 So, let's start where we ended up last time. Like I said, hopefully I'll get the test back to you guys, for those of you guys who are in the on-campus class. If you're on the online class, you can be guaranteed that your test scores will be put online by Friday. But hopefully for the on-campus class I'll be able to return your tests on Wednesday. So, that would be nice. That being said, let's get started from where we ended up last time. So, the last thing we were talking about was Vesper theory. And we were talking about essentially three different molecules. Let's see if I still have one built here. Methane, ammonia, and water. Passes around in a second. Remember, the electronic structure of methane, ammonia, and water are very similar to each other. They all have four separate pairs of electrons around the central atom. So, with methane, all four of those pairs are locked up in chemical bonds or covalent bonds. So that has four pairs of electrons around it. It has four pairs of electrons around it, right? This one has three chemical bonds, one long pair. So, that also has four pairs? That long pair of things has messed me up on that. I would count them and I was wrong. I would say it had two long pairs of electrons and it had two long pairs. For some reason, it's going to tell me I'm wrong that it had this many long pairs. Bring it to me and we'll check it out and we'll figure out what I was saying. So long, maybe long pairs and now they're saying it's like individual electrons. So, maybe they're saying you've got four long pair of electrons as opposed to two long pairs. So, you do have two long pairs, but if you counted up the number of long pair of electrons, I know it's kind of a nuanced way of saying things, but maybe they're saying, well, you've got four of them. And four bonding electrons here as opposed to two bonds. But either way, bring it to me after class or whatever and we'll look at it. But anyways, we see central atoms, central atoms, central atoms. Around all those central atoms, there's eight electrons or four pairs of electrons. Some of them are tied up in bonds, others are tied up in lone pairs like what we were talking about. But if you look at the structures of all of these, they're all different. This is due to the fact that if you want to think of it this way, if you were like a little guy or gal that was small enough to be able to see an individual molecule, you would be able to see the different atoms that make up the molecule, but you wouldn't be able to see the lone pair electrons. So we would be able to see something like this for methane, something like this for nitrogen with the top cut off there, no lone pair at the top because we can't see it, and something like this for oxygen without those two lone pairs. But notice all the bond angles are very similar. The bond angles are due to the electronic structure of the molecule. So why don't we test those three guys around, and you can look at them and see them. Remember the bond angle here is the number I want you to remember. 109.5 degrees. The bond angle here, you guys remember that one? So they're all very similar. They're really close to about 109.5. And what you find here is that when you add a lone pair, it decreases the bond angle. Add a second lone pair decreases the bond angle even more. This is due to the fact that these lone pairs aren't really kind of as heavily pulled as these bonding electrons. So they're kind of more like a nebulous or ameval-like. So they're kind of more just like spread out. So these ones, since all electrons are negative charged, they don't like to be around each other. So they kind of squeeze down. So that's what's happening in the structures of those various molecules. So let's move on. Let's consider another type of molecule, a linear molecule. So notice the difference between, well, let's build CO2 first, and then we can talk about it. So let's build carbon dioxide. So we recall to build molecules. We got to build them from using their Lewis structures first. So we got to ask ourselves, well, which is going to be the central, which is going to be the outer atoms? How do we figure out which is going to be the central? Okay, well, you can look at it that way, but the rule is that the least electronegative atom is going to be in the center. So carbon, if you remember, and you'll always have a table of electronegativity values just like you have a periodic table that's given to you on the test. So from now on, you'll be able to look at the table of electronegativities and ask yourselves, well, which one's more electronegative, oxygen or carbon? And when you do that, you'll figure out that carbon is less electronegative, so it's going to be in the middle. So we'll write the central atom first. Then we write the Lewis stop structure. Remember, for carbon, it's one, two, three, four, like that. And then we have the two oxygens. Well, where do they go relative to the carbon? Of course, on the outside. And they're exactly the same. So you don't have to say, well, one oxygen, oxygen one has to go here, or oxygen two has to go here. You can put them either way. So oxygen and oxygen. And then what we find is oxygen, of course, has six day-once electrons. One, two, three, four, five, six. Then the other oxygen also has six day-once electrons, like that. And if we look, carbon needs four more day-once electrons to fully fill its outer day-once shell, and oxygen needs two. So, of