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If NaCl dissolves in water, it is said that it happens because the hydration energy is greater than the lattice energy. If so, the enthalpy of solution attained is about +3.88kJ/mol, implying that it is endothermic.

From what I have understood, the hydration enthalpy is exothermic and thus its effect should be noticeable. That is, when we add up the lattice and hydration, the negative sign should have come in our final answer. Why does this not happen? Where have I gone wrong?

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    $\begingroup$ You don't add up. You destroy the lattice, so the lattice energy should be subtracted. If you think of it, the answer (and even sign thereof) is not all that obvious. $\endgroup$ – Ivan Neretin Dec 28 '16 at 10:20
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    $\begingroup$ Spontaneity of a process (dissolution in this case) is determined by Gibbs free energy, not enthalpy. So your first statement is categorically wrong. There are very many salts with endothermic dissolution enthalpies that still dissolve. I don't know off-hand if NaCl is one of them, but the point still stands. See: Why does a substance with an endothermic heat of solution dissolve? $\endgroup$ – orthocresol Dec 28 '16 at 10:20
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For any salt dissolving in water, its not compulsory that enthalpy change needs to be negative. Yes, if it is negative then its better dissolving for sure, But it doesn't mean the enthalpy change has to be compulsorily negative for salt to dissolve in water.

Enthalpy change when we dissolve Nacl in water is positive. This is because breaking of ionic bonds in the crystal lattice of Nacl requires more energy than that of provided by hydration of Nacl. Yes, hydration energy for Nacl is less than the energy required to break the Nacl apart. Hence the extra energy which can add up with the hydration energy to break the bonds of Nacl apart comes from the heat of the water. So the extra amount of energy which hydration could not provide comes from the heat of the water. Hence water looses heat, becomes slightly cold and also hydration energy (exothermic process) is less than the required energy to break the bonds, hence the net enthalpy change is positive.

So till now we learnt that for Nacl hydration enthalpy is not greater than enthalpy required to break the bonds and hence the extra energy required to satisfy the bond breaking comes form the heat of the water, and as hydration enthalpy is less than required energy to break the bond of Nacl, the net enthalpy change is positive, hence endothermic process.

Now the question arises how comes then Nacl dissolve in water? As already said its not compulsory that enthalpy change needs to be negative for a salt to dissolve in water. Remember my friend, dissolving depends upon ΔG. ΔG is the most important factor which decides weather the reaction will happen or not. If ΔG is negative and ΔH is positive, salt will for sure dissolve because dissolving of salt (or any reaction) depends on ΔG. So if ΔG is negative for the dissolving of salt then salt would dissolve. Infact for Nacl ΔG is negative. Lets see how.

ΔG= ΔH −TΔS.

In dissolving Nacl in water ΔS is very much positive because is dissolving Nacl, Na and Cl gets apart, increasing the disorderedness in the system, hence ΔS is highly positive. So practically TΔS term is higher than ΔH term and hence we end up having ΔG negative.

Hence as ΔG is negative, Nacl dissolve. Remember dissolving of Nacl depends on ΔG. Negative ΔG makes Nacl dissolve in water.

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  • $\begingroup$ If the enthalpy of dissosiation was influenced by the heat from the water, then why doesn't it add up to the effect of the hydration enthalpy? $\endgroup$ – KV18 Feb 28 '17 at 16:35
  • $\begingroup$ I do feel that your answer is correct, although when we add up the lattice and hydration energy, the hydration energy is still lesser than the lattice energy. But then if the heat was being used, then the overall energy should surpass the magnitude of the lattice energy, right? In that way, the answer would have been negative! However, I do feel that the equation in my textbook that says the enthalpy of solution equals enthalpy of hydration and lattice is incomplete. I mean, it should have mentioned the heat from the water right? $\endgroup$ – KV18 Feb 28 '17 at 16:37
  • $\begingroup$ sorry for my delayed reply! Ya in textbooks they don't go through detailed information. However what I mentioned is what exactly happens. Rather hydration enthalpy plays a very important role. For example in reduction potential of Co, Ni,Mn,Cr are all +ve. For reduction we require ionisation. Ionisation enthalpy of these transitive elements is very high. So think from where do they get the energy? Its from hydration energy. Infact hydration enthalpy is much higher than the ionisation enthalpies (1st and 2nd) and hence making electrode potential positive. So hydration enthalpy is imp factor. $\endgroup$ – Yb609 Mar 14 '17 at 16:01

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