Molecular Orbital Theory: Study with Video Lessons, Practice Problems & Examples

Molecular orbital theory explains how atomic orbitals combine to form molecular orbitals, influencing bonding and stability in compounds like alkenes. Constructive interference leads to bonding molecular orbitals, enhancing electron density between atoms, while destructive interference results in antibonding molecular orbitals, reducing stability. Understanding hybridization, such as sp², is crucial for predicting molecular behavior and reactivity, particularly in reactions involving conjugation and resonance. This foundational knowledge is essential for grasping complex organic reactions and mechanisms.

Review of Atomic Orbitals

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 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 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 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 mean? Resonance means that you're sharing electrons from one atom to another. You guys remember that? You can resonate electrons, etcetera. Well, one of the technical ways that you can talk about resonance is that resonance happens from nonbonding orbitals to adjacent nonbonding orbitals. What do I mean by nonbonding? That it's 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 you can 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 1s orbitals, 2s orbitals, etcetera. 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 2nd 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. What 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 1. This is one of the first things you learned in organic chemistry 1, 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 3 bonds, so an alkene carbon has 3 bonds attached to it, which would mean what we call them 3 groups or 3 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 it have attached to it right now? It has a hydrogen, a hydrogen, and a carbon, meaning that there are 3 bond sites, meaning that this mean that this equates to an sp2 hybridization. Remember that? That you're supposed to know that however many bond sites there are, that's how many, that's what your hybridization is. And 3 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 6. Carbon has an atomic number of 6, which means that at its neutral state, how many protons does it have? 6. How many electrons does it have? 6. So when we build the electron configuration of carbon, we need to figure out where all those 6 electrons are going to go. Right? And remember that you always start with Aufbau 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 6 electrons, where should those electrons go? Well, 2 of them should go into the 1s orbital because that's the lowest energy state orbital possible. Then another 2 of them should go into the 2s orbital because that's the next highest, right? And then we have Hund's rule. Remember 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 2 kids on one seat and 0 kids on another seat. So in this case, notice that the p orbitals are all the same energy state, right? So that means that I would then get one electron in the 2px, one electron in the 2py, and now I ran out of electrons. I just put all my 6, meaning that there's no electrons for the 2pz. 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 of just like, hey, there's 6 electrons and we need to put the put them into the electron configuration, this is what it should look like. But remember that 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 4 bonds, right? But right now, the way that you have the carbon set up, the 2s 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 2 orbitals could make a bond because they could accept 1 electron. And then this one has no electrons, so it's not very good at making a bond because it would have to accept 2 electrons, not just 1. 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 3 groups or 3 bond sites, what happens is that the 2s orbital blends with 2 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 3 bond sites equals sp2, which means that 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 sp2sp2, sp2. Okay. Now you might be wondering Johnny, why did you put 2 sp2? Well, because technically you can just call it sp2, but you can also call it 2 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 3 bonds and and they can all blend together in this way. Notice that now what happens is instead of getting 2 electrons in a lower energy orbital and then 2 electrons in higher, now we have is 3 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 sp 2 23 bond sites, that means that there's a 4th bond that's not being made. That means that 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 2pz. 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, potentially make a bond or to interact with other orbitals, okay? So this is gonna be what we keep, what we're gonna call our nonbonding orbital and this is gonna be the one that's gonna be the really interesting one for us for the rest of this section. We're gonna be talking 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 2 sp2s and talk about what they're doing. Okay? Well, remember that we said that it's making 3 bonds, right? So what we could do is we could show where those 3 bonds are happening. 1 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? Hydrogens have an atomic number of 1, 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 2 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 if you just have a hydrogen. What it actually is, is it's another sp2 orbital, another sp2 hybridized orbital because notice that it has 3 bond sites. So you're going to have 1s and 2p's blend together and give you an sp2. And what's going to happen is that those 2 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. So 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. Okay. They're all overlapping in one place, giving us 3 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 3 sigma bonds, 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 in a in 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 that you started with, where it has one electron and basically it's overlapping with another, basically 2pz. Does that make sense? It's overlapping with... **[MESSAGE CUTS OFF]**

Review of Molecular Orbitals

When adjacent non-bonded atomic orbitals overlap with each other or are next to each other, they create more favorable molecular orbitals. A molecular orbital is simply the overlap of a few atomic orbitals. If you want to know what the molecular orbital is going to look like, we can use a system very common in organic chemistry called the linear combination of atomic orbitals (LCAO), which helps you predict what the molecular orbital will look like.

Now that I've hopefully convinced you that atomic orbitals like to share electrons, I want to talk about what those electrons look like after they share. We're going to take the example of ethene. Recall that one orbital has one electron, another orbital has another electron—these are the atomic orbitals, and this is typically how we represent them. Each of the conjugated atoms will receive one atomic orbital. Notice that I have 2 conjugated atoms, so I draw 1, 2 atomic orbitals next to each other and place however many free electrons there are into those orbitals. Atom 1 is donating one electron, which is why I put one electron there. Atom 2 is donating another, which is why I add another electron.

