Molecular Orbital Theory

Hey, everyone, in this video, we're going to start our journey down a path of organic chemistry that's very important called molecular orbital theory. Let's go ahead and get started, guys. So first of all, I want to just give you a disclaimer. This topic is called The Basics of Molecular Orbital Theory. But to be honest, there's nothing basic about molecular orbital theory. It's one of the most widely misunderstood parts of organic chemistry, and many students just avoid it entirely because they're so confused andare very few resources out there that provide very clear explanations that they just try to ignore it and try to get through organic chemistry without it. Unfortunately, there are some reactions that we're going to need to really understand molecular orbital theory so that we can learn those reactions. And without a good understanding of Emma theory, you're just going to get lost. So what I'm gonna try to do in the next 15 minutes or so is I'm going to try to tell you a really smooth story based on what you already know on how to understand Molecular Orbital's, and I actually worked really hard on this to try to build a good flow based on what I believe you already know and what you need to know by the end of this topic. So please let me know if this story made sense to you at the end. I'm totally down to redo this if it's confusing, but I'm gonna try to just take my time. This isn't about getting through this quickly. It's about making sure that all of you guys get it at the end. So if it seems like I'm going slow, that's on purpose. Because there are very few videos you could go to online that explain molecular orbital theory thoroughly. And I'm gonna try to build that video now. Okay, cool. So let's start with what we know as previously discussed. There's this idea called conjugation. You guys remember what conjugation means. Conjugation means that you have the ability to resonate. OK, what does resonance mean? Resonance means that you're sharing electrons from one Adam toe. Another. You guys remember that you can resonate. Electrons etcetera. Well, one of the technical ways that you can talk about residents is that residents happens from non bonding. Orbital's two adjacent non bonding orbital's what Talking by non bonding that it's not making a bond toe another atom. Now, why are Onley non bonding? Orbital's involved in residence because if you're making a bond to an atom, it's stuck to that Adam. And remember that residents structures, you can't move atoms around. Remember that. Remember, the only thing you could move is bonds like pi bonds and electrons. You can't move atoms. So when we're going to talk about the idea of conjugation, we're always gonna talk about the non bonding orbital's meaning, ones that don't have in an Adam attached to them. Okay, quote. So now, since we're gonna be talking about non bonding, orbital's a lot today, I want to remind you that non bonding on Lee takes place in the outermost shell oven Adams Electron configuration. So remember that you have, like, one s orbital's two s orbital's etcetera. You would on Lee be dealing with last shell. And since we're in organic chemistry, that last shell is usually going to be the second shell, meaning that the electrons that are in the first shell that one s are not involved in any of the things we're going to talk about today. You can pretty much ignore. The one s we're gonna talk about is the two s. Okay, So what I wanna dio is show you guys are very basic example off hybridization from organic chemistry. One. This is one of the first things you learned in organic chemistry. One. And I want to remind you how these electrons behave, how these valence electrons behave. Okay, so here is an AL Keen, and we know that we learned a long time ago that Al Keane's have three bonds. So an alky in carbon has three bonds attached to it, which would mean what we call them three groups or three bond sites, Remember? Ah, bond site or a group is just anywhere that you haven't Adam attached or that you have a lone pair attached. Okay, So if we were to look at this carbon right here that I already have circled, how many atoms does it have attached to it right now? It has ah, hydrogen. Ah, hydrogen and carbon. Meaning that there are three bond sites meaning that this means that this equates to an SP two hybridization. Remember that That you're just supposed to know that However, many bond sites there are. That's how many That's what you're hybridization is, and three always means sp two. But let's go a little bit deeper into the electron configuration to remember how this hybridization works. And by the way, I have videos on all of this so far. Um, but I'm just here to remind you thes air the highlights. So remember that carbon is in which, you know, is what's the atomic number of carbon? It's six. Carbon has an atomic number of six, which means that at its neutral state, how many protons does it have? Six. How many electrons does it have? Six. So when we build the electron configuration of carbon, we need to figure out where all those six electrons we're gonna go right. And remember that you always start with a foul principle from the lowest energy orbital. So you have to start filling your orbital's in ascending order of energy. So that means that out of the six electrons, where should those electrons go? Well, two of them should go into the one s orbital because that's the lowest energy state orbital possible. Then another two of them should go into the two s orbital because that's the next highest right. And then we have Huns rule. Where that hunts rule says that if you have, um, a bunch of seats on the bus, you need to fill them equally. You can't just have two kids on one seat and zero kids on another seat. So in this case, notice that the P orbital's are all the same. Energy state, right? So that means that I would then