Lecture 08. Chemical Bonds
Brindley, Amanda
Chem 1A is the first quarter of General Chemistry and covers the following topics: atomic structure; general properties of the elements; covalent, ionic, and metallic bonding; intermolecular forces; mass relationships. 
 Okay, so today we're going to start chapter two. And chapter two is going to start off with a lot of review, mostly because I thought you needed to know a lot of this before we really even got started with class and so I talked about it earlier. But we should go over it again and just make sure it's cemented in there before we start talking about the more complicated aspects of chapter two. So we're going to go through it and we're going to remind you about ionic and covalent bonding so that you remember exactly what that is and what the differences are because it becomes really important here. All your naming is still really important here and knowing how to identify between each one is very important. Now we're going to introduce Lewis Dot symbols in a little bit more of a formal nature. Before we kind of threw them at you and said this is some really simple ones, this is how we represent them, now we're going to go through and we're going to do a bunch of very complicated ones that go way more into depth on what exactly is going on which will lead us into trying to talk about how molecules look like and doing geometries and things like that. Then we're going to do some more Lewis Dot structures that are even more complicated and more Lewis Dot structures. Really chapter two is kind of all about Lewis Dot structures and getting good at them. Once we've really got that down packed, we're going to go through and do some bond strength and length comparisons so that we can see how they're related to each other. Again some review from last time. Bonding is all about trying to figure out how to complete octets or make things more stable than they were. Now you can do that by in something like this case sharing electrons. If two different electrons have six valence electrons, they can go share them. They can take two of their electrons, share them with each other and then form some bonds there so that they each can have eight. There's our three main types of bonding. We have our ionic bonding, our covalent bonding and our metallic bonding. Remember these are important for knowing how to name them because we use different names for our ionic than we do covalent. They're very different from each other. So in ionic you're trading an electron. Sodium says, hey I have an extra electron and chlorine says, well I don't have enough and so the sodium gives its electrons to chlorine. Sodium becomes a 1 plus, chlorine becomes a 1 minus. They share the electrons or they switch the electrons. Now that plus and minus, now they're attracted to each other. They're like little magnets, right? So they're going to be attracted to each other and that's how that bond forms. Now something that's covalent on the other hand, they're going to share the electrons. So something like carbon monoxide, carbon and oxygen, both of them don't have quite enough electrons. So if they share some of them, now they can each have what they need. And so they're kind of stuck together by the fact that they're sharing the electrons. Now we need to know how to differentiate between the two, right? So the way that you decide whether something's ionic or whether something's covalent is for right now, seeing where they are in the periodic table. So if something's on the right hand side and the left hand side are two different things like sodium and chloride, you have a metal and a non-metal, that means that you're going to have an ionic bond. Something that's covalent will be two non-metals, so everything up on that top right hand corner of the periodic table. If it's made up of things over there, that's going to be covalent. Now this is a little bit of an oversimplification and toward the end of the chapter we will start talking about how this isn't quite exactly how things are and that really there's a gradient. Some things are more ionic than others. You can have a purely covalent compound and then there's sort of a gradient all the way up until the ionic ones. And so we'll talk about that later. And then we had the metallic which we aren't talking a lot about now. You'll get into that more in Chem 1b where, but we did want to talk about it a little bit. The main difference here is that all of your electrons are delocalized. It means that one particular electron doesn't really belong to any particular atom. And so there's this kind of sheet of atoms and sheet of electrons all bonded together, which means that you can pass electrons through it. If you put an electron in one side, you can make it come out the other. And so that's why they're conductive. Now this is sort of, you know, we didn't quite get into this before. And there's a little bit of background here that you sort of need to make sure you know from the fundamentals that maybe you glazed over a bit. So all of the processes that consume energy are called endothermic. So think back to your high school or one of the P experiences and think about these words coming back to you again, right? For taking in energy, that's endothermic, bringing in energy. If you're releasing energy, you're throwing energy out, that's exothermic, right? I think exoskeleton on the other side is putting energy out into the environment. We're going to go really far into that in Chem 1b. Here we just need to be able to remember what's endothermic, what's exothermic in the signs that are related to those. So we want to talk about the energy of an ionic bond. So think about before I show you the next slide, which has a lot on it, what I said before about how we form an ionic bond. An electron goes away from one atom. It goes to the other atom and then the plus and the minus are attracted to each other. So we have three kind of steps there. We have an ionization energy, the electron being pulled off an atom. We have an electron affinity, an electron coming into an atom, the other one. And then we have this other energy that comes from the attraction of a plus and a minus. And we have that energy. So three energies. And for it to be favored, we want it to be exothermic. We want it to give off energy. If we have to put a bunch of energy into form something, that's not going to be favorable. So this is sort of the background for the next slide. Okay. So this is what I just said all kind of laid out formally. First this one. We have sodium losing an electron. We have an electron going away from the sodium. Now we have a word for that, right? We talked about it just last lecture. We called that ionization energy. An electron is leaving. So the only real difference here is instead of talking about a trend, I just put a number on it. And I put a plus number on it. So it takes energy to be able to lose an electron. So this endothermic reaction, in quotes. Now C.L. is going to gain an electron. It's going to take an electron in. So that's, in other words, an electron affinity, right? Once again what we talked about last time. And again we'll just kind of give it a number now. So that takes a little bit of energy. And so that becomes 340, or negative 349. So these are just numbers for these specific things. Obviously not something that you would ever be expected to memorize. What I really want you to get out of this is the magnitudes. So 1B will go into more how to calculate energies and how to deal with them on a very mathematical form. I want you to kind of think about the magnitudes here. So we have a plus energy. We have a minus energy. So we have an endothermic. We have an exothermic taking in energy, giving off energy. OK. So once we do that, now we've switched to the electrons, right? We've taken away the electron from one and we've put it on the other. One took in energy, one released energy. Now we have sodium chloride. We have sodium plus and chloride minus attracting each other. And that's favorable too. That's exothermic as well. So we get a negative 787. So you may want to add some little notes here. If you've forgotten what designates an endothermic and an exothermic for signs. Pluses are endothermic, minuses are exothermic. So in order for this to be energetically favorable, you want it to be exothermic, which means do you want a positive or a minus? Remember, endothermic, exothermic. And now we want this whole thing to be exothermic. We want the whole thing to be negative. Well what that means then is we need all three of these numbers to add up to be a negative number. So if we do that here and we add up this, this, and this, you'll see that we get a negative 642 kilojoules per mole. So that's definitely exothermic, right? We have a negative charge. So it's lower. We're losing energy. We're releasing energy into the environment. And so it's going to be stable. So the idea here is that you want your electron affinity and you want this value to be negative. It would be bigger than this one. If this is too large or these are too small negatively, you know, very, it's close to zero, you're not going to be able to form this. In this case we do though, because these added to this make a negative number. So your ionization energy has to be low enough to make this exothermic. And these have to be large enough to make it, excuse me, these have to be negatively large enough. They have to be negative enough to make the whole thing negative. Now this goes into why we tend to only form ionic compounds with metals. Something needs to have a low enough ionization energy to actually be able to do this. If this is too high, it's going to overwhelm these. If this was like 2000, then these wouldn't, it wouldn't be exothermic anymore. These wouldn't be able to get it into the negative line. So because of this, it's typically just your metals, right? Think back to last class. As we were on the left-hand side of the periodic table where all the metals were, those were where our low ionization energies were. So this is where these numbers are kind of low. And so these numbers are able to overwhelm them. They're able to say, no, I don't care if you take a little bit of energy to do, we give off so much energy that it's still energetically favorable. So it's mostly just metals that do this because it's metals that have a low ionization energy. The ones on the right-hand side of the periodic table, carbons, oxygens, fluorines, all of those, those have such a high ionization energy. This number is so high. And go look at it. You know, pull out your table of numbers and see they're so high that it's going to overwhelm it and it's not going to happen. Okay. So now some loose dot symbols. And there's a little more formal introduction of them. They're going to go, they're going to be how we represent our elements when we draw them out. And there's lots of different ways to represent them. Loose dot symbols is one of the simplest and most used. So you put the symbol for the element inside and then you put your dots around the outside. So for iodine, for instance, you have seven valence electrons. You're not going to draw all your electrons, of course, just your valence ones. And then whenever you need to draw a bond, you do a line. You'll sometimes see people represent these, the, instead of a line, they'll put two dots instead and then just squish them close together. Like this. That's fine. It's not wrong. And then here I like to do it to show you how we're switching electrons that when we talk about ionic bonds, you can see it. When you're actually drawing these, it gets a little bit messy. And I would suggest keeping it like this and just putting the bonds. Just remember that every time you draw one of these lines, it's comprising two electrons. Okay. So now let's do some ionic bonds with these. If we have potassium and we have iodide, we are going to take the electron from potassium and give it to iodine, which forms potassium iodide with a plus and a minus charge. Lewis structures get a little boring with ionic compounds. They're a little bit better, a little more entertaining to use with covalent compounds. With magnesium, remember this was our example where we said, well, sometimes ionic compounds aren't going to form a one to one ratio. Sometimes magnesium might need to give an electron to two different atoms in order to make everybody happy. So magnesium says, well, I have two electrons that I need to get rid of and fluoride says, okay, well, I'm only going to need one. So they give one to fluoride and one to the other fluoride, or fluorine forming magnesium fluoride. And the simple way of thinking about that was always, well, we need to balance the charges, right? If we know what our ions are, we know that magnesium likes to be a two plus. And if you forgot why, look at your periodic table, it's in the second column. So you go to your periodic table, you see it's in the second column, you know it needs to be a two plus. You go to fluorine on the periodic table and you say, okay, well, that has seven electrons, so it needs to have a minus charge. So now, if I have an Mg2 plus and an F minus, I need to balance those charges. So that was a faster way of looking at it. Now there was a trick to doing that as well. And it's to crisscross the charges. So if you had aluminum three plus and oxygen two minus, you're going to take the one charge and bring it down, the other charge and bring it down. And then that'll give you a neutral compound because two aluminum is two times six, or two times three is six. And then two oxygens, so, excuse me, three oxygens, two minus, so that's a negative six. We have a plus six and a minus six, so we have a neutral charge. And whenever you're forming ionic compounds, you need it to be a neutral charge. Now there is a caveat to that trick, right? And that caveat was, well, what if you have the exact same charge of a higher number? Really at that point, your best bet is if you have something that's a two plus and a two minus, just know to put them together. Just know that it's a one to one thing, or if it's a three plus and a three minus, it's one to one. But if you don't remember that and you don't pick up on it right away and you do the crisscross trick anyways, and you end up with MG202, you have to reduce that down. Now remember, only four ionic compounds do you reduce this down. For this, you have to make it MGO. You have to reduce it down. That's not true of covalent compounds, right? Covalent compounds might be N2H4, and you wouldn't reduce that down. Ionic compounds always have to be lowest whole number ratio. Covalent compounds don't. Okay, so now let's talk about some covalent compounds. So this is when you're going to be sharing electrons. So for instance, with chlorine, you can go ahead and you can take, and you can share each one, this will share one of its electrons, this will share one of its electrons, and you'll form a single bond. If each element is sharing one electron, that's a single bond. And you can draw it like this, or this, I highly suggest doing it this way though. And so that's how we're going to be representing all our covalent compounds. So if I ask you, draw this, that's how you want to draw it. Here's a little bit more complicated example, where now sulfur is needing to share electrons with multiple atoms in order to get what it wants. So it's going to share two, one oxygen, two with another oxygen, to form sulfur dioxide. And so you can go ahead and whenever you are sharing four electrons total, two from each element, that's a double bond. And so you can write it like this, which gets a little messy, so I much prefer that you write it like this and actually form the double bonds. Now for these simpler ones, that's fine to just kind of be able to write them out, but all those structures get pretty complicated. And so let's go through some general rules. And this is a good slide to keep with you as we go through all of the examples, because we have a whole bunch of examples to do. And we'll follow through these guidelines as we go. Step one, and this is the step that when you start getting good at them, everybody skips and it's a terrible idea, it's where you start making your mistakes. And you'll see I'm kind of lazy about it sometimes too, and I'll have to go back and do it at the end just to make sure I did it right. So add up all your valence electrons. In the beginning, this will help get you started. In the end, it'll check for your mistakes, so it's very important to do. Make a skeletal structure. So what I mean by that is you put your central atom inside, and then you line, you put everything else around it, and you just draw single bonds. You connect them together, you kind of play connect the atoms. And you don't worry too much about lone pairs and in double bonds and all of that, that just gets you started. So you know every time that you draw that line, that's going to be two electrons. Now you're going to have some electrons left over. So figure out how many electrons you have left over. Start with the number that you started with. Subtract out how many you had to use to form your single bonds, and now you know you have a certain amount of electrons left over. Put those around where you need to put them around. Now octet rule, right? We're trying to get eight. So if you can't get eight, then you're going to have to do something else. So at this point, you do want to kind of be thinking, certain things don't follow the octet rule, and we're going to talk about what those are. These are some common exceptions, but there's going to be some other ones too, and we'll talk about why they are in a minute. But you're going to want to check for those so that you know, hey, this is an atom that typically does something that's a little bit funny. So let's go ahead and make sure that, you know, I'm at least aware that it doesn't have to follow the rules. Maybe it will, maybe it won't. Now if you can't distribute the electrons, if you try to put them around and you say, I have two atoms that I don't have enough electrons, where do I put them? That means that you need to start forming more bonds. Start putting double bonds, start putting triple bonds. A little note once you get good at them, if you're short by two electrons, that's going to, you're going to have to add one more bond. If you're short by four electrons, you're going to have to add two bonds. So that's sort of a little trick to, when you start getting good at them and you want to move faster. However many electrons you have left out that you're missing, you need six, you're probably going to need to form three bonds. Now this next part, and again we don't know, we don't know what all of this means yet. We're going to go through each of these in more detail as we do the examples. But this is just a good slide to keep with you while you're doing these. So you have, you don't look at them and say, what do I do? I'm stuck. I don't even know how