# Why Ionic compounds dissolve into water

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?

• Does this help? Mar 23 '14 at 23:18

2. Each ion can participate in multiple ion-dipole interactions when dissolved, and these multiple ion-dipole interactions could potentially store more energy than the ion-ion interactions, thus making the dissolving of ionic compounds more thermodynamically favored. You can see proof of this through a few exceptions in solubility (such as $$\mathrm {Li_3PO_4}$$ not being soluble) and some types of ionic compounds never being soluble (ionic compounds with $$\mathrm{OH^-}$$ are almost never soluble), where some ionic compounds will not dissolve due to it not being more thermodynamically favored. In fact, this also explains why some ionic compounds (such as $$\mathrm{CaS}$$) decompose in water: because it is more thermodynamically favored.