Learn about Advanced Placement Chemistry, Chemical bond 4, in this comprehensive video by bannanaiscool.
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The last of the loose diagram set of shapes are the six effective parallels. So if you have PCl6 negative here’s your Lewis diagram. So there’s your PCl6, now you're going to have to put six bonds around the central phosphorous and there’s your shape—four sides here and a total of eight sides an octahedral that’s what it’s called. It looks the same everywhere you turn it. So now you’ve got your octahedral shape for six effective pairs and that’s your seed shape as it were. So what do you get when you have another one that’s going to—SF5 negative. S is going to have with a long pair six effective pairs around it when you do the Lewis diagram so that’s going to be octahedral but you need a long pair in there. Where’s the long pair going to go? Let’s put it right on the top. So therefore what do you get out of this? You have a pyramid that has four sides. So we call it square pyramidal because the base is square instead of like trigonal as before. So square pyramidal when you have six effective pairs but one of them is a lone pair. And then when you have the zenon tetraflouride—yeah you can bond noble gases like that they do bond you know. And six effective pairs but two of them are lone pairs where are you going to stick these next lone pair? Where do you think? Farthest away from the other lone pair as possible and now you’ve got yourself what? Well it’s plainer isn’t it? And it’s a square. Square plainer is the shape. Now you’ve got all the shapes for all of the molecules that we can do.