 So here's another good one. This one actually is a polyatomic ion on I3 minors. And so it's pretty easy, I think, in this one to distinguish what's the central, right, because there's only I. The only thing you want to know about this is that one of the I minors, the central one, is expanding its bay lines, and the other one's off. So everybody OK with that? So you've got I. We're just going to put all of the Lewis dot structures. OK, so when you look at this, well, we still got to add another electron. But already when you look at this, you probably are like, if you didn't realize one of them has to expand its bay lines, you look at this, and you're like, I can't make single bonds to all of these things. So at that point, you should appreciate that one of them has to expand its bay lines, OK? So now we have to add our electron to it. So you've got to add the electron, especially in this case, the one that's expanding its bay lines, because if you added it to over here, it couldn't make a bond. You added it to over here, it can't make a bond. Does everybody see that? We're going to put that electron there. OK, so that's I minus I. It's going to happen here. Well, let's just use these two electrons to bond just so we don't have to perturb the structure here too much. And remember, what I'm drawing here is just the Lewis structure, so it doesn't really have any perspective information. So again, I can move these electrons all around if I want to, too. So I'm just going to do that because I'm going to draw it in such a way that gives us kind of a perception as to what it looks like. I already know that formal charge is on that iodine. And that's the iodine, because you put the electron in the iodine. If you didn't know that, you can calculate it using your formal charge. We're going to be erasing this. Well, hopefully you guys can tell me how many electron groups are around the central atom because I guess I gave it away. What does that determine? Can you tell me the number of electron geometry? The electronic geometry, right? So let's figure out what the electronic geometry is. What is the electronic geometry? How many groups does it have around it? 5 on octahedral. Trigonal bipyramidal, this is the trigonal bipyramidal structure. So you can see the two iodines that are making the sigma bonds are here and here, just the non-hypervalent iodines. And the middle one is the hypervalent. Of course, these, in this case, are electron groups, which you don't see. Is everybody OK with that? So I mean, we could attempt to draw this structure emphasizing the lone pairs. Let's try to do that in such a way. It's not going to be the easiest thing to do. I would accept this as actually the perspective structure as well, you know, so the best for structure. But you can think of it as fully you are. One of those lone pairs will be in the plane, and one of them will be coming out towards us. So if that's OK for your eyes or whatever. And then the other, what you're going to, is that. And of course, the negative formal charge on that iodine. Is everybody OK with that structure? So hopefully you can see the trigonal bipyramidal nature of the electronic geometry. What would we say that the molecular geometry is? It could fit. Linear. Linear. And this is non-polar because it's also standard. Yeah, non-polar. So. It wouldn't form like a resonance structure because those can't make up bonds. Oh yeah, so we're not talking. So I mean, somebody may say you can do non-bombing resonance, but this is general chemistry. We're not talking anything about the advanced trigonometry. So yeah, no. Can't break single bonds for resonance. Let's just make that rule of thumb, OK, for this bond. Oh, you can only go from, like let's say there were double bonds. There were double bonds, then you can move it. You know, or triple bonds, you can move it, OK? So that's resonance, if you're moving a double bond to an atom or to another bond, you know? So we'll, I mean, we saw one earlier. I think we've gone over a few. Any other questions on this?