 I did call some names to give back quiz, some quiz ones and quiz twos, if you haven't gotten them, if you haven't gotten them, maybe I called your name. So come up after class and come and get that stuff from me. Okay, so the last thing we talked about last time was writing Lewis structures of simple compounds, simple covalent compounds. Today we're going to start with drawing Lewis structures of polyatomic ions. Okay, so the way that you do this is fairly similar to the way that you build covalent molecules. So let's just do a little review from last time. Let's try to build CH4. Okay, so most compounds, few rules apply. You want to put the most electronegative atom in the middle and hydrogens and halogens around the outside. Okay, so in this case, of course, we only have a C and four H's. So you're going to put the H's all around the edges and the C in the middle. Okay, the first thing you do is you draw the chemical symbol and then you put your valence electron around it. So if you recall, carbon has four valence electrons. So the way we do this is draw one, two, three, four like that. Again, hydrogen can only make one bond so it has to be on the outside because you can't have a hydrogen atom looking like this because that would be making two bonds. Okay, so all the hydrogen atoms have to be on the outside and conveniently enough there's four of them and this carbon atom needs four more valence electrons to fill its octet for the sharing of the electrons, right? So what I would do is just go ahead, draw hydrogen there, hydrogen there, like that, and then recall the types of arrows I want to make. They're a little bit different than the arrows that are shown in the examples in the book. What I would like you to do is one side of arrows so that electron goes there and combines with that electron. Through this process of sharing, both of the valence shells of each of these atoms is now full. Remember when two electrons are paired up in between two atoms, we call that a bond and we make a new symbol for that bond. It's just a line between those two atoms, so it would be something like that. But in this case we've got one, two, three, four shared pairs, so we're going to go. So does everybody recall how to do that? Okay, so I remember, this is a review, this is what we did last time. So let's take a step forward and try to do, let's do polyatomic ions now. Okay, so let's try to do this first polyatomic ion up there on the left. And in fact, in the book they usually have brackets around these things and show the positive charge as the whole thing has the positive charge. I'd like you to think about which atom, because that positive charge is actually located on a specific atom. I'd like you to start thinking about which atom that positive charge is actually located on. And I'm going to show you exactly how to do that. Through similar rules that I give you on the lecture slides, but I think I'd like to describe it a little differently, so you can get a little different perspective on it. Okay, so for NH4, H4 plus, you're going to do the exact same analysis that you did for CH4. Remember, the H's just have to be, they can't be in the center, they can't be in the center atoms, so they're all around the perimeter, okay? So when we put our atoms down on the piece of paper, the first thing you want to do is write the central atom down, so that's the end. Of course, if you look up at the periodic table, it'll indicate to you that nitrogen has five valence electrons, so we're going to draw our five valence electrons, just like we would for carbon. One, two, three, four, five, and now we need to place our four hydrogens around that nitrogen, just like we did for carbon. Okay, so notice there's a little bit of a difference between this compound here and this compound over here, right? The carbon central atom has four valence electrons, but the nitrogen has five, but this is going to be taken care of because of course you can't have three electrons in a bond, that's going to be our problem here when we're looking at the structure, we're like, okay, well this one's okay, right, we can do that, so on and so forth, but when we get to here, we can't have three electrons in a bond because a bond is two electrons, okay? So what we know about the structure of the ammonium ion is that it's got an overall positive charge on it, so it's less than one, and if you recall, electrons are negatively charged, so what this implies is that one electron has been removed from this group of atoms, okay? So that electron is always going to be removed ironically enough from the most electronegative atom, so the most electronegative atom will always have the charge on it, whether it's positive or negative. So if we look at these five atoms here, and look at our electronegativity chart, which I don't have in front of me, but you'll realize that nitrogen is more electronegative than hydrogen is, okay? So in order to get that positive one charge on there, what we need to do is take one of these electrons on hydrogen and remove it, okay? So just cross it out, or if you prefer, shoot it out over there, okay? It's gone. If we do that, we get a positive charge, so what's going on here is this nitrogen now has, is now positively charged like that, okay? Because we've removed an electron, one of its valence electrons, and now we can bond it like we did with CH4 or methane. And we'll notice we've got the right electrons to make the four bonds that we want, plus now our positive charge, because we've lost that valence electron from the nitrogen, okay? So let's go ahead and draw our structure. I know we're kind of on the edge here. And so where it puts it in overall brackets, I think you guys, I think you guys can do this well enough to be