get one electron in the two p x one electron in the two p y. And now I ran out of electrons. I just put all my six, meaning that there's no electrons for the two PC. Does that make sense so far? So this is the way that we would fill these orbital's based on what we know from Gen. Camp, based on what we know of just like, hey, there's six electrons and we need to put the put them into the election configuration. This is what it should look like. But remember that basically, in the first chapter of organic chemistry, what we learned is that this is not favored, and the reason is because guys remember that Carbon always wants to be able to make four bonds, right? But right now, the way that you have the carbon set up, the two s already has a filled orbital, right? This is already filled. So can that orbital make a bond? No. And then we have I'm just gonna use different colors. These two orbital's could make a bond because they could accept one electron, and then this one has no electrons. So it's not very good at making a bond because it would have to accept two electrons, not just one. So that limits the amount of things that it could make bonds with. So what I'm trying to say is that it's not very favored toe have these electrons scattered like this. What's more favored is to spread the electrons out evenly throughout all the orbital's, so that all the orbital's have a chance to make a bond. And this is the process that we call hybridization. Remember that when you have specifically three groups or three bond sites, what happens is that the two s orbital blend with two of the P orbital's to give you an SP two hybridized blended orbital. Okay, and that's what we have happening here in this gray box. Notice that remember what you're supposed to memorize is that three bond sites equals s p two, which means that the two s the yellow from the two s blended with the two with the two p orbital's two of the P orbital's to give us three new orbital's called SP two sp two s, p two. Okay, now you might be wondering, Johnny, why did you put two SP too? Well, because, technically, you could just call it s p two. But you can also call it to SP two because they're in the second shell, and anything that's in the second shell can get a two behind it. Okay, so remember that s p three means that there are three bonds and they can all blend together in this way. Notice that now what happens is instead of getting two electrons in a lower energy orbital and then two electrons in higher now we have is three electrons evenly spaced out between this, uh, more averaged out energy level. Okay, But there's also one more thing, which is that when you have sp two and three bond sites, that means that there's 1/4 bond. That's not being made. That means that there's an extra electron that is just gonna be in an extra orbital. And that extra orbital does not hybridized. So that is gonna be here. My two pz notice that this one didn't hybridize. So it's actually a little bit higher and energy because it didn't blend with the other ones. And it has one electron that's free Thio interact with either to make a potentially make a bond or to interact with other orbital's. Okay, so this is gonna be what we keep. We're gonna call our non bonding orbital, and this is gonna be the one that's gonna be the really interesting one for us. For the for the rest of this section. We're going to talk about the non bonding orbital a lot, Okay, But let's put this on hold for a second because I want to go back to the two sp two and talk about what they're doing. Okay, We'll remember that. We said that it's making three bonds, right? So what we could do is we could show where those three bonds air happening. One of them is happening. I'm just going to use different colors for this One of them is happening to an electron from a one s orbital in the hydrogen. So that's this guy right here. I'm gonna call him a And then this is a What's happening is that the hydrogen has one electron, right? Hydrogen have an atomic number of one. So it has one electron, and it's sharing that electron with the one electron from the S P two. And what that's doing is it's making a bond. So when I drew this blue area here, this actually means that we're making a new bond between those two electrons that are now gonna be shared in one orbital. Does that make sense so far? Cool. Notice that this is also happening on the bottom. I have another one. This is HB. Let's say this one is also overlapping with the electron from the SP too. And it's making a bond cool. And then lastly, guys noticed that the carbon is also making a bond to another carbon. Right. Let's call this See here. So it's making a bond to that carbon. Okay, But that carbon isn't just a one s orbital because one s is what you have. If you just have a hydrogen. What it actually is is it's another SP two orbital, another sp two hybridized orbital because notice that it has three bond sites. So you're gonna have 11 s and two piece blend together and give you an SP two. And what's gonna happen is that those two SP two s are gonna overlap in one place and give us a new Sigma Bond, and that's our new signal on. So, by the way, all of these Air Sigma Bonds, this is Sigma Oops, Sigma and sigma because they're all overlapping in just one place, basically, like you could think of it like this. Like this tip is going to overlap with this tip like this. They're all overlapping in one place, giving us three new Sigma bonds. Okay, so now this brings us to the interesting part. What is happening with the non bonding orbital? It's left over, but it has one electron left, so its able to interact with something. But what is it going to do? What it actually looks like guys, is it looks more like just to show you it looks more like this. Okay, where you have your three Sigma bonds. So Sigma one, Sigma two, Sigma three. We already talked about how that's