to start this problem. This is your guidelines. They aren't hard rules, they change a little bit. Once you're done with all of that and you theoretically have everything with octets, you theoretically have all of your atoms happy, now you check your formal charges. And we'll talk about exactly what that is. We want them small. So when we get there, we want to minimize them. Now the only way to really do these is to go through and do a ton of examples. And so that's what we're going to do. We're going to do a ton of examples at which point we'll talk about formal charges. And then we're going to do some one that break octet rules too. Okay. So the first, we'll start with one that isn't too bad. We'll start with nitrogen. So our first step, and this is, you know, our error checking and making sure that we have somewhere to start from is to count our valence electrons. So you pull out your periodic table and you say, well, how many valence electrons does nitrogen have? Well, each one of them has five. So we have two nitrogens, each one with five valence electrons, and so we have ten electrons. We don't have a charge, so we don't have to add or subtract any electrons. And that's important to make sure you remember too. We'll do some of those later as well. So our first, our next step then is to draw our skeletal structure. So we go through and we do that. And now we need to add in our electrons so that we satisfy the octet rule. Well, let's say we try to do this with nitrogen and we say, okay, well, I used up two electrons making that bond. So I have eight left. So one, two, seven, eight. Okay. Well, I'm short electrons, right? I can't give everything an octet. Nitrogen would need two more, or four more here, and I can't do that. So that means I'm not sharing enough electrons. I'm short four electrons, which probably means I'm going to need to do two bonds as sort of a good rule in your head to think about. So there's only one place to put those bonds, so I'm just going to put them there. Which means now I need, I have four electrons left because I have two, four, six. So I've used up six and I started with ten. So four electrons left. So now I'm set. Two, four, six, eight, ten. Always go back at the end and double check and make sure your electrons add up. And this is why it's important to do this in the beginning, even once you start getting good at them and you start just kind of whipping them down. At the very end, if you didn't do this, go back and do it real quick and double check it. Make sure that you've done it right. Okay. Next one. Now we've made it a little bit more complicated. So now we have to figure out which one's going to be our central atom. Now, good rules of thumb for this, and there can be more than one central atom as well. The least electronegative element, that's going to be your central atom. And then everything else kind of goes around it. Hydrogen, how many bonds is hydrogen allowed to form? So all it has filled in its ground state is it has one S electron, one S electron. And so it can only ever form one bond. Hydrogen's not allowed to form any more than one bond. And so because of that, you're going to, that's never going to be a central element. Oxygen can, but it's usually only if it's with things that only form one bond. So with all of those rules that puts carbon in our center. Okay, so let's add up our electrons. So we have four from carbon. We have two from hydrogen, well, let's write that this way. We have two hydrogen atoms, and each one of them has one electron. And from oxygen we have six. So we have a total of 12 electrons. So we put carbon in our center, and we put everything else around it. Now we try to distribute our electrons. So hydrogen doesn't form an octet, right? Hydrogen is the duet rule. It only has two electrons. So we have two electrons here, and two electrons here. So that leaves us with, and two electrons here. So that leaves us with six electrons to work with. So let's go through and see if we can do this. Well we have two, four, six, eight, ten, twelve. So we're set for electrons, everything has an octet. So if you just walk away at this point, you're not going to get your points. And the reason for that we'll talk about in just a minute. But let's for right now see if there's another way we can do it. And after we talk about formal charges, we'll talk about why one is better than the other. So the other way that you can do this is by doing this. Forming a double bond here, and putting your two oxygens here. Now let's look at all of our octets and see what we have. Here we have two, four, six, eight, we're good. Two, four, six, we're not good. And then our hydrogens are good. Here we have two and two, so we're good. Our oxygen has two, four, six, eight, so we're good. And now let's look at our carbon. Two, four, six, eight, so we're good with carbon. Here we didn't have enough electrons to do it. Here we do. So double bonds and triple bonds are a way to give ourselves extra electrons because we're sharing them. So this one's wrong. And this one is the one that we want to be using. Okay, next example. And when you're doing these, you'll notice I've given myself an entire sheet of paper to do these on. That's on purpose. We're going to be using these for the entire rest of the quarter, pretty much, all the way up until at least gas loss. So make sure that you give yourself just as much room as I did, and we're going to be adding to them. When we get to geometries, we're going to go through and we're going to put all of our geometries on. When we get to hybridization, we're going to go through and put all of our hybridization on. We'll do bond angles. We'll do polarity. We're going to just spend a lot of time with these sheets of paper. So one per sheet is about right, any more than that, and you start running out of room near the end of the quarter. Okay, now the next one. And this one gets a little complicated because now we have two central atoms. Now luckily, people usually write these in a way that makes it a little bit easy to see what's going on. So you can kind of see here that we have BH3 and H3, and that's written that way on purpose. So you can sort of see that you have like a BH3 group bonded to an NH3 group. So first, count our electrons. We go to boron. We have three. We go to nitrogen, and we have five. We go to hydrogen. There's six hydrogens, and each one has one electron. So we have 14 electrons to work with. So now our skeletal structure. We can kind of see that boron and nitrogen are the only ones who can form multiple bonds. So they're going to need to be bonded to each other, as well as all the hydrogens. And so at this point, we say, OK, well, how many electrons do we have? 2, 4, 6, 8, 10, 12, 14. So we're good. Does everything have its octet filled, or for hydrogen, it's duet? Well, all the hydrogens are set. And then boron has 2, 4, 6, 8, and nitrogen has 2, 4, 6, 8. And so we're set there. Now later on, we're going to talk more about exceptions. And we'll see that boron's actually one that a lot of times will have less than eight electrons. A lot of times, it'll just have six, and it's fine with that. But here's an example where it actually does what most atoms do, and it has its eight. OK, next one. Now we need to do an ion. So you have some practice doing ions. So we have XCF4, 4 plus. So conopolar electrons. So xenon has eight. It's a noble gas. And then fluorine has seven each. So we have four fluorines, and each one of them has seven. So that would be 36. But it's a four plus charge, which means that we need to take away four electrons to give us 32 electrons. So it's a plus charge. We have less electrons than the neutral atom, and so we minus the four electrons. So we have 32 electrons to work with. So we put xenon in our center. Fluorine never allowed to break the octet rule, and we'll talk more about that in a minute as well. So fluorine is ever only going to have one bond. So whenever you see a hydrogen or a fluorine, only one bond, it can't be a central atom. So when we go through and we draw these out, we know fluorine can't have more than one bond since I just told you. So really all we can do at this point is add in our lone pairs and hope everything works. If not, we'll have to add some things to xenon. So if you count all of these, 2, 4, 6, 8, all the way around, what you see is that we do in fact have 32 electrons. So we're done. We don't have to do anything else with this one. OK. Now our last and most complicated one for right now. This one we won't actually even be able to finish until we talk about some other things. OK. So N2O. So this one's kind of not as simple as the other ones, right? And we'll see why in a minute. So first of all, most electronegative element here is oxygen. So nitrogen's going to be in our center, one of our nitrogens. So we can set up our central atom as being nitrogen. And we can put a nitrogen on each side, or excuse me, a nitrogen on one side, an oxygen on the other one. And then we can draw our skeletal structure. Now if we try to give everything an octet here, first we're going to need to count our electrons. So you'll see that we have five electrons from each nitrogen and we have two nitrogens. We have six electrons from our oxygen leaving us with 16 electrons. If you try to make your octet here with just going and filling in your elements or your electrons, and we just kind of fill in, you notice we can't do it. We don't have enough here, right? We would need four more electrons. So if we're short four electrons, how many bonds do we need to form? We need to form two bonds. Okay. Well, where do we put those bonds? Now we have some options. When we had to form two bonds with that N2 group, we didn't really have any options. So we just sort of put them in between the two nitrogens. Here we have options. So I have two bonds that I need to form. So I could just put one bond on each group. I could put both bonds to the oxygen, or I could put both bonds to the nitrogen. And let's go through and fill in our octets at this point. So we have two here, so we need two more. This nitrogen set, two, four, six, eight. This oxygen needs four more. We count up our electrons, two, four, six, eight, ten, twelve, fourteen, sixteen, we're set. We move down to this one. Now this nitrogen needs six. This nitrogen again is all set. And this oxygen is going to need one. This nitrogen needs two. And this oxygen needs three. We go through and count, two, four, six, eight, ten, twelve, fourteen, sixteen, we're good. Two, four, six, eight, ten, twelve, fourteen, sixteen, we're good. Okay. Well, which one is it? So all of them follow all of the octet rules. All of them have the least electronegative element in the center. So which one do we do? This is where that very last step on the slide comes in. We don't really know how to do that yet. So let's learn how to do that and then we'll come back. So there's an equation for this. And once you start getting good at it, you're going to want to do it more sort of internally and be able to quickly check all of your formal charges. You don't necessarily always want to be filling into this equation. But for the moment, it's a good thing to start with. So your valence electrons. So you start there. So for nitrogen, it's five. Oxygen, six. If you ever forget, you just go to your periodic table and count over until you get to the element where you're at. Now that's your starting point. Then you subtract the number of electrons that you're, the atom that you're talking about in your, whatever you're looking at. What I kind of always think of it is which ones they actually own. So which ones are sort of their electrons? So the way that that's calculated is the number of electrons, the number of electrons that are actually in, that it has all by itself. It's lone pairs. So you're subtracting away all the electrons that are inside the lone pairs. And then half of the bonding electrons. So if you have a bond, one of those, one of those electrons belongs to one element and one of those electrons belongs to the other element. So for accounting purposes, you always think of half of the bonded electrons. And we'll do some examples that'll make that make more sense. So let's go through and assign formal charges to all of the Lewis structures that we did. And then when we get done with them, we'll be able to actually go through that last harder example that we did and decide which one it actually is, which Lewis structure is actually correct. So just sort of starting over on our Lewis structures. Again you'll want to keep them handy and do this quite a bit throughout this next little bit of the quarter. Okay. So our first one. So what is our formal charges on each nitrogen? So we go and we say, well, and formal charges are going to be atom specific. So it's for each atom in the molecule. Now in this case, they're going to be the same, but we can still calculate them for both. So how many electrons does nitrogen have, valence electrons does nitrogen have? You go to the periodic table, you come over and you see that nitrogen has five. Now you have to decide how many are its own, how many does it have in this structure. And you're going to subtract those away. So anything in a lone pair is its own. So it has two from the lone pairs. From the bonds, it only gets one from each bond. So one, two, three. So you have five valence electrons minus five is zero. So this nitrogen's formal charge is zero. Now if you go to the other nitrogen, the exact same thing happens. It has five valence electrons. It has two electrons that are in lone pairs and it gets one from each bond. And so that nitrogen's going to be plus zero as well. Next one. So now we can go through and we can talk about each one individually. So let's start with the oxygen. So how many valence electrons does oxygen have? It has six. And we're going to subtract away anything that are in lone pairs. So one, two, three, four. And any of the bonding electrons but only one from each bond. One, two. So we have six minus six. So the oxygen's charge is zero. Carbon. We go to the periodic table. We say how many valence electrons does it have? And it has four. So we subtract away anything. In this case, there's no lone pairs so we don't have to count any of those. And it gets one from each bond. So one, two, three, four. So we have four electrons there. So its formal charge is zero. Now we need to talk about hydrogen. Hydrogen has one valence electron and gets one from each bond. So its formal charge is zero as well. And that would be true of the other hydrogen as well. So those are all our formal charges for that one. Next one. So now this one when we go through, it's going to be a little bit different. We'll start actually getting some formal charges. So let's start with the hydrogens. Hydrogen gets one valence electron. And it's the same for each of these hydrogens. It'll get one from the bond. So every one of the hydrogens on here follow the same rules and so they're all going to be plus zero. Now let's talk about the boron and the nitrogen. For boron, how many valence electrons does it have? So you go to the periodic table and you see that it has three. Now if you come here and we count electrons, no lone pairs, so don't worry about those, gets one from each bond. One, two, three, four. So boron has a minus one charge. Now nitrogen, if you go to its period, it on the periodic table, you'll see that it has five valence electrons. And you count around one, two, three, four. And so you have a plus one charge on nitrogen. Now the xenon tetrafluoride. So for each of the fluorines, because again they're all going to follow this numerals, so we're only going to do it once, they get seven electrons from the periodic table, valence electrons. Now here each one of them has two, four, six that are in lone pairs. So those are the fluorine lone pairs. It's not sharing them, it's not doing anything with them. So it owns all of those. Now for the bonds, it just gets one. So we have seven valence electrons, minus six electrons that are in lone pairs and one electron that is in the bond. So each of the fluorines are going to have a zero charge. Now we move on to the xenon. It's a noble gas, so it's in that eighth group. Has eight valence electrons. And we subtract one from each bond. One, two, three, four. So we subtract four electrons. So it has a plus four charge. So as we're going through these, you should notice that we're counting electrons very differently depending on what we're doing. When we were counting for octets, we counted everything, right? For xenon we went through and said two, four, six, eight. Xenon has an octet. Fluorine has two, four, six, eight. Fluorine has an octet. That's very different from when we do formal charges. The bonds only count as one for each atom. So fluorine, we only got to count that as one electron. For the sake of octets, you counted as two. Okay, so now we can move on to our last one, which was a bit more complicated. And we weren't actually able to finish. Okay, so we had this. And we said, well, I don't know which of these three to do. Now we can actually go through and we can figure this out. So let's do that. Let's go through and put formal charges on everything. And I'm going to do this a little bit quicker this time and do it a little bit more like how you should be doing it on an exam and how after maybe about a week of practice and a week of homework you should be doing it. So we look at oxygen and we say, how many valence electrons does oxygen have? Well, it has six of them. How many does it have in lone pairs? Four. Okay, so we're down to two electrons at that point and then two here. So oxygen has six valence electrons and here it has one, two, three, four, five, six. So six minus six is zero. Let's actually switch colors just to make it a little bit more obvious. Plus zero. Now nitrogen. Nitrogen has five valence electrons and here we have one, two, three, four. So we have five minus four is equal to one and it's a positive one. Now here nitrogen has four valence electrons and here it has one, two, three, four, five, six. So five minus six is minus one. So again, when you're practicing at home, do it the way that in the beginning, do it the way that I've been doing it up until now. When you get a little bit on a roll with it, this is kind of the fast way to do it. This is how you want to be able to do it for the sake of, you know, checking your work quickly without having to waste too much time. Okay, let's do the next one. Here we have oxygen normally has six valence electrons. Here it has one, two, three, four, five. So six minus five is a plus one. Over here, nitrogen normally has five and here it has one, two, three, four. Five minus four is plus one. Here nitrogen normally again has five but here it has one, two, three, four, five, six, seven. Five minus seven being minus two. Okay. So just going through, putting formal charges on everything and then this is going to