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Why do certain solutes dissolved in water release heat and other solutes absorb heat?

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To understand this phenomenon first we must point out that there are very different solid substances and solvents. Different combinations of them lead to different interactions and properties.

Anyway for the regular case - ionic or soluble compound in water - there is at least two steps to realize the solution:

Edited Answer:

  1. Break the solid structure (eg. the crystal lattice). This step generally absorbs energy (endothermic, see lattice energy);
  2. Solvate the compound. Because there are interactions between matter, the solvent might organize itself around the compound (solvation shell). Those molecules involved are not any more available, they are tied to the compound, they are not any more free water. This step generally releases energy (exothermic, see solvation energy). You can think about this step as a reaction of the solvent to recover form the intrusion by coating the intruder with one or several layers of its molecules. Then it 'feels' if there were only water. You will find details of the complex solvation process in good Physical Chemistry book.

Succeeding these two unit steps leads to dissolution. Solubility is the thermodynamic limit beyond that you cannot dissolve further molecule of a substance. Off course this is related with every substance present in the solution and available free water. Solubility are generally expressed for pure substance only.

When you realize your solution, speaking of heat exchanged with the environment, you experienced the balance between those two reactions. This is why the heat balance can lead to a global endothermic or exothermic reaction.

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  • $\begingroup$ Breaking a crystal lattice is endothermic. Solvation is exothermic - talking about enthalpy. Entropy terms then play a role as well. Your answer is good, but it says that breaking the crystal lattice is exothermic, and solvating is endothermic. Isn't it the other way round?? $\endgroup$ – Swedish Architect May 10 '14 at 9:50
  • $\begingroup$ Thanks to point out this mistake. I double checked it in Atkins Physical Chemistry. Answer has been edited. $\endgroup$ – jlandercy May 10 '14 at 10:34
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Enthalpy of hydration is exothermic, Lewis acid-base neutralization. Entropy of dissolution can be either positive or negative. Negative - ordering effect of ion on solvent is greater than the entropy increase of the crystal (highly ordered) lattice breaking down. Positive - increase in entropy because the solvent hydrogen bonding is disrupted. Dissolving KOH is a very large exotherm, Dissolving urea in water is a very large endotherm. This is descriptive and qualitative,

http://www1.lsbu.ac.uk/water/kosmos.html

We invite the physical chemists to put a quantitative mathematical model on it.

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    $\begingroup$ "Entropy of dissolution can be endothermic" is quite a poor way to word it. Endo and exothermicity are related to enthalpy only, so perhaps you could clarify your sentence? And I'd like to reinforce that the linked site is exceptional. It is definitely one of the most complete descriptions of a single substance I have ever seen. $\endgroup$ – Nicolau Saker Neto Apr 9 '14 at 22:30
  • $\begingroup$ Entropy at J/K-mole is ~1000-fold smaller than enthalpy at J/mole (KJ/mole typical), but temperature takes up a lot of slack. Water is extremely ordered re its large specific heat via hydrogen bonding, offering a large free energy sink and temperature drop when disordered other than by adding energy as heat. $\endgroup$ – Uncle Al Apr 9 '14 at 23:27
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    $\begingroup$ Entropy cannot be qualified as endo or exothermic, those terms are for Enthalpy only. Please rewords for the sake of your answer and the exactness of thermodynamic. $\endgroup$ – jlandercy May 10 '14 at 7:42

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