OK. All right. So, what's our first step when we're figuring out VSEPR geometry? What do we have to figure out? Steric number, right? We need to figure out our steric number and one other sort of number.

What is that? Coordination number. Good. So let's look at XeF4. So we drew this Lewis structure out before. We have it set up. Now what is our steric number here? Remember what steric number is. We're counting bonds or I should say things that are bonded to the central atom or whichever

atom we're talking about along with lone pairs. So our steric number for this would be what? Well how many things do we have bonded to the central atom? Four. And how many lone pairs do we have? Two. So our steric number would be?

Six. Good. Now, what is our coordination number? That's the number of things that are bonded to it, right? So it is? Four. Good. So since we have a steric number of six, what does that mean our electron geometry would be?

And remember, we're calling that EG, so just so you remember, that's kind of my shortcut for it. Six things bonded to it, so that makes it octahedral, right? So that's our electron geometry.

That's acting as if we can kind of see the electrons. We're counting that as part of our geometry. Now remember, even though there's six things bonded to it, we call it octahedral because that's the three-dimensional shape that it makes, all right? OK. So then we have to move on to what kind of geometry? Molecular geometry.

And remember my shortcut for that has just been MG for the sake of, you know, speed and all that. OK. So what shape do we have there then? So you're picturing sort of six things coming off a central atom and we take away one of them. And then do we take away the one directly across from it or one of the side ones?

Directly across from it, right? There seems to be a lot of debate about that one, so let's just make this thing real fast. The level of debate to it is high enough I would like to make it.

We're going to do it shortcut way this time since we've already done it out with the molecules. So octahedral shape, six things bonded to it, right? So it's like three sticks straight across from each other. Now, two of those are electron pairs though.

So let's take away one. Let's declare this a lone pair. Which one's the other one are we going to take? Are you going to take one of these pointing at you guys or this one? The top one, right? And why do we do that? We want the electrons to push on each other equally, right? We want the electrons to be as far away as part as possible. So if we took one from here and then one from here, these two electron pairs would

be next to each other. Which is OK if we don't have a choice. But here, we can put one here and one here. So what shape is that? Square planar, right? A square base and it's all on one plane. So square planar.

Next one. So we have boron trihydride.

So we have our central atom and we have three things bonded to it. So again, sort of taking the shortcut with just sticks instead of the full molecules. We have this shape. So what is that? Trigonal planar, right? So first of all, what is our steric number?

Three. What is our coordination number? Three. Is our molecular and electron geometry the same or different? Same. And it's trigonal planar.

OK. So as you're going through your notes, actually we can do this one.

So now this one. So we have a few different things that we can talk about on this one. So let's start with the sulfur geometry. What is the steric number on sulfur? Be really careful.

Is it four or is it six? It's four. Good. Right? Remember, we are not counting the number of bonds. We are counting the number of atoms that are directly bonded to it. So it's one, two, three, four. So the steric number and the coordination number are the same and they're four. So what geometry does that make?

This one actually kind of matches with the prefix you're used to, right? Catrihedral? Now is that electron geometry, molecular geometry or both for the sulfur? It's both, right? There's no lone pair so they're the same.

Now I wrote sulfur geometry because I also want to talk about the oxygen's geometry. So I'm going to talk about this oxygen and this oxygen. They have the same geometry.

OK? These are just bonded to one thing so we don't talk about the geometry there. But what about this oxygen? What is the steric number? Good. The steric number is four, right? We have two lone pairs and two things bonded to it. Now what is the coordination number? Two. Because we have this sulfur that's bonded there and we have this hydrogen that's bonded

there. So now we have an electron geometry and molecular geometry that are different, right? So what would the electron geometry be? Careful. Electron geometry is counting the electrons in. That's part of the geometry here.

So, it would be tetrahedral. Four things. Now, what's the molecular geometry? Good. Bent.

All right? Now, remember for all of these, I haven't been writing in the bond angles for all of them but maybe we should do that for some to make sure that you were remembering to do that. So what would your bond angles on this sulfur be? 109.5. Good.

