6.5.4 Factors Affecting Enthalpy of Hydration

This topic explains the factors that affect the enthalpy of hydration. It describes how ionic charge and ionic radius influence the energy released when ions are hydrated.

Definition

Enthalpy of Hydration (ΔH_hyd): The energy released when one mole of gaseous ions is dissolved in water to form one mole of aqueous ions.

Charge Density: The ratio of ionic charge to the ionic radius of an ion.

Charge density = ionic chargeionic radius

Main Idea

• Ionic Charge
• Ionic Radius

Core Concept

• Ionic charge: As the ionic charge increases, the hydration enthalpy also increases (becomes more exothermic).
• Ionic radius: As the ionic radius decreases, the hydration enthalpy increases (becomes more exothermic).
• Charge density: Higher charge density → stronger attraction to water molecules → higher (more exothermic) hydration enthalpy.

Detailed Explanation

What is Charge Density?

The magnitude of enthalpy of hydration is influenced by the charge density of an ion. The charge density is defined as the ratio of ionic charge to the ionic radius of an ion.

Charge density = ionic chargeionic radius

Therefore, two factors affect enthalpy of hydration.

1. Ionic Charge

• The ionic charge of an ion affects the electrostatic interaction between the ion and water molecules.
• Ions with a higher charge exert stronger attractive forces on water molecules through ion-dipole interactions, resulting in a higher enthalpy of hydration.
• Therefore, as the ionic charge increases, the hydration enthalpy also increases, and the hydration process becomes exothermic.
• Example: The hydration enthalpy of MgO is more exothermic than that of NaCl. These compounds are ionic compounds and consist of Mg²⁺, O²⁻ and Na⁺, Cl⁻ ions, respectively. Both Mg²⁺ and O²⁻ have a greater charge than Na⁺ and Cl⁻, so the hydration enthalpy of MgO is more exothermic than NaCl.

2. Ionic Radius

• The size of the ion also affects the enthalpy of hydration.
• Smaller ions have a higher charge density and exert a stronger attraction on water molecules.
• Therefore, more energy is released when they hydrate, and the enthalpy of hydration becomes exothermic.
• On the other hand, larger ions spread out the charge distribution and exert a weaker attractive force on water molecules. Thus, less energy is released when they hydrate.
• Example: The hydration enthalpy of MgSO₄ is greater than that of CaSO₄. Sulphate ions are common in these compounds; their hydration enthalpy difference must be due to Mg²⁺ and Ca²⁺ ions. Both Mg and Ca belong to group 2 of the periodic table. Mg is in the group above Ca, so the Mg²⁺ ion is smaller than the Ca²⁺ ion. Thus, the attraction of water molecules to Mg²⁺ ions is stronger than to Ca²⁺ ions. As a result, the hydration enthalpy of MgSO₄ is more exothermic than that of CaSO₄.

Summary Table