Hybridization occurs when an atom bonds using electrons from both the s and p orbitals, creating an imbalance in the energy levels of the electrons. To equalize these energy levels, the s and p orbitals involved are combined to create hybrid orbitals. Hybridization is a key concept in valence bond theory, but alternate models are proposed in molecular orbital theory.
Alright so we're going to talk about hybridization of atomic orbital and just like you'd think when something is hybridized that you're familiar with let's hybrid cars or hybrid brakes. They're 2 ideas of, 2 separate types of cars 1 being electric and one being gas and kind of combine them together to make a new hybrid car it's a completely new different type of a car. Or bikes for example a mountain bike and a road bike they come together to form a new type of bike it's called a hybrid bike. We're going to do the same thing with atomic orbital, we're going to bring 2 different orbitals together to make a completely new orbital we're going to call that process hybridization. This is a process in which atomic orbitals are mixed to form new identical orbitals this is, are identical in energy.
Alright so let's talk about why we're actually going to do this, why is this even necessary. So let's start with an example carbon tetrachloride, alright so carbon tetrachloride is a single carbon atom bonded to 4 chlorine atoms. So we agree we have to agree that all these bonds, these 4 bonds are equal in energy. One is not more energetic than the other, they all have, one is not more special or anything than the other one. So how do we make sure that, that's the case? When you look at that orbital, that carbon has in the valence shell we have a 2S orbital and we have 3 equal in energy P orbitals. Okay so we know carbon has 4 electrons, we're going to denote them.
Okay so we want to have 4 equal places where chlorine come in and bond with this carbon. So we're going to hybridize all these orbitals to make 4 equal in energy orbitals. So we're going to have 4 new orbitals and we're going to call them the 1S and the 3 of them are P so we're going to call it SP3, 1 from S 3 from P. And we're going to spread these out just like the rule tells us to and we're going to say okay we have 4 electrons which gives us 4 equal places for chlorine to come in and actually bond with that carbon. So chlorine is coming here, here, here and here and make 4 of those bonds like you see in the picture.
Alright so what actually would get hybridized? What do we create to actually mix it so that these equal orbitals are necessary? So we know all single bonds are going to be hybridized because a single bond there's not one that's more energetic than the other. So all single bonds are going to be hybridized. Because they're hybridized bonds we're going to now call single bonds sigma bonds, this is just the way they overlap, the way that orbitals overlap we're going to call them, denote them sigma bonds. And we also want to say that low in pairs are also going to be hybridized because they're not higher or lower in energy than those bonds either. So let's look at ammonia as an example, ammonia if you look at nitrogen within ammonia it has these 2 lone pair of electrons. So ammonia before had the same thing, ammonia has 5 valence electrons so 2, 3, 4, 5, this should be the same I'm sorry they're kind of uneven, they should actually be the same in energy and we have the 5 electrons.
If we're going to hybridize all of them we need to have 1, 2, 3 of these are the same along with this fourth one so we need to have all 4 of these is the same, so we're goingto have again 4 equal in energy we're going to call it SP3, 1 from S, 3 from P 1, 2, 3, 4, 5 here's our lone pair and here's the hydrogens that are going to come in and bond with them all equal in energies so we have this new hybrid orbitals. Okay about when we have multiple bonds? So in different cases we may have multiple bonds, double bonds and triple bonds. So what happens in those guys? Well one of those bonds within a multiple bond is called a sigma bond and again don't forget sigma bonds are hybridized so one of those bonds is going to be hybridized.
The rest of those bonds are called pi bonds, those pi bonds are just P orbitals overlapping each other, they're only P orbitals they're a little higher in energy, they actually are different in energy. So we're going actually keep then separated, so we have 1 sigma and 1 pi. So let's look at carbon dioxide here, we have a double bond alright of these is going to be a sigma bond and we're going to denote that with a sigma and one of the bonds is going to be a pi bond we'll denote it with pi. These are only P orbitals, these are hybridized orbitals we're just talking about the carbon right now. Okay so carbon we already know looks like this we're going to save 2 of the P orbitals and I don't care which 2 I save it doesn't really matter, they're all the same in energy I donât care.
I'm just going to save this just for practical purposes, these are going to be the ones to use in pi bonding, so I'm going to save those so they hybridize 1S and the other P. SO instead of being SP3 this time it's going to be SP, 1S and 1P SP orbitals. Let's look at the oxygen, oxygen also has a sigma bond, a pi bond but notice it has lone pair. So the sigma bond and the lone pair are going to be hybridized but not the pi bond we're going to leave this lone. So we need 3 hybrid orbital, 1 from S and 2 from P, so it's going to be SP2. So this guy is going to be SP2, and this guy was SP don't forget it's orbitals were SP. It doesn't matter that these guys have different types of orbitals, we just want to make sure that the orbitals within the atom itself are the same.
So let's look at triple bond, triple bond one is sigma don't forget. So we're going to say this is sigma and these 2 are pi. Okay, so these 2 are just P we're going to ignore them and this nitrogen has, needs 2, 1 for the lone pair and 1 for the sigma hybridized orbitals. So we need 2, 1 from S, 1 from P so we're going to call this SP, this also is just SP. Okay look at Ozone, O3 all 3 of these are a little bit different so this guy has 1, 2, 3 lone pair and 1 sigma bond so it needs 4, 1 from S, 3 from P so it's going to be SP3. Oxygen is going to need, it has 2 sigma bonds 1 from here, 1 from here and lone pair so it needs 3 so it's going to be SP2 1 from S, 2 from P. This guy over here has 1 sigma bonds and 2 lone pairs again 3 hybridized orbitals, so it's going to be SP2, 1 from S, 2 from P. So hopefully that made it a little bit easier for you to figure out hybridized orbitals.