Now, what about the oxygens? It's tetrahedral electron geometry so is it 109.5? No. What is it? Good. Less than. All right.

We're going to skip POCl3 so that one's the same as what we've been doing so I want to skip that one and go on.

OK. ClF4 minus. So this one is one with an ion but we've already drawn the Lewis structure. So once you get the Lewis structure drawn, you don't have to worry about whether it's an ion or not too much. You're just basing it off the Lewis structure. And we counted that in already. So we have a steric number of what?

6. And we have a coordination number of what? 4. All right. So steric number comes from 1, 2, 3, 4 and then each lone pair is 1 so 6. Coordination number comes from the fact that we have 4 things bonded to it.

So what is our electron geometry? So we have 1, 6 things bonded or excuse me, steric number of 6. So octahedral, good. So we're again sitting with this sort of structure.

So when we tear in two of those bonded atoms into lone pairs, what happens? What does it become? Square planar. When you're doing these problems, I would really suggest sitting there with like some toothpicks or at least, you know, if you don't want to carry them around with you ink pens, whatever.

I don't care if you bring toothpicks with you on the exam and are kind of playing around with them. It doesn't bother me at all. OK. We'll end there for the Lewis structures.

Don't put your Lewis structures away. We have one slide and then we're going to go back through them all and do something else with them. Getting a lot of run out of those few Lewis structures. OK. So the next thing that we need to spend some time talking about is dipole moment. So this is going to be largely based on that video that I sent you guys.

So hopefully you all watched it over the weekend. Otherwise, you might be a little lost. So if you didn't, go back and read it or watch it rather. I'll give you a second to get your work out.

OK. So your work should be out now? OK. All right. So with a dipole moment, this has to do with a vector addition of magnetic moments of polar bonds. Now, what does that mean? That's sort of the technical definition for it. So if you've had a lot of math and a lot of physics, you want to think of this as vector

addition because that's what it is. If you haven't, I'm going to give you a sort of more general feel for how it works and what it is. The idea behind this dipole moment has to do with that polarity that I talked about in the video and the difference in electronegativity.

We go to look at these. I had talked about the difference between a bond on electronegativity. Now when we had talked about it or whether you had watched it in a bond, what did we say a polar bond was? We said if you have an atom that's very electronegative and one that's not very electronegative or at

least isn't as electronegative as the other one, that you would have a polar bond. So it's something, what's an electronegative element that you can think of? Maybe the most electronegative. Yeah, fluorine is a good one. And a non-electronegative one, maybe hydrogen? So if you had something like HF, would that be a polar bond?

Two atoms, a molecule like HF, if that bond is polar, the atom is polar. Because what's happening there is the fluorine is taking all that electron density and it's pulling the electron density toward it. And that's making that bond polar. Now you have lots and lots of different bonds, now you have to look at the geometries

a little bit. You still have to go through and decide, are there polar bonds? Once you decide if there's polar bonds, then you have to actually see, well, how do these

polar bonds add up? And that's what I mean by vector addition. But you can just sort of look at it in general and see if they cancel. If you have a bond going, an electronegative bond going that way and an electronegative bond going that way, they're going to cancel each other out. This atom's going to be pulling the electron density that way, this atom's going to

be pulling the electron density that way, and they're going to completely cancel. Now if you have bonds that are in the same direction, now this doesn't have to be the exact same direction. Obviously I can't have two bonds going exactly the same way to two different atoms. But if you have two bonds that are in the same general direction, they're going to add up.

So something like HF, that's a good example of a one bond where you just have the fluorine taking all the electron density to it. There's nothing to cancel it out, there's no other atoms that you have to worry about whether that particular bond is polar or not. And so you're just going to get a polar bond and a polar molecule.

Now this is how we write this. So there's two different ways, and you can pick one or the other, it doesn't really matter, just so you don't have to write them both, I did just to show you. But sometimes you'll see one written as a delta positive, so that delta meaning just a partial positive, it's not like it actually has a positive charge, but it has a partial

positive charge. And then a negative over here saying, well this is a partial negative. The other way that you can write it is you put this little plus sign arrow on the positive side and you draw an arrow to the other side. Now what if we have something more complicated? So for instance, ClF3, so I've sort of taken my structure apart, let me put it

together real fast. Something like this is a little bit harder to see. First of all, is a bond between a Cl and an F polar?

