 Okay guys, let's build the vestibular structure of ICL3, iodine trifluoride. So first off, what did we say, period 3 and below can do what? They can expand their valence, right? So that means that even though initially when we saw this we thought, wow, that's so weird because all of them only have one spot for an electron, right? So which one would be the central atom? Now we know that period 3 and below can expand their valence or use all of their valence electrons to grab another atom to make a bond, okay? So since that's the case, well, we look up at the periodic table and we see, well, chlorine's period 3 and iodine's period 3, which one is the central atom? So we're still not clear on it. But what did we say? How do we identify? All is to electronegativity, right? So iodine has that, chlorine's got a bigger electronegativity, so iodine's going to be the central atom, okay? So let's go ahead and draw the iodine as our central atom and then draw it, so the Lewis structure, then draw its valence electrons around it, and we'll draw the chlorine and I'm going to draw that space next to the iodine there. Ready to get to this point? So now we're going to show our fish hook arrows. So notice, this we would normally think of as a long pair, right? But what can happen is it can split those two up, okay? That might be the money fall, guys. It's like a good call, but of course I might not be hanging out with you guys very much longer. Okay, so that one? Sad, I don't know. Like y'all can tell me how much you love me after class. Don't do it already. Is everybody okay with that? So notice, the ones that I didn't use, I still want to keep them as a pair, okay? Because electrons prefer that. So now we're going to get something that looks like this, one going back like that. If you look here on my little list, right, you're going to say how many electron groups do I have around that thing? One, two, three, four, five. Is everybody okay with that? So sadly enough, you're going to have to memorize all of this stuff, okay? So notice I have a five there. Is everybody okay with that? So it's going to be in that group, okay? So now we look and say, well, how many atoms do I have? One, two, three, right? And how many long pairs? One, two. So I look for that thing. One, two, three, one, two, it's a T-shaped molecule, okay? So even a better way to draw this would be with the orbitals kind of pointed to the sides, okay? But you can't really show that very well on the board without showing the orbitals themselves. So oftentimes you'll see this molecule drawn like this, but it kind of looks weird. So this long pair is coming out towards us and that one's back away. Does that make sense? Okay, so we call this T-shaped, okay? And if you look, you should figure out that these bond angles are 90 degrees there and 180 degrees, respectively, okay? Could you guys do that on your own? So we've got to remember these things are spherical, okay? I'm not going to go over every one of them because there's a lot, but I have recorded a bunch of them, okay? And they're not all embedded, unfortunately, in the lecture, okay? So what you want to do if you, I think I have a playlist, but just type in this word in the search bar and you'll get a bunch of them, okay? So just type in that word to get a bunch of them and then some of them will be T-shaped, some of them will be octahedral, some of them will be square planar, some of them will be trigonal bipyramidal, okay? Go through all of them. So again, we don't have time to go through every one of them in here and you guys probably don't have the patience to go through every one of them. Everybody's sleeping the other time might be done. Okay, can I, about, this video probably won't be posted, but I'm going to kill it unless you guys got questions.