Double and Triple Bond – Chemical Bonding

00:01 So let's have a quick look at what I mean.

00:05 Let's take, for example, our carbon. I've shown here an sp2 hybridized orbital (you should recognize that: big fat green lobe, very small orange lobe), and I've shown another carbon on the other side on the board. And as you can see, these two hybrid orbitals can overlap to form a sigma molecular orbital. And, as I said to you before, pi bonds are formed after sigma bonds are formed. And as you can see, pi bonding is possible because we also have the unhybridized py atomic orbital from each carbon able to overlap as long as those lobes are in the right phase. This gives rise to a rather interesting sort of hot dog / sausage structure that you can see there, where you have a sigma bond shown in green in the center and you have a pi bond which is stacked on one either side.

01:00 So the last part of this is where I want to talk to you a little bit about alkynes, or triple bonds. And I explained to you that the shapes of some organic molecules can be explained by hybridization. And I want to draw to your attention one final example, and that is ethyne. Now if we look at ethyne, which is otherwise known as acetylene (the actual use of which is mostly in welding torches, but there are a wide variety of other things that can be done with it) is that we actually have three bonds in the center. Now we know how we can get two: sigma and pi bonds. However, ethyne is a linear molecule. It is not trigonal planar; it is not tetrahedral. All of its bonds are in a line, and why is this? So let's go back to our hybridization. Let's look at carbon. Prior to hybridization, we have two electrons in the 2s, two unpaired in the 2p. However, when it comes to triple bond formation, we only hybridize two of those orbitals, and this gives rise to so-called sp hybridization.

02:13 Looking across, you can see you've got two unpaired electrons in the hybridized orbitals, and you have two unpaired electrons in the non-hybridized orbitals, the 2p. The sp orbitals have a linear geometry with the p orbitals at right angles. So instead of being one over the top and one along, you have one bond over the top, another in this direction, and also one along. Hopefully, you can see that in the diagram shown here. Here I've shown the z and I've shown the y axes. Those z and y axes are where you would see the p orbitals engaging in one pi bond and then a second pi bond 90 degrees from the first.

03:02 Sp orbitals will make sigma bonds either to carbon or hydrogen, and the p orbitals will overlap to give two perpendicular pi bonds. And this is the origin of the triple bond. And in a triple bond, of course, that means we have six electrons in total: two in the sigma and two each in each of the pi bonds that are formed. Hybridization and bonding in molecules doesn't just affect the geometry, that as you might imagine, it also affects bond energies and bond lengths. And here is a table, shown on the board, which indicates how compounds which are similar to each other have different strengths in bonds and different bond lengths, depending on the degree of hybridization and the types of bonds that are formed. In the first instance, we see ethane with a single sigma bond. The strength of this bond shown there is 376 kJ per mole. The second double bond, higher in energy of course, but it is not twice the strength of the original sigma bond. Equally, when we add a further pi molecular bond to this in the form of an alkyne, such as ethyne shown here, we again see an increase in the strength of that bond. But notice the gap between those two types—so between the single sigma, the double, and the triple bond. Notice the difference in strength and the difference in length. They do not go up as you would expect. And this is because as you add additional electron density across that system, you're weakening the bond and leaving it more prone to being broken open via reactions—reactions which we'll come on to in future lectures.