London dispersion forces and dipole-dipole forces are subsets of Van der Walls forces, which themselves are the weakest intermolecular forces. Intermolecular forces directly determine properties like melting point, boiling point, and viscosity. London dispersion is most easily noticed in nonpolar molecules, where stronger IMFs don’t drown out the London dispersion.
London dispersion forces are the weakest of the three types of intermolecular forces. London dispersion forces are actually what hold many substances together. Real-life examples include gasoline and petroleum jelly, which are mostly make of alkanes of varying lengths. Alkanes are non-polar and exhibit no hydrogen bonding, so Van der Waals London dispersion are the only IMFs they exhibit.
Temporary polarity resulting from unsymmetrical distribution of electrons creates an attractive force in otherwise nonpolar molecules like Br2.
Let’s say we’ve got a single dibromide molecule like the one above. The two atoms in the molecule are identical, so it might seem at first that the electrons will have a symmetrical distribution, but that’s an oversimplification.
Electron density will actually oscillate—this happens because electrons aren’t point-objects; it might help to think about them as clouds of probability. Those partial charges will oscillate as electron density moves around; those partial positive charges result from lack of electron density.
Of course, molecules don’t exist in isolation. Let’s look at what happens when there are more dibromide molecules in the area. Those temporary dipoles will actually induce dipoles in neighboring molecules. That makes sense if you think about it in terms of electron density. Electrons will repel each other, so a partial negative in one bromine will push electron density away in a neighboring dibromide. This results in a relatively stable attractive force between the partial negative and induced partial positive.
The dibromides will actually line up with each other with opposite partial charges facing each other. This trend can continue indefinitely and isn’t limited to a single dimension like in the diagram above.
The strength of London dispersion forces is influenced by molecular size and molecular shape. Basically, the larger the molecule is, the more surface area there generally is for partial charges to develop. Based on that idea, which of the following molecules would you expect to have the higher boiling point? Remember that boiling point is directly determined by intermolecular forces.
Ding ding! The answer is pentane! The boiling point of propane is -42ºC and the boiling point of n-pentane is 36.1ºC.
So that tackles the size… but what about shape? Let’s look at three molecules with 5 carbons: cyclopentane, n-pentane, and 2-methylbutane. Try to order them in order of increasing boiling point. Think about the ways that the molecules can arrange themselves to maximize contact.
Cyclic compounds can stack pretty easily since their sigma bonds don’t really rotate too much, but chain compounds do rotate around a lot. Branched chain hydrocarbons also rotate, but they have the extra substituents that prevent overlap.
The general order of increasing London dispersion forces, and therefore boiling points, is ring > straight chain > branched chain. The boiling points of the above compounds are as follows:
Boiling Point (ºC)
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