course, they're going to combine in a fashion that we're already familiar with, making two pairs of double bonds between them. Notice the two pairs of arrows here, the two pairs of arrows here, indicating that we're making two double bonds. If we look, we've still got one, two, three, four day-once electrons around each of the oxygens. So we'll put those. Hopefully everybody's able to build carbon dioxide, compared to those that we have passing around right now. Let's look at carbon dioxide. Here's carbon dioxide here. Notice when we look here, when we look at methane and compare it to carbon dioxide, in methane we say we've got four bonds, right? Four bonds from the C to the H's. In carbon dioxide, we actually can say we only have two bonds, two double bonds, okay? Even though those double bonds are composed of four electrons each, okay? So, a double bond you can think of as two single bonds that are kind of overlapping with each other, just kind of like what this structure is depicting here, okay? So if you don't mind just a second. So if you can see, right, kind of the two double bonds overlapping with each other. So, remember here we have these electrons trying to get as far apart from each other as possible along this sphere of carbon atom, right? So remember a carbon atom is a spherical thing. So those electrons are trying to get as far apart from each other as possible. So they can do that by going to 109.5 degrees. But with the double bonds, since we only got two bonds here, the furthest they can get apart from each other is a total of 180 degrees, okay? So that's what that structure is showing there, right? So on a sphere, the furthest away you can get from each other is to be on opposite sides, right? Just kind of like the poles on the earth or something like that. So this angle is going to be due to the fact that these electrons don't like each other, so they want to get as far apart from each other as possible. So you're going to have a 180 degree angle there. And this is known as a linear molecule. So this is going to be the most common way a linear molecule is formed when you've got these two bonds on opposite sides. So that's what I was just talking about. So if you look here, right, with the band structure, you've got the four electrons, right? Yeah, so the electrons are kind of trying to get as far away from each other as possible. So when we're looking at the linear structure here, we've only got these two bonds, right, or these two areas of electron density if you want to think about it that way, it might make more sense, right? So since you've got only two areas of electron density here and here, right, they're trying to get as far away from each other as possible. But when you've got, look at this or this or this, right, we've got these four areas of electron density that are trying to get away from each other as far as possible, okay? So like I was saying, usually you'll find that when you've got these kind of two double bonds around the central atom, it's going to be a linear molecule. If you've got two single bonds with the lone pairs, it's going to be a band molecule, okay? So you really want to think of things as they're, if you will, electronic structure or this regions of electron density, and it will really help you out, okay? Let's look at one more type of molecule. This one is going to be a trigonal planar molecule, trigonal because you see that it's got three bonds, okay, three bonds, and planar because all the atoms are in the same plane, okay? So trigonal planar. Notice one of those three bonds is a double bond, okay? This molecule here is called formaldehyde. I'll keep it for a second. So remember, this is a tetrahedral. This is trigonal. I'm going to look at a trigonal planar. Okay, can I erase this stuff? Or CO2? A molecule is known as formaldehyde. This is the stuff you stick dead things in and it keeps them preserved or whatever. So if I were to ask you to build the structure, or the Lewis structure of formaldehyde and show me the structural formula, you would have to decide, well, what's the central atom first? So hopefully you would decide that it was carbon, right? Because the other thing you want to know is that carbon makes the most bonds, okay? Anything in group 4 or 4A can make up to four bonds out of all the non-metals. So usually what you'll find is that carbon is almost always the central atom when you got that as an option. Of course, hydrogens can only make one bond so you know they're around the perimeter. Oxygen, of course, has to be bonded to that carbon. So let's just go ahead and put the carbon in first. Draw the Lewis structure of carbon. One, two, three, four. Then we'll put the oxygen. Five, six. Imagine that you could try to bond this in a number of ways. Maybe the H's to the O, H's to the C, O to the C, so on and so forth. But what you'll find is I suggest that you try to do it that way. What you'll find is that there's only one way to bond these atoms together and we'll just go over that. You can go over on the others on your own. Of course, the carbon's going to bond to the oxygen. So notice we've got a single bond being formed between that carbon and that hydrogen. Single bond being formed between that carbon and that hydrogen and two single bonds being formed between that carbon and that oxygen, so that would be considered a double bond. Let's draw the Lewis structure of that and what we just drew. So notice the Lewis structure