That's our first question. Are these going to be polar bonds? They're not going to be super polar, but they'll be polar. Fluorine is more electronegative than chlorine. So there's going to be, each of these fluorines will steal the electron density away from the chlorine. So now the question becomes, is it a polar molecule?

So the easiest way to look at it is what cancels and what adds together? So if this one's pulling up, this one's pulling down, are they going to add, cancel, or do something else? Yeah, I do. Sorry, I made it more complicated than I needed to.

We'll do that one next. Okay, so yeah, these will cancel. So for the sake of our little discussion on polarity, we can pretend they're not having any impact in here. Is this one going to have anything to cancel with? No. So is this going to have a dipole? Yeah. Build it back again.

In which direction is that dipole going to be? That way. Yeah, we'll point. That way. Right? Because this one and this one cancels. The fluorine is more electronegative. So your partial positive would be here, the plus sign of your arrow, and your delta negative would be over here.

Alright, so you can sort of just think of it like that without having to think of it as vector addition, or if, you know, you really like the physics and the math aspect of it and you want to think of it as full vector addition, you can do that too. Okay, now, let's think of some more examples.

Let's look at XeF2Cl2. Now yeah, no, it would be polar.

So you can say it a couple of different ways, you can say have a dipole or is polar, either way means the same thing, and it would have a dipole, and it would be in the direction where everything doesn't cancel. So when I had it drawn like this, the top and the bottom would cancel out, and so it

would have just been in that direction. Yeah, you would use your periodic trends. So you know, your chapter one material didn't quite go away, unfortunately for you guys. So you have to remember that. So yeah, you'd refer to the periodic table.

Oh, good question. Let's go through that in a minute. So what shape is this? So best way to start when you're talking about these is to start thinking about the electron

geometry and then move to the molecular geometry, keeps you from making mistakes. So our electron geometry would have been like this, but two of those would have

been replaced by lone pairs. So what shape is this? Let's turn it on its side. What shape is this? T-shaped. So having three things bonded to it doesn't make it trigonal planar because you have to pay attention to the lone pairs, and the lone pairs make a difference.

Now how did you know about the lone pairs there? Well you would have had to draw out the full Lewis structure, which I didn't necessarily do here, but you would draw out the full Lewis structure. Before you can ever decide on a geometry, and therefore before you can ever decide on whether something is polar or not, you have to go back and you have to draw the Lewis structure. Because otherwise you might think, well, there's three things on it, there's three

things bonded to it, why wouldn't it be trigonal planar? And you might draw it like this. Now if it was shaped like this, would it have a dipole? No, right? All three cancel. They're all going completely opposite directions and so they all cancel. So if you had a trigonal planar molecule, you'd say, well no, it doesn't have a dipole.

Which is very different from a T-shaped, where you would have a dipole. Make sense? So, you know, Lewis structure is by far and away. If you can't draw a Lewis structure, and you can't draw them properly, you're going to have a really rough time on this next exam. So you always start at the Lewis structure, that's why we've been going back to our

pages, and you start with your Lewis structure and then you work your way all the way through. Good question though. Okay. Anything else on this one? No. Okay. So now this one. XeF2Cl2.

So this is one we haven't drawn out yet. So when we go to draw this, there's a couple of different arrangements. We could have it like this, or we can have it like this. Okay? Now what are the differences between those?

So look at them kind of closely. The one sort of arrow means that it's coming out at you, the other means that it's going into the board. So you need to picture these to be able to see what's going on. So I'm going to draw them out for you, or build them for you rather.

So what shape is this? What shape of electron geometry is it? Octahedral. So there's six things bonded to it. What shape is it electron, or excuse me, molecular geometry? Square planar.

