Enthalpies of Solution and Hydration

• Hess' law states that total enthalpy change for a reaction is independant of the route by which the chemical change takes place.
• When might you not be able to determine an enthalpy change experimentally?

  • Trying to measure the temperature change of a solid
  • Trying to measure the enthalpy change of a reaction that requires heating (e.g. decomposition reaction)
  • Adding a specific amount of substance to another, when several products are possible
  • If the reaction rate is too slow
• Hess' law is used to work out the enthalpy change to form a hydrated salt from an anhydrous salt as this cannot be done experimentally. It is impossible to add the exact amount of water and it is not easy to measure the temperature change of a solid.
• When an ionic lattice dissolves in water the lattice breaks down and the ions are separated. This is a very endothermic process (energy needed to break bonds).
• The enthalpy change of solution is the enthalpy change when 1 mole of solid is dissolved in sufficient solvent to give an infinitely dilute solution.
• Infinitely dilute solution:

  • ions are fully separated and do not interact with each other
  • further dilution does not cause any further heat change
• Cations and anions dissociate (lattice dissociation enthalpy).

  Ions then become surrounded by water molecules (enthalpy of hydration).
• Strong ion-dipole forces act between:

  • positive cations and the delta - oxygen atoms in the water molecule
  • negative anions and the delta + hydrogen atoms of the water molecules
• The process by which an ion is surrounded by water is called hydration.

  Hydration is highly exothermic - this compensates for the endothermic break-up of the lattice.
• The hydration enthalpy of an ion is the enthalpy change when 1 mole of gaseous ions is completely surrounded by water molecules to produce an infinitely dilute solution.
• The hydration energy depends on the charge and size of the ions.
• The larger the charge and the smaller the size, the larger the hydration energy.
• The larger the charge and the smaller the size, the greater the charge density and the larger the hydration energy.
• The cations increase in size down the group with the same charge.

  The enthalpy of hydration becomes less exothermic.
• Charge density decreases as ionic radius increases, therefore decreasing attraction for delta - oxygen atom of the water molecules.

  Hence less energy is released when the ion-dipole attractions form.
• Charge density decreases as ionic radius increases, therefore decreasing attraction for the delta + hydrogen atoms of the water molecules.

  Hence less energy is released when the attractions form.
• The more exothermic (or less endothermic) the enthalpy of solution, the more likely the compound is to dissolve.

  Not all ionic compounds are soluble due to great attraction between ions in the solid state (lattice energy). The energy required to break up an ionic lattice is known as the lattice dissociation enthalpy.
• Solubility of ionic compounds in water varies widely. It depends on the balance between the hydration energy of the ions and the lattice energy of the compound.

  Most ionic compounds are at least partially soluble in water.