happening, but then we have this extra electron that's just floating in an orbital, waiting to do something. So what is it going to do? Well, guys, it's going to be able to congregate if you can put an end. If you can put another non bonding orbital next to it, it will congregate. And what conjure it means is that it can share its electrons between them freely so the electrons can actually resonate or jump around from here to here from here to here, and they can blend together. Okay, now, the type of residence that you get depends on what type of orbital is the second one that you're interacting it with. Okay, now, in a in when you make a pie bond, remember that a pie bond has to do with making a double bond, right? Ah, pie bond would just mean that you have another non bonding orbital. That's exactly like the one that you started with. Where has one electron and basically it's overlapping with another um, basically two p Z. Does that make sense? It's overlapping with another two PC and what it's going to do is it's going to make what we consider to be a pie bond. So a pie bond would mean that one electron from here is sharing with one electron from here. And they're making a new region of overlap here and here. Okay, that would be what we call a pi bon. And that would be a form of conjugation because now those electrons can be shared between them. Okay, By the way, you might be wondering, What does this thing mean? That I drew? It just means that this is the thing that it can make that it can interact with. This is the This is the unknown non bonding orbital. If it happens to be the same thing that we started with, then it's gonna make a pie bond. Okay. But that's not the only possibility. There are other types of non bonding orbital's that we could put here. Another type would be just an empty orbital. So when empty orbital, there are lots of examples, but examples could be something like aluminum or boron. Thes are atoms that have, uh, just an empty orbital that you can share electrons with That's non bonding. Another one would be like an orbital that's empty with a positive charge, which is the same thing. Just the formal charges. Different. So that would be basically like a cat ion. Okay, Another example would be a non bonding orbital with just one electron in it. That that would be a radical. Okay, another one would be a lone pair. Remember, Lone pairs. Once again, they are non bonding because they're not making a bond to anything. And then the last one would be like an anti, which is just simply, it's a It's a lone pair with a negative charge on it. So it's just it's just a the formal charges different. So what happens is guys, even though the example that I've given here is of a double bond, and that's what we're doing over with their al cane. If we wanted to, we could have also conjugated this non bonding orbital with one electron toe, any of those other atomic orbital's that are non bonding. Okay, and when we do that, what we're gonna do is we're going Thio instead of having two different atomic orbital's. What's gonna happen is that they're gonna make one new molecular orbital. So what we're gonna do is we're going to make molecular Orbital's out of where there were atomic Orbital's before. Okay, so what I'm gonna do is in the next video, I'm going to remind you guys how atomic Orbital's become molecular orbital's by overlapping with any of these non bonding orbital's. Okay, so let's go ahead and go to the next video.

when adjacent non bonded atomic orbital's overlap with each other or a next to each other. What they do is they create these more favorable molecular orbital's. So what a molecular orbital is is It's just the overlap of a few atomic orbital's. Okay, now, if you want to know what the molecular orbital is gonna look like, we can use a system that's very common in organic chemistry, called the linear combination of atomic orbital's L C A O. And what that does is it helps you to predict what the molecular orbital is gonna look like. Okay, so what I want to do now is now that I have kind of hopefully convinced you that atomic orbital's like to share electrons. Now what I want to do is talk about what do those electrons look like after their share. And we're gonna take the example of ethane again. Remember that with ethane, What did we just say? We said that one orbital has one electron, another orbital has another electron, and that's basically what we've done right here. These are the atomic orbital's, and usually this is the way that we represent it. What we do is all of the conjugated Adams will get one atomic orbital, so notice that I have to conjugated atoms. So here I draw 12 atomic orbital's next to each other, and I put, however, many free electrons there are into those orbital's. So notice that Adam one is donating one. That's why I put one electron Adam to is donating another one. That's why I put another one. Okay, now these air again, what we called the atomic orbital's. But remember that what is an orbital even like? What is the definition of an orbital? Remember that in orbital is just a region of space that is statistically probable. Toe have electrons in it, so it's like a cloud of electron density where there's a high chance will find electrons in it. But it's not actually a particle. It's not actually tangible. So when you bring these atomic orbital's close together, remember that let's say that they're about to touch what happens? Do they collide like tennis balls? Would if we brought them together? No, Remember that they don't collide. Remember that what they do is they interfere with each other like a wave. They interfere like waves. They don't collide like particles. So what that means is that there's different types of interference that can happen when you bring these to a owes close enough to interact. Okay, there's one type of interference called constructive interference. What