Yes. So these ones, the thick ones, means that they're coming out at you, and these means that they're going into the board at you, or away from you I guess. Okay. So this is what we have for the one thing. Now if, which one is this one? Is it this one or this one?

So there's a couple different ways we can make this. One is to have them directly across from each other, like this, where the chlorines and the fluorines are directly across from the same atom. And one is to put them next to each other, like this, okay?

Now there's a difference on what these are on whether you're going to have a dipole or not. So if we look at this first one, we have the chlorines across from each other and the fluorines across from each other, right? So we'd have this.

Okay. Question then. Does this one have a dipole? Right. So this chlorine and this chlorine will directly cancel with each other, and this fluorine and this fluorine would directly cancel with each other. And so you would get a direct cancellation of everything and you'd have no dipole.

Now, are one of these more polar than the other, or excuse me, more electronegative than the other? Yeah. Which one? Which one's more electronegative? Fluorine. So let's say we do this. Now you think we're going to have a dipole now? Yeah, because what will happen is sure the chlorine and this will cancel a little bit.

All of these bonds are polar, right? A chlorine and a xenon is still going to be a polar bond, but it's not going to be quite as polar. Meanwhile, the xenon fluorine bond is going to be a lot more polar. And so while they'll cancel a little bit, these fluorines are going to overpower the

chlorines and you're going to have a dipole. Now let's say I just, for, you know, the sake of consistency, let's say I hold it like this. Which way is the dipole? Straight up, right? Why isn't it like to this side or this side? The side to side cancel, right?

So even though this is going to the left or right a little bit, and that's going that way a little bit, the side to side part cancels out. But they're both going up. And so they're both going to pull the electron density up. And so your dipole is going to be straight up if I'm holding it this direction.

So that arrow is sort of supposed to be going into the board. So well this one will be polar, but if we drew it like this and put them directly across from each other, is that one going to be polar?

No. Because this one and this one will directly cancel and this one and this one will directly cancel. So this one, this isn't one I could give you on an exam. I would have to draw it for you or something like that. Yeah, this isn't one that I could just say XeF2Cl2, is it polar or not?

You wouldn't know. Well, we're not going to have to draw like this. I'm not going to ask you to draw it like that. You should know what it means. So if I were to draw these two and say which one is polar, you could circle it, something

like that. But don't worry too much about having to actually draw them. Correct. Having a dipole makes something polar. But if something has no dipole because everything would cancel, then they wouldn't be polar. But to continue, so one other thing, now if I were to say does this have polar bonds

though, what's the answer to that? Yes. Right? It has polar bonds, it just doesn't have a dipole. So be careful with that. Okay. So I'll give you a second to write that. So this sort of shows you how the geometry really matters.

Now sort of a lesson in drawing Lewis structures properly and what will happen if you don't.

Let's take these two examples. CF4, or let's just say CLF4, sorry about that, CLF4, and CLF4 minus.

So, let's first look at this one. So how many, what's the steric number on this? Five. So what's the electron geometry?

We have five things, okay well let's start this one from scratch then. So let's put the first two directly across from each other.

Put this here. If you can't see these without building them, you really shouldn't feel bad at all.

Seeing in 3D is sort of one of those skills people can develop but some people have it more than others to begin with. So we have five things bonded to it. So what does that, what does this look like if we put it in a triangle? We have a triangle base and we would have like a pyramid, right? And if we did that, we'd have two pyramids, so what shape would that be?

Yeah. Trigonal bipyramidal. Okay, now the question becomes, what's the molecular geometry? So I have to remove one of these atoms and put a lone pair there for visualization purposes.

Am I going to remove this one or that one? What are this and that called? What are these called? Equatorial, right? You kind of think of it as a globe and the things going around the equator, and what are these called? Good, axial. So are we going to remove an equatorial one or an axial one?

Okay, let's look at bond angles. We're not in agreement yet. So what is the bond angle right here? We have 360 divided by 3, so it's 120. And what's the bond angle here? 90, right? 180 divided by 2, so it's 90. So are we going to want to give the electrons lots of room and give it 120 on each side?