help us decide something when we're all done. Down here, oxygen normally has six. Now we have one, two, three, four, five, six, seven. We have minus one. Here nitrogen normally has five. One, two, three, four. So five minus four is one. Over here, nitrogen has five valence electrons and here it has one, two, three, four, five. Five minus five being zero. Now as a good check, we need to go, this is something you kind of want to do in your head for everything. You don't necessarily need to write them down. If I ask you to write down a Lewis structure, you don't have to put informal charges. If you have something that's an ion, actually like this, if you have something like an ion, it's not actually correct unless you put the charge somewhere but you don't have to put the charge on Xenon. You can just put brackets and put the charge and that's fine. That's just as correct. Something like this that doesn't have a charge, you don't actually have to put the formal charges unless I ask for it. You should do it in your head anyways though. Now we have to decide which one's correct. So if we look at these, a good check first is to make sure that all of our formal charges add up to whatever the charge is here. So charge is zero. So this one adds up to be zero. This one adds up to be zero. So we've done our formal charges correct. We've done our Lewis structures correct. Now we have to decide which set of formal charges is the best set of formal charges. And so if we go through and we look, we can say, well, this one has a negative 2, a 1, and a 1. This one has a negative 1, 1, 0, negative 1, 1, 0. This one has some really high formal charges and there's no reason for that. 2, 1, 1 is acceptable if there's no other option. You will have formal charges of plus 4 sometimes. We saw that with Z non. But if there's another option, you don't want to do that. So we're going to get rid of that one right away. That one's not going to be right. So now we have to decide between these two. So they're basically the same. Our central nitrogen has the same. But now here we have a negative one on the nitrogen. Here we have a negative one on the oxygen. So we have to figure out which one's better. So for this, we need to kind of backtrack to what we talked about earlier today or last class, where we were talking about electronegativities. So when we're talking about electronegativities, we're talking about the amount that an atom is willing to pull the electron density toward it inside of a bond. So this is kind of exact in this situation. And something like oxygen is more electronegative than something like nitrogen. So if the electron density is going to be shifted one way or the other, it's going to be shifted so that the more electronegative element, the one that's pulling the electron density toward it, is going to be the negative one. Because well, it's more electronegativity. It has more electronegativity. It's pulling those electrons toward it more. It's going to take the negative charge, not nitrogen. So this one is going to be correct because oxygen is more electronegativity, or more electronegative. So you always, if it's possible, and keep this in mind, if it's possible, you want to put your negative charge on your most electronegative element. That's not always possible. For instance, this one again. In this one, we had a plus. Oh, I guess we didn't have a negative charge. Never mind. So draw out CO for fun sometime. Draw out carbon monoxide. And you'll see that you're forced to do the opposite of this. You don't have a choice. And that's fine. If you have a choice, you don't have a choice. But in something like this, we do. We can choose to put it on the nitrogen by drawing it that way, or we can choose to put it on the oxygen by drawing it that way. So this is the more correct way. OK. So now, moving on a bit. More in detail about breaking the octet rule. I've been sort of alluding to this and talking about it, but now we're going to go into it in a little bit more detail and talk about exactly when we can and can't do this. So your second period elements, they're not going to do this. So that second row of the periodic table, well, and of course, hydrogen and helium as well. Hydrogen, helium, and then lithium up through neon. They're not going to do it. Do not ever have 10 electrons off fluorine or an oxygen or nitrogen. You will just get it wrong. There will be no partial credit for that. You'll just have it wrong. Now third period and lower, you can do that and it's fine. That's not a big deal. There's some that are going to do it more than others, but not all of them. Now the third period element can do it where the second one can't because of the fact that they have some d orbitals sitting right there. So the d orbitals are right there. They're able to be used for bonding. You can use them for bonding. You can use them for hybridization, which we'll talk about later in the next chapter. And so anything below, third row and below can do it. Second period ones do not do it ever. Some of them are going to have a lot less than eight and they don't have to, but they can. Theoretically, anything could have less than eight electrons, but that's not necessarily normal. For beryllium, aluminum, and boron, it's pretty normal. If you see one of these and you can't fill the octet, don't worry about it. Just move on. Anything like carbon and oxygen, you're going to probably fill the octet. Sell them going to break the octet rule. It's not something that you're going to ever see in this class. Where these elements, they do it all the time, so get used to these elements breaking the octet rule on the low side and having less. As far as having more, that falls into the category of sort of a lot of these that I talked about. There's a lot of phosphorus xenon, chlorine does it a fair amount. Anything third row and below can do it, but these are sort of your common ones that actually do it all the time. We'll start and do some examples of this now and then do some more next time. Okay, so we'll start with SF6. Our central one is sulfur, since we can't really do anything else, we can't put fluorine in the center. When we go through and do this, sulfur here, and we have six things, so we just spread them all out and put fluorines on all of them. Because it's a Lewis structure, we need to go through and we need to add in all of our lone pairs. We count up how many electrons we started with just to make sure that everything is still okay. We have six from the sulfur, seven from each fluorine, and we have six of them. We have 48 electrons. If we go through and count all of these up, it adds up to be 48, and so we're all set. Sulfur is an example where it's okay with having six bonds. Usually with sulfur, you want to kind of get in the habit of seeing, well, it'll either have two bonds and fill its octet, or it'll have four bonds, or excuse me, six bonds in this sort of configuration. So it'll either act like oxygen, which is right below on the periodic table, or it'll form six. Okay, next one. So this is kind of similar to one that we've done before, only now it's not an ion. Now it's XCF4 without any 4 plus. So we count up our electrons. We had eight from xenon. We have seven from each fluorine, and we have four of them. So we have 36 electrons. It's not a charge, so we don't have to worry about adding or subtracting any. Now when we go to draw this, this is a good example of one where you can get yourself into trouble if you don't count the electrons first. So up until now, a lot of them have been, once you get really good at them, you may say, well, I don't need to count electrons. I know what it looks like. This is one where you may get yourself into trouble if you don't. So if we go through and we kind of set up our skeletal structure, and we know fluorine only gets to have one bond, so we're just going to fill those in. At this point, you might be tempted to say, okay, everything has an octet. I'm done. And walk away. In which case, you'd get it wrong. And then all the other questions following it that we're going to learn how to do later, like hybridization and geometries and bond angles and things like that. So we look at this and we say, okay, how many electrons have we used up? If we count them all, we've only used up 32. So we have two more, or four more electrons that we have to deal with. So we can't put those electrons on fluorine, because fluorine only gets to have eight. Xenon can have more if need be. So we're going to give them the xenon. So that gives xenon more, it has two, four, six, eight, ten, twelve electrons. So that's okay. So that's one where we're giving xenon twelve electrons. And so we're breaking the octet rule there. We can't break it with fluorine, but xenon can break it. Okay, so now let's do BH3. So we add up our electrons. We say, well, we have three from boron. We have three from each hydrogen, one from each hydrogen. So we have six electrons total. We go to draw this. And we have this. And at this point, there's nothing else we can do. We only have six electrons to work with. Well, if we only have six electrons total, there's absolutely no way to give boron eight. We can't just make two electrons out of malware. So we're done. That's all there is to it. We can't really do anything else. Now as good outside practice on this, go through and calculate all your formal charges. We haven't done any formal charges yet, because it hasn't really come up where we've had an option where we needed to decide between things. But it's a good practice to go through and make sure that you can do it and get in the habit of doing it. And so there's lots of places here for you to practice it. Now a more complicated one. And these ones tend to get people a lot. For all of these polyatomic ions where we turn them into acids, so they're called oxoacids. So in this case, this one would be called what? Sulfuric acid. So, or phosphoric acid or carbonic acid. Any of these, you want to get in the habit of knowing sort of the general way that these look like. Because otherwise you might be tempted to take sulfur and bond one to each oxygen and to the hydrogens. You don't want to do that. Whenever you see these, you need to know that these hydrogens are actually bonded to the oxygen. So when we go to do this, we're going to put sulfur in the center. We're going to put four oxygens around. And then the hydrogens go on the oxygens. And that should just, you should know that as part of your skeletal structure from now on. And if you come to me later and say, well, how was I supposed to know that? I'm going to point you to this slide. So this is how you know it. And this goes for sulfuric acid, carbonic acid, any of those oxoacids. So now let's go through. We can fill our octet in here because there's not really much else we're going to probably do with oxygen. And at this point in practicing, you may get in the habit of saying, okay, well, I know oxygen pretty much always has two, two bonds. So I'm just going to give it two bonds in its lone pairs. Now if we add in our electrons and we try to add 2,4,6 here and 2,4,6 here, we'll see that we can't do it. And so what we need to do, we'll see that we're short for electrons. So go home, practice it, do it. We'll see we're short for electrons, so we need two bonds. And so we're going to put our two bonds here, or excuse me, our two lone pairs here and our two lone pairs here. And that's where we'll finish. Now you'll notice I forgot to add in all of my electrons and make sure that we have the right amount. So let's do that real quickly, just as a double check. If you forget to do it in the beginning, especially when you get good at them, you may get out of the habit and you may forget. Just do it at the end and make sure that you're right. And when you go through and do this, you'll see that, hey, we do have 32 electrons and we are all set. So at this point, we've gone through and we've done a bunch of ones that break the octet rule. We have a bunch more that we're going to do because Lewis structures are really, really important. You will not be able to do all of the rest of the material for this chapter and chapter 3 and chapter 5 if you can't do these. So we're going to do a bunch more examples next time. So we'll just kind of go through them at that point and keep going through and get into all of that. So we'll end there for today though and start up on the rest of them next time.
 Okay, so today we're going to start chapter two.  And chapter two is going to start off with a lot of review, mostly because I thought  you needed to know a lot of this before we really even got started with class and so  I talked about it earlier.  But we should go over it again and just make sure it's cemented in there before we start  talking about the more complicated aspects of chapter two.  So we're going to go through it and we're going to remind you about ionic and covalent  bonding so that you remember exactly what that is and what the differences are because  it becomes really important here.  All your naming is still really important here and knowing how to identify between each  one is very important.  Now we're going to introduce Lewis Dot symbols in a little bit more of a formal nature.  Before we kind of threw them at you and said this is some really simple ones, this is how  we represent them, now we're going to go through and we're going to do a bunch of very complicated  ones that go way more into depth on what exactly is going on which will lead us into trying  to talk about how molecules look like and doing geometries and things like that.  Then we're going to do some more Lewis Dot structures that are even more complicated  and more Lewis Dot structures.  Really chapter two is kind of all about Lewis Dot structures and getting good at them.  Once we've really got that down packed, we're going to go through and do some bond strength  and length comparisons so that we can see how they're related to each other.  Again some review from last time.  Bonding is all about trying to figure out how to complete octets or make things more  stable than they were.  Now you can do that by in something like this case sharing electrons.  If two different electrons have six valence electrons, they can go share them.  They can take two of their electrons, share them with each other and then form some bonds  there so that they each can have eight.  There's our three main types of bonding.  We have our ionic bonding, our covalent bonding and our metallic bonding.  Remember these are important for knowing how to name them because we use different names  for our ionic than we do covalent.  They're very different from each other.  So in ionic you're trading an electron.  Sodium says, hey I have an extra electron and chlorine says, well I don't have enough  and so the sodium gives its electrons to chlorine.  Sodium becomes a 1 plus, chlorine becomes a 1 minus.  They share the electrons or they switch the electrons.  Now that plus and minus, now they're attracted to each other.  They're like little magnets, right?  So they're going to be attracted to each other and that's how that bond forms.  Now something that's covalent on the other hand, they're going to share the electrons.  So something like carbon monoxide, carbon and oxygen, both of them don't have quite  enough electrons.  So if they share some of them, now they can each have what they need.  And so they're kind of stuck together by the fact that they're sharing the electrons.  Now we need to know how to differentiate between the two, right?  So the way that you decide whether something's ionic or whether something's covalent is  for right now, seeing where they are in the periodic table.  So if something's on the right hand side and the left hand side are two different things  like sodium and chloride, you have a metal and a non-metal, that means that you're going  to have an ionic bond.  Something that's covalent will be two non-metals, so everything up on that top right hand corner  of the periodic table.  If it's made up of things over there, that's going to be covalent.  Now this is a little bit of an oversimplification and toward the end of the chapter we will  start talking about how this isn't quite exactly how things are and that really there's a gradient.  Some things are more ionic than others.  You can have a purely covalent compound and then there's sort of a gradient all the way  up until the ionic ones.  And so we'll talk about that later.  And then we had the metallic which we aren't talking a lot about now.  You'll get into that more in Chem 1b where, but we did want to talk about it a little  bit.  The main difference here is that all of your electrons are delocalized.  It means that one particular electron doesn't really belong to any particular atom.  And so there's this kind of sheet of atoms and sheet of electrons all bonded together,  which means that you can pass electrons through it.  If you put an electron in one side, you can make it come out the other.  And so that's why they're conductive.  Now this is sort of, you know, we didn't quite get into this before.  And there's a little bit of background here that you sort of need to make sure you know  from the fundamentals that maybe you glazed over a bit.  So all of the processes that consume energy are called endothermic.  So think back to your high school or one of the P experiences and think about these words  coming back to you again, right?  For taking in energy, that's endothermic, bringing in energy.  If you're releasing energy, you're throwing energy out, that's exothermic, right?  I think exoskeleton on the other side is putting energy out into the environment.  We're going to go really far into that in Chem 1b.  Here we just need to be able to remember what's endothermic, what's exothermic in the signs  that are related to those.  So we want to talk about the energy of an ionic bond.  So think about before I show you the next slide, which has a lot on it, what I said  before about how we form an ionic bond.  An electron goes away from one atom.  It goes to the other atom and then the plus and the minus are attracted to each other.  So we have three kind of steps there.  We have an ionization energy, the electron being pulled off an atom.  We have an electron affinity, an electron coming into an atom, the other one.  And then we have this other energy that comes from the attraction of a plus and a minus.  And we have that energy.  So three energies.  And for it to be favored, we want it to be exothermic.  We want it to give off energy.  If we have to put a bunch of energy into form something, that's not going to be favorable.  So this is sort of the background for the next slide.  Okay.  So this is what I just said all kind of laid out formally.  First this one.  We have sodium losing an electron.  We have an electron going away from the sodium.  Now we have a word for that, right?  We talked about it just last lecture.  We called that ionization energy.  