Remember, an orbital is just a region of space that is statistically probable to have electrons. It's like a cloud of electron density where there's a high chance we'll find electrons, but it's not actually a particle. When you bring atomic orbitals close together, they don't collide; they interfere with each other like waves, not like particles. This interference can be constructive or destructive.

Constructive interference means that the waves of those atomic orbitals build on each other, increasing their amplitude and the chances of finding electrons between them. This is called an in-phase overlap, forming a bonding interaction, which means that the chances of finding electrons between these two atoms is unusually high.

On the other hand, destructive interference occurs when atomic orbitals are out of phase, causing the waves to cancel each other out, creating a node where there is no mathematical chance of finding electrons. This leads to an antibonding interaction, making the atoms unstable and likely to repel each other.

It is important to note that the positive and negative lobes of an atomic orbital do not relate to electrical charges; it's just a way to think about orbitals. During constructive overlap, the whites (or positives) and grays (or negatives) are on the same side, aligning properly; during destructive overlap, they are opposite each other.

According to the Pauli exclusion principle, you can only put two electrons in each orbital. When we make our new molecular orbitals, molecular orbital pi 1 and molecular orbital pi 2, based on the overlapping atomic orbitals, both electrons can fill the lowest energy orbital, creating a bonding interaction with no electrons in the antibonding region. If we had an extra electron, it would go into the higher energy, antibonding orbital, destabilizing the bond.

In conclusion, the reason alkenes can form such good double bonds is that they have exactly two electrons to share constructively in one molecular orbital, effectively making it look like a single, low-energy molecular orbital that promotes bonding between the two. I hope this is a good start to understanding molecular orbital theory, and I will follow up with more videos explaining exactly what you need to know so you can apply this theory to solve problems.

Here’s what students ask on this topic:

What is molecular orbital theory in organic chemistry?

Molecular orbital theory (MO theory) in organic chemistry explains how atomic orbitals combine to form molecular orbitals, which are spread over the entire molecule. These molecular orbitals can be bonding, antibonding, or nonbonding. Bonding molecular orbitals result from constructive interference of atomic orbitals, increasing electron density between atoms and stabilizing the molecule. Antibonding molecular orbitals result from destructive interference, decreasing electron density between atoms and destabilizing the molecule. Understanding MO theory is crucial for predicting molecular behavior, stability, and reactivity, especially in reactions involving conjugation and resonance.

How do atomic orbitals combine to form molecular orbitals?

Atomic orbitals combine to form molecular orbitals through the process of linear combination of atomic orbitals (LCAO). When two atomic orbitals overlap, they can interfere constructively or destructively. Constructive interference leads to bonding molecular orbitals, where electron density between the nuclei increases, stabilizing the molecule. Destructive interference leads to antibonding molecular orbitals, where electron density between the nuclei decreases, creating a node and destabilizing the molecule. The resulting molecular orbitals can accommodate electrons from the combining atomic orbitals, influencing the molecule's overall stability and reactivity.

What is the difference between bonding and antibonding molecular orbitals?

Bonding molecular orbitals result from constructive interference of atomic orbitals, where the electron density between the nuclei increases, leading to a stable interaction. These orbitals are lower in energy compared to the original atomic orbitals. Antibonding molecular orbitals result from destructive interference, where the electron density between the nuclei decreases, creating a node with zero electron probability. This interaction is unstable and higher in energy compared to the original atomic orbitals. Electrons in bonding orbitals stabilize the molecule, while electrons in antibonding orbitals destabilize it.

What is the significance of hybridization in molecular orbital theory?

Hybridization is significant in molecular orbital theory as it explains the mixing of atomic orbitals to form new hybrid orbitals, which can then combine to form molecular orbitals. For example, in sp² hybridization, one s orbital and two p orbitals mix to form three sp² hybrid orbitals. These hybrid orbitals can overlap with other atomic or hybrid orbitals to form sigma bonds, while the remaining unhybridized p orbital can form pi bonds. Understanding hybridization helps predict the geometry, bonding, and reactivity of molecules, particularly in organic compounds like alkenes.

How does molecular orbital theory explain the stability of alkenes?

Molecular orbital theory explains the stability of alkenes through the formation of bonding and antibonding molecular orbitals. In alkenes, the sp² hybridized carbons form sigma bonds, while the unhybridized p orbitals overlap to form a pi bond. The pi bond results from the constructive interference of the p orbitals, creating a bonding molecular orbital with increased electron density between the carbons, stabilizing the molecule. The absence of electrons in the antibonding molecular orbital further enhances stability. This delocalization of electrons in the pi bond contributes to the overall stability and reactivity of alkenes.