and the structural formula are going to be different. Why is that? Because if you think about that in like what we were talking about earlier in terms of areas of electron density, we notice around this carbon atom here, we've got one, two, three areas of electron density. So in order to get as far away from each other as possible, they're going to be in a 120 degree angle. Who's got, oh, I still got it. So notice at the 120 degree angle that they're at. So this is 120 degrees. The furthest you can be away from each other. On a sphere, if you've got three spokes or whatever, if you will, three bonds is 120 degrees. Where are the other ones that are out there? You have all of them? Like I said, this is the Lewis structure, not the structural formula. If we were to draw the structural formula, it would be those hydrogen. I guess here we're just going over again the basic procedure to determine the shape of the molecule. First, you're going to write the Lewis structure. I don't know, this is another way of determining it. Not really the way that I taught you, but you can read over this, and if it makes more sense to you to do it that way, you can do it this way. I think honestly drawing the Lewis structures of the particular atoms first and then putting the arrows together really is the easiest way to do it. Yeah, this is trigonal planar. Again, here's a number, these are the different structures, the molecular structures I'd like you to remember, plus trigonal planar. I don't think trigonal planar is listed on this plot here. So make sure you put something that's trigonal planar in there. Again, you can think of them as areas of electron density. This has 4, 4, 4. This is really a good way to remember it. So let's erase this and put that trigonal planar up there and compare them this way. Electron density, we notice we've got 4, 4, and 4. So all those have 4. And notice their bond angles are really close to each other. So 4, you've got 4 is about 109.5 degrees, because that's ideal. Because 703 is about 109. 104 is about 109, especially when you compare it to 180 and 120. So these are the three numbers you really want to remember. One on 90, 180, and 120. One you've got. That one is also known as the angular. Angular, yeah. I was looking for it on here. Yeah, I'll usually call it van. Because that's easier to say. Get van, or whatever. It works. Yeah, so just remember that. So we go from 4, then we go to 3, 3 areas of electron density. And we go up from 109.5 to about 120 degrees. And then when you go to 2, you go up again from 120 to 180. So this is kind of the way you want to remember how to do it. Okay, so if we look, of course, not all molecules just consist of a couple of atoms. Let's see if I can find a molecule that's not enough. So this molecule here consists of more than just a couple of atoms. In fact, it consists of about a dozen atoms. The molecules didn't consist of a lot more atoms. But this is supposed to be like DNA strands stuck together. You can imagine that all the reds are oxygens, the blues are nitrogen, the blacks are carbons, the whites are hydrogens. So you notice that you can't just say this is a bent structure, right? Because you can see kind of part of it's bent there. You can see the tetrahedral is there, a little bit of the tetrahedral, so for those of you who are watching at home, right? Tetrahedral part, bent part, more tetrahedral stuff, kind of a trigonal pyramidal, okay? So there's the bigger and bigger type molecules you get. So more and more variation you're going to have at different particular centers, okay? So let's build a molecule that instead of just has one central atom, let's build a molecule that has three central atoms, okay? And that's the molecule that's being shown on the slides right now. Can I erase all of this stuff here? Hopefully. Say no. Fast. Okay, so, really like I was saying, it really helps for you to think. Four areas of electron density, three areas of electron density, two. And that will really help you out in trying to figure out where, what the structure is around that central portion. So let's look at this molecule. It's called, it's up there. It's called dimethyl ether, okay? We can try another one, diethyl ether. The first diethyl ether would be the first anesthetic that was used after, you know, either knocking somebody out or, you know, making them drink a lot of booze, right? Before you, like, pull their tooth out or something like that. The next one would be, like, diethyl ether to kind of get you, if you smell it, it gets you really loopy. It's got a structure that's kind of like that one. In fact, it's got two more carbon atoms on it. But let's look at dimethyl ether to begin with. Notice the central oxygen atom has two bonds and two lone pairs. So this looks, though, let's draw the other ones. This looks very similar if we're looking just at that central oxygen, like water, right? So if we take these CH3s, these are called methyl groups, because methane is CH4. So methyl just means you took off a hydrogen and stuck something else on it, okay? So if we look at just this central portion, it looks very similar to water. Does everybody agree with me? That looks very similar, right? So what you would imagine, and if you did, you would be correct that the central region of this molecule, dimethyl ether, looks structurally very similar to water, okay? And in fact, we would, if I were to ask you what's the angle