Okay, on with part three from chapter one, and I'm not going through every slide. I'm just hitting some highlights here. When it comes to covalent bonds, you won't be able to put together molecules showing the presence or absence of double bonds or triple bonds. Whenever those bonds show up, they are important, and so we want to know when they're called for. If you've only got two or three atoms, like in the previous examples, you can draw them individually and just kind of connect the dots literally. But when you have several atoms, and when there is the possibility of these extra bonds, the double and triple, it's not always obvious where they go. So these rules will kind of step you through any problem that we would encounter when it comes to building molecules. As it says in rule three, the connectivity has to be given. But apart from that, we can use these rules of valence electrons to come up with correct structures. The octet rule is our constant guide, but there's always going to be a set number of electrons we can use to get us there. So examples like this one for formaldehyde, off to the right is the calculation of how many valence electrons the whole molecule brings to the table. I really don't care in my final answer which electrons came from carbon or oxygen. I just need the grand total. And here that total is 12, six of which are already drawn. So if I build on top of this structure, as you see here, we can't get carbon and oxygen satisfied just by putting in pairs of dots. We end up having to make a double bond between the carbon and the oxygen. And like I say, we want to know when that's called for. So by knowing we can't use more than 12 electrons for this molecule. And by knowing that I have to get carbon and oxygen to satisfy the octet, double bonds, the only thing that works. Hydrogens are an easy case. They never make double bonds. They never have dots around them. They just make one single bond to get their two electrons. This is another example where you don't have a whole lot of electrons to work with, but both that carbon and nitrogen need to satisfy the octet. So this molecule only has 10 total. And if I consider the two bonds already shown, those have to be there. Plus putting those six dots around the nitrogen. Well, almost done, but this carbon in the middle still isn't happy until we move those electrons off of nitrogen. We're going to have to make a triple bond. Notice that gives eight electrons around both carbon and nitrogen. And it does that using a grand total of 10. So we're constantly counting electrons and then distributing them. Polyatomic ions like nitrate work the same way with the exception that negative charge, as you see, is an extra electron. We have to put in our total. So whereas I would normally count 23 electrons for nitrate, that negative charge means there's 24. If there were a positive charge up top here, we would have taken away an electron. And so atoms will make those adjustments when they can do something that's more stable. And as I put in those electrons, I find that with 24 grand total, I'm going to have to put a double bond between nitrogen and one of those oxygens in order to get the octet. Rule to be satisfied for all atoms. And that's what I'm trying to figure out. If I need a double bond, where can it go? And in this case, as it says at the bottom, we've got a choice of where we make the double bond. So there's really three different answers. And that leads into this next topic of resonance. Because when we have more than one way to draw the octet rule structure, we generally draw all the possibilities. Because nitrate here really has three different answers that would satisfy our requirements. Resonance is this phenomenon where we can draw more than one structure that obeys the octet rule. And it just changes the location of things like double bonds or triple bonds. Ozone is a simple example of this because you can see those two structures are very similar, but the location of the double bond is different. And as it says, we really want to draw both of these because the real answer to what ozone looks like is a hybrid of these two. This next slide explains that as well. So being able to recognize the existence of resonance and when it occurs and how many structures we can draw, that makes for good practice, but it also represents the reality of how these molecules are really made. In real life, ozone doesn't have a short double bond and a long single bond. The real answer is somewhere in the middle. So we want to draw a structure to reflect that.