Enthalpy of Solution, Hsoln

The enthalpy of solution, enthalpy of dissolution, or heat of solution is the enthalpy change associated with the dissolution of a substance in a solvent at constant pressure resulting in infinite dilution.
The enthalpy of solution is one of the three dimensions of solubility analysis. It is most often expressed in kJ/mol at constant temperature. Just as the energy of forming a chemical bond is the difference between electron affinity and ionization energy, the heat of solution of a substance is defined as the sum of the energy absorbed, or endothermic energy, expressed in positive values and unit kJ/mol, and energy released, or exothermic energy (negative value).
Because heating decreases the solubility of a gas, dissolution of gases is exothermic. Consequently, as a gas continues to dissolve in a liquid solvent, temperature will decrease, while the solution continues to release energy. This is an effect of the increase in heat or of the energy required to attract solute and solvent molecules—in other words, this energy outweighs the energy required to separate solvent molecules. When the gas is completely dissolved (this is purely theoretical as no substance can infinitely dissolve), the heat of solution will be at its maximum.
Dissolution can be viewed as occurring in three steps:

1. Breaking solute-solute attractions (endothermic), see for instance lattice energy in salts.
2. Breaking solvent-solvent attractions (endothermic), for instance that of hydrogen bonding
3. Forming solvent-solute attractions (exothermic), in solvation.

The value of the overall enthalpy change is the sum of the individual enthalpy changes of each of these steps. For example, dissolving ammonium nitrate in water decreases the temperature of the solution. Solvation does not compensate energy spent in breaking down the crystal lattice, while adding potassium hydroxide will increase it.
Solutions with negative enthalpy changes of solution form stronger bonds and have lower vapor pressure.
The enthalpy of solution of an ideal solution is zero since the attractive and repulsive properties of ideal fluids are equal, irrespective of the compounds, thus mixing them does not change the interactions.

Enthalpy change of solution for some selected compounds
hydrochloric acid	-74.84
ammonium nitrate	+25.69
ammonia	-30.50
potassium hydroxide	-57.61
caesium hydroxide	-71.55
sodium chloride	+3.88
potassium chlorate	+41.38
acetic acid	-1.51
sodium hydroxide	-44.51
Change in enthalpy ΔHo in kJ/mol in water at 25°C[1]