Or less room and only give it 90? Lots of room, right? Electrons get the most room. So we'll put it here. Okay, so now we have our molecular geometry. And what would that be called? Yeah, tip it on the side if you forget, right?

Seesaw? Okay. So, now, what do we think? Dipole or no dipole? Well, let's look at these first for a second. You think these are going to contribute to the dipole at all? No, they cancel, right? This one's going straight, almost straight up. This one's going almost straight down. So what about these two? So let's say I'm going to hold it just like this.

What way would the dipole be then? Directly that way, right? Because this part that's coming out at you and going into the board, those parts are going to cancel. Meanwhile, since they're both going that direction, they're going to go that way. So does it have polar bonds first?

We'll start there. Yes. Does it have a dipole? Yes. So is it a polar molecule? Okay. Next one. ClF4-. Now, that gives us an extra two electrons, which means that now we have a different steric

number all together, right? What's our steric number now? Six. So what's our molecular, or excuse me, let's start with electron, what's our electron geometry? Octahedral.

So we have this shape. So that's for our electron geometry.

Now we have two lone pairs, so does it matter which one I remove first? Are the bond angles all the same or different? They're all the same, right? They're all 90. So it doesn't matter which one I remove to begin with. So now that I've removed that one, which one do I have to remove for the second one? Top one. So now I have a lone pair here and a lone pair here.

So what shape do I have? Square planar, right? It's a square and it's all in one plane. All right, so what do we think, dipole or no? No, because this one and this one cancel and this one and this one cancel. So even though we have polar bonds, we're not going to have a dipole.

If you had the most electronegative element in the center, it would, but in general, that's not how molecules get set up. So for the most part, that your least electronegative one is going to be in the center.

And so all the ones on the outside will be more and it'll pull it that direction. But I mean, theoretically, if you had a molecule that was the most electronegative in the center, then yeah, yeah. For this one, yes.

What do you mean individual molecule? Oh, so for the bond, that's where you're looking at the difference in electronegativity. So if there's a difference in electronegativity, then that's going to be polar unless it's so big that it's ionic.

Anything else? Yeah, so if we wanted to switch out two of these for something that didn't have the same electronegativity at that level, then we would be back to our xenon example

where we had like fluorine and chlorine. And if this was something that wasn't as electronegative, say hydrogen or something like that, then we would have a dipole.

So we have octahedral, square planar, and in that case, no dipole because they all cancel. Now, because I know it's been like five minutes since we've seen those Lewis structures we did and you've been missing them, back to the Lewis structures. Okay, so we have one spoiler up there, but that's alright.

So let's look at nitrogen. So do we think that N2 would have a dipole? First of all, does it have any polar bonds? No, right? Nitrogen and nitrogen, they have the exact same electronegativity because they're

the exact same atom. So your first step is always, well, do they have polar bonds? And they don't. So nitrogen isn't going to have a dipole. Now, what do we think about this one? N and O.

So that was our big resonance structure one that, or the structure that we drew out with our three different structures and determined the best one, that this was the best one. So, first of all, do we have any polar bonds? Yeah, which one? The N to the O, right? Sure, they're next to each other, but they still have a difference in electronegativity

and so it's going to have a dipole. Which direction? Toward what atom then? Okay, well, double spoiler on that one. Alright, let's go to this one. So, we have a dipole, that which we now know, right? Where's our polar bond? Which one?

Or our main polar bond here? The carbon to oxygen. So, which way is our dipole going to be? It's going to be completely up. And just so that you have the other way of writing it as well. Okay, now what about this one?

So, what's the shape of that one? What's the electron geometry? Octahedral, right? That's what I've been doing a lot of making of today. And then, so what's the molecular geometry? Now that we've taken two of those atoms away and put on lone pairs? Square planar. And they're all the same atom that are bound, so they all have the same electronegativity.

So, first of all, is there polar bonds? All of them, right? It doesn't have a dipole. No, they all cancel, right? It's this kind of same example where we have them all directly across from each other. And the two that are this way cancel and the two that are this way cancel.