able to indicate which atom actually has the charge, and in this case, of course, you figure that out in the previous step that's going to be the nitrogen, okay? So hopefully everybody understands that. We're going to try another one that's a little bit more difficult. We're going to try this CO3 2 minus 1, okay? So I'll leave that one up there, and we can compare that one to the one we're about to do. Again, you know, I want you guys to start thinking about which atom actually has that charge on it, because it really does help you in thinking about chemistry altogether, you know? So... So on a test, I'll indicate whether I want you to put the charge on the atom that it actually has the charge, okay? I'll definitely indicate to you on the test or in a quiz or anything what I really want, okay? And if it's not indicated, you can ask me and then I'll tell you. Okay, but let's go ahead and compare this to CO3 2 minus. And you've got different atoms that can make multiple bonds. Of course, we don't have any halogens or hydrogens in this compound, okay? Or in this polyatomic ion. So we're not really certain which is the central atom, okay? The rule is that you want to put your least electronegative atom in the center, okay? So if we look at our electronegativity tables, we'll realize that carbon is less electronegative than oxygen, okay? So we're going to put carbon, then we're going to draw the valence electrons around carbon. How many does it have? Four. Around that carbon. Of course, there's only three oxygens, so we can only put it on three sides. And we have to ask ourselves how many valence electrons does oxygen have? Six, right? All three of them, each has six valence electrons. So let's go ahead and draw those. Recall what we said before that the most electronegative atom will have the charge on it, okay? The most electronegative atom. So in this case, we said, well, oxygens are more electronegative than carbon, but there's three oxygens here, okay? So which one's going to have these two charges? Well, what it really comes down to is that no nonmetal will have more than one charge if it's in a covalent compound, okay? So if it's in an ionic compound, it can have more than one charge, like S2 minus or P3 minus, okay? Or O2 minus and 3 minus, okay? But since this is a covalent compound, this is a covalent compound, right? We're making covalent bonds. That means that none of these atoms can have more than a single charge on it, okay? So what that really is indicating to us is that since we've got 2 minus charge, two of these atoms have a single charge on them, okay? So there's three oxygen atoms here. They're all equivalent, so we can just choose two of them to add an electron to, okay? So I'm just going to choose this one here for adding this electron and this one here. So does everybody get what we're doing so far? So now we can see, right, that this one here, right, that's going to fill the octet of this, oh, sorry, let's step back for a second. Remember, electrons are negatively charged, electrons are negatively charged, so when you add the electrons to those particular atoms, they're going to have a negative charge overall, okay? So since we added a negative charge to this oxygen, it's going to have the negative charge there. Does that make sense? And then since we're adding the electron to this oxygen, it also has a negative charge, okay? So that's where your two negatives are. Okay, now let's go back and try to fill these octets, okay? So this oxygen's octet can be filled by combining with that electron from the carbon. This oxygen's octet can also be filled by using that one of carbon's valence electrons. And if you'll notice, this oxygen needs how many more electrons to fill its octet? Hopefully everybody says two, right? Why? Because it has six here. And how many does carbon still need to fill its octet? So it's got one, two, three, four, five, six, right? So it needs two more as well. So what's going to happen is you're going to make what's called a double bond here. And let's just draw a structure over here. So we've got a carbon that has two bonds to this top oxygen. Notice that oxygen still has its two lone pair electrons on it in a single bond to this oxygen here. And this oxygen has one, two, three lone pair of electrons plus a negative charge on it. So we've got to indicate all that. The same is true for this oxygen over here. One single bond, three lone pairs, and a negative charge. So that's the ion that goes CO3 two months. And clearly indicating which atoms have the charge. Okay? Does that make sense to everyone? Hopefully that makes sense. There's any questions we can talk about right now. Okay, just like I can stand on my head, you know, these molecules can flip over, you know. So I could have drawn, so for example, if this was what we were drawing out there, I could draw it like this, or like this, or like this, or like this. And it's the same molecule. Just like if I stand on my head, right, I'm still heat, you know, I'm just standing on my head, okay? Or if I'm turned around, you know, I'm not a different person. I'm still heat, but I'm just turned around. So these molecules, in fact that's what they do is they like spin and spin and spin and spin, you know. So what you'll find is there's different representations of them that emphasize the things that you want to emphasize. So if I wanted to show this kind of edge on, I could turn it like that if I wanted to, or something like that. Okay? You'll get more of a feel for that in a little bit when we start drawing structures. Any other questions about drawing polyatomic ions? Here are some kind of rules that I explained a little differently. And the slides and what I explained it up here, I really prefer the way that