here, hopefully you would say 104.5, because you would be right if you did, okay? How do you know it's 104.5? Well, you got there the four electron groups around that central oxygen, right? One, two, three, four, and two of those electron, areas of electron density are long pair electrons. So they're going to force that ideal 109.5 degree angle downwards a little bit, making it 104.5, okay? So the central area around this oxygen atom is actually a van region, okay? So let's kind of look at that. So look what we're doing here. We're taking water, right? Now we'll take one of those hydrogens off and put a methyl group on, and that's essentially what we got there, right? Okay, let's take the other hydrogen off, put another methyl group on, if I can find it, but you almost done with all that stuff. Hey, can I just see water? I've got water. So if we look here, right, this guy is that, okay, and this is dimethyl ether, right? Everybody hopefully can agree with me. Notice the similarity between that central oxygen, right? Just to remove these hydrogens and put what we call these methyl groups there. So the central oxygen there is van, and we looked at the carbon here. Hopefully you would say, well that carbon's got four regions of electron density around it, or four electron groups. We have a more lone pair, so the ideal bond angle would be 109.5. So that's the bond angle there, 109.5 degrees. And if you were to guess what the molecular geometry around that central atom was, hopefully you would say tetrahedral, right? Because it's a problem. Tetra means four, yeah, tetra means four, yeah. Yeah, so like a common name for four-legged creatures is like tetrapods, like horses and dogs, you know. Well, now you know. Pod means foot, you know. So if we're to analyze in more detail this structure, notice we've got three areas of structural identification to do, okay? So we would say, well around this central atom here, this first carbon, you've got a tetrahedral structure, right? With the bond angle being the ideal bond angle of 109.5 degrees. Then if we move to the central oxygen here, well since it's got the two lone pairs and two bonds, well we know that's 104.5 degrees because those lone pairs are squishing it down a little bit, making this event structure. And then we come over to this second carbon and ask ourselves, so what's the molecular structure around that? Well since it's got four bonds, it's also got to be 109.5 degrees, okay? Which is the tetrahedral structure, okay? So hopefully, I wouldn't be at, so some, of course, molecules sort of get bigger and bigger and bigger and bigger, right? So it's very hard to figure out, well this has to be combined in this particular fashion, okay? So once you get bigger and bigger molecules, you can actually start combining them in different ways that the molecular formula doesn't give you enough information to show which way the atoms are actually attached to each other. So on there, are you saying that each angle in between... So this angle here would be 109.5. 109.5, what's that angle? That should be 109.5. 109.5? Just want to make sure that I'm... 109.5, this one, 109.5, 109.5, 109.5, 109.5. Just because another thing that's closing in on the fact of the way that they've done it. Yeah, well it's a perspective. Again, you know, if I could draw on the board this thing here, then I would, you know? But I'm not like some sort of, you know, Picasso or... Well, probably Picasso would draw a very interesting way of looking at it, but whatever some artist that might be able to draw things in reality. Okay, so I'd like you to try on your own to determine the molecular geometry of these particular structures, okay? These would be good ones for you to try for the test, or try the study for the test, and... So let's talk about polarity now. We've talked kind of about it already, but let's go into more depth about it. So let's look at, say, water, for example. This molecule up here. So we know the structure of water. We know why the structure of water is bent. Now, let's look... Well, we've talked about bonds. Bonds can be polar, right? If the electronegativity is different between the two atoms that are in the bond, okay? What you'll find is the electronegativity, of course, is hydrogen and oxygen are quite different, okay? I think hydrogen is 2.4 and oxygen is 3.1, if I'm not mistaken. But there's something like that, or oxygen is 3.5. But oxygen is very, very electronegative. Hydrogen is not very electronegative at all. So what you find is that when you look here, you've got a delta minus up there and a delta plus there and a delta plus there, okay? So this molecule is a polar molecule, okay? It's due to the fact that it's like this little magnet that's got these oppositely charged regions of it, okay? So if it were to interact with another molecule of water, in fact, it would interact in such a way that would make these two molecules orient their positive and negative ends in such a way that they would want to, where there's the delta positive, the delta negative wants to be delta positive, delta negative. And in fact, the way that they would actually orient themselves is that these long pair of electrons would be kind of associating with that electron-poor hydrogen. Remember, these electrons here are being shared but not equally, right? So more electron density is here than there is here. So that hydrogen kind of wants to have a little more electron density. So