An electron is leaving.  So the only real difference here is instead of talking about a trend, I just put a number  on it.  And I put a plus number on it.  So it takes energy to be able to lose an electron.  So this endothermic reaction, in quotes.  Now C.L. is going to gain an electron.  It's going to take an electron in.  So that's, in other words, an electron affinity, right?  Once again what we talked about last time.  And again we'll just kind of give it a number now.  So that takes a little bit of energy.  And so that becomes 340, or negative 349.  So these are just numbers for these specific things.  Obviously not something that you would ever be expected to memorize.  What I really want you to get out of this is the magnitudes.  So 1B will go into more how to calculate energies and how to deal with them on a very mathematical  form.  I want you to kind of think about the magnitudes here.  So we have a plus energy.  We have a minus energy.  So we have an endothermic.  We have an exothermic taking in energy, giving off energy.  OK.  So once we do that, now we've switched to the electrons, right?  We've taken away the electron from one and we've put it on the other.  One took in energy, one released energy.  Now we have sodium chloride.  We have sodium plus and chloride minus attracting each other.  And that's favorable too.  That's exothermic as well.  So we get a negative 787.  So you may want to add some little notes here.  If you've forgotten what designates an endothermic and an exothermic for signs.  Pluses are endothermic, minuses are exothermic.  So in order for this to be energetically favorable, you want it to be exothermic, which means  do you want a positive or a minus?  Remember, endothermic, exothermic.  And now we want this whole thing to be exothermic.  We want the whole thing to be negative.  Well what that means then is we need all three of these numbers to add up to be a negative  number.  So if we do that here and we add up this, this, and this, you'll see that we get a negative  642 kilojoules per mole.  So that's definitely exothermic, right?  We have a negative charge.  So it's lower.  We're losing energy.  We're releasing energy into the environment.  And so it's going to be stable.  So the idea here is that you want your electron affinity and you want this value to be negative.  It would be bigger than this one.  If this is too large or these are too small negatively, you know, very, it's close to  zero, you're not going to be able to form this.  In this case we do though, because these added to this make a negative number.  So your ionization energy has to be low enough to make this exothermic.  And these have to be large enough to make it, excuse me, these have to be negatively  large enough.  They have to be negative enough to make the whole thing negative.  Now this goes into why we tend to only form ionic compounds with metals.  Something needs to have a low enough ionization energy to actually be able to do this.  If this is too high, it's going to overwhelm these.  If this was like 2000, then these wouldn't, it wouldn't be exothermic anymore.  These wouldn't be able to get it into the negative line.  So because of this, it's typically just your metals, right?  Think back to last class.  As we were on the left-hand side of the periodic table where all the metals were, those were  where our low ionization energies were.  So this is where these numbers are kind of low.  And so these numbers are able to overwhelm them.  They're able to say, no, I don't care if you take a little bit of energy to do, we give  off so much energy that it's still energetically favorable.  So it's mostly just metals that do this because it's metals that have a low ionization energy.  The ones on the right-hand side of the periodic table, carbons, oxygens, fluorines, all of  those, those have such a high ionization energy.  This number is so high.  And go look at it.  You know, pull out your table of numbers and see they're so high that it's going to overwhelm  it and it's not going to happen.  Okay.  So now some loose dot symbols.  And there's a little more formal introduction of them.  They're going to go, they're going to be how we represent our elements when we draw  them out.  And there's lots of different ways to represent them.  Loose dot symbols is one of the simplest and most used.  So you put the symbol for the element inside and then you put your dots around the outside.  So for iodine, for instance, you have seven valence electrons.  You're not going to draw all your electrons, of course, just your valence ones.  And then whenever you need to draw a bond, you do a line.  You'll sometimes see people represent these, the, instead of a line, they'll put two dots  instead and then just squish them close together.  Like this.  That's fine.  It's not wrong.  And then here I like to do it to show you how we're switching electrons that when we  talk about ionic bonds, you can see it.  When you're actually drawing these, it gets a little bit messy.  And I would suggest keeping it like this and just putting the bonds.  Just remember that every time you draw one of these lines, it's comprising two electrons.  Okay.  So now let's do some ionic bonds with these.  If we have potassium and we have iodide, we are going to take the electron from potassium  and give it to iodine, which forms potassium iodide with a plus and a minus charge.  Lewis structures get a little boring with ionic compounds.  They're a little bit better, a little more entertaining to use with covalent compounds.  With magnesium, remember this was our example where we said, well, sometimes ionic compounds  aren't going to form a one to one ratio.  Sometimes magnesium might need to give an electron to two different atoms in order to  make everybody happy.  So magnesium says, well, I have two electrons that I need to get rid of and fluoride says,  okay, well, I'm only going to need one.  So they give one to fluoride and one to the other fluoride, or fluorine forming magnesium  fluoride.  And the simple way of thinking about that was always, well, we need to balance the charges,  right?  If we know what our ions are, we know that magnesium likes to be a two plus.  And if you forgot why, look at your periodic table, it's in the second column.  So you go to your periodic table, you see it's in the second column, you know it needs to  be a two plus.  You go to fluorine on the periodic table and you say, okay, well, that has seven electrons,  so it needs to have a minus charge.  So now, if I have an Mg2 plus and an F minus, I need to balance those charges.  So that was a faster way of looking at it.  Now there was a trick to doing that as well.  And it's to crisscross the charges.  So if you had aluminum three plus and oxygen two minus, you're going to take the one charge  and bring it down, the other charge and bring it down.  And then that'll give you a neutral compound because two aluminum is two times six, or  two times three is six.  And then two oxygens, so, excuse me, three oxygens, two minus, so that's a negative six.  We have a plus six and a minus six, so we have a neutral charge.  And whenever you're forming ionic compounds, you need it to be a neutral charge.  Now there is a caveat to that trick, right?  And that caveat was, well, what if you have the exact same charge of a higher number?  Really at that point, your best bet is if you have something that's a two plus and a  two minus, just know to put them together.  Just know that it's a one to one thing, or if it's a three plus and a three minus, it's  one to one.  But if you don't remember that and you don't pick up on it right away and you do the crisscross  trick anyways, and you end up with MG202, you have to reduce that down.  Now remember, only four ionic compounds do you reduce this down.  For this, you have to make it MGO.  You have to reduce it down.  That's not true of covalent compounds, right?  Covalent compounds might be N2H4, and you wouldn't reduce that down.  Ionic compounds always have to be lowest whole number ratio.  Covalent compounds don't.  Okay, so now let's talk about some covalent compounds.  So this is when you're going to be sharing electrons.  So for instance, with chlorine, you can go ahead and you can take, and you can share  each one, this will share one of its electrons, this will share one of its electrons, and  you'll form a single bond.  If each element is sharing one electron, that's a single bond.  And you can draw it like this, or this, I highly suggest doing it this way though.  And so that's how we're going to be representing all our covalent compounds.  So if I ask you, draw this, that's how you want to draw it.  Here's a little bit more complicated example, where now sulfur is needing to share electrons  with multiple atoms in order to get what it wants.  So it's going to share two, one oxygen, two with another oxygen, to form sulfur dioxide.  And so you can go ahead and whenever you are sharing four electrons total, two from each  element, that's a double bond.  And so you can write it like this, which gets a little messy, so I much prefer that you  write it like this and actually form the double bonds.  Now for these simpler ones, that's fine to just kind of be able to write them out, but  all those structures get pretty complicated.  And so let's go through some general rules.  And this is a good slide to keep with you as we go through all of the examples, because  we have a whole bunch of examples to do.  And we'll follow through these guidelines as we go.  Step one, and this is the step that when you start getting good at them, everybody skips  and it's a terrible idea, it's where you start making your mistakes.  And you'll see I'm kind of lazy about it sometimes too, and I'll have to go back and  do it at the end just to make sure I did it right.  So add up all your valence electrons.  In the beginning, this will help get you started.  In the end, it'll check for your mistakes, so it's very important to do.  Make a skeletal structure.  So what I mean by that is you put your central atom inside, and then you line, you put everything  else around it, and you just draw single bonds.  You connect them together, you kind of play connect the atoms.  And you don't worry too much about lone pairs and in double bonds and all of that, that just  gets you started.  So you know every time that you draw that line, that's going to be two electrons.  Now you're going to have some electrons left over.  So figure out how many electrons you have left over.  Start with the number that you started with.  Subtract out how many you had to use to form your single bonds, and now you know you have  a certain amount of electrons left over.  Put those around where you need to put them around.  Now octet rule, right?  We're trying to get eight.  So if you can't get eight, then you're going to have to do something else.  So at this point, you do want to kind of be thinking, certain things don't follow the  octet rule, and we're going to talk about what those are.  These are some common exceptions, but there's going to be some other ones too, and we'll  talk about why they are in a minute.  But you're going to want to check for those so that you know, hey, this is an atom that  typically does something that's a little bit funny.  So let's go ahead and make sure that, you know, I'm at least aware that it doesn't  have to follow the rules.  Maybe it will, maybe it won't.  Now if you can't distribute the electrons, if you try to put them around and you say,  I have two atoms that I don't have enough electrons, where do I put them?  That means that you need to start forming more bonds.  Start putting double bonds, start putting triple bonds.  A little note once you get good at them, if you're short by two electrons, that's going  to, you're going to have to add one more bond.  If you're short by four electrons, you're going to have to add two bonds.  So that's sort of a little trick to, when you start getting good at them and you want  to move faster.  However many electrons you have left out that you're missing, you need six, you're probably  going to need to form three bonds.  Now this next part, and again we don't know, we don't know what all of this means yet.  We're going to go through each of these in more detail as we do the examples.  But this is just a good slide to keep with you while you're doing these.  So you have, you don't look at them and say, what do I do?  I'm stuck.  I don't even know how to start this problem.  This is your guidelines.  They aren't hard rules, they change a little bit.  Once you're done with all of that and you theoretically have everything with octets,  you theoretically have all of your atoms happy, now you check your formal charges.  And we'll talk about exactly what that is.  We want them small.  So when we get there, we want to minimize them.  Now the only way to really do these is to go through and do a ton of examples.  And so that's what we're going to do.  We're going to do a ton of examples at which point we'll talk about formal charges.  And then we're going to do some one that break octet rules too.  Okay.  So the first, we'll start with one that isn't too bad.  We'll start with nitrogen.  So our first step, and this is, you know, our error checking and making sure that we  have somewhere to start from is to count our valence electrons.  So you pull out your periodic table and you say, well, how many valence electrons does  nitrogen have?  Well, each one of them has five.  So we have two nitrogens, each one with five valence electrons, and so we have ten electrons.  We don't have a charge, so we don't have to add or subtract any electrons.  And that's important to make sure you remember too.  We'll do some of those later as well.  So our first, our next step then is to draw our skeletal structure.  So we go through and we do that.  And now we need to add in our electrons so that we satisfy the octet rule.  Well, let's say we try to do this with nitrogen and we say, okay, well, I used up two electrons  making that bond.  So I have eight left.  So one, two, seven, eight.  Okay.  Well, I'm short electrons, right?  I can't give everything an octet.  Nitrogen would need two more, or four more here, and I can't do that.  So that means I'm not sharing enough electrons.  I'm short four electrons, which probably means I'm going to need to do two bonds as  sort of a good rule in your head to think about.  So there's only one place to put those bonds, so I'm just going to put them there.  Which means now I need, I have four electrons left because I have two, four, six.  So I've used up six and I started with ten.  So four electrons left.  So now I'm set.  Two, four, six, eight, ten.  Always go back at the end and double check and make sure your electrons add up.  And this is why it's important to do this in the beginning, even once you start getting  good at them and you start just kind of whipping them down.  At the very end, if you didn't do this, go back and do it real quick and double check  it.  Make sure that you've done it right.  Okay.  Next one.  Now we've made it a little bit more complicated.  So now we have to figure out which one's going to be our central atom.  Now, good rules of thumb for this, and there can be more than one central atom as well.  The least electronegative element, that's going to be your central atom.  And then everything else kind of goes around it.  Hydrogen, how many bonds is hydrogen allowed to form?  So all it has filled in its ground state is it has one S electron, one S electron.  And so it can only ever form one bond.  Hydrogen's not allowed to form any more than one bond.  And so because of that, you're going to, that's never going to be a central element.  Oxygen can, but it's usually only if it's with things that only form one bond.  So with all of those rules that puts carbon in our center.  Okay, so let's add up our electrons.  So we have four from carbon.  We have two from hydrogen, well, let's write that this way.  We have two hydrogen atoms, and each one of them has one electron.  And from oxygen we have six.  So we have a total of 12 electrons.  So we put carbon in our center, and we put everything else around it.  Now we try to distribute our electrons.  So hydrogen doesn't form an octet, right?  Hydrogen is the duet rule.  It only has two electrons.  So we have two electrons here, and two electrons here.  So that leaves us with, and two electrons here.  So that leaves us with six electrons to work with.  So let's go through and see if we can do this.  Well we have two, four, six, eight, ten, twelve.  So we're set for electrons, everything has an octet.  So if you just walk away at this point, you're not going to get your points.  And the reason for that we'll talk about in just a minute.  But let's for right now see if there's another way we can do it.  And after we talk about formal charges, we'll talk about why one is better than the other.  So the other way that you can do this is by doing this.  Forming a double bond here, and putting your two oxygens here.  Now let's look at all of our octets and see what we have.  Here we have two, four, six, eight, we're good.  Two, four, six, we're not good.  And then our hydrogens are good.  Here we have two and two, so we're good.  Our oxygen has two, four, six, eight, so we're good.  And now let's look at our carbon.  Two, four, six, eight, so we're good with carbon.  Here we didn't have enough electrons to do it.  Here we do.  So double bonds and triple bonds are a way to give ourselves extra electrons because  we're sharing them.  So this one's wrong.  And this one is the one that we want to be using.  Okay, next example.  And when you're doing these, you'll notice I've given myself an entire sheet of paper  to do these on.  That's on purpose.  We're going to be using these for the entire rest of the quarter, pretty much, all the  way up until at least gas loss.  So make sure that you give yourself just as much room as I did, and we're going to be  adding to them.  When we get to geometries, we're going to go through and we're going to put all of our  geometries on.  When we get to hybridization, we're going to go through and put all of our hybridization  on.  We'll do bond angles.  We'll do polarity.  We're going to just spend a lot of time with these sheets of paper.  So one per sheet is about right, any more than that, and you start running out of room  near the end of the quarter.  Okay, now the next one.  And this one gets a little complicated because now we have two central atoms.  Now luckily, people usually write these in a way that makes it a little bit easy to see  what's going on.  So you can kind of see here that we have BH3 and H3, and that's written that way on purpose.  So you can sort of see that you have like a BH3 group bonded to an NH3 group.  So first, count our electrons.  We go to boron.  We have three.  We go to nitrogen, and we have five.  We go to hydrogen.  