Learn about Advanced Placement Chemistry, Chemical bond 3, in this comprehensive video by bannanaiscool.
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Rob Lederer: Now that you know that you can exceed the octet rule, well, we can have more than the two, the three, or the four effective pairs around the central atom to give a shape for a molecule. Look at the formula for PCL5. That should have impressed you right away before when we've mentioned that molecule, you should say, well, how do you build that, because if you do a Louis diagram for it, phosphorous has two, four, six, eight, ten; it's already exceeded the octet rule, yeah, but it can. By the way, the formula chart for phosphorous here, it's in group five normally; one, two, three, four, five bonds and no lone pairs, gives it a formal charge zero, that's the best Louis diagram we can do for PCL5. So, PCL5 has five bonds around the phosphorous. So you know, one better than a tetrahedral, what we are going to be looking at here in three-dimensional space? Here is the shape for five effective pairs and it is the seed shape for five effective pairs. Everything is arranged in something called, 'ready for this', trigonal because there is three in the plane here that you can see, right? See that plain through in there. But if you make this into a three-dimensional shape with size; one, two, three, four, five, six size, you'd have like a three-sided pyramid on top and a three-sided pyramid below. The name trigonal bipyramidal because of the two pyramids that it makes, okay, that's the name. So, this is a trigonal bipyramidal shape and you can tell that all of these being in the same plane, looks to me like they are going to be a 120 degrees apart and 90 degrees away from this one at the top, 90 degrees away from the bottom on here and these guys are 180 degrees away from each other, that gives a very symmetrical, non-polar type of shape. Try, go and buy a pyramid. Now, PCL4 negative also exists as a molecules and it's got a phosphorous in the center, 4CL surrounded by an extra lone pair. So, we had still got five as the effective pair or the seed shape of trigonal bipyramidal, but one of the chlorines has to go here and we've got a lone pair in. Where do we take it off? We don't take it off of the top or the bottom. The lone pairs like to be in the plane because then they can be farther away from everything else, because lone pairs, VSEPR theory tell you that test they occupy the more space. They want more space; they go in the 120 degree plane always. So, we pluck off this chlorine. We understand that the lone pair is out here, what's the name of this shape? Believe it or not, this is called seesaw. It's a seesaw shape when you have five effective pairs, but only four bonds; seesaw shape. Now, CLF3 has one, two, three bonds surrounded to the central chlorine with two lone pairs around the chlorine to give five effective pairs. So, trigonal bipyramidal arrangement of electrons in the five effective pairs, but you got a lone pair that you got here and another lonely pair, where do you going to take that lone pair? From top or bottom? No, no, from the plane boss, the plane. So, there are lone pairs here and here in the plane, so what does that giving as a shape? T-shape. This is called T-shape when you have five effective pairs, but two lone pairs in the plane; T shape. The last one for five effective pairs; how about I3 negative? That has in the central atom one, two, three, four, five effective pairs. When you have extra lone pairs to put into the Louis diagram, because this has 22 valence electrons, you need to put them in the center. So, now we've got an eye with five effective pairs around it, but it's only got two bonds; the lone pairs in the plane, all, one, two, three of them, and that leaves you with linear.
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