 In the last video I introduced you to three basic molecular shapes predicted by VSEPR theory, linear, trigonal planar and tetrahedral shapes. In this video we're going to look at some variations on those shapes. So let's start with ammonia. First pause the video and draw the Lewis structure. You would have come up with this. So how do you think this molecule would arrange its bonds? Well it has three of them, so it might seem logical to say it's like the boron trichloride that we looked at in the last video. It's trigonal planar. But when you do crystallography on ammonia you find out that actually the shape of its molecules is like this. So what's going on? Well the main thing that differs between boron trichloride and ammonia is that the ammonia has a lone pair on the central atom. Could that make a difference? The answer is yes it can. Non-bonding electrons are still negatively charged and can exert a repulsive force just like bonding electrons. So if a central atom has one or more lone pairs those electrons will compete for space around the atom as well. In effect they act like a phantom bond. So rather than thinking of ammonia as being based on a central atom with three bonds around it, which would make us think of a trigonal planar shape, we need to think of it as having a central atom surrounded by three bonds and a lone pair, a total of four electron groups. And this makes it more like methane with a tetrahedral geometry. But although the lone pair here is drawn like a giant balloon it actually doesn't stick out like that. The lone pair repels the other three genuine bonds pushing them closer together but it doesn't take up space like the hydrogen atoms do. This means that the shape of the molecule is like a tetrahedral molecule with the top bond chopped off. So like a little pyramid. There's one more thing. If you look at the bond angles in ammonia there are 107 degrees. Do you remember what the bond angles were in methane? An undistorted tetrahedral molecule like methane will have bond angles of 109.5 degrees. Instead ammonia has slightly smaller bond angles 107 degrees which means the hydrogens have somehow been drawn closer together. What could cause this? Well the inference that we make is that a lone pair exerts a greater repulsive force than a bonding pair. This means that if you had two lone pairs repelling each other the force would be greater than if you had a lone pair versus a bond as we have here. And that would be greater than the force between two bonds. So here the three NH bonds lose out a bit and they're pushed closer together by the lone pair which makes the bond angles smaller. So ammonia with four electron groups around the central atom is based on a tetrahedral geometry but one of the groups is a lone pair so it doesn't show up in the final molecular shape which leaves us with a sort of truncated tetrahedron which we call a trigonal pyramidal molecule. And to draw that VSEPR structure you use a wedge and a dash to show the 3D shape and the lone pair so that you don't forget what's making it that shape in the first place.