There's six hydrogens, and each one has one electron.  So we have 14 electrons to work with.  So now our skeletal structure.  We can kind of see that boron and nitrogen are the only ones who can form multiple bonds.  So they're going to need to be bonded to each other, as well as all the hydrogens.  And so at this point, we say, OK, well, how many electrons do we have?  2, 4, 6, 8, 10, 12, 14.  So we're good.  Does everything have its octet filled, or for hydrogen, it's duet?  Well, all the hydrogens are set.  And then boron has 2, 4, 6, 8, and nitrogen has 2, 4, 6, 8.  And so we're set there.  Now later on, we're going to talk more about exceptions.  And we'll see that boron's actually one that a lot of times will have less than eight electrons.  A lot of times, it'll just have six, and it's fine with that.  But here's an example where it actually does what most atoms do, and it has its eight.  OK, next one.  Now we need to do an ion.  So you have some practice doing ions.  So we have XCF4, 4 plus.  So conopolar electrons.  So xenon has eight.  It's a noble gas.  And then fluorine has seven each.  So we have four fluorines, and each one of them has seven.  So that would be 36.  But it's a four plus charge, which means that we need to take away four electrons to give  us 32 electrons.  So it's a plus charge.  We have less electrons than the neutral atom, and so we minus the four electrons.  So we have 32 electrons to work with.  So we put xenon in our center.  Fluorine never allowed to break the octet rule, and we'll talk more about that in a  minute as well.  So fluorine is ever only going to have one bond.  So whenever you see a hydrogen or a fluorine, only one bond, it can't be a central atom.  So when we go through and we draw these out, we know fluorine can't have more than one  bond since I just told you.  So really all we can do at this point is add in our lone pairs and hope everything works.  If not, we'll have to add some things to xenon.  So if you count all of these, 2, 4, 6, 8, all the way around, what you see is that we  do in fact have 32 electrons.  So we're done.  We don't have to do anything else with this one.  OK.  Now our last and most complicated one for right now.  This one we won't actually even be able to finish until we talk about some other things.  OK.  So N2O.  So this one's kind of not as simple as the other ones, right?  And we'll see why in a minute.  So first of all, most electronegative element here is oxygen.  So nitrogen's going to be in our center, one of our nitrogens.  So we can set up our central atom as being nitrogen.  And we can put a nitrogen on each side, or excuse me, a nitrogen on one side, an oxygen  on the other one.  And then we can draw our skeletal structure.  Now if we try to give everything an octet here, first we're going to need to count our  electrons.  So you'll see that we have five electrons from each nitrogen and we have two nitrogens.  We have six electrons from our oxygen leaving us with 16 electrons.  If you try to make your octet here with just going and filling in your elements or your  electrons, and we just kind of fill in, you notice we can't do it.  We don't have enough here, right?  We would need four more electrons.  So if we're short four electrons, how many bonds do we need to form?  We need to form two bonds.  Okay.  Well, where do we put those bonds?  Now we have some options.  When we had to form two bonds with that N2 group, we didn't really have any options.  So we just sort of put them in between the two nitrogens.  Here we have options.  So I have two bonds that I need to form.  So I could just put one bond on each group.  I could put both bonds to the oxygen, or I could put both bonds to the nitrogen.  And let's go through and fill in our octets at this point.  So we have two here, so we need two more.  This nitrogen set, two, four, six, eight.  This oxygen needs four more.  We count up our electrons, two, four, six, eight, ten, twelve, fourteen, sixteen, we're  set.  We move down to this one.  Now this nitrogen needs six.  This nitrogen again is all set.  And this oxygen is going to need one.  This nitrogen needs two.  And this oxygen needs three.  We go through and count, two, four, six, eight, ten, twelve, fourteen, sixteen, we're good.  Two, four, six, eight, ten, twelve, fourteen, sixteen, we're good.  Okay.  Well, which one is it?  So all of them follow all of the octet rules.  All of them have the least electronegative element in the center.  So which one do we do?  This is where that very last step on the slide comes in.  We don't really know how to do that yet.  So let's learn how to do that and then we'll come back.  So there's an equation for this.  And once you start getting good at it, you're going to want to do it more sort of internally  and be able to quickly check all of your formal charges.  You don't necessarily always want to be filling into this equation.  But for the moment, it's a good thing to start with.  So your valence electrons.  So you start there.  So for nitrogen, it's five.  Oxygen, six.  If you ever forget, you just go to your periodic table and count over until you get to the  element where you're at.  Now that's your starting point.  Then you subtract the number of electrons that you're, the atom that you're talking  about in your, whatever you're looking at.  What I kind of always think of it is which ones they actually own.  So which ones are sort of their electrons?  So the way that that's calculated is the number of electrons, the number of electrons that  are actually in, that it has all by itself.  It's lone pairs.  So you're subtracting away all the electrons that are inside the lone pairs.  And then half of the bonding electrons.  So if you have a bond, one of those, one of those electrons belongs to one element and  one of those electrons belongs to the other element.  So for accounting purposes, you always think of half of the bonded electrons.  And we'll do some examples that'll make that make more sense.  So let's go through and assign formal charges to all of the Lewis structures that we did.  And then when we get done with them, we'll be able to actually go through that last harder  example that we did and decide which one it actually is, which Lewis structure is actually  correct.  So just sort of starting over on our Lewis structures.  Again you'll want to keep them handy and do this quite a bit throughout this next little  bit of the quarter.  Okay.  So our first one.  So what is our formal charges on each nitrogen?  So we go and we say, well, and formal charges are going to be atom specific.  So it's for each atom in the molecule.  Now in this case, they're going to be the same, but we can still calculate them for  both.  So how many electrons does nitrogen have, valence electrons does nitrogen have?  You go to the periodic table, you come over and you see that nitrogen has five.  Now you have to decide how many are its own, how many does it have in this structure.  And you're going to subtract those away.  So anything in a lone pair is its own.  So it has two from the lone pairs.  From the bonds, it only gets one from each bond.  So one, two, three.  So you have five valence electrons minus five is zero.  So this nitrogen's formal charge is zero.  Now if you go to the other nitrogen, the exact same thing happens.  It has five valence electrons.  It has two electrons that are in lone pairs and it gets one from each bond.  And so that nitrogen's going to be plus zero as well.  Next one.  So now we can go through and we can talk about each one individually.  So let's start with the oxygen.  So how many valence electrons does oxygen have?  It has six.  And we're going to subtract away anything that are in lone pairs.  So one, two, three, four.  And any of the bonding electrons but only one from each bond.  One, two.  So we have six minus six.  So the oxygen's charge is zero.  Carbon.  We go to the periodic table.  We say how many valence electrons does it have?  And it has four.  So we subtract away anything.  In this case, there's no lone pairs so we don't have to count any of those.  And it gets one from each bond.  So one, two, three, four.  So we have four electrons there.  So its formal charge is zero.  Now we need to talk about hydrogen.  Hydrogen has one valence electron and gets one from each bond.  So its formal charge is zero as well.  And that would be true of the other hydrogen as well.  So those are all our formal charges for that one.  Next one.  So now this one when we go through, it's going to be a little bit different.  We'll start actually getting some formal charges.  So let's start with the hydrogens.  Hydrogen gets one valence electron.  And it's the same for each of these hydrogens.  It'll get one from the bond.  So every one of the hydrogens on here follow the same rules and so they're all going to  be plus zero.  Now let's talk about the boron and the nitrogen.  For boron, how many valence electrons does it have?  So you go to the periodic table and you see that it has three.  Now if you come here and we count electrons, no lone pairs, so don't worry about those,  gets one from each bond.  One, two, three, four.  So boron has a minus one charge.  Now nitrogen, if you go to its period, it on the periodic table, you'll see that it  has five valence electrons.  And you count around one, two, three, four.  And so you have a plus one charge on nitrogen.  Now the xenon tetrafluoride.  So for each of the fluorines, because again they're all going to follow this numerals,  so we're only going to do it once, they get seven electrons from the periodic table, valence  electrons.  Now here each one of them has two, four, six that are in lone pairs.  So those are the fluorine lone pairs.  It's not sharing them, it's not doing anything with them.  So it owns all of those.  Now for the bonds, it just gets one.  So we have seven valence electrons, minus six electrons that are in lone pairs and one  electron that is in the bond.  So each of the fluorines are going to have a zero charge.  Now we move on to the xenon.  It's a noble gas, so it's in that eighth group.  Has eight valence electrons.  And we subtract one from each bond.  One, two, three, four.  So we subtract four electrons.  So it has a plus four charge.  So as we're going through these, you should notice that we're counting electrons very  differently depending on what we're doing.  When we were counting for octets, we counted everything, right?  For xenon we went through and said two, four, six, eight.  Xenon has an octet.  Fluorine has two, four, six, eight.  Fluorine has an octet.  That's very different from when we do formal charges.  The bonds only count as one for each atom.  So fluorine, we only got to count that as one electron.  For the sake of octets, you counted as two.  Okay, so now we can move on to our last one, which was a bit more complicated.  And we weren't actually able to finish.  Okay, so we had this.  And we said, well, I don't know which of these three to do.  Now we can actually go through and we can figure this out.  So let's do that.  Let's go through and put formal charges on everything.  And I'm going to do this a little bit quicker this time and do it a little bit more like  how you should be doing it on an exam and how after maybe about a week of practice and  a week of homework you should be doing it.  So we look at oxygen and we say, how many valence electrons does oxygen have?  Well, it has six of them.  How many does it have in lone pairs?  Four.  Okay, so we're down to two electrons at that point and then two here.  So oxygen has six valence electrons and here it has one, two, three, four, five, six.  So six minus six is zero.  Let's actually switch colors just to make it a little bit more obvious.  Plus zero.  Now nitrogen.  Nitrogen has five valence electrons and here we have one, two, three, four.  So we have five minus four is equal to one and it's a positive one.  Now here nitrogen has four valence electrons and here it has one, two, three, four, five,  six.  So five minus six is minus one.  So again, when you're practicing at home, do it the way that in the beginning, do it  the way that I've been doing it up until now.  When you get a little bit on a roll with it, this is kind of the fast way to do it.  This is how you want to be able to do it for the sake of, you know, checking your work  quickly without having to waste too much time.  Okay, let's do the next one.  Here we have oxygen normally has six valence electrons.  Here it has one, two, three, four, five.  So six minus five is a plus one.  Over here, nitrogen normally has five and here it has one, two, three, four.  Five minus four is plus one.  Here nitrogen normally again has five but here it has one, two, three, four, five, six,  seven.  Five minus seven being minus two.  Okay.  So just going through, putting formal charges on everything and then this is going to help  us decide something when we're all done.  Down here, oxygen normally has six.  Now we have one, two, three, four, five, six, seven.  We have minus one.  Here nitrogen normally has five.  One, two, three, four.  So five minus four is one.  Over here, nitrogen has five valence electrons and here it has one, two, three, four, five.  Five minus five being zero.  Now as a good check, we need to go, this is something you kind of want to do in your head  for everything.  You don't necessarily need to write them down.  If I ask you to write down a Lewis structure, you don't have to put informal charges.  If you have something that's an ion, actually like this, if you have something like an ion,  it's not actually correct unless you put the charge somewhere but you don't have to put  the charge on Xenon.  You can just put brackets and put the charge and that's fine.  That's just as correct.  Something like this that doesn't have a charge, you don't actually have to put the formal  charges unless I ask for it.  You should do it in your head anyways though.  Now we have to decide which one's correct.  So if we look at these, a good check first is to make sure that all of our formal charges  add up to whatever the charge is here.  So charge is zero.  So this one adds up to be zero.  This one adds up to be zero.  So we've done our formal charges correct.  We've done our Lewis structures correct.  Now we have to decide which set of formal charges is the best set of formal charges.  And so if we go through and we look, we can say, well, this one has a negative 2, a 1,  and a 1.  This one has a negative 1, 1, 0, negative 1, 1, 0.  This one has some really high formal charges and there's no reason for that.  2, 1, 1 is acceptable if there's no other option.  You will have formal charges of plus 4 sometimes.  We saw that with Z non.  But if there's another option, you don't want to do that.  So we're going to get rid of that one right away.  That one's not going to be right.  So now we have to decide between these two.  So they're basically the same.  Our central nitrogen has the same.  But now here we have a negative one on the nitrogen.  Here we have a negative one on the oxygen.  So we have to figure out which one's better.  So for this, we need to kind of backtrack to what we talked about earlier today or last  class, where we were talking about electronegativities.  So when we're talking about electronegativities, we're talking about the amount that an atom  is willing to pull the electron density toward it inside of a bond.  So this is kind of exact in this situation.  And something like oxygen is more electronegative than something like nitrogen.  So if the electron density is going to be shifted one way or the other, it's going to  be shifted so that the more electronegative element, the one that's pulling the electron  density toward it, is going to be the negative one.  Because well, it's more electronegativity.  It has more electronegativity.  It's pulling those electrons toward it more.  It's going to take the negative charge, not nitrogen.  So this one is going to be correct because oxygen is more electronegativity, or more  electronegative.  So you always, if it's possible, and keep this in mind, if it's possible, you want  to put your negative charge on your most electronegative element.  That's not always possible.  For instance, this one again.  In this one, we had a plus.  Oh, I guess we didn't have a negative charge.  Never mind.  So draw out CO for fun sometime.  Draw out carbon monoxide.  And you'll see that you're forced to do the opposite of this.  You don't have a choice.  And that's fine.  If you have a choice, you don't have a choice.  But in something like this, we do.  We can choose to put it on the nitrogen by drawing it that way, or we can choose to put  it on the oxygen by drawing it that way.  So this is the more correct way.  OK.  So now, moving on a bit.  More in detail about breaking the octet rule.  I've been sort of alluding to this and talking about it, but now we're going to go into it  in a little bit more detail and talk about exactly when we can and can't do this.  So your second period elements, they're not going to do this.  So that second row of the periodic table, well, and of course, hydrogen and helium as  well.  Hydrogen, helium, and then lithium up through neon.  They're not going to do it.  Do not ever have 10 electrons off fluorine or an oxygen or nitrogen.  You will just get it wrong.  There will be no partial credit for that.  You'll just have it wrong.  Now third period and lower, you can do that and it's fine.  That's not a big deal.  There's some that are going to do it more than others, but not all of them.  Now the third period element can do it where the second one can't because of the fact that  they have some d orbitals sitting right there.  So the d orbitals are right there.  They're able to be used for bonding.  You can use them for bonding.  You can use them for hybridization, which we'll talk about later in the next chapter.  And so anything below, third row and below can do it.  Second period ones do not do it ever.  Some of them are going to have a lot less than eight and they don't have to, but they  can.  Theoretically, anything could have less than eight electrons, but that's not necessarily  normal.  For beryllium, aluminum, and boron, it's pretty normal.  If you see one of these and you can't fill the octet, don't worry about it.  Just move on.  Anything like carbon and oxygen, you're going to probably fill the octet.  Sell them going to break the octet rule.  It's not something that you're going to ever see in this class.  Where these elements, they do it all the time, so get used to these elements breaking the  octet rule on the low side and having less.  As far as having more, that falls into the category of sort of a lot of these that I  talked about.  There's a lot of phosphorus xenon, chlorine does it a fair amount.  Anything third row and below can do it, but these are sort of your common ones that actually  do it all the time.  We'll start and do some examples of this now and then do some more next time.  