I explained it up here, but this might click with you too, okay? So I'm going to let you go through those on your own. And again, here are some covalent compounds. None of these are polyatomic ions, but you can see methane here. The one thing I want to emphasize is that your arrows, I said double-headed arrows like the book does. I wanted my single-headed arrows because that actually depicts the motion of a single electron as opposed to two electrons. But you can see a number of structures that you might want to try on your own, just doing this valence electron sort of analysis. And remember, the valence electron is really the only thing that matters in a bull reactive. There's some more. And then we talked about bond energy last time, how a single bond is longer but weaker than a double bond, which is longer and weaker than a triple bond. Or in other words, a triple bond is much stronger and shorter than a double bond, which in turn is much stronger and shorter than a single bond. So hopefully now, you could draw all of these different Lewis structures. Of course, four and five are the two that we just went over in class. I would suggest you try to draw those. There's some other polyatomic ions that you guys are familiar with or should be familiar with that I've given you as a list to memorize. Try to go over with those and try to build these types of structures because I could only imagine that I would put a couple of these on the next exam. So if you know how to do them, then those are like three points, right? Okay. So we've been drawing, okay, let me erase here. Can I erase? Has everybody got this? So we've been drawing these structures, well, these Lewis structures, if you will. We've been drawing these Lewis structures on the board and the way we're representing them is as if they were like pieces of paper or something, like flat objects that like, so like this, right? So we can see it all like this, but if I were to turn it like this, I don't see anything, right? But that's not the way that molecules are, okay? In fact, let's build methane. So this representation of methane shows its structure more accurately than what the Lewis structure depicts. And I will draw another representation, one is a structural formula that adequately represents this or at least somewhat adequately on a two-dimensional surface. Okay, so hopefully you can see this little plastic thing that's in my hand, okay? This is representing methane just like this is representing methane, CH4, CH4, okay? Notice the difference here, right? Whereas that seems as if it's all in the same plane of the board, right? If you pass this around, if you don't mind, you can see that those four hydrogens aren't actually in the same plane, okay? And in fact, we've represented, so there's a problem with the way that we can represent things on a two-dimensional surface, okay? Because it's two dimensions, it's hard to represent things like a three-dimensional shape. So we have these rules that we have to follow that kind of emphasize different directional types of bonds, okay? So let's draw this as if we were looking at it like that, okay? So the central atom there, and if you'll notice, look at that compound that's going around, you'll see that two of the atoms actually are in the same plane, or actually three of the atoms, but two of the hydrogens are actually in the same plane. So if we look at this, so if I turn it like this, hopefully you can see that this atom, this atom, and this atom, right? This atom, this atom, and this atom are all in the same plane, with this one sticking towards you and this one sticking towards me, okay? Does everybody see that? Okay, so there's one, what we call one going forward and one coming back, that's what we say, okay? Let's draw this representation on the board, okay? So straight lines, like we've been drawing, mean that they're in the plane of the board, like that, okay? So let's draw the other one, like this. Notice, notice what this shows is it's also not at a 90-degree angle, right? So if you look at that, it's a little bit more than a 90-degree angle, right? Specifically, it's 109.5 degrees angle, okay? That's the number you're going to have to know, okay? 109.5. So let's draw that second one, like that. In fact, what you find is all the bond angles, if you will, okay? So now let's represent the going forward bond and the going backwards bond, okay? So the forward bond is represented as a wedge, as a dashed line or a hashed line. If you've got this in your hand, you might want to look at this and look at the picture that we've drawn on the board and try to see how we're actually representing them. All of the bond angles are 109.5. This one, wedge, no wedge, wedge bond. That means it's going forward. So why is this? Why do these bonds go into these particular bond angles in 109.5 and not 90 degrees, like the Lewis structure is telling us? Well, it's because of this idea, of course, electrons are negatively charged, right? If I'm an electron, I'm negatively charged. If there's another electron, that electron is also negatively charged, okay? Well, there's this theory, vet for theory, and it works pretty well, that describes the shapes of these bonds, of these molecules of the bond angles. That says, well, since there's electrons in a bond and electrons in the other bond, well, they don't like to be around each other because they repulse each other. So they try to get as far away from each other as possible, okay? So if we imagine an atom to be a sphere, you guys are going to have around the central atom as four, okay? Most bonds you'll have around the central atom as four. But if you imagine an atom to be a sphere and not like a box, kind of like what we represented over there, you'll realize that the furthest apart these bonds could be is not 90 degrees, but 109.5 degrees. So instead of sticking here and here and here and here, right? They can get a little farther away, right? 