First, I would like to say that I understand the basic idea of disassociation and solubility. However, as I study intermolecular forces I feel like this basic Idea is being challenged. My understanding from basic chemistry is that ionic bonds are much stronger than covalent bonds, due to the actual exchange of electrons and the formation of cations and anions, causing an attractive force between the opposite charges to "hold" the compound together. When explaining solubility with IMF theory, however, I fail to understand how the IM forces present in water--dipole-dipole, H-bonds, and dispersion--are able to overcome that attraction between the cation and anion present in the ionic bond; because, I was under the impression that actual bonds are always stronger than IMF bonds. So, for instance, does it not infer that if, say, Sodium Chloride is placed into water, and the attraction of the h-bonds in water cause the NaCl to essentially break apart into ions, that the inherent attractions between cations and anions in an ionic compound are weaker than the attraction force of the h-bond water molecule?
It's not that ionic bonds are weaker than hydrogen bonds, but that the energy cost of breaking the ionic bonds is well compensated by the gain in solvation energy when you put NaCl in solution. Solvation energy drives the spontaneity of this dissolution process, because of entropic considerations.
An individual IMF bond is weaker than an ionic bond. But when a salt dissolves, at least 4 IMF bonds are formed for each ionic bond. This is why some salts (e.g. CaCl2) release heat when they dissolve.