Okay, so we'll start with SF6.  Our central one is sulfur, since we can't really do anything else, we can't put fluorine  in the center.  When we go through and do this, sulfur here, and we have six things, so we just spread  them all out and put fluorines on all of them.  Because it's a Lewis structure, we need to go through and we need to add in all of our  lone pairs.  We count up how many electrons we started with just to make sure that everything is  still okay.  We have six from the sulfur, seven from each fluorine, and we have six of them.  We have 48 electrons.  If we go through and count all of these up, it adds up to be 48, and so we're all set.  Sulfur is an example where it's okay with having six bonds.  Usually with sulfur, you want to kind of get in the habit of seeing, well, it'll either  have two bonds and fill its octet, or it'll have four bonds, or excuse me, six bonds in  this sort of configuration.  So it'll either act like oxygen, which is right below on the periodic table, or it'll  form six.  Okay, next one.  So this is kind of similar to one that we've done before, only now it's not an ion.  Now it's XCF4 without any 4 plus.  So we count up our electrons.  We had eight from xenon.  We have seven from each fluorine, and we have four of them.  So we have 36 electrons.  It's not a charge, so we don't have to worry about adding or subtracting any.  Now when we go to draw this, this is a good example of one where you can get yourself  into trouble if you don't count the electrons first.  So up until now, a lot of them have been, once you get really good at them, you may  say, well, I don't need to count electrons.  I know what it looks like.  This is one where you may get yourself into trouble if you don't.  So if we go through and we kind of set up our skeletal structure, and we know fluorine  only gets to have one bond, so we're just going to fill those in.  At this point, you might be tempted to say, okay, everything has an octet.  I'm done.  And walk away.  In which case, you'd get it wrong.  And then all the other questions following it that we're going to learn how to do later,  like hybridization and geometries and bond angles and things like that.  So we look at this and we say, okay, how many electrons have we used up?  If we count them all, we've only used up 32.  So we have two more, or four more electrons that we have to deal with.  So we can't put those electrons on fluorine, because fluorine only gets to have eight.  Xenon can have more if need be.  So we're going to give them the xenon.  So that gives xenon more, it has two, four, six, eight, ten, twelve electrons.  So that's okay.  So that's one where we're giving xenon twelve electrons.  And so we're breaking the octet rule there.  We can't break it with fluorine, but xenon can break it.  Okay, so now let's do BH3.  So we add up our electrons.  We say, well, we have three from boron.  We have three from each hydrogen, one from each hydrogen.  So we have six electrons total.  We go to draw this.  And we have this.  And at this point, there's nothing else we can do.  We only have six electrons to work with.  Well, if we only have six electrons total, there's absolutely no way to give boron eight.  We can't just make two electrons out of malware.  So we're done.  That's all there is to it.  We can't really do anything else.  Now as good outside practice on this, go through and calculate all your formal charges.  We haven't done any formal charges yet, because it hasn't really come up where we've had  an option where we needed to decide between things.  But it's a good practice to go through and make sure that you can do it and get in the  habit of doing it.  And so there's lots of places here for you to practice it.  Now a more complicated one.  And these ones tend to get people a lot.  For all of these polyatomic ions where we turn them into acids, so they're called oxoacids.  So in this case, this one would be called what?  Sulfuric acid.  So, or phosphoric acid or carbonic acid.  Any of these, you want to get in the habit of knowing sort of the general way that these  look like.  Because otherwise you might be tempted to take sulfur and bond one to each oxygen and  to the hydrogens.  You don't want to do that.  Whenever you see these, you need to know that these hydrogens are actually bonded to the  oxygen.  So when we go to do this, we're going to put sulfur in the center.  We're going to put four oxygens around.  And then the hydrogens go on the oxygens.  And that should just, you should know that as part of your skeletal structure from now  on.  And if you come to me later and say, well, how was I supposed to know that?  I'm going to point you to this slide.  So this is how you know it.  And this goes for sulfuric acid, carbonic acid, any of those oxoacids.  So now let's go through.  We can fill our octet in here because there's not really much else we're going to probably  do with oxygen.  And at this point in practicing, you may get in the habit of saying, okay, well, I know  oxygen pretty much always has two, two bonds.  So I'm just going to give it two bonds in its lone pairs.  Now if we add in our electrons and we try to add 2,4,6 here and 2,4,6 here, we'll see  that we can't do it.  And so what we need to do, we'll see that we're short for electrons.  So go home, practice it, do it.  We'll see we're short for electrons, so we need two bonds.  And so we're going to put our two bonds here, or excuse me, our two lone pairs here and  our two lone pairs here.  And that's where we'll finish.  Now you'll notice I forgot to add in all of my electrons and make sure that we have the  right amount.  So let's do that real quickly, just as a double check.  If you forget to do it in the beginning, especially when you get good at them, you may get out  of the habit and you may forget.  Just do it at the end and make sure that you're right.  And when you go through and do this, you'll see that, hey, we do have 32 electrons and  we are all set.  So at this point, we've gone through and we've done a bunch of ones that break the octet  rule.  We have a bunch more that we're going to do because Lewis structures are really, really  important.  You will not be able to do all of the rest of the material for this chapter and chapter  3 and chapter 5 if you can't do these.  So we're going to do a bunch more examples next time.  So we'll just kind of go through them at that point and keep going through and get into  all of that.  So we'll end there for today though and start up on the rest of them next time.
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Okay, so today we're going to start chapter two. And chapter two is going to start off with a lot of review, mostly because I thought you needed to know a lot of this before we really even got started with class and so I talked about it earlier. But we should go over it again and just make sure it's cemented in there before we start talking about the more complicated aspects of chapter two. So we're going to go through it and we're going to remind you about ionic and covalent bonding so that you remember exactly what that is and what the differences are because it becomes really important here. All your naming is still really important here and knowing how to identify between each one is very important. Now we're going to introduce Lewis Dot symbols in a little bit more of a formal nature. Before we kind of threw them at you and said this is some really simple ones, this is how we represent them, now we're going to go through and we're going to do a bunch of very complicated ones that go way more into depth on what exactly is going on which will lead us into trying to talk about how molecules look like and doing geometries and things like that. Then we're going to do some more Lewis Dot structures that are even more complicated and more Lewis Dot structures. Really chapter two is kind of all about Lewis Dot structures and getting good at them. Once we've really got that down packed, we're going to go through and do some bond strength and length comparisons so that we can see how they're related to each other. Again some review from last time. Bonding is all about trying to figure out how to complete octets or make things more stable than they were. Now you can do that by in something like this case sharing electrons. If two different electrons have six valence electrons, they can go share them. They can take two of their electrons, share them with each other and then form some bonds there so that they each can have eight. There's our three main types of bonding. We have our ionic bonding, our covalent bonding and our metallic bonding. Remember these are important for knowing how to name them because we use different names for our ionic than we do covalent. They're very different from each other. So in ionic you're trading an electron. Sodium says, hey I have an extra electron and chlorine says, well I don't have enough and so the sodium gives its electrons to chlorine. Sodium becomes a 1 plus, chlorine becomes a 1 minus. They share the electrons or they switch the electrons. Now that plus and minus, now they're attracted to each other. They're like little magnets, right? So they're going to be attracted to each other and that's how that bond forms. Now something that's covalent on the other hand, they're going to share the electrons. So something like carbon monoxide, carbon and oxygen, both of them don't have quite enough electrons. So if they share some of them, now they can each have what they need. And so they're kind of stuck together by the fact that they're sharing the electrons. Now we need to know how to differentiate between the two, right? So the way that you decide whether something's ionic or whether something's covalent is for right now, seeing where they are in the periodic table. So if something's on the right hand side and the left hand side are two different things like sodium and chloride, you have a metal and a non-metal, that means that you're going to have an ionic bond. Something that's covalent will be two non-metals, so everything up on that top right hand corner of the periodic table. If it's made up of things over there, that's going to be covalent. Now this is a little bit of an oversimplification and toward the end of the chapter we will start talking about how this isn't quite exactly how things are and that really there's a gradient. Some things are more ionic than others. You can have a purely covalent compound and then there's sort of a gradient all the way up until the ionic ones. And so we'll talk about that later. And then we had the metallic which we aren't talking a lot about now. You'll get into that more in Chem 1b where, but we did want to talk about it a little bit. The main difference here is that all of your electrons are delocalized. It means that one particular electron doesn't really belong to any particular atom. And so there's this kind of sheet of atoms and sheet of electrons all bonded together, which means that you can pass electrons through it. If you put an electron in one side, you can make it come out the other. And so that's why they're conductive. Now this is sort of, you know, we didn't quite get into this before. And there's a little bit of background here that you sort of need to make sure you know from the fundamentals that maybe you glazed over a bit. So all of the processes that consume energy are called endothermic. So think back to your high school or one of the P experiences and think about these words coming back to you again, right? For taking in energy, that's endothermic, bringing in energy. If you're releasing energy, you're throwing energy out, that's exothermic, right? I think exoskeleton on the other side is putting energy out into the environment. We're going to go really far into that in Chem 1b. Here we just need to be able to remember what's endothermic, what's exothermic in the signs that are related to those. So we want to talk about the energy of an ionic bond. So think about before I show you the next slide, which has a lot on it, what I said before about how we form an ionic bond. An electron goes away from one atom. It goes to the other atom and then the plus and the minus are attracted to each other. So we have three kind of steps there. We have an ionization energy, the electron being pulled off an atom. We have an electron affinity, an electron coming into an atom, the other one. And then we have this other energy that comes from the attraction of a plus and a minus. And we have that energy. So three energies. And for it to be favored, we want it to be exothermic. We want it to give off energy. If we have to put a bunch of energy into form something, that's not going to be favorable. So this is sort of the background for the next slide. Okay. So this is what I just said all kind of laid out formally. First this one. We have sodium losing an electron. We have an electron going away from the sodium. Now we have a word for that, right? We talked about it just last lecture. We called that ionization energy. An electron is leaving. So the only real difference here is instead of talking about a trend, I just put a number on it. And I put a plus number on it. So it takes energy to be able to lose an electron. So this endothermic reaction, in quotes. Now C.L. is going to gain an electron. It's going to take an electron in. So that's, in other words, an electron affinity, right? Once again what we talked about last time. And again we'll just kind of give it a number now. So that takes a little bit of energy. And so that becomes 340, or negative 349. So these are just numbers for these specific things. Obviously not something that you would ever be expected to memorize. What I really want you to get out of this is the magnitudes. So 1B will go into more how to calculate energies and how to deal with them on a very mathematical form. I want you to kind of think about the magnitudes here. So we have a plus energy. We have a minus energy. So we have an endothermic. We have an exothermic taking in energy, giving off energy. OK. So once we do that, now we've switched to the electrons, right? We've taken away the electron from one and we've put it on the other. One took in energy, one released energy. Now we have sodium chloride. We have sodium plus and chloride minus attracting each other. And that's favorable too. That's exothermic as well. So we get a negative 787. So you may want to add some little notes here. If you've forgotten what designates an endothermic and an exothermic for signs. Pluses are endothermic, minuses are exothermic. So in order for this to be energetically favorable, you want it to be exothermic, which means do you want a positive or a minus? Remember, endothermic, exothermic. And now we want this whole thing to be exothermic. We want the whole thing to be negative. Well what that means then is we need all three of these numbers to add up to be a negative number. So if we do that here and we add up this, this, and this, you'll see that we get a negative 642 kilojoules per mole. So that's definitely exothermic, right? We have a negative charge. So it's lower. We're losing energy. We're releasing energy into the environment. And so it's going to be stable. So the idea here is that you want your electron affinity and you want this value to be negative. It would be bigger than this one. If this is too large or these are too small negatively, you know, very, it's close to zero, you're not going to be able to form this. In this case we do though, because these added to this make a negative number. So your ionization energy has to be low enough to make this exothermic. And these have to be large enough to make it, excuse me, these have to be negatively large enough. They have to be negative enough to make the whole thing negative. Now this goes into why we tend to only form ionic compounds with metals. Something needs to have a low enough ionization energy to actually be able to do this. If this is too high, it's going to overwhelm these. If this was like 2000, then these wouldn't, it wouldn't be exothermic anymore. These wouldn't be able to get it into the negative line. So because of this, it's typically just your metals, right? Think back to last class. As we were on the left-hand side of the periodic table where all the metals were, those were where our low ionization energies were. So this is where these numbers are kind of low. And so these numbers are able to overwhelm them. They're able to say, no, I don't care if you take a little bit of energy to do, we give off so much energy that it's still energetically favorable. So it's mostly just metals that do this because it's metals that have a low ionization energy. The ones on the right-hand side of the periodic table, carbons, oxygens, fluorines, all of those, those have such a high ionization energy. This number is so high. And go look at it. You know, pull out your table of numbers and see they're so high that it's going to overwhelm it and it's not going to happen. Okay. So now some loose dot symbols. And there's a little more formal introduction of them. They're going to go, they're going to be how we represent our elements when we draw them out. And there's lots of different ways to represent them. Loose dot symbols is one of the simplest and most used. So you put the symbol for the element inside and then you put your dots around the outside. So for iodine, for instance, you have seven valence electrons. You're not going to draw all your electrons, of course, just your valence ones. And then whenever you need to draw a bond, you do a line. You'll sometimes see people represent these, the, instead of a line, they'll put two dots instead and then just squish them close together. Like this. That's fine. It's not wrong. And then here I like to do it to show you how we're switching electrons that when we talk about ionic bonds, you can see it. When you're actually drawing these, it gets a little bit messy. And I would suggest keeping it like this and just putting the bonds. Just remember that every time you draw one of these lines, it's comprising two electrons. Okay. So now let's do some ionic bonds with these. If we have potassium and we have iodide, we are going to take the electron from potassium and give it to iodine, which forms potassium iodide with a plus and a minus charge. Lewis structures get a little boring with ionic compounds. They're a little bit better, a little more entertaining to use with covalent compounds. With magnesium, remember this was our example where we said, well, sometimes ionic compounds aren't going to form a one to one ratio. Sometimes magnesium might need to give an electron to two different atoms in order to make everybody happy. So magnesium says, well, I have two electrons that I need to get rid of and fluoride says, okay, well, I'm only going to need one. So they give one to fluoride and one to the other fluoride, or fluorine forming magnesium