109.5 is bigger than 90 degrees, okay? So that's what's going on, okay? And that's all vet for theory is saying. So let's go ahead and read some of these slides out to you and maybe I can explain a little bit better what we're talking about. So anyways, molecular shape plays a large part in determining the properties and the shape of a molecule. This is known as vest for theory, which means valence, shell, electron, pair, repulsion, okay? Just like the name implies. It's the valence, shell, electrons that are repulsing each other, okay? That's why they're far apart away as possible. They repulse each other because they're negatively charged and don't like to be around each other. So this is used to predict the shape of molecules and all electrons around the atom, central atom, arrange themselves so they can be as far away from each other as possible to minimize this electron repulsion. So here's three shapes that you want to become familiar with. This one here at the end is the tetrahedron, tetrahedral structure. And in fact, that's the structure of this molecule here. There's a couple other ones. This one's linear. This one, it says triangular, but what it really should be saying is trigonal planar, okay? I don't know how many trigonal planar molecules we'll come into contact with in this class, but definitely linear molecules we will, okay? And then there's another one, a couple other ones that I'd like you to know. Let's see if they come up soon. But for right now, let's just keep going along with these ones. So in a covalent bond, the bonding electrons are localized around the nucleus, and the covalent bond is directional, right? Just like these. They have a specific orientation in space, specifically 109.5 degrees away from each other in this particular molecule. So ionic bonds, they're inherently different. They're these electrostatic forces that are sticking these, like, kind of spheres together, okay? And they have no specific orientation. They just go in a three-dimensional array from the center, as far as you can imagine. So, again, remember the terms bonding pair, two electrons shared by an atom, non-bonding pair or lone pair are two electrons belonging to one atom, so the pair is not shared, okay? And the maximal separation between electron pairs is going to be four bonds, and that's the tetrahedron, okay? So here's methane, this molecule that we've just described up here. There are four shared electron pairs around the central carbon, right? So hopefully everybody can see that, right? One, two, three, four. Or eight electrons shared around that central carbon, right? Four electron pairs. Minimal repulsion is when the electrons are placed at the four corners of the tetrahedron, okay? This structure here, like this square pyramid, that's the tetrahedron, okay? So when the place in that orientation, they're as far away from each other as possible, okay? And in that case, you get to be 109.5 degrees. All of these, of course, are going to be showing a full octet. Remember the wedge, the hash bond and the straight lines, and what they mean. So let's consider a new, a different type of molecule, a molecule that shapes a little differently but still a common molecule, ammonia, is very, very similar to methane structure-wise or electronic structure-wise. But when you look at the actual structure, it's a little bit different, okay? Too significant to show you the difference. Ammonia, and there's methane, okay? It looks like ammonia just got kind of, it's top cut off, right? They look kind of the same, there's just the top cut off. Okay, so this is a new structure that you're going to have to memorize the name to. It's called trigonal pyramidal, okay? Because it looks like a pyramid, right? Kind of like a pyramid with three sides or three points on its face, okay? So the three kind of like a little pyramid there. But trigonal, of course, implies three, okay? So let's look at the electronic structure of ammonia and compare it to that of methane and then indicate to ourselves why ammonia and methane actually looked different, okay? So the first thing we want to do is build the Lewis structure of ammonia. How do we do that? Well, hydrogens must be on the outside because they only can make one bond. Nitrogen must be the central atom. Nitrogen has five valence electrons, like that. So it only needs three more valence electrons to fill its octet. So let's put those three in the form of hydrogen plus its electrons, like that. Let's build it. The structure of this would be nitrogen with its long pair because nitrogen still has that long pair that hasn't done anything. Then hydrogen, hydrogen, hydrogen. Notice again these bond angles aren't correct, right? So we're going to have to depict this in such a way that represents this structure here because this is what ammonia looks like, not like a T. Okay? It doesn't look like a T. So let's draw this in the same sort of fashion that we drew methane. Remember, the hashed lines are going back, the wedge lines are going forward, and the straight lines are in the plane, right? So if we look at this, how many straight lines, or how many forward bonds are there? Just one. How many back bonds are there? One. And how many straight bonds are there? One, right? So let's draw that. So we got a nitrogen jet with one straight bond there. We can put its long pair. One wedge bond. That implies the trigonal pyramidal structure. The bond angle, I don't know, does it give you that effect? 