Intermolecular Forces

So, when we discussed Intramolecular Forces, we discussed the forces that keep molecules together. But there are also forces that keep molecules attracted to other molecules within a substance. This is illustrated above in the diagram.

These are known as intermolecular forces(IMFs) and, believe it or not, they are the reason why certain compounds are solids at room temperature and others are liquids or gases at the same temperature. This is expanded on in the page for States of Matter.

London Dispersion Forces

So, the first IMF to cover are London Dispersion Forces(LDFs), which are found between every molecule, regardless of their chemical or physical nature. So, how do LDFs work? They work because of the "random" motion of electrons around atoms. Remember that atoms are located(generally) within orbitals and don't have fixed circular orbits around atomic nuclei.

Because of their "random" motion, this means the electrons within a molecule have a chance of being distributed asymmetrically, with there being a few more electrons on one side of an atom than another. This can be represented by the diagram above, where the electrons happen to be distributed so that there are more electrons on the right side of both molecules. If there's more electrons on one side, that means the molecule is a temporary dipole and there will be attraction between the positive end of one molecule and the negative end of another molecule.

LDFs are very weak because they are temporary, meaning that once the molecules move away from each other, the LDFs are pretty much negligible.

However, LDFs can be made stronger if you increase the polarizability of the molecules involved. So, how do you increase the polarizability? Well, it's quite simple really. You have to increase the number of electrons in the molecule. This is because if you increase the number of electrons, there is a better chance that those electrons will become polarized. If there's a better chance of polarization, the electrons will form a stronger dipole. So, how do you increase the electrons of a molecule? Well, heavier atoms have more electrons since if they're heavier, they're likely to have more protons and thus more electrons. This means that molecules with higher molar mass have greater LDFs. For two molecules that only exhibit LDFs, the molecule with a higher molar mass will then exhibit stronger LDFs.

Dipole-Dipole Attractions

Recall that polar molecules have a positive end and a negative end, forming an electric dipole. Dipole-dipole(DD) attractions form between two polar molecules because the positive end of one molecule attracts the negative end of the other, causing them to stick very close.

DD attractions are stronger than London Dispersion Forces(LDFs) because they are more permanent than LDFs. However, in principle, they are no different from LDFs except for the fact that DD attractions involve permanently polarized molecules instead of temporarily polarized molecules.

Hydrogen Bonding

So, it is established that Dipole-Dipole(DD) attractions are stronger than London Dispersion Forces(LDF), but there is a special type of DD attraction that's even stronger than regular DD attractions known as Hydrogen Bonding.

A hydrogen bond involves hydrogen(H) bonded with either Fluorine(F), Oxygen(O), or Nitrogen(N). Why these 3 atoms and why hydrogen? Well, Hydrogen has a very low electronegativity while these 3 atoms have high electronegativity, making the electronegativity difference great, causing a very strong electric dipole between H and these 3 atoms. Also, Hydrogen is the smallest atom so it gets attracted even closer to the highly electronegative atom, making the bond extremely strong. The more hydrogen bonds a molecule has, the stronger its hydrogen bonding.

So, the general rule of thumb to find out if a molecule has hydrogen bonding is if its has H bonded to either F, O, or N. This can be seen in the diagram above showing the hydrogen bonds in water, one of the most notable compounds that exhibit hydrogen bonding.

Ion-Dipole Attractions