fluoride. And the simple way of thinking about that was always, well, we need to balance the charges, right? If we know what our ions are, we know that magnesium likes to be a two plus. And if you forgot why, look at your periodic table, it's in the second column. So you go to your periodic table, you see it's in the second column, you know it needs to be a two plus. You go to fluorine on the periodic table and you say, okay, well, that has seven electrons, so it needs to have a minus charge. So now, if I have an Mg2 plus and an F minus, I need to balance those charges. So that was a faster way of looking at it. Now there was a trick to doing that as well. And it's to crisscross the charges. So if you had aluminum three plus and oxygen two minus, you're going to take the one charge and bring it down, the other charge and bring it down. And then that'll give you a neutral compound because two aluminum is two times six, or two times three is six. And then two oxygens, so, excuse me, three oxygens, two minus, so that's a negative six. We have a plus six and a minus six, so we have a neutral charge. And whenever you're forming ionic compounds, you need it to be a neutral charge. Now there is a caveat to that trick, right? And that caveat was, well, what if you have the exact same charge of a higher number? Really at that point, your best bet is if you have something that's a two plus and a two minus, just know to put them together. Just know that it's a one to one thing, or if it's a three plus and a three minus, it's one to one. But if you don't remember that and you don't pick up on it right away and you do the crisscross trick anyways, and you end up with MG202, you have to reduce that down. Now remember, only four ionic compounds do you reduce this down. For this, you have to make it MGO. You have to reduce it down. That's not true of covalent compounds, right? Covalent compounds might be N2H4, and you wouldn't reduce that down. Ionic compounds always have to be lowest whole number ratio. Covalent compounds don't. Okay, so now let's talk about some covalent compounds. So this is when you're going to be sharing electrons. So for instance, with chlorine, you can go ahead and you can take, and you can share each one, this will share one of its electrons, this will share one of its electrons, and you'll form a single bond. If each element is sharing one electron, that's a single bond. And you can draw it like this, or this, I highly suggest doing it this way though. And so that's how we're going to be representing all our covalent compounds. So if I ask you, draw this, that's how you want to draw it. Here's a little bit more complicated example, where now sulfur is needing to share electrons with multiple atoms in order to get what it wants. So it's going to share two, one oxygen, two with another oxygen, to form sulfur dioxide. And so you can go ahead and whenever you are sharing four electrons total, two from each element, that's a double bond. And so you can write it like this, which gets a little messy, so I much prefer that you write it like this and actually form the double bonds. Now for these simpler ones, that's fine to just kind of be able to write them out, but all those structures get pretty complicated. And so let's go through some general rules. And this is a good slide to keep with you as we go through all of the examples, because we have a whole bunch of examples to do. And we'll follow through these guidelines as we go. Step one, and this is the step that when you start getting good at them, everybody skips and it's a terrible idea, it's where you start making your mistakes. And you'll see I'm kind of lazy about it sometimes too, and I'll have to go back and do it at the end just to make sure I did it right. So add up all your valence electrons. In the beginning, this will help get you started. In the end, it'll check for your mistakes, so it's very important to do. Make a skeletal structure. So what I mean by that is you put your central atom inside, and then you line, you put everything else around it, and you just draw single bonds. You connect them together, you kind of play connect the atoms. And you don't worry too much about lone pairs and in double bonds and all of that, that just gets you started. So you know every time that you draw that line, that's going to be two electrons. Now you're going to have some electrons left over. So figure out how many electrons you have left over. Start with the number that you started with. Subtract out how many you had to use to form your single bonds, and now you know you have a certain amount of electrons left over. Put those around where you need to put them around. Now octet rule, right? We're trying to get eight. So if you can't get eight, then you're going to have to do something else. So at this point, you do want to kind of be thinking, certain things don't follow the octet rule, and we're going to talk about what those are. These are some common exceptions, but there's going to be some other ones too, and we'll talk about why they are in a minute. But you're going to want to check for those so that you know, hey, this is an atom that typically does something that's a little bit funny. So let's go ahead and make sure that, you know, I'm at least aware that it doesn't have to follow the rules. Maybe it will, maybe it won't. Now if you can't distribute the electrons, if you try to put them around and you say, I have two atoms that I don't have enough electrons, where do I put them? That means that you need to start forming more bonds. Start putting double bonds, start putting triple bonds. A little note once you get good at them, if you're short by two electrons, that's going to, you're going to have to add one more bond. If you're short by four electrons, you're going to have to add two bonds. So that's sort of a little trick to, when you start getting good at them and you want to move faster. However many electrons you have left out that you're missing, you need six, you're probably going to need to form three bonds. Now this next part, and again we don't know, we don't know what all of this means yet. We're going to go through each of these in more detail as we do the examples. But this is just a good slide to keep with you while you're doing these. So you have, you don't look at them and say, what do I do? I'm stuck. I don't even know how to start this problem. This is your guidelines. They aren't hard rules, they change a little bit. Once you're done with all of that and you theoretically have everything with octets, you theoretically have all of your atoms happy, now you check your formal charges. And we'll talk about exactly what that is. We want them small. So when we get there, we want to minimize them. Now the only way to really do these is to go through and do a ton of examples. And so that's what we're going to do. We're going to do a ton of examples at which point we'll talk about formal charges. And then we're going to do some one that break octet rules too. Okay. So the first, we'll start with one that isn't too bad. We'll start with nitrogen. So our first step, and this is, you know, our error checking and making sure that we have somewhere to start from is to count our valence electrons. So you pull out your periodic table and you say, well, how many valence electrons does nitrogen have? Well, each one of them has five. So we have two nitrogens, each one with five valence electrons, and so we have ten electrons. We don't have a charge, so we don't have to add or subtract any electrons. And that's important to make sure you remember too. We'll do some of those later as well. So our first, our next step then is to draw our skeletal structure. So we go through and we do that. And now we need to add in our electrons so that we satisfy the octet rule. Well, let's say we try to do this with nitrogen and we say, okay, well, I used up two electrons making that bond. So I have eight left. So one, two, seven, eight. Okay. Well, I'm short electrons, right? I can't give everything an octet. Nitrogen would need two more, or four more here, and I can't do that. So that means I'm not sharing enough electrons. I'm short four electrons, which probably means I'm going to need to do two bonds as sort of a good rule in your head to think about. So there's only one place to put those bonds, so I'm just going to put them there. Which means now I need, I have four electrons left because I have two, four, six. So I've used up six and I started with ten. So four electrons left. So now I'm set. Two, four, six, eight, ten. Always go back at the end and double check and make sure your electrons add up. And this is why it's important to do this in the beginning, even once you start getting good at them and you start just kind of whipping them down. At the very end, if you didn't do this, go back and do it real quick and double check it. Make sure that you've done it right. Okay. Next one. Now we've made it a little bit more complicated. So now we have to figure out which one's going to be our central atom. Now, good rules of thumb for this, and there can be more than one central atom as well. The least electronegative element, that's going to be your central atom. And then everything else kind of goes around it. Hydrogen, how many bonds is hydrogen allowed to form? So all it has filled in its ground state is it has one S electron, one S electron. And so it can only ever form one bond. Hydrogen's not allowed to form any more than one bond. And so because of that, you're going to, that's never going to be a central element. Oxygen can, but it's usually only if it's with things that only form one bond. So with all of those rules that puts carbon in our center. Okay, so let's add up our electrons. So we have four from carbon. We have two from hydrogen, well, let's write that this way. We have two hydrogen atoms, and each one of them has one electron. And from oxygen we have six. So we have a total of 12 electrons. So we put carbon in our center, and we put everything else around it. Now we try to distribute our electrons. So hydrogen doesn't form an octet, right? Hydrogen is the duet rule. It only has two electrons. So we have two electrons here, and two electrons here. So that leaves us with, and two electrons here. So that leaves us with six electrons to work with. So let's go through and see if we can do this. Well we have two, four, six, eight, ten, twelve. So we're set for electrons, everything has an octet. So if you just walk away at this point, you're not going to get your points. And the reason for that we'll talk about in just a minute. But let's for right now see if there's another way we can do it. And after we talk about formal charges, we'll talk about why one is better than the other. So the other way that you can do this is by doing this. Forming a double bond here, and putting your two oxygens here. Now let's look at all of our octets and see what we have. Here we have two, four, six, eight, we're good. Two, four, six, we're not good. And then our hydrogens are good. Here we have two and two, so we're good. Our oxygen has two, four, six, eight, so we're good. And now let's look at our carbon. Two, four, six, eight, so we're good with carbon. Here we didn't have enough electrons to do it. Here we do. So double bonds and triple bonds are a way to give ourselves extra electrons because we're sharing them. So this one's wrong. And this one is the one that we want to be using. Okay, next example. And when you're doing these, you'll notice I've given myself an entire sheet of paper to do these on. That's on purpose. We're going to be using these for the entire rest of the quarter, pretty much, all the way up until at least gas loss. So make sure that you give yourself just as much room as I did, and we're going to be adding to them. When we get to geometries, we're going to go through and we're going to put all of our geometries on. When we get to hybridization, we're going to go through and put all of our hybridization on. We'll do bond angles. We'll do polarity. We're going to just spend a lot of time with these sheets of paper. So one per sheet is about right, any more than that, and you start running out of room near the end of the quarter. Okay, now the next one. And this one gets a little complicated because now we have two central atoms. Now luckily, people usually write these in a way that makes it a little bit easy to see what's going on. So you can kind of see here that we have BH3 and H3, and that's written that way on purpose. So you can sort of see that you have like a BH3 group bonded to an NH3 group. So first, count our electrons. We go to boron. We have three. We go to nitrogen, and we have five. We go to hydrogen. There's six hydrogens, and each one has one electron. So we have 14 electrons to work with. So now our skeletal structure. We can kind of see that boron and nitrogen are the only ones who can form multiple bonds. So they're going to need to be bonded to each other, as well as all the hydrogens. And so at this point, we say, OK, well, how many electrons do we have? 2, 4, 6, 8, 10, 12, 14. So we're good. Does everything have its octet filled, or for hydrogen, it's duet? Well, all the hydrogens are set. And then boron has 2, 4, 6, 8, and nitrogen has 2, 4, 6, 8. And so we're set there. Now later on, we're going to talk more about exceptions. And we'll see that boron's actually one that a lot of times will have less than eight electrons. A lot of times, it'll just have six, and it's fine with that. But here's an example where it actually does what most atoms do, and it has its eight. OK, next one. Now we need to do an ion. So you have some practice doing ions. So we have XCF4, 4 plus. So conopolar electrons. So xenon has eight. It's a noble gas. And then fluorine has seven each. So we have four fluorines, and each one of them has seven. So that would be 36. But it's a four plus charge, which means that we need to take away four electrons to give us 32 electrons. So it's a plus charge. We have less electrons than the neutral atom, and so we minus the four electrons. So we have 32 electrons to work with. So we put xenon in our center. Fluorine never allowed to break the octet rule, and we'll talk more about that in a minute as well. So fluorine is ever only going to have one bond. So whenever you see a hydrogen or a fluorine, only one bond, it can't be a central atom. So when we go through and we draw these out, we know fluorine can't have more than one bond since I just told you. So really all we can do at this point is add in our lone pairs and hope everything works. If not, we'll have to add some things to xenon. So if you count all of these, 2, 4, 6, 8, all the way around, what you see is that we do in fact have 32 electrons. So we're done. We don't have to do anything else with this one. OK. Now our last and most complicated one for right now. This one we won't actually even be able to finish until we talk about some other things. OK. So N2O. So this one's kind of not as simple as the other ones, right? And we'll see why in a minute. So first of all, most electronegative element here is oxygen. So nitrogen's going to be in our center, one of our nitrogens. So we can set up our central atom as being nitrogen. And we can put a nitrogen on each side, or excuse me, a nitrogen on one side, an oxygen on the other one. And then we can draw our skeletal structure. Now if we try to give everything an octet here, first we're going to need to count our electrons. So you'll see that we have five electrons from each nitrogen and we have two nitrogens. We have six electrons from our oxygen leaving us with 16 electrons. If you try to make your octet here with just going and filling in your elements or your electrons, and we just kind of fill in, you notice we can't do it. We don't have enough here, right? We would need four more electrons. So if we're short four electrons, how many bonds do we need to form? We need to form two bonds. Okay. Well, where do we put those bonds? Now we have some options. When we had to form two bonds with that N2 group, we didn't really have any options. So we just sort of put them in between the two nitrogens. Here we have options. So I have two bonds that I need to form. So I could just put one bond on each group. I could put both bonds to the oxygen, or I could put both bonds to the nitrogen. And let's go through and fill in our octets at this point. So we have two here, so we need two more. This nitrogen set, two, four, six, eight. This oxygen needs four more. We count up our electrons, two, four, six, eight, ten, twelve, fourteen, sixteen, we're set. We move down to this one. Now this nitrogen needs six. This nitrogen again is all set. And this oxygen is going to need one. This nitrogen needs two. And this oxygen needs three. We go through and count, two, four, six, eight, ten, twelve, fourteen, sixteen, we're good. Two, four, six, eight, ten, twelve, fourteen, sixteen, we're good. Okay. Well, which one is it? So all of them follow all of the octet rules. All of them have the least electronegative element in the center. So which one do we do? This is where that very last step on the slide comes in. We don't really know how to do that yet. So let's learn how to do that and then we'll come back. So there's an equation for this. And once you start getting good at it, you're going to want to do it more sort of internally and be able to quickly check all of your formal charges. You don't necessarily always want to be filling into this equation. But for the moment, it's a good thing to start with. So your valence electrons. So you start there. So for nitrogen, it's five. Oxygen, six. If you ever forget, you just go to your periodic table and count over until you get to the element where you're at. Now that's your starting point. Then you subtract the number of electrons that you're, the atom that you're talking about in your, whatever you're looking at. What I kind of always think of it is which ones they actually own. So which ones are sort of their electrons? So the way that that's calculated is the number of electrons, the number of electrons that are actually in, that it has all by itself. It's lone pairs. So you're subtracting away all the electrons that are inside the lone pairs. And then half of the bonding electrons. So if you have a bond, one of those, one of those electrons belongs to one element and one of those electrons belongs to the other element. So for accounting purposes, you always think of half of the bonded electrons. And we'll do some examples that'll make that make more sense. So let's go through and assign formal charges to all of the Lewis structures that we did. And then when we get done with them, we'll be able to actually go through that last harder example that we did and decide which one it actually is, which Lewis structure is actually correct. So just sort of starting over on our Lewis structures. Again you'll want to keep them handy and do this quite a