What is the molecular orbital theory? How are molecular orbitals determined? Who created the molecular orbital theory? How do you find hybridization orbitals? 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.
All right, so once again we have four SP three hybrid orbitals, and each one of these hybrid orbitals is gonna have an electron in it, so we can see that each one of these SP three hybrid orbitals has one electron in there, like that, and so the final orbital, the final hybrid orbitals here contain 25 percent S character. Let me go ahead and write this down here: So 25 percent S character, and 75 percent P character, in this new hybrid orbital.
Once again, that's because we started out with one S orbital and three P orbitals for our hybridization. All right, let's go back to methane, and let's go ahead and draw in a picture, because now we know that this carbon is SP three hybridized, so let's go ahead and draw a picture of that hybridized carbon, here. So we're gonna go ahead and draw in our carbon, and we know that it has four SP three hybrid orbitals, and once again, when we draw the orbitals, we're gonna ignore the smaller back lobe here, so it doesn't confuse us.
So we go ahead and draw in, here's one of our orbitals, for carbon, so that's an SP three hybrid orbital. Here's another SP three hybrid orbital, here's another one, and then, finally, a fourth one. So let's go back up here, to this picture, 'cause once again, we need to show that each of these hybrid orbitals has one valence electron in it, so I can go ahead and put in my one valence electron, in each of my hybrid orbitals, like that, so here's our valence electron.
If we're talking about methane, so carbon is bonded to four hydrogens, each hydrogen has an un-hybridized S orbital, and each hydrogen has one electron in that, so I'm gonna go ahead and sketch that in; let's go ahead and use blue here.
So here's an un-hybridized S orbital, I'm gonna go ahead and draw these in, so an un-hybridized S orbital; each one of these un-hybridized S orbitals for hydrogen has one valence electron, so I'm gonna go ahead and put in those one valence electrons, in here, like that, so same for here, and then, finally, for here.
So this is just one picture of the methane molecule, so this is hydrogen, these are all the hydrogens right here, like that. All right, let's think about this bond that we formed right here, so here we have an overlap of orbitals, an overlap of an SP three hybrid orbital form carbon, with a un-hybridized S orbital from hydrogen here, and so this is a head-on overlap, so we're sharing electrons here, in this head-on overlap.
And a head-on overlap, in chemistry, is called a sigma bond, so this is a sigma bond, sigma bond here, a head-on overlap, and this happens three more times in the methane molecule. So here's a head-on overlap, here's a head-on overlap, and here's a head-on overlap. Orbital configuration for carbon atom.
Figure 2. Promotion of carbon s electron to empty p orbital. Figure 3. Carbon sp 3 hybrid orbitals. Figure 4. Summary Electrons hybridize in order to form covalent bonds.
Nonequivalent orbitals mix to form hybrid orbitals. Practice Use the link below to answer the following questions. What contributes to these unexpected bond angles? What happens to the lone pair electrons in ammonia when hybridization occurs? Does the same thing happen with water? Review Why is carbon expected to form only two covalent bonds? How many covalent bonds does carbon actually form? What needs to happen to allow carbon to form four bonds? Learning Objectives Describe the formation of sp and sp 2 orbitals.
How do you open the closed circle? Figure 5. Promotion of Be 2s electron. Figure 6. Be hybrid orbitals. Figure 8. Promotion of 2s electron. Figure 9. Formation of sp 2 orbital.
0コメント