 Hi, well, I'm Stephen Nesheba, and I want to tell you a little bit about electron delocalization and resonance, and the set of ideas that I want to touch on here, listed here of the x, y, z pattern of the lone pairs and so forth, and to do that, I'll just, I mean, I have a number of examples to show you. So the x, y, z pattern, what does that mean? Well, what it means is I imagine that I have three atoms, there's a double bond between the first two and the third, it's got a single bond to it, but it has a lone pair of electrons, and whenever you see that x, y, z pattern, one thing that you can think is, oh, I could imagine another Lewis structure that involves those electrons coming down there forming a double bond, and those electrons in here leaving that double bond and forming a lone pair of electrons on x. And that's what I've drawn here, that lone pair of electrons, there they are, the electrons that form that extra bond, the pi bond between y and z is right there. So let's do some examples, here's o-zone, okay, does it satisfy the x, y, z pattern? Well, I think so, I've got a double bond here, single bond, and a third atom that's got a lone pair, it actually has three lone pairs, and so I can imagine now this pair of electrons coming down here, those pair of electrons leaving to form a lone pair, and here's that double bond. Now, a little bit about the nomenclature here, what we've drawn here are different resonance forms, okay, and so there's the word resonance, and the pair of electrons that came down off z to form the double bond, that's called the allelic lone pair of electrons. I have this phrase equivalent or unequivalent, what we have to decide here is do I have any reason to prefer this resonance form over that one, and the answer here is no, these oxidants are all identical to each other, so I would say these are equivalent resonance forms, okay, and when you have equivalence resonance forms, it means that I have to, in my mind, kind of think that reality is some average of those. That average gives rise to this idea of fractional bond order and partial charge, here's how it works. In this resonance form, I have a double bond, in this one I have a single bond, what's the truth? Well, it must be some sort of average, so I would say that the bond order between the first two oxidants is not one, not two, but one and a half, one and a half bond order, same thing here, one and a half. What about the charges? Well, I've drawn in the formal charges for you here, and both of these equivalent resonance forms, the central oxygen has a plus one charge, but notice here, there's a zero formal charge here, a minus one here, what's the average? Well, I would say the average must be minus and a half on that oxygen, same thing here, right, a minus one and a zero, so that must have a partial charge of minus and a half. So, whenever we have equivalent resonance forms, we typically get fractional bond orders and partial charges. One more thought here before I go into the next example, it's useful to think about how many electrons are delocalized, that is, how many are playing in this XYZ pattern. Well, it's kind of obvious here that there's those two that did that and then the two that left. So, we would say that the number of delocalized electrons in any XYZ pattern is two pairs, or, you know, there are four delocalized electrons that are, that ozone has, we would say, ozone has four delocalized electrons, pi electrons. Okay, how about NO2? Similar kind of story here, I have a double bond, single bond, and the third atom, the Z type atom has a lone pair, so I can do the same thing down they go, up you go, and so that lone pair is there now and that those electrons are not going to go, yeah, the second bond, the pi bond, partial charges looks like minus a half, plus one, minus a half again. Fractional bond orders, sure, I've got a bond order of two, bond order of one, the average two, one and a half. How about this one, this is CH3, C double bond O, CH2 with a net minus charge, and so I've got a lone pair of electrons on that carbon, and now I can say, oh yeah, that looks like a Z, down they go, up those electrons go to form now this oxygen with a minus one formal charge. This negative formal charge in that carbon has now gone away because it's formed a double bond. It's still XYZ, so we would still say that there are four de-localized electrons in this situation, but what the difference is that these are no longer equivalent residence forms. How do I know that? Well, in this case I have a formal charge of minus one on a carbon, in this case I have a formal minus charge on the oxygen, those are clearly not equivalent situations. However, we can also say when we have any equivalent residence forms, which one we like better, which one is predominant, and generally the way that we do that is we say, oh well, I think the more electronegative element oxygen will claim that minus one charge, so we generally prefer, we'd say this is a more predominant residence form. Okay, here's one last example. It doesn't satisfy the XYZ pattern, but there are still de-localized electrons. I just wanted to show you an example of this. So I have a carbon over here, okay, this is the allylic cation, it's called, and there's a, this carbon here doesn't even have a complete octet, and there's no lone pair floating around here, but this carbon could claim that pair of electrons by stealing it basically from that carbon, and that's what I've drawn here. Now this carbon has lost a pair of electrons, so it's got a plus one charge, and this carbon here has a complete octet, so it's happy. We have a similar thing that we would say, oh, the bond order must be one and a half, the partial charges must be plus a half, plus a half, and zero on that carbon. But the big difference now, of course, is that besides the fact that it's not XYZ de-localization, we also have just two de-localized electrons that are better at playing a role in that molecule. Okay.