So, anything where we don't cancel out all of the dipole or all of the polar bonds, they aren't directly across from each other. Those ones have a dipole. The other ones don't. Yeah.

Can you talk a little bit louder? I can't hear you. How do the, oh, the H's? So, you actually, for the most part, can ignore the H's here because the difference between carbon and hydrogen electronegativity wise is really, really tiny. I think it's like, if you were to look up the numbers, I think it's like .3 or something like that.

Where the difference between carbon and oxygen is much higher. And so, you don't really have to worry about the carbon hydrogen issue at all because it's such a small difference as compared to the carbon oxygen, which is a really big difference.

No, so I was just doing that mostly to show you where the polar bonds were that you had to worry about. Yeah, so if I were to ask you if this has a dipole, I wouldn't want you to, I would just want you to say no dipole. Because otherwise, it looks like you're trying to make four dipoles. So, this one just kind of shows you where the polar bonds are.

Good question. Yeah. Yeah, and hydrogen is a little, hydrogen is actually sort of an exception to the trend on some level. So, that's why it's actually pretty close to carbon is because it's an exception because it only has that one electron.

So, it's helpful, if you want to look at the carbon hydrogen issue, it's helpful to actually pull out one of the slides on electronegativity and look and see that they're pretty much identical. Just a tad off. Yeah.

You mean like if there was like three O's here? So, with the, yeah. So, let's say it was three things of all equal electronegativity.

So, we'll not make up a real example off the top of my head in case I, you know, don't pull the right one, but something like this. Right? So, what shape is this? Okay, so I would like more consensus. You're right, but more consensus. So, what shape do we have first? What, like, triangle, right?

And is it on a plane? Is it a pyramid? What is it? It's a plane, so it's a trigonal planar. So, if these all had the same, if each of these bonds were polar, and they were all polar equally, right? They were all the same atom. Would this have a dipole? No, they would all cancel. Now, let's say that, like, we had two things that were kind of polar and one thing that was super polar.

Then would it have a dipole? It kind of depends, really. So, if these were just a little bit electronegative, or a little bit polar, but, you know, a fair amount, maybe something like a chlorine,

and then this was something that was, you know, more very electronegative, like a fluorine, it would be really tough to tell. You'd actually have to do the bond addition, like the vector addition, and I wouldn't give you that. It would be something really obvious, like a carbon and hydrogen with an oxygen or something like that.

Yeah. Oh, and in my fake example here? It's hard to tell because these two would be adding that direction, and this one would be adding that direction, so these two are also sort of canceling with each other, but they're also adding that way,

so there's sort of a lot going on that becomes a lot harder to tell. You know what I mean? That's when you have to actually do some vector addition, and actually add up the dipl- the de-bias, the differences between the two, and then add them. Make sense? You mean going down a periodic table?

So what do we think? Going down a periodic table, what happens to the electronegativity in general? Gets bigger or smaller going down the periodic table?

It gets smaller, right? Okay, next ones. So what about these? So let's start here. SF6. Dipl? Yes, no? No, right? They all cancel. They're all going in different directions, they all cancel.

Alright, next one. BH3. Dipl? No, they're all going in opposite directions, and there's not much of a difference in electronegativity there anyways. Alright, next one.

So, this one, this actually falls into the category of ones I probably wouldn't ask you, for the exact reason that I was talking before. So, if you look it up, I looked it up at some point and it happens to have one,

but this goes on the example of things I wouldn't ask you, okay? So you can actually just write that down if you want. Because what happens is these are all pointing in one direction, we have a tetrahedral geometry,

and that's pointing in the opposite, and so it becomes a little tough to tell. So this kind of goes in the category of things that you'd actually have to do all the vector addition for.

Now, to kind of switch gears a bit, but not too much. I want to talk a moment about greenhouse gases. Because now we know some stuff that we can talk about greenhouse gases with. So, I hope everyone has heard of these.

These are kind of responsible for what important thing going on right now. Global warming, right? So, you know of a lot of greenhouse gases, right? You know of ones that are talked about, such as, yeah, CO2, things of that sort.

