Aqueous Solutions and Heats of Hydration

Water and other polar molecules are attracted to ions. The electrostatic attraction between an ion and a molecule with a dipole is called an ion-dipole attraction. These attractions play an important role in the dissolution of ionic compounds in water.

When ionic compounds dissolve in water, the ions in the solid separate and disperse uniformly throughout the solution because water molecules surround and solvate the ions, reducing the strong electrostatic forces between them. This process represents a physical change known as dissociation. Under most conditions, ionic compounds will dissociate nearly completely when dissolved, and so they are classified as strong electrolytes. Even sparingly, soluble ionic compounds are strong electrolytes, since the small amount that does dissolve will dissociate completely.

Consider what happens at the microscopic level when solid KCl is added to water. Ion-dipole forces attract the positive (hydrogen) end of the polar water molecules to the negative chloride ions at the surface of the solid, and they attract the negative (oxygen) ends to the positive potassium ions. The water molecules surround individual K⁺ and Cl^? ions, reducing the strong interionic forces that bind the ions together and letting them move off into solution as solvated ions. Overcoming the electrostatic attraction permits the independent motion of each hydrated ion in a dilute solution as the ions transition from fixed positions in the undissolved compound to widely dispersed, solvated ions in solution.

This text is adapted from Openstax, Chemistry 2e, Section 11.2: Electrolytes.

Ionic solutes are held together by attractive interactions called Coulombic forces.

When ionic solutes are dissolved in water, the hydrogen bonds between the water molecules break and disrupt the Coulombic forces between the ions.

Breaking up an ionic crystal lattice into its constituent ions in this manner requires a large energy input and, thus, the enthalpy of the solute is an endothermic process.

The energy released when one mole of ionic solid is formed from the constituent gaseous ions is called the lattice energy and is always exothermic. Thus, the enthalpy to break one mole of the solute into its components is equal and opposite to the lattice energy.

However, when an ionic lattice is broken in an aqueous solution, each ion is surrounded and stabilized through ion-dipole interactions with the oppositely charged end of the water dipole.

This phenomenon is called hydration. The enthalpy change associated with the dissolution of one mole of ions in water is known as the heat of hydration. It is a combination of the enthalpy of the solvent and the enthalpy of mixing.

Since the ion-dipole interactions between a hydrated ion and the water molecules are much stronger than the hydrogen bonds in water alone, hydration is always an exothermic process.

The overall enthalpy of the solution is a sum of the endothermic enthalpy of the solute and the exothermic heat of hydration and therefore depends on the relative magnitudes of these two terms.

If the enthalpy of the solute is less than the heat of hydration, the enthalpy of solution will be negative, and dissolution will be exothermic—as seen in a sodium hydroxide solution.

If the enthalpy of the solute is greater than the heat of hydration, the enthalpy of  solution will be positive and dissolution will be endothermic—as in an ammonium chloride solution.

If the enthalpy of the solute is much greater than the heat of hydration, the solute will remain insoluble in water—as seen in the case of calcium sulfate.

If the two terms are close to equal, the enthalpy of solution will be nearly zero, as it is for sodium chloride. Such solutes do not change the temperature of the solution.