what it'll do is, since these long pairs are so, these bunny ears huge things, right, they'll kind of associate themselves close to them, okay? And in fact, if you looked at the crystal structure of ice, it would look very similar to this with another, with the, of course this is delta minus, delta plus, delta minus interaction. These interactions here, specifically when you've got hydrogen interacting with the three most electronegative elements, N, O, and S. Of course, this is not a bond, right? It's an interaction because a bond is when you actually have them stuffed together, okay? These are in different molecules so it's not a bond, it's an interaction. But unfortunately, it's named incorrectly, it's called a hydrogen bond, okay? So this thing, this interaction is called a hydrogen bond. It occurs when the very electron-pore hydrogen that's attached to one of these three atoms, nitrogen, oxygen, or fluorine is kind of interacting with another one of these three atoms. Oxygen, oxygen, or fluorine, you get this kind of, what we say, a partial kind of connection, okay? And this is known as a hydrogen bond. So if you will, a misnomer, it's misnamed as a bond. But it's because they came up with this name before they really understood the whole concept of bonding, okay? So, oh, there is one thing I wanted to show you, the polarity of the molecules, some demonstration I'll have to wait until Wednesday to do it, but it's pretty interesting. I'll have to remember, it's really a school. Okay, so remember, let's go back over what we just talked about. Electronegativity, the most electronegative elements are found in the upper right corner of the periodic table. The least electronegative are in the lower left. Look at the electronegativity, the difference between hydrogen and oxygen, 2.1 to 3.5. So 2.1 here, 3.5 here. So that means that you're going to have bar negative charge associated with the oxygen. Excuse me, and then the hydrogen. Okay, so if we look at HF here, this is a polar covalent bond. We call it a polar covalent bond because, of course, it's a covalent bond, sharing of the two electrons. It's polar because the electronegativity difference of the two atoms that are sharing those electrons is, there is a difference, okay? So again, when you have a polar covalent bond, what happens is the molecule acts like a little magnet, giving one part of the molecule delta minus and the other part of delta plus. We can represent that in a couple of ways. Like we've been representing it up here with the delta minus and the delta plus, like that. Or we can represent it with this arrow here. This arrow indicates the same thing as delta minus and delta plus, but it makes it a lot easier to draw than delta, delta, delta. So when you draw the arrow, of course, arrow means negative charge in chemistry. Remember, electron arrows are like that. So when you draw the arrow, you put a little positive sign on one hand of it and put the arrow going towards the negative sign. So in water, it would be like this. So that would show the polarity of the two bonds. Okay, notice that the little cross shows the positive side of the bond. That's the less electronegative atom and the arrow goes towards the more electronegative atom. You can look at this, same stuff we've been talking about. Okay, and then you can ask yourself, which would be more polar? How do I determine that? Well, I look at the table of electronegativity. Hf is 4, Cl is 3, H, of course, is 2.1. So the polarity of the Hf bond is more than the polarity of the HCl bond. You can see the polarity in water here. Okay, this is a different sort of depiction of it. Blue indicating positive charge, red indicating negative charge. Yeah, the center of the partial positive charge is midway between those two atoms there. Notice here you've got a molecule that has two polar covalent bonds. Okay, but they're opposite of each other, exactly opposite of each other and they're the exact same magnitude since they're between carbon and oxygen, both of them. So what they do is cancel each other out. Okay, so this actually, even though it's got two polar covalent bonds, is a nonpolar molecule. Okay, so it's a nonpolar molecule even though it's got two polar covalent bonds. Okay, and then you can see here, so polar covalent bond, nonpolar bond is zero change of electronegativity of course, polar covalent is all the way up to 2.1. That's the HF bond, or a little bit more than that. And then after 2.1 you go to ionic. Okay, then you get straight up electron transfer because the electronegativity is so great. The difference is that one just rips it off the other one. It doesn't share it anymore. Okay. Yeah, we could talk about more intermolecular forces next time. The one thing I want to say is that you should be expecting a quiz sometime soon. Probably we'll do it either Friday or Monday. Okay, so Friday or Monday. I'll tell you more about it on Wednesday. I just want to cover parts. What was that? So it'll cover everything to what we talked about before the quiz. So if it's on Friday, then it'll be everything that we covered up, so on Wednesday. So probably the chapter 4 is what it'll be. Chapter 4, we can do a little bit of chapter 5 and it'll be that one. So I'm saying that stuff from chapter 1 could be on there also. No, well, potentially, you need to know everything with chemistry to keep going. So I'm just saying that's a possibility. Okay, let's turn this video off.