107 degrees. It's actually 107.3 degrees, that bond angle there. Okay. So why is it 107.3 degrees if there's only three bonds and we're thinking of this thing as a sphere, right? And those electrons are repulsing each other. Why isn't it 120 degrees? That's what you would expect, right? If they were as far away from each other as possible, it should be something like this, right? As far away from each other as possible. Even on a sphere, 120 degrees is bigger, of course, than 107.3, right? So you know you can make that structure. Why isn't that the case? Yeah. Well, it's because you got the long pair there, right? This long pair is still up there at the top. But if we were to take a picture of this molecule, right? We couldn't see it because we can't see electrons, right? We could only see the atoms, okay? So the structure of this molecule, so the physical structure of the molecule is trigonal pyramidal, but the electronic structure is still tetrahedral, okay? But if it's tetrahedral still, then why the smaller bond angle than 109.5? It's because the long pair electrons aren't like kind of stretched in space, kind of taught if you can imagine a difference between like a rope being stretched by two people and a rope just kind of hanging, you know? That hanging rope is going to have more kind of area than the stretched type rope. So what happens is this long pair is actually all around here. We think of electrons as little particles, but they're kind of like, I don't know, some weird kind of force type thing, okay? So you can kind of think of them as like a force field type thing that's associated with all of this area here, okay? So when those electrons kind of are long pairs, they're not stretched to be right in the middle, right? So they kind of just wander around, and that kind of makes these other bonds here get away from them even more, okay? So this is like the contraction due to the long pair electrons up here. So long pairs need a lot of space, yeah? Not necessarily, but probably, okay? Almost always, but not necessarily. And again, always depends on the way I'm looking at this molecule. So let's try to draw, let's draw from this way. It's like this. Like I can draw like, I can stand like this, right? And all my hands and legs will be in the plank, right? But if I stand like this, then they won't be. But I can stand like this, and one of them would be and one of them wouldn't be, you know? So these molecules can kind of do the same thing, right? We could draw it, let's draw it to where we're looking at it like this. So how many bonds are going to be going forward if we're looking at zero, right? And how many are in the plank? Zero, two, right? So all of them are actually going to be hashed bonds going back. Let's draw that kind of like this, what we've drawn here. It's kind of looking, it's like we're looking from down on top of it so when we do that, we like that. And then the lone pair would kind of be sticking out like here. Yeah, you'd have to wedge your lone pair, which is not very easy. So the reason why we draw them in the orientation we normally do is because that lone pair is not very easy to show the wedge coming out of the plane of the board. You can imagine it would be something like, or something like that, you know? Coming out. But it's not very easy to depict on a two-dimensional surface. So can everybody see the difference between that molecule and that molecule and why the bond angles are the way they are? Okay, cool. So let's go down to the next one, which would be oxygen and two hydrogens. Oh, I guess let's read what this slide says real quick. So consider ammonia, there are four electron pairs around the central atom, three shared pairs, and one lone pair. The lone pair is more electronegative with a greater electron repulsion. So it takes up more area. The lone pair takes one of the corners of the tetrahedron without being visible, and thus it distorts the arrangement of the other electron pairs, giving ammonia a trigonal pyramidal structure with a 107.3 degree bond angle. So let's go down to water now. So notice what we're doing. We're going from carbon to nitrogen to oxygen, okay? So this will be the last thing we talked about today. So I'm going to go through this a little more quickly because we only got three minutes. Next one would be OH2, right? Oxygen has to be the central atom. One, two, three, four, five. I'm just going to draw it like this. Make our bonds. You might think that linear, 180 degrees would be the best orientation for these bonds, but of course there's these two lone pairs that we have to contend with. So in actuality, water looks like this, okay? Kind of looks like a little boomerang. Okay? So in fact we call this the bent structure, okay? Or angular, you might hear it a little, but bent is far more common. So what we would show in this structure is the oxygen being the central atom. Of course it's got these two lone pairs that are like big bunny ears up here, right? Like that and like that. And then, so in order to show those in the plane, right, we're going to have to put it like this. And then that means that there's going to be a wedge and a dash, okay? So if we're emphasizing those kind of bunny ears, we want to draw it like this. And the bond angle here is actually, would you imagine it to be bigger or smaller than 107.3 degrees? Smaller, right? And it is about three degrees, 104.5 degrees. These things here, remember these are called Lewis structures. Things here, all of these guys, these are called structural formulas. If you weren't going to emphasize those lone pair electrons, you could draw water. And this is the most normal way of drawing water, just with those two bonds in the plane, just showing it like this, okay? Not emphasizing those lone pair electrons. Okay? Does everybody see that? Okay, cool. Thank you guys. Bye.