and we're just drawing these ions as particles. Would it look something like this, where maybe the positive ion is all on one side, and then the negative ion is on the other side, is on the bottom if we were to take a slice? Would something like this make sense? Or maybe it's random. Maybe you have a positive there, and then you have some negatives right over there. And then, maybe you have a positive and a positive, and then a positive right over there. And then maybe you have some negatives right over there. Would this be a reasonable configuration, as they form these ionic bonds? Well, when we think about Coulomb forces, we know that like charges repel each other and unlike charges, or opposite charges, attract each other. And so, when these ionic solids form, they're unlikely to form in this way, or even in this way, because they're gonna form in a way that maximizes the attractive forces and minimizes the repulsive, the repelling forces. And so what would be an arrangement that does that? Pause this video and think about it. Well, all the positive charges are gonna try to get as close as possible to the negative charges and as far as possible from other positive charges. And the same thing is going to be true of negative charges. They're gonna try to get as far away from other negative charges as possible, and as close to other positive charges as possible. So the arrangement that you are likely to see is going to look something more like a checkerboard pattern. So it may be a positive there, a positive there, a positive there, a positive there, and a positive there. These are all the same ion, I'm not drawing it perfectly, they'd be the same size. And when you do these two-dimensional representations, these particulate models, it is important to get the size right, 'cause we're gonna think about that in a second. And then the negative charges would be in between. So notice. In this configuration, every negative is surrounded by positives, and every positive is surrounded by negatives. So it's maximizing the attractive forces and it's minimizing the repulsive forces. And if you were to think about it in three dimensions, you would have a lattice structure that looks something like that. And we have seen this in other videos. Now another interesting thing to think about is the size of the ions that form that ionic solid. Let's say we wanted to deal with rubidium bromide. Rubidium bromide. What would this look like if I were to draw it in a two-dimensional particulate model like this, and I wanted to make the size roughly comparable to what we would see between the rubidium and the bromide? Pause this video and think about that, and I'll give you a little bit of a hint. It might be useful to look at this periodic table of elements. All right, if we were to separate this out into its ions, it is a rubidium cation, and a bromide anion. Now a rubidium cation, it has lost an electron. So even though it still has 37 protons, its electron configuration now looks like that of krypton. Now, the bromide anion, even though it only has 35 protons, it's going to gain an electron to become a bromide anion, and it also has an electron configuration of krypton. So both of these have the same number of electrons, but rubidium has two more protons than bromide does. And so the rubidium is going to attract that outer shell of electrons, that fourth shell of electrons, more than the bromide nucleus is going to. And so, the rubidium in this example is going to be smaller than the bromide. And so if I were to draw one of these diagrams, it would look something like this. Let me draw the bromide first. So I have a bromide anion, I have another bromide anion, another bromide anion, maybe I have a bromide anion right over here, bromide anion over there, maybe a few more. Make 'em a little bit, if I was doing this with a computer, I would make them all the same size. So these are our bromide anions. And then your rubidium cations would be a bit smaller. And so, our particulate model right over here might look something like this. We wanna make it clear that the cation is a bit smaller than the anion. It would arrange, it would likely arrange in a pattern that looks like this. And notice, I am trying to make the sizes roughly accurate, to show that the cation is indeed smaller than the anion. Although it wouldn't be dramatically smaller. Remember, they have the same number of electrons. And they don't have that dramatically different number of protons. And this is just a very rough drawing. If they were dramatically different, you might show that in the sizes on this diagram.
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