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 and there are very few resources out there that provide very clear explanation 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 MO Theory you're just going to get lost. So, what I'm going to 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 orbitals 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 going to 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 can go to online that explain Molecular orbital Theory thoroughly and I'm going to 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, okay? What does resonance means? resonance means that you're sharing electrons from one atom to another you guys remember that? you can resonate electrons etc. Well, one of the technical ways you can talk about resonance is that resonance happens from nonbonding orbitals to adjacent nonbonding orbitals, what do I mean by nonbonding, that is not making a bond to another atom. Now, why are only nonbonding orbitals involved in resonance? because if you're making a bond to an atom it's stuck to that atom and remember that resonance structures you can't move atoms around, remember that? Remember, the only thing to 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 going to talk about the nonbonding orbitals, meaning ones that don't have an atom attached to them okay, cool? So, now since we're going to be talking about nonbonding orbitals a lot today I want to remind you that nonbonding only takes place in the outermost shell of an atom's electron configuration. So, remember that you have like 1's orbital 2's orbitals etc. you would only be dealing with the 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 1s, we're going to talk about is the 2s, okay? So, what I want to do is show you guys a very basic example of 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 alkene and we know that, we learned a long time ago, that alkenes have three bonds. So, an alkene carbon has three bonds attached to it, which would mean what we call them three groups or three bond sites. Remember, a bond site or a group is just anywhere that you have an atom 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 that have attached to it right now? it has a hydrogen, a hydrogen and a carbon, meaning that there are three bond sites, meaning that this means that this equates to an sp2 hybridization. Remember that? that you're just supposed to know that how many bond sites there are, that's how many, that's what your hybridization is and three always means sp2.

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 but I'm just here to remind you, these are 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 of that 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 are going to go, right? And remember that you always start with principle from the lowest energy orbital. So, you have to start filling your orbitals 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 1s orbital because that's the lowest energy state orbital possible then another two of them should go into the 2's orbital because that's the next highest, right? And then we have Hund's rule where that Hund's rule says that if you have a bunch of seats on the bus you need to fill them equally, you can't just have two kids on one seat and 0 kids on another seat. So, in this case, notice that p orbitals are all the same energy state, right? So, that means that I would then get one electron in the 2p x, one electron in the 2p Y and now I ran out of electrons I just put all my six, meaning there's no electrons for the 2p Z, does that make sense so far? so this is the way that we would fill these orbitals based on what we know from gen chem based on what we know just like, hey there's six electrons and we need to put the put them into the electron configuration, this is what it should look like, but remember that pretty 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 2'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 getting different colors, these two orbitals could make a bond because they could accept one electron and then this one has no electron. 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 can make bonds with. So, what I'm trying to say is that it's not very favored to have these electrons scattered like this, what's more favored is to spread the electrons out evenly throughout all the orbitals so that all the orbitals 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 2's orbital blends with two of the p orbitals to give you an sp2 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 equal sp2, which means the 2s, the yellow from the 2s blended with the 2 with the 2p orbitals 2 of the p orbitals to give us 3 new orbitals called SP2, SP2, SP2, okay? Now, you might be wondering, Johnny why did you put two SP2? Well, because technically you can just call it SP2 but you can also call it two SP2 because they're in the second shell and anything that's in the second shell can get a 2 behind it, okay? So, remember that SP3 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 three electrons evenly spaced out between this more averaged out energy level, okay? But there's also one more thing, which is that when you have SP2 and three bond sites that means that there's a fourth bond that's not being made that means there's an extra electron that is just going to be in an extra orbital and that extra orbital does not hybridize so that is going to be here, my 2p z. Notice that this one didn't hybridize. So, it's actually a little bit higher in energy because it didn't blend with the other ones and it has one electron, that's free to interact with either to make a potential to make a bond or to interact with other orbitals, okay? So, this is going to be what we're going to call our non bonding orbital and this is going to be the one that's going to be the really interesting one for us for the for the rest of this section, we're going to talk about the nonbonding orbital a lot, okay? But, let's put this on hold for a second because I want to go back to the two SP2's and talk about what they're doing okay? Well, remember that we said that it's making three bonds, right? So, what we could do is we could show where those three bonds are 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 1s orbital in the hydrogen. So, that's this guy right here, I'm going to call him a and then this is a, what's happening is that the hydrogen has one electron, right? Hydrogen's have an atomic number of one, so it has one electron and it's sharing that electron with the one electron from the sp2 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 going to 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 sp2 and it's making a bond, cool? And then lastly guys. Notice that the carbon is also making a bond to another carbon, right? Let's call this C here. So, it's making a bond to that carbon, okay? But that carbon isn't just a 1s orbital because 1s is what you have a few just have a hydrogen, what it actually is, is it's another sp2 orbital, another sp2 hybridized orbital because notice that it has three bond sites. So, you're going to have one p, one S and two P's blend together and give you an sp2 and what's going to happen is that those two SP2's are going to overlap in one place and give us a new Sigma bond and that's our new Sigma bond. So, by the way all of these are Sigma bonds, this is Sigma, 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, okay? 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 nonbonding orbital? it's left over but it has one electron left. So, it's 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 bond. So, Sigma 1, Sigma 2, Sigma 3, 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 conjugate, if you can put an N, if you can put another nonbonding orbital next to it, it will conjugate and what conjugate 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 resonance that you get depends on what type of orbital is the second one that you're interacting it with, okay?

