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I understand the idea that the solvation enthalpy (ΔH-) must exceed the lattice enthalpy (ΔH+) for a solution to be more energetically stable, but what type of bond is formed between the disassociated ions and the the polar solvent molecules? Surely it must be even stronger than the ionic bonding in the lattice if it releases more energy than the energy needed to break the ionic bonds in the ionic compound.

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Firstly, There is no such bond existing inbetween the polar solvent molecule and the solute ions. It is ion-dipole interaction which exists between the polar solvent molecules and the ions of the solute.

I hope you got your answer.

I wanted to add something more which I think you got wrong.

I agree that if hydration enthalpy is greater than lattice enthalpy then solute will dissolve but it is not compulsory for hydration enthalpy to be greater than lattice enthalpy for dissolving. What i mean is: For dissolving its important to break the ionic bonds between the ions of the solute molecule but that energy required to break the ionic bond is not necessarily from hydration only. That is sometimes hydration enthalpy couldn't meet the energy required to break the ionic bonds (lattice energy). In those cases the extra amount of energy can be from heat of the solvent. In the cases where hydration enthalpy is larger than lattice energy, ΔH is negative. In cases where hydration enthalpy is lesser than lattice energy, in those cases the more energy needed is driven from the heat of the solvent and in this case ΔH becomes positive. Negative ΔH is preferred to be better for dissolving because it helps in making ΔG more negative. But it doesn't mean solute with positive ΔH doesn't dissolve. If ΔS is highly negative, then even though ΔH is positive still it can make ΔG negative and hence dissolving possible. Dissolving of any solute depends on ΔG. For dissolving, ΔH doesn't matter as long as ΔG is negative. Say for example in case of Nacl, net solvation enthalpy is positive, but still it dissolves in water.

Nacl hydration enthalpy is not greater than enthalpy required to break the bonds in between Na+ and Cl-. As hydration enthalpy is not sufficient to release energy which is required for bond breaking incase of Nacl, hence the extra amount of energy required comes from the heat of the water ,which in addition to the hydration enthalpy is able to break the bonds in between Nacl. As energy released is less than net energy absorbed, hence net enthalpy (solvation enthalpy) is positive.The reason behind the dissolving of the salt is ΔS being highly positive which in turn ends up making ΔG negative as ΔG= ΔH −TΔS.

So for a well dissolving ionic solute, hydration enthalpy is always not greater than the energy required to break the ions apart in the solute molecule. An ionic solute with positive ΔH would still dissolve if ΔG is negative.

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