Enthalpy change of solution
- Enthalpy change of solution
The enthalpy of solution (or enthalpy of dissolution) is the enthalpy change associated with the dissolution of a substance in a solvent at constant pressure.
The heat 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" kJ/mol), and energy released, or exothermic energy (expressed in "negative" kJ/mol).
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:
# Breaking solute-solute attractions (endothermic), see for instance lattice energy in salts.
# Breaking solvent-solvent attractions (endothermic), for instance that of hydrogen bonding
# Forming solvent-solute attractions (exothermic), in solvation.
The value of the overall enthalpy change is sum of the individual enthalpy changes of each step. For example dissolving ammonium nitrate in water will decrease the temperature of water (solvation does not weigh up against energy spent in breaking down the crystal lattice) while adding potassium hydroxide will increase it.
Solutions with negative heats of solution form stronger bonds and have lower vapor pressure.
ee also
* Heat
* Solution
* Thermometric titration
References
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2010.
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