Factors Affