London Dispersion Forces

Hi. It's Mr. Andersen and this is chemistry essentials video 16. It's on London dispersion forces which were first identified by Fritz London. That's where they get their name. After this whole idea of quantum mechanics had been developed. Can you find him in this picture? He is right here. Before we move on and talk about it, think about the forces in this picture. Which are greater? The forces within the individual people holding their atoms together? Or the forces between them holding them together in this photo? Well clearly it's the forces within your body. And so this is the first video in a set a videos that talk about bonding between molecules and within molecules. And so if we think of them all as molecules, and we move them apart a little bit, there are going to be the intermolecular forces. Those are going to be the forces between the molecules. Between this molecule and this molecule. And between this molecule and that molecule. And these are generally going to be really, really low compared to the intra molecular, within the molecule itself. And we'll talk more about those in later videos. But London dispersion forces are going to be attractive forces. These are intermolecular, between molecules. And they're going to be found in all atoms and molecules. Everything is attracted through London dispersion forces. It's really important. What really develops them are uneven electrons. As the electrons are distributed unevenly they create what are called a temporary dipole. Those are going to be a molecule that has a charge. It has a positive and a negative charge. These are going to increase as we increase the contact area between the molecules. Also as we increase the polarizability, which is going to be the squashiness of a molecule. Generally, the more electrons you have, the higher this is going to be. And so as we increase molecular size we're going to increase these dispersion forces. And finally pi bonding, which is going to be where orbitals overlap. If we have more pi bonding, we're going to increase these. And so let's look at it. Why are London dispersion forces important? Well they're in everything. And if it weren't for them we'd never have noble gases in a liquid phase. And we know we can do that if we can cool it down. And so what holds helium together when it gets into this liquid phase, is going to be these London dispersion forces. So right here I have two helium atoms. And so they're going to have two protons on the inside. Two electrons on the outside. And so as we move them close together, watch what happens to their electrons. Let's watch that again. As they move close together the electrons are going to migrate. What's doing that? Well the electrons on the outside of one helium are going to be attracted, just Coulomb's law, to the protons on the inside of that other helium. There's also going to be repulsion between those electrons on the outside. And so what this really is creating are two, what are called temporary dipoles. They have positive and negative ends. And so a London dispersion force is simply going to be the connection between those two. It's fairly simple. It's going to hold them together. Now it's just instantaneous. It's temporary. And then it's going to go away. But if we have enough of these atoms in an area and it's cool enough, then we can get that. So what's going to increase these London dispersion forces? Well contact area is one thing. The more electrons we can have closer to each other, the more of these forces we're going to find. So if we look over here at n-pentane & neopentane, they're both going to be made of the same atoms. But these ones right here are going to be looking more like a column and these more like spheres. And as a result, these ones are going to have more surface area, more contact area. So we're going to have greater forces. And so the boiling point of n-pentane is going to be higher than that of neopentane based on these London dispersion forces. What else? The polarizability. And polarizability remember is the squashiness. The more electrons we have, the more forces we're going to have. And so this right here is a protein. Proteins are massive in size. And a lot of that 3-dimensional shape comes from these London dispersion forces. Also it's going to explain why as we go down in the periodic table, let's look at the halogens right here. Halogens are going to have the same valence electrons, but as we go to fluorine, chlorine, bromine and iodine, we're actually going to go from gases, fluorine, chlorine, to liquid bromine and then solid iodine. And that's all going to be based on these London dispersion forces. As the atoms get bigger they have more electrons. We're going to have more of this uneven electron. And then finally pi bonding. That's going to have orbitals that overlap. It forms something called a pi bond. If we increase the number of pi bonds inside those molecules, we're going to increase these London dispersion forces. So again, what are they? They are attractive in nature. Found in everything. It's uneven electrons create these temporary dipoles. Remember what increases it, our list would be contact area or surface contact. Also the polarizability which is going to be based on the molecule size. And pi bonding. Can you do that? Could you explain these trends? If you can remember those three things, you've got it. And I hope that was helpful.