 All right, let's talk about polarity, the last of the big, big topics that we have to talk about in covalent bonding. Polarity is the result of atoms in a covalent bond not sharing electrons equally. Ideally, you'd have a situation like you do in N2, where the two nitrogen atoms are going to share a total of six electrons. It makes a triple bond. Each bond representing two electrons that are shared between the two atoms. These electrons will move back and forth through the hybrid orbitals between the two Nitrogens. This is a perfect scenario because when you look at the electronegativity for the two elements, the electronegativity will be exactly the same. Nitrogen is 3.0. So this nitrogen has an electronegativity of 3.0, this one has an electronegativity of 3.0. There is no difference. They're both equally good at attracting electrons to them in the covalent bond. So theoretically, these electrons will spend half their time over here at this atom and half their time over there at the other. So I mean to really drive this point home. Let's say we're talking about 10 minutes of time. The electrons will spend five minutes over here and they'll spend five minutes over here. In the end, that means that this atom will spend half its time being positive, half its time being negative. This one will spend half its time being positive, half its time being negative. All those charges cancel out and what we get is what we call a non-polar bond. That doesn't always happen though. There are situations where the elements that are involved in the bond have significantly different electronegativities and those differences in electronegativities cause the shearing to be kind of out of balance. That would happen in hydrogen fluoride, HF. Hydrogen fluoride is a two atom molecule, so we know it's going to be linear. Hydrogen and fluorine need to only make one bond a piece, so we know that it's just going to be a single bond between the two. So again, that dash represents the two electrons are going to be shared back and forth between these two atoms, but unlike here, it's not going to be a nice 50-50 thing. Hydrogen has an electronegativity of 2.1 and fluorine has an electronegativity of 4.0. There are significantly different electronegativities, a difference of 1.9. Because of that, we know the higher electronegativity atom, the fluorine, is going to keep the electrons most often, most of the time. And the hydrogen is going to be without the electrons most of the time, meaning that most of the time fluorine is a negative charge. And most of the time hydrogen is going to be a positive charge. They will flip. It's not permanent charge. This is not ionic bonding. But most of the time fluorine will be negative and hydrogen free positive. So again, if we want to oversimplify this just to drive the point home, if we're talking about 10 minutes worth of time, the electrons might spend eight minutes over here and only two minutes over here. And again, I'm just making these numbers up to drive the point home. That means almost the whole time fluorine is sitting there being a negative charge and almost all the time hydrogen is sitting here being a positive charge. This develops some weak charges in the molecule and creates what we call a polar bond. So that's what polarity is all about. It's all about this sharing of electrons, whether or not the atoms share them equally as they do at N2 and there are no charges produced in the molecule. We get a non-polar bond. Or whether there's some unequal sharing and the electrons spend more of their time around one atom than the other and we end up developing weak positive negative charges. Now if this polarity extends to the entire molecule, in other words, the polar bonds that are produced produce polarity in the molecule, which means the molecule develops a positive and negative side to it, then that's going to affect the behavior of the particle altogether because of what we call intermolecular forces. Non-polar bonds, like we had at N2, automatically produce non-polar molecules. Non-polar bonds automatically produce non-polar molecules. Polar bonds, while required, don't necessarily mean you have a polar molecule. To have a polar molecule, you need two things. First thing you need is polar bonds. Again, if you have non-polar bonds, it's automatically a non-polar molecule because of that. You have to have polar bonds to have a polar molecule, but that's not the only criteria. There's a second thing you have to have and that is separation of charge. The charges have to be on opposite sides or opposite ends of the molecule to have any real effect. Now, we've already talked about how you figure that out. You're going to look at electronegativity differences and you're going to use this scale. Zero to point four is non-polar bonds. Zero to point four is non-polar bonds. So when you go through this and you get a difference that falls between zero and point four, you're going to say non-polar bonds automatically non-polar molecule. The range for polarity is point five to two point zero. So again, you're going to go through the subtracting process and if it's point five to two point zero, you're going to have the polar bond, so you're going to meet the first criteria and then you'll move on to the second. So as you see here, in our hydrogen fluoride, we had a difference of one point nine. That's why I called that a polar bond. Now you might be thinking, why the upper cap? Why is there this upper limit? And that's because if you go over two point zero, then the sharing stops and when the sharing stops, you have ionic bonds. So anything over two point zero is technically an ionic bond and there's no sharing that takes place whatsoever. What you get there is permanent electrical charge, a transfer of electrons, where the one with a low electronegativity loses the electron and is a permanent positive charge, the one that takes the electron because it's high electronegativity becomes a permanent negative charge. So that's ionic bonding at that point. This is almost ionic. It goes right up to that higher limit. So the question is, how do we figure out if this has separation of charge? And the answer is you try to draw a straight line through it. If you can put all the positive charges on one side of that straight line and all the