Now, when you make a pi bond. Remember, that a pi bond has to do with making a double bond, right? A pi bond would just mean that you have another nonbonding orbital that's exactly like the one you started with, where it has one electron and basically it's overlapping with another basically 2p z does that make sense it's overlapping with another 2P z and what it's going to do is going to make what we consider to be a pi bond. So, a pi 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 bond 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 unknown nonbonding orbital, if it happens to be the same thing that we started with then it's going to make a pi bond, okay? But, that's not the only possibility, there are other types of nonbonding orbitals that we could put here, another type would be just an empty orbital. So, an empty orbital, there are lots of examples, but examples could be something like aluminum or boron these are atoms that have a just an empty orbital that you can share electrons with, that's nonbonding, another one would be like an orbital that's empty with a positive charge, which is the same thing, just the formal charge is different so that would be basically like a cation, okay? Another example would be a nonbonding 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 nonbonding because they're not making a bond to anything and then the last one would be like an anion, which is just simply it's a, it's a lone pair with a negative charge on it. So, it's just a, the formal charge is 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 our alkene, if we wanted to, we could have also conjugated this nonbonding orbital with the one electrons to any of those other atomic orbitals that are nonbonding, okay? And, when we do that, what we're going to do is we're going to instead having two different atomic orbitals, what's going to happen is that they're going to make one new molecular orbital. So, what we're going to do is we're going to make molecular orbitals out of where there were atomic orbitals before, okay? So, what I'm going to do is in the next video, I'm going to remind you guys how atomic orbitals become molecular orbitals by overlapping with any of these nonbonding orbitals, okay? So, let's go ahead and go to the next video.

When adjacent non-bonded atomic orbitals overlap with each other or are next to each other, what they do is they create these more favorable molecular orbitals. So, what a molecular orbital is, is it's just the overlap of a few atomic orbitals, okay? Now, if you want to know what the molecular orbitals going to look like, we can use a system, it's very common in organic chemistry called the linear combination of atomic orbitals LCAO, and what that does is it helps you to predict what the molecular orbitals going to look like, okay? So, what I want to do now, is now that I kind of, hopefully convinced you that atomic orbitals like to share electrons. Now, what I want to do is talk about what do the letters look like after they share and we're going to take the examples ethene again. Remember, that with Ethene and 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 orbitals and usually this is the way that we represent it, what we do is all of the conjugated atoms will get one atomic orbital, so notice that I have two conjugated atoms. So, here I draw one, two atomic orbitals next to each other and I put however many free electrons there are into those orbitals, so notice that atom one is donating one. So, why I put one electron, atom two is donating another one, that's why I put another one, okay? Now, these are again what we call the atomic orbitals, but remember that, what is an orbital even, like, what is the definition of an orbital? Remember that an orbital is just a region of space that is statistically probable to have electrons in it. So, it's like a cloud of electron density, where there's a high chance we'll find electrons in it but it's not actually a particle, it's not actually tangible. So, when you bring these atomic orbitals close together. Remember that, like, 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 AO's close enough to interact, okay? There's one type of interference called constructive interference, which I'm going to write down here, constructive interference. Now, what constructive interference is in a nutshell, it means that the two, the waves of those atomic orbitals build on each other. So, then the waves between them get actually increase, they get higher, 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 an in phase overlap, because it means that the atomic orbitals 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 the 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 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 orbitals, when your atomic orbitals are out of phase, meaning that your positives and your negatives don't align properly. So, what ends 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 0, it actually hits a mathematical limit of zeros, we call a node, 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 orbitals perfectly cancel each other out, okay? When this happens, it actually is the opposite of a bond, it's very unstable because your atoms actually want to repel from each other because they're not sharing electrons at all and this is what we call an antibonding interaction, okay? Now, I want to just point out one thing, which is that the positive and negative lobes of an atomic orbital have nothing to do with positive and negative charges. So, this doesn't, I'm not saying if 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 orbitals that they have different signs, but you could also just call it the white side in the gray side, if you want it, it doesn't matter, all I'm saying is that the whites has 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, you guys remember what Pauli exclusion says? this is just comes from Chapter one of organic chemistry, it says that you can only put two electrons in each orbital and once you have two electrons that's it, you can't go any more, you have to then fill the next higher energy orbital, okay? So, what this means is that you actually prefer not to put any electrons into the anti-bonding molecular orbital, we only do it if we have no other choice because you have too many electrons and we need to put one up there. So, how many electrons are being shared in the atomic orbitals? two, so that means that when we make our new molecular orbital, this is molecular orbital pi-1, this is molecular orbital pi-2, these are the potential pi bonds that we could make based on the atomic orbitals that are overlapping and what we know 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 going to have 0 electrons up in the antibonding region, okay? But, let's say that one of these orbitals 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 to Aufbau principle. 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 antibonding orbital, which means that it's not going to promote bonding between the two carbons, it's actually going to make them less stable and going to want to make them break apart, okay? So, basically, when we build out our molecular orbitals, we're going to be using this system and I'm going to teach you how to write out all your atomic orbitals, how to write out your molecular orbitals but it's important that you guys know that we're going to always be starting from the lowest energy state and then only 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 alkenes 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 orbitals, 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 going to follow this up with more videos explaining exactly what you need to know so you can apply molecular orbital theory to solve whatever problems you have."