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- [Voiceover] We've already seen that formal charge is equal to the number of valence electrons in the free atom, minus the number of valence electrons in the bonded atom, and another way of saying that is the formal charge is equal to the number of valence electrons the atom is supposed to have, minus the number of valence electrons the atom actually has in the drawing, so let's assign a formal charge to oxygen in this molecule. Remember that each bond is made up of two electrons, so this bond right here is made up of two electrons. This bond over here is made up of two electrons, and our goal is to find the formal charge on oxygen, so the formal charge on oxygen is equal to the number of valence electrons in the free atom, so the number of valence electrons that oxygen is supposed to have, we know that's six, right, oxygen is supposed to have six valence electrons, minus the number of electrons that oxygen actually has in our drawing. Now remember when you have a bond with two electrons, we give one electron to one atom, and the other electron to the other atom. So from these two electrons, oxygen gets one of those electrons. All right, same thing for this other bond here, oxygen gets one of those electrons, and the other electron goes to hydrogen. So now we have a total of six electrons around our oxygens, so I'll highlight those. We have one, two, three, four, five, six, so there are six valence electrons around the oxygen in our drawing, so six minus six is equal to zero, so the formal charge on oxygen is equal to zero. And let me go ahead and write down this pattern that we've just seen. When oxygen has two bonds, and two lone pairs of electrons, one oxygen has two bonds and two lone pairs of electrons, the formal charge is equal to zero. And sometimes, the lone pairs are just left off for convenience reasons, right, so you could draw this with your oxygen, and your hydrogen like that, or you could even go like this, and all those are just different ways of representing the same molecule. So leaving off lone pairs of electrons. Let's look at some other examples, where the formal charge on oxygen is equal to zero, and we'll look at the one on the left first, so the formal charge is equal to zero on this oxygen and we can see we have two bonds here to oxygen, here's one of the bonds to oxygen, and here's the other bond to oxygen. The lone pairs of electrons have been left off this dot structure, but we know, since the formal charge is zero, we already have our two bonds here. There should be two pairs of electrons on that oxygen, so you can put them in there or you could leave them off. I'll go ahead and put them in. So we have a total of eight electrons around our oxygen, so oxygen's following the octet rule here, so I'll highlight them, two, four, six, and eight. On the right, we have another example, where oxygen has a formal charge of zero, and this oxygen has two bonds to it so here's one of the bonds, and here's the other bond, and this oxygen would also have to have two lone pairs of electrons on it, and again, I didn't draw them in here. Sometimes you don't draw them in here for convenience, but I could go ahead and add them, so it's easier to see that that oxygen has a formal charge of zero. And a total of an octet of electrons around it, so let me highlight those, two, four, six, and eight. So these patterns are important, and for oxygen, two bonds, and two lone pairs of electrons give us a formal charge of zero. Let's move on to another formal charge situation for oxygen, let's find the formal charge on oxygen here, so we start by drawing in the electrons in our bonds. Each bond consists of two electrons, so I draw those in. What is the formal charge on oxygen this time? So the formal charge is equal to the number of valence electrons oxygen is supposed to have, which is six, minus the number of valence electrons oxygen actually has in our drawing. So we divide up our electrons again, so oxygen gets one from this bond, and one from this bond, and one from this bond, so how many electrons are around oxygen now? This would be one, two, three, four, and five. So six minus five is equal to plus one, so it's like oxygen has lost an electron here. So oxygen has a formal charge of plus one. I could re-draw that, let me go ahead and do that over here. So I could re-draw that over here, so we have oxygen with our bond to hydrogens, and oxygen has a lone pair of electrons on it, and this oxygen has a plus one formal charge. And so we can come up with another pattern here. Here oxygen has three bonds, let me highlight those bonds. I'll use red this time, so here's one bond, two bonds, and then three bonds and then one lone pair of electrons, so the pattern of three bonds plus one lone pair of electrons for oxygen will give you a formal charge of plus one, and again, it's good to recognize these patterns. You should be able to do the calculation, and then after you do enough of these problems, you can just look at it and figure out what the formal charge is. Let's look at some more examples where oxygen has a formal charge of plus one, and the lone pairs were left off of this one, again, for convenience reasons, so we'll start with this example on the left, we can see that the oxygen with a plus one formal charge has three bonds to it. Here's one, here's two, and then here's three, so three bonds, in order for that oxygen to have a plus one formal charge, it must also have one lone pair of electrons on it, so you could just leave them off and know that they're there, or you could go ahead and draw them in, and I'll draw them in on that oxygen. Now to our other example over here on the right. This is the oxygen with a plus one formal charge, so that oxygen must have three bonds and one lone pair. So here are the three bonds, here's one, two, and three, and again, I didn't draw in the lone pair of electrons on the oxygen, but the lone pair is there, so I'll go ahead and put it in like that. So recognize this pattern, three bonds, plus one lone pair for oxygen, gives us a formal charge of plus one. You could have also figured out how many electrons are necessary. Let's use this example, by, let me go ahead and re-draw it here. So how else could we figure out how many electrons are on that oxygen, if that oxygen has a plus one formal charge. Well you could say that, all right, the calculation for formal charge would be six minus x is what we don't know, but we do know the formal charge is plus one, so let me just put this in a little box over here. So six minus x is equal to plus one. Obviously, x would have to be equal to five, meaning that oxygen would have five electrons around it, let's think about how many electrons we see right now around oxygens. So let me draw in some electrons here and I'll use red, so this bond is two electrons, and then we have a bunch of electrons in here, so how many electrons around oxygen do we have so far? Well, we would have three, just these three, and we need five, which means we need two more, which means we need a lone pair of electrons on the oxygen. So that's a little bit too complicated, I think, for figuring it out, but you could use that method, or you could just learn this pattern, right, and eventually, you'll have to have this pattern down pretty well. Let's look at another example for assigning formal charge to oxygen. So our goal is to find the formal charge on oxygen in this example, and we put in our electrons in this bond. Each bond represents two electrons, so the formal charge on oxygen, is equal to the number of valence electrons that oxygen is supposed to have which is six, minus the number of valence electrons that oxygen actually has, and in this example, right, we would take one of these electrons from this bond, and how many electrons is that total around oxygen? This would be one, two, three, four, five, six, seven, so six minus seven gives us a formal charge of negative one. So I could re-draw that over here, I could say oxygen has three lone pairs of electrons, and a negative one formal charge, so our pattern, this time, our pattern is one bond. Here's the one bond, and then three lone pairs of electrons. So let me write that down. So the pattern is one bond, when oxygen has one bond, and three lone pairs of electrons the formal charge is negative one, just like we saw up here with the calculation. So we could leave those electrons off if you wanted to save some time, we could just say, oh, this is oxygen with a negative one formal charge, and we should know that there must be three lone pairs of electrons on that oxygen. Let's look at one more example, where formal charge is negative one. So right here, this oxygen has a negative one formal charge, and we can see it already has one bond to it. And so the pattern, of course, is one bond plus three lone pairs of electrons. So we already have the one bond. In order for that oxygen to have a negative one formal charge, we need three lone pairs of electrons. So we could re-draw this, so that is one way to represent that ion, and we could also represent it like this, with putting three lone pairs of electrons on that oxygen with a negative one formal charge. So again, become familiar with these patterns.