you're gonna write down here, constructive interference. Now what constructive interference is? In a nutshell, it means that the two waves of those atomic orbital's build on each other. So then the waves between them get actually increased. They get hired, they increase in amplitude, and they increase the chances of finding electrons inside of them in between them. Okay, so when there's a constructive overlap, that's what we call in in phase overlap. Because it means that the atomic orbital's are aligned in such a way that they agree with each other. The waves at the top are adding together. The waves at the bottom are adding together, and they're increasing the chances of finding electrons on both sides of the orbital. So what they do is they form this type of interaction called a bonding interaction. And what a bonding interaction means is that the chances of finding electrons between these two atoms is unusually high. It's higher than normal because these waves blended together in a really positive in a good way. But that's not the only type of interference that can happen. There's another type of interference called destructive interference. Okay, destructive interference happens when your orbital's when you're atomic orbital are out of fate, meaning that your positives and your negatives don't align properly. So what winds up happening is that instead of the waves blending together in a good way to make the chances of finding electrons higher, they actually perfectly cancel out. So one positive cancels out one negative, and the chances of finding electrons between them is actually hits a limit of zero. It actually hits a mathematical limit of zero. That's what we call unnoted. A node is a region of space in between the atoms that there's actually no mathematical chance that there could be electrons there because the orbital's perfectly canceled each other out. Okay, when this happens, it actually is the opposite of a bond. It's very unstable because you're Adams actually want to repel from each other because they're not sharing electrons at all. And this is what we call an anti bonding interaction. Okay, now I want to just point out one thing, which is that the positive and negative lobes, often atomic orbital, have nothing to do with positive and negative charges. So this doesn't I'm not saying there are negative or positive charges like you're thinking maybe with acids and bases. This has to do with just the nomenclature or the way that we think about Orbital's, that they have different sides. But you could also just call it the white side and the gray side if you wanted. It doesn't matter. All I'm saying is that the whites have to be on the same side and the grays have to be on the same side. You don't want them to be opposite to each other. Now you might be saying, Okay, Johnny, I understand that mathematically constructive overlap is possible, and destructive overlap is possible. But why in the hell would electrons ever go into the destructive overlap? Why would you ever do that? And guys, the reason is because we have these rules of atomic of electron configuration, and one of them is called the Pauli exclusion principle. Do you remember what Polly Exclusion says? This is just comes from Chapter one of organic chemistry It says that you can Onley put two electrons in each orbital. And once you have two electrons, that's it. You can't go anymore. You have to Then fill the next higher energy orbital. Okay, So what this means is that you actually we prefer not to put any electrons into the anti bonding molecular orbital. We Onley do it if we have no other choice because we have too many electrons and we need to put one up there. So how many electrons are being shared in the atomic orbital's too. So that means that when we make our new molecular orbital, this is molecular orbital pie one. This is molecular orbital pie to these air, the potential pi bonds that we could make based on the atomic orbital's that are overlapping. And what we know is that both of these electrons can come down and fill the lowest energy orbital and fill it completely, meaning that two electrons is a great number because we're going to get a bonding interaction and we're gonna have zero electrons up in the anti bonding region. Okay, but let's say that one of these orbital's instead of donating one electron, let's say that it was donating two electrons. So let's say there was an extra electron that we had to deal with that Third Electron would then according toe off Bell principal, remember building up. It would need to get kicked up here, which would be really bad, because this means that this electron is in an anti bonding orbital, which means that it's not gonna promote bonding between the two carbons. It's actually gonna make them less stable, and it's gonna wanna make them break apart. Okay, so basically, when we build out our our molecular orbital's, we're gonna be using this system, and I'm gonna teach you how to write out all your atomic orbital is how to write out your molecular orbital's. But it's important that you guys know that we're going toe always be starting from the lowest energy state and then Onley going up to the higher energy states when you need to, because you have extra electrons that you need to get rid of. Okay, So, basically, in conclusion, the reason that Al Keane's can make such a good double bond is because they have exactly two electrons that they can share constructively in one molecular orbital. So, really, instead of having two separate atomic orbital's, what it really looks like is one molecular orbital of low energy that promotes bonding between the two. Okay, so I hope that this is a good start to molecular orbital theory. And now I'm gonna follow this up with more videos explaining exactly what you need to know so that you can apply molecular orbital theory to solve whatever problems you have.