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Let's use the VSEPR theory to predict the structure of this molecule-- so phosphorus pentachloride So the first thing we need to do is draw a dot structure to show our valence electrons. We find phosphorus in Group 5, So 5 valence electrons Chlorine in Group 7, So 7 valence electrons and I have 5 of them So 7 times 5 is 35 Plus 5 gives us a total of 40 valence electrons that we need to show in our dot structure. So phosphorus goes in the center because it is not as electronegative as chlorine. And we have five chlorines. So we go ahead and put our five chlorines around our central phosphorus atoms like that. If we see how many valence electrons we've drawn so far, this would to be 2, 4, 6, 8, and 10. So 40 minus 10 gives us 30 valence electrons left over. And remember, you start putting those leftover electrons on your terminal atom. So we're going to put those on the chlorines Each chlorine's going to follow the octet rule. So that means each chlorine needs 6 more electrons. Now, each chlorine is surrounded by 8 valence electrons like that. So if I'm adding 6 more electrons to 5 atoms, 6 times 5 is 30. So I have now represented all of my valence electrons on my dot structure. Notice that phosphorus is exceeding the octet rule here. There are 10 valence electrons around phosphorus. And it's Ok for the phosphorus to do that because it's in Period 3 on the periodic table. I like to think about formal charge. And so if you assign a formal charge to phosphorus, you'll see it has a formal charge of 0. And that helps to explain- for me, anyway- the resulting dot structure. Now, step two We're going to count the number of electron clouds that surround our central atom. Remember, an electron cloud is just a region of electron density. So I could think about these bonding electrons in here as a region of electron density around my central atom. I could think about these bonding electrons, too. So here's another electron cloud. And you can see we have a total of five electron clouds around our central atom. The next step is to predict the geometry of the electron clouds. Those valence shell electrons are going to repel each other. All right. So that's a VSEPR theory-- Valence Shell Electron Pair Repulsion Since they're all negatively charged, they're going to repel and try to get as far away from each other as they possibly can in space. When you have five electron pairs, it turns out the furthest they can get away from each other in space is a shape called a trigonal bipyramidal shape. So let me see if I can draw our molecule in that shape. We're going to have our phosphorus in the center, and we're going to have three chlorines on the same plane. So let me attempt to show three chlorines on the same plane here. These are called the equatorial positions because they're kind of along the equator, if you will. So three chlorines in the same plane, one chlorine above the plane, and one chlorine below the plane. Those are called axial positions. All right. So there's a quick sketch Let me see if I can draw a slightly better shape of a trigonal bipyramial shape here. So let me see if I can draw one over here so you can see what it looks like a little bit better. So we could have one pyramid looking something like that. And then, down here, let's see if we can draw another pyramid in here like that So that's a rough drawing, but we're trying to go for a trigonal bipyramidal shape here So let's focus in on those chlorines that are on the same plane first. If I'm looking at these three chlorines and I go over here to my trigonal bipyramidal shape, you could think about those 3 chlorines as being at these corners here. So it's a little bit easier to see. They're in the same plane. So those are the equatorial chlorines. When I think about the bond angle for those- so those chlorines being in the same plane, you have these three bond angles here. And so when we did trigonal planar, we talked about 360 degrees divided by 3- giving us a bond angle of 120 degrees. So you could think about that as being a bond angle of 120. All right. So same idea, Those bonding electrons are going to repel each other. When we focus in on our axial chlorines- so this one up here and this one down here. You could think about those as being here and here on your trigonal bipyramidal shape like that. And if you draw the axis, if you draw a line down this way connecting those, it's easy to see those are 180 degrees from each other. So you could think about a bond angle of 180 degrees between your chlorines like that. And then, finally, if we think about the bond angle between, let's say, this axial chlorine up here at the top and then one of these green chlorines right here, I think it's a little bit easier to see that's 90 degrees here. So this bond angle right here would be 90 degrees. And so those are your three ideal bond angles for a trigonal bipyramidal situation here. It's important to understand this trigonal bipyramidal shape because all of the five electron cloud drawings that we're going to do are going to have the electron clouds want to take this shape. So it's important to understand those positions. For step four, ignore any lone pairs and predict the geometry of the molecule. Well, there are no lone pairs on our central phosphorus. So the electron clouds take a trigonal bipyramidal shape and so does the molecule. Let's go ahead and do another example. Sulfur tetrafluoride here. So we're going to start by drawing the dot structure, and we need to count our valence electrons, of course. So sulfur's in Group 6, so 6 valence electrons. Fluorine is in Group 7, so 7 valence electrons. I have 4 of them. 7 times 4 is 28. 28 plus 6 is 34 valence electrons. We know sulfur is going to go in the center because fluorine is much more electronegative. We put sulfur in the center here. We know sulfur is bonded to 4 fluorines. So we put our fluorines around like that. And let's see how many valence electrons we've shown so far- 2, 4, 6, and 8. So 34 minus 8 gives us 26 valence electrons we still need to account for on our dot structure.