negative charges on the other, then it's a polar molecule. And for this one, I would have to draw my line vertically like that. I would have my positive charge on one side of the line, my negative charge on the other, so this is a polar molecule. I'm going to go through the three examples that I did in class just to make sure that you have them for reference whenever you need them. And the three examples that I did in class was phosphorus triiodide, PI3. I did SiO2 and then I did NH3 ammonia. Silicon dioxide and ammonia. Those are the three I did in class. So let's take a look at how these things work out. First, it's best if we go ahead and figure out if we have polar bonds or not because if the bonds are polar, the molecule is not polar. And we can stop looking at anything that falls into that category. It's only when the bonds are polar that we have to go on and we have to go deeper. So we look up our electronegativities. Phosphorus is a 2.1 and iodine is a 2.5. For an electronegativity difference of 0.4, this falls into the non-polar range. That is a non-polar bond, so automatically we have a non-polar molecule. Again, that's why it's good to look at these electronegativities first. We can figure out where we have to do more work and where we don't. This one, we don't have to do any more work. We're done with that. We know that. Silicon is a 1.8. Oxygen is a 3.5. 1.8 and 3.5 is the difference of 1.7. That falls into the polar range. So we know we have polar bonds here. We know how you have to work further. I have to get a shape figured out for that. So then I can figure out if I have a polar molecule. Let's go ahead and do the ammonia next. Nitrogen is 3.0. Hydrogen is 2.1. For a difference of 0.9, again, the polar range is 0.5 to 2.0. So this is in that polar range. So we do have polar bonds. So I know I have to do more work with that one as well. I need a structural formula. I need to figure out if my charges are separated or not. So let's go ahead and continue on with the SiO2, the silicon dioxide. Figuring out its shape. Silicon is in group 14. It has four single dots in its Lewis structure. It can make four covalent bonds. Oxygen is in group 16. It has six valence electrons. It can make two covalent bonds because of those two single dots. And I have two of them. I can bond that up there. I can bond that up there. Bond that there. Bond that there. Now this illustration tells me a couple of very important things. When I look at the silicon, the center atom, I've used up all the electrons. There are no unshared electrons there, so I know the shape is linear. So it's going to be O, SiO. Linear. Everything in a straight line. The other thing it tells me is because I circled two sets of electrons between the oxygen and the silicon, I know it's a double bond there. And again, because I circled two sets of electrons over here between the silicon and oxygen, I know it's a double bond on that side as well. So those Lewis structures modeling the bonding allows us to figure out what the shape is going to be and what the bonding is going to be. That's why they're so important. Now that's the structure. The next thing I got to do is figure out what's positive and what's negative. For that, I look at the electronegativities. Again, the one that has the higher number has a stronger attractive force in the electrons. The one with the higher number is going to be negative. So oxygen ends up being negative. The one with the lower number ends up being positive. So I come back to my structure. I'll put a little negative sign there. Negative sign there next to my oxygens. And I'll put a little plus sign next to my silicon. Then I got to look for that separation of charge. And I got to ask myself, is there any way I can draw a straight line through that, put all the pluses on one side, all the negatives on the other. It can be any orientation, horizontal, vertical, diagonal. But no matter what I try to do here, I can't separate the charge. So I fail on the second criteria. While I have polar bonds, it is a non-polar molecule. And again, that's because there's no separation of charge there. Now I move on to the ammonia, the NH3, and I'll do the same thing with it. So I'll model the bonding. Electrogens in group 15, so it has five valence electrons. That is one pair in three singles. I know because of that pair there, it's going to be one of my deformed shapes. It's either going to be bent or pyramidal because of that. Whenever I see those double dots on my middle atom, and again, nitrogen is my middle atom for a couple reasons. One, there's only one of them. So it has to be that center attaching point. And two, it's got the most single dots. When it's got the most single dots, it has the most bonding points. You can attach the most stuff to it. When I see those paired up electrons on the middle atom, I know that my three atom ones are going to be bent and my four atom ones are going to be pyramidal. This is a four atom one, so I already know what the shape is going to be. I already know it's going to be pyramidal. This will help me figure out if my bonding is single or double. Hydrogens in group one has one valence electron. I can see because of the single loop between all the nitrogens and the hydrogens that it's all single bonding. So to draw the pyramidal, I put my center atom in first and then I attach the other three underneath it. Like the legs on a tripod. If you were viewing it from the side, all the hydrogens would be below the nitrogen. Now I'm going to assign my positive and negative charges. Again, I go back to here. The one with the higher charge, the nitrogen, that's negative. The one with the lower electronegativity, that's the positive one. So I do a little negative sign by the nitrogen and I put little positives by the hydrogens. And I ask myself, is there any way I can draw a straight line through that? Put all those positive signs on one side and the negative sign on the other and I can if I draw my line horizontally. Again, the line can be horizontal, vertical, or diagonal. It doesn't matter what the line's orientation is. You just have to have a straight line in just one of them. Positives on one side, negative on the other. So I do have separation of charge here. So I can say not only does this have polar bonds, it is a polar molecule.