Pi Bonds – Chemical Bonding

00:01 Okay. Now I want to bring you on to another type of molecular orbital. And this is the pi molecular orbital, okay? So this is where we've got pi bonds formed. Pi bonds are formed from the overlap of 2p orbitals which is, as you can see, are in line with each other.

00:21 Remember what we said about the fact things need to be in phase, or electrons need to be in phase, and that the lobes of a p orbital show out and in phase. And in this particular case, what we're seeing here is that the overlap between the two green lobes and the overlap between the two orange lobes of the p atomic orbitals successfully forms a pi molecular bond, a pi molecule orbital. Imagine for a second, if you will, I was to have taken one of those p atomic orbitals and flipped it upside down. You would not get a pi orbital formed. In fact, you get something called an antibonding orbital (which is, for the purposes of this course, beyond our terms of reference but something you can look up if you wish). So the geometry and the phase of the bonding orbitals from the atom must be correct in order for a bonding molecular orbital to form. And as you can see, you get orbitals perpendicular to the bond axis—not in line with, but perpendicular.

01:30 Pi bonds almost always occur between atoms which are already bonded via a sigma molecular orbital. And that's pretty much the rule. If we take, for example, a look at this very simple alkene molecule, and of course, we'll come on to alkenes in more depth later on in the course when we look at homologous series, but for the moment, let's try and consider it in isolation. Here, we have ethene. Note two carbons bound together with four hydrogens, two binding to each carbon. Unlike the methane structure, which is tetrahedral, ethene itself is flat. All of its atoms are on a single plane. And this is what happens when you have a so-called double bond. A double bond almost exclusively is formed from a sigma molecular orbital and then a pi molecular orbital being formed on top. So how can this be accounted for? Let's have a look at our hybridization again and look at pi bonds. For each carbon atom, two p orbitals and one s orbital hybridize to give three sp2 orbitals, but the remaining 2p orbital on the carbon remains untouched. It is not hybridized. And so as a direct result, you have only three uniform hybridized orbitals in an angle around that single carbon, and this gives rise to sp2 hybridization and a so-called trigonal planar arrangement, where each of the molecular orbitals is separated by approximately 120 degrees, as shown in the diagram. So let's have a quick look at what I mean.