London Dispersion Forces

My next few posts will discuss intermolecular forces. The textbook put this information in the same chapter as functional groups, because functional groups play such a huge role in the type and strength of intermolecular forces a compound has. I don't know if this is the best approach, but I'm too lazy to figure out a different one and whatever, this works.

Hopefully the word "intermolecular" tips you off to the fact that these are forces that occur between molecules. If not, what is wrong with you. Intermolecular forces are much weaker than normal covalent or ionic bonds. If they weren't, they wouldn't be intermolecular because they'd be binding atoms as tightly as the bonds within the molecules and the molecules wouldn't really have distinct identities at all.

London dispersion forces are the weakest and most ubiquitous ones we'll look at. In this textbook and in the college chemistry classes I took, they're more often referred to as van der Waals forces. But from what I can tell, that term is actually more inclusive, so I'll be calling them "London dispersion forces" or "dispersion forces" or maybe "London forces" until I have some reason not to. Do note that it seems to be common to refer to these as "van der Waals" forces/interactions. They're also known as induced dipole–dipole forces.

London dispersion is so ubiquitous in organic chemistry in part because it's the basis for intermolecular attraction in hydrocarbons. The other forces don't come into play unless there are functional groups. It would be nice to have the picture in my textbook to convey what I'm talking about, because it really does seem better than what I'm finding on the web, but this ball & stick picture that uses hydrogen molecules should serve well enough...
Because of electromagnetic repulsion, positive and negative charge gather in regions of the molecule. The whole molecule is neutral in charge, but some parts of it are more positive and others are more negative. I like the methane example used in my textbook better, because it shows more randomness, but the principle is the same: the negative portion of one molecule is attracted to the positive portion of another molecule. And this interaction, magnified over massive numbers of molecules, gives the whole thing some cohesion. And when I say some, in the case of hydrogen at normal temperatures and pressures, it's not enough. That's why pure hydrogen is usually a gas. The same is true with methane. If it gets cold enough, as happens in that atmosphere of Titan, one of Saturn's moons, methane is liquid.

Not all hydrocarbons are gases at room temperature. The bigger they are, the higher the temperature has to be in order to excite the movements of the molecules enough to overcome London dispersion forces and liberate the molecules from each other. But it's not all because the molecules are simply bigger. They're also longer. And that means more surface area. My textbook points out that n-pentane has stronger dispersion forces than neopentane. Condensed structures should be enough to show why...

n-pentane: CH₃CH₂CH₂CH₂CH₃
neopentane: C(CH₃)₄

The long, straight chain of n-pentane has more surface area than the bunched up neopentane. They have the exact same quantities of each atom, but this difference has significant effects on the physical properties of these molecules.

Finally, dispersion is also affected by the polarizability of atoms. In a large atom, the electrons are further from the nucleus and are more subject to disruption and fluctuation in electromagnetic forces than in a small atom, where the electrons are more tightly held.

Oh, and a fun trivia fact that everyone loves is that geckos hold onto surfaces using these forces (the bristles on their toes are small enough to do this and well distributed enough that enough surface area is there for the forces to be strong enough to hold the lizard's whole body even when walking upside-down). Seriously. Pretty cool, huh?