bit throughout this next little bit of the quarter. Okay. So our first one. So what is our formal charges on each nitrogen? So we go and we say, well, and formal charges are going to be atom specific. So it's for each atom in the molecule. Now in this case, they're going to be the same, but we can still calculate them for both. So how many electrons does nitrogen have, valence electrons does nitrogen have? You go to the periodic table, you come over and you see that nitrogen has five. Now you have to decide how many are its own, how many does it have in this structure. And you're going to subtract those away. So anything in a lone pair is its own. So it has two from the lone pairs. From the bonds, it only gets one from each bond. So one, two, three. So you have five valence electrons minus five is zero. So this nitrogen's formal charge is zero. Now if you go to the other nitrogen, the exact same thing happens. It has five valence electrons. It has two electrons that are in lone pairs and it gets one from each bond. And so that nitrogen's going to be plus zero as well. Next one. So now we can go through and we can talk about each one individually. So let's start with the oxygen. So how many valence electrons does oxygen have? It has six. And we're going to subtract away anything that are in lone pairs. So one, two, three, four. And any of the bonding electrons but only one from each bond. One, two. So we have six minus six. So the oxygen's charge is zero. Carbon. We go to the periodic table. We say how many valence electrons does it have? And it has four. So we subtract away anything. In this case, there's no lone pairs so we don't have to count any of those. And it gets one from each bond. So one, two, three, four. So we have four electrons there. So its formal charge is zero. Now we need to talk about hydrogen. Hydrogen has one valence electron and gets one from each bond. So its formal charge is zero as well. And that would be true of the other hydrogen as well. So those are all our formal charges for that one. Next one. So now this one when we go through, it's going to be a little bit different. We'll start actually getting some formal charges. So let's start with the hydrogens. Hydrogen gets one valence electron. And it's the same for each of these hydrogens. It'll get one from the bond. So every one of the hydrogens on here follow the same rules and so they're all going to be plus zero. Now let's talk about the boron and the nitrogen. For boron, how many valence electrons does it have? So you go to the periodic table and you see that it has three. Now if you come here and we count electrons, no lone pairs, so don't worry about those, gets one from each bond. One, two, three, four. So boron has a minus one charge. Now nitrogen, if you go to its period, it on the periodic table, you'll see that it has five valence electrons. And you count around one, two, three, four. And so you have a plus one charge on nitrogen. Now the xenon tetrafluoride. So for each of the fluorines, because again they're all going to follow this numerals, so we're only going to do it once, they get seven electrons from the periodic table, valence electrons. Now here each one of them has two, four, six that are in lone pairs. So those are the fluorine lone pairs. It's not sharing them, it's not doing anything with them. So it owns all of those. Now for the bonds, it just gets one. So we have seven valence electrons, minus six electrons that are in lone pairs and one electron that is in the bond. So each of the fluorines are going to have a zero charge. Now we move on to the xenon. It's a noble gas, so it's in that eighth group. Has eight valence electrons. And we subtract one from each bond. One, two, three, four. So we subtract four electrons. So it has a plus four charge. So as we're going through these, you should notice that we're counting electrons very differently depending on what we're doing. When we were counting for octets, we counted everything, right? For xenon we went through and said two, four, six, eight. Xenon has an octet. Fluorine has two, four, six, eight. Fluorine has an octet. That's very different from when we do formal charges. The bonds only count as one for each atom. So fluorine, we only got to count that as one electron. For the sake of octets, you counted as two. Okay, so now we can move on to our last one, which was a bit more complicated. And we weren't actually able to finish. Okay, so we had this. And we said, well, I don't know which of these three to do. Now we can actually go through and we can figure this out. So let's do that. Let's go through and put formal charges on everything. And I'm going to do this a little bit quicker this time and do it a little bit more like how you should be doing it on an exam and how after maybe about a week of practice and a week of homework you should be doing it. So we look at oxygen and we say, how many valence electrons does oxygen have? Well, it has six of them. How many does it have in lone pairs? Four. Okay, so we're down to two electrons at that point and then two here. So oxygen has six valence electrons and here it has one, two, three, four, five, six. So six minus six is zero. Let's actually switch colors just to make it a little bit more obvious. Plus zero. Now nitrogen. Nitrogen has five valence electrons and here we have one, two, three, four. So we have five minus four is equal to one and it's a positive one. Now here nitrogen has four valence electrons and here it has one, two, three, four, five, six. So five minus six is minus one. So again, when you're practicing at home, do it the way that in the beginning, do it the way that I've been doing it up until now. When you get a little bit on a roll with it, this is kind of the fast way to do it. This is how you want to be able to do it for the sake of, you know, checking your work quickly without having to waste too much time. Okay, let's do the next one. Here we have oxygen normally has six valence electrons. Here it has one, two, three, four, five. So six minus five is a plus one. Over here, nitrogen normally has five and here it has one, two, three, four. Five minus four is plus one. Here nitrogen normally again has five but here it has one, two, three, four, five, six, seven. Five minus seven being minus two. Okay. So just going through, putting formal charges on everything and then this is going to help us decide something when we're all done. Down here, oxygen normally has six. Now we have one, two, three, four, five, six, seven. We have minus one. Here nitrogen normally has five. One, two, three, four. So five minus four is one. Over here, nitrogen has five valence electrons and here it has one, two, three, four, five. Five minus five being zero. Now as a good check, we need to go, this is something you kind of want to do in your head for everything. You don't necessarily need to write them down. If I ask you to write down a Lewis structure, you don't have to put informal charges. If you have something that's an ion, actually like this, if you have something like an ion, it's not actually correct unless you put the charge somewhere but you don't have to put the charge on Xenon. You can just put brackets and put the charge and that's fine. That's just as correct. Something like this that doesn't have a charge, you don't actually have to put the formal charges unless I ask for it. You should do it in your head anyways though. Now we have to decide which one's correct. So if we look at these, a good check first is to make sure that all of our formal charges add up to whatever the charge is here. So charge is zero. So this one adds up to be zero. This one adds up to be zero. So we've done our formal charges correct. We've done our Lewis structures correct. Now we have to decide which set of formal charges is the best set of formal charges. And so if we go through and we look, we can say, well, this one has a negative 2, a 1, and a 1. This one has a negative 1, 1, 0, negative 1, 1, 0. This one has some really high formal charges and there's no reason for that. 2, 1, 1 is acceptable if there's no other option. You will have formal charges of plus 4 sometimes. We saw that with Z non. But if there's another option, you don't want to do that. So we're going to get rid of that one right away. That one's not going to be right. So now we have to decide between these two. So they're basically the same. Our central nitrogen has the same. But now here we have a negative one on the nitrogen. Here we have a negative one on the oxygen. So we have to figure out which one's better. So for this, we need to kind of backtrack to what we talked about earlier today or last class, where we were talking about electronegativities. So when we're talking about electronegativities, we're talking about the amount that an atom is willing to pull the electron density toward it inside of a bond. So this is kind of exact in this situation. And something like oxygen is more electronegative than something like nitrogen. So if the electron density is going to be shifted one way or the other, it's going to be shifted so that the more electronegative element, the one that's pulling the electron density toward it, is going to be the negative one. Because well, it's more electronegativity. It has more electronegativity. It's pulling those electrons toward it more. It's going to take the negative charge, not nitrogen. So this one is going to be correct because oxygen is more electronegativity, or more electronegative. So you always, if it's possible, and keep this in mind, if it's possible, you want to put your negative charge on your most electronegative element. That's not always possible. For instance, this one again. In this one, we had a plus. Oh, I guess we didn't have a negative charge. Never mind. So draw out CO for fun sometime. Draw out carbon monoxide. And you'll see that you're forced to do the opposite of this. You don't have a choice. And that's fine. If you have a choice, you don't have a choice. But in something like this, we do. We can choose to put it on the nitrogen by drawing it that way, or we can choose to put it on the oxygen by drawing it that way. So this is the more correct way. OK. So now, moving on a bit. More in detail about breaking the octet rule. I've been sort of alluding to this and talking about it, but now we're going to go into it in a little bit more detail and talk about exactly when we can and can't do this. So your second period elements, they're not going to do this. So that second row of the periodic table, well, and of course, hydrogen and helium as well. Hydrogen, helium, and then lithium up through neon. They're not going to do it. Do not ever have 10 electrons off fluorine or an oxygen or nitrogen. You will just get it wrong. There will be no partial credit for that. You'll just have it wrong. Now third period and lower, you can do that and it's fine. That's not a big deal. There's some that are going to do it more than others, but not all of them. Now the third period element can do it where the second one can't because of the fact that they have some d orbitals sitting right there. So the d orbitals are right there. They're able to be used for bonding. You can use them for bonding. You can use them for hybridization, which we'll talk about later in the next chapter. And so anything below, third row and below can do it. Second period ones do not do it ever. Some of them are going to have a lot less than eight and they don't have to, but they can. Theoretically, anything could have less than eight electrons, but that's not necessarily normal. For beryllium, aluminum, and boron, it's pretty normal. If you see one of these and you can't fill the octet, don't worry about it. Just move on. Anything like carbon and oxygen, you're going to probably fill the octet. Sell them going to break the octet rule. It's not something that you're going to ever see in this class. Where these elements, they do it all the time, so get used to these elements breaking the octet rule on the low side and having less. As far as having more, that falls into the category of sort of a lot of these that I talked about. There's a lot of phosphorus xenon, chlorine does it a fair amount. Anything third row and below can do it, but these are sort of your common ones that actually do it all the time. We'll start and do some examples of this now and then do some more next time. Okay, so we'll start with SF6. Our central one is sulfur, since we can't really do anything else, we can't put fluorine in the center. When we go through and do this, sulfur here, and we have six things, so we just spread them all out and put fluorines on all of them. Because it's a Lewis structure, we need to go through and we need to add in all of our lone pairs. We count up how many electrons we started with just to make sure that everything is still okay. We have six from the sulfur, seven from each fluorine, and we have six of them. We have 48 electrons. If we go through and count all of these up, it adds up to be 48, and so we're all set. Sulfur is an example where it's okay with having six bonds. Usually with sulfur, you want to kind of get in the habit of seeing, well, it'll either have two bonds and fill its octet, or it'll have four bonds, or excuse me, six bonds in this sort of configuration. So it'll either act like oxygen, which is right below on the periodic table, or it'll form six. Okay, next one. So this is kind of similar to one that we've done before, only now it's not an ion. Now it's XCF4 without any 4 plus. So we count up our electrons. We had eight from xenon. We have seven from each fluorine, and we have four of them. So we have 36 electrons. It's not a charge, so we don't have to worry about adding or subtracting any. Now when we go to draw this, this is a good example of one where you can get yourself into trouble if you don't count the electrons first. So up until now, a lot of them have been, once you get really good at them, you may say, well, I don't need to count electrons. I know what it looks like. This is one where you may get yourself into trouble if you don't. So if we go through and we kind of set up our skeletal structure, and we know fluorine only gets to have one bond, so we're just going to fill those in. At this point, you might be tempted to say, okay, everything has an octet. I'm done. And walk away. In which case, you'd get it wrong. And then all the other questions following it that we're going to learn how to do later, like hybridization and geometries and bond angles and things like that. So we look at this and we say, okay, how many electrons have we used up? If we count them all, we've only used up 32. So we have two more, or four more electrons that we have to deal with. So we can't put those electrons on fluorine, because fluorine only gets to have eight. Xenon can have more if need be. So we're going to give them the xenon. So that gives xenon more, it has two, four, six, eight, ten, twelve electrons. So that's okay. So that's one where we're giving xenon twelve electrons. And so we're breaking the octet rule there. We can't break it with fluorine, but xenon can break it. Okay, so now let's do BH3. So we add up our electrons. We say, well, we have three from boron. We have three from each hydrogen, one from each hydrogen. So we have six electrons total. We go to draw this. And we have this. And at this point, there's nothing else we can do. We only have six electrons to work with. Well, if we only have six electrons total, there's absolutely no way to give boron eight. We can't just make two electrons out of malware. So we're done. That's all there is to it. We can't really do anything else. Now as good outside practice on this, go through and calculate all your formal charges. We haven't done any formal charges yet, because it hasn't really come up where we've had an option where we needed to decide between things. But it's a good practice to go through and make sure that you can do it and get in the habit of doing it. And so there's lots of places here for you to practice it. Now a more complicated one. And these ones tend to get people a lot. For all of these polyatomic ions where we turn them into acids, so they're called oxoacids. So in this case, this one would be called what? Sulfuric acid. So, or phosphoric acid or carbonic acid. Any of these, you want to get in the habit of knowing sort of the general way that these look like. Because otherwise you might be tempted to take sulfur and bond one to each oxygen and to the hydrogens. You don't want to do that. Whenever you see these, you need to know that these hydrogens are actually bonded to the oxygen. So when we go to do this, we're going to put sulfur in the center. We're going to put four oxygens around. And then the hydrogens go on the oxygens. And that should just, you should know that as part of your skeletal structure from now on. And if you come to me later and say, well, how was I supposed to know that? I'm going to point you to this slide. So this is how you know it. And this goes for sulfuric acid, carbonic acid, any of those oxoacids. So now let's go through. We can fill our octet in here because there's not really much else we're going to probably do with oxygen. And at this point in practicing, you may get in the habit of saying, okay, well, I know oxygen pretty much always has two, two bonds. So I'm just going to give it two bonds in its lone pairs. Now if we add in our electrons and we try to add 2,4,6 here and 2,4,6 here, we'll see that we can't do it. And so what we need to do, we'll see that we're short for electrons. So go home, practice it, do it. We'll see we're short for electrons, so we need two bonds. And so we're going to put our two bonds here, or excuse me, our two lone pairs here and our two lone pairs here. And that's where we'll finish. Now you'll notice I forgot to add in all of my electrons and make sure that we have the right amount. So let's do that real quickly, just as a double check. If you forget to do it in the beginning, especially when you get good at them, you may get out of the habit and you may forget. Just do it at the end and make sure that you're right. And when you go through and do this, you'll see that, hey, we do have 32 electrons and we are all set. So at this point, we've gone through and we've done a bunch of ones that break the octet rule. We have a bunch more that we're going to do because Lewis structures are really, really important. You will not be able to do all of the rest of the material for this chapter and chapter 3 and chapter 5 if you can't do these. So we're going to do a bunch more examples next time. So we'll just kind of go through them at that point and keep going through and get into all of that. So we'll end there for today though and start up on the rest of them next time.
we're going to remind you about ionic and covalent bonding and our metallic bonding. calculated is the number of electrons, the number next little bit of the quarter. Okay. So our