But what makes something a greenhouse gas? Well, what happens is you get either a permanent dipole or what we call an induced dipole, which I haven't really talked about yet. So, we already talked about a permanent dipole, right? Difference in electronegativity, and they don't cancel each other out. Now we have something like an induced dipole.

So, what an induced dipole is. So, a great example of this would be, well, let's actually just go through some examples and I'll show you one. So, let's go through those and see which ones actually are going to be a greenhouse gas and which ones aren't. So, let's look at N2.

So, first of all, does it have a permanent dipole? No. Now, is there any way that we can make some sort of dipole there by bending it or stretching it or doing something like that? You know, you can't, right? There's only two atoms. You can't change anything about that. So, is N2 going to be a greenhouse gas? No.

What about NO2? So, we have this sort of shape going on with N in the center. So, we have polar bonds going this way, polar bonds going that way. So, which, if I'm holding it like this, which way would the dipole be?

Down. Just, you know, since they are the same thing. Does that double bond matter? No, right? Because is it really a double bond? No, right? Why? What's the word here you're looking for? Resonance, right? Really, that double bond is sort of split between here and here.

We say those electrons are delocalized, right? So, they're kind of going in between. So, yeah, this one's going to be a greenhouse gas. It has a dipole. What about oxygen? Does that one have a dipole? Can we make a dipole out of it? So, that one's not going to be.

What about CO? Yeah, that's carbon monoxide, right? That one has a dipole as well. We have a difference in electronegativity. Alright, N2O. So, we sort of talked about this one already. We have a dipole, right? Which direction is it going again?

Yeah, toward the oxygen. Okay, now there's one I didn't put up here. What's another greenhouse gas you know of? CO2. So, what's CO2 shaped like?

It's linear, right? If you don't believe me, draw it. It's a carbon with a double bond to one oxygen and a double bond to the other oxygen. No lone pairs, so it's linear. So, is CO2 a greenhouse gas? Just use general knowledge right now.

In your life, is CO2 a greenhouse gas? Yes. Now, I told you, does it have a dipole? No, it does not have a permanent dipole. Now, how is it a greenhouse gas then? Yeah. You can induce it. Yeah, you can induce it, good. So, what if I take this and I bend it like that?

Would it have a dipole? It would be going which direction? Up. What if I bent it like that? Yeah, then it would be down, right? So, you can make a dipole in it by bending the molecule. You can't do that with something like, you know, N2 or O2 or any sort of just two atoms bonded together. But you can do that when you have this central atom that you can bend.

So, that CO2 is a greenhouse gas because you can induce a dipole. Yeah. So, it's not actually that you're changing the bond. It's that you're just sort of stretching it with electromagnetic radiation. So, you can just sort of bend this by putting energy into it.

And you can also stretch the bonds. The only way to actually do the stretching part is with my hands. So, if my body is carbon and these are oxygen, you can kind of stretch it like that. And then one dipole would be different from the other. Or you can bend it.

Anytime you have a bond angle, you can bend it at that angle. OK, so in the last two minutes, don't pack up on me yet. Alright, I want to talk about the fact that we have these two different theories of bonding. Bond theory, and we'll talk about that as hybridization.

And then we have something called MO theory. So, what we'll start with next time is the valence bond theory and the hybridization. And then we'll go on to talk about MO theory. They both really have their benefits and their flaws. Something like hybridization is really quick and simple to look at.

And we can look at really large complicated molecules and yet you can still go through them very quickly. MO theory is a little bit better at predicting things, but it's very, very complicated. So, with MO theory, I'm going to sort of say, hey, the computer said this and show you some pictures and explain to you why.

But MO theory becomes very complex very quickly. We'll only make it up to second row diatomics in this class with MO theory. With valence bond theory, we'll start looking at, you know, large drug molecules fairly quickly because it just expands out really nicely. So, those are sort of your two different theories and why you would use each one. But you should kind of keep in mind that they are separate to some extent.

That you can use the one or you can use the other or you could use both of them, but they are very different. And again, if you didn't watch that video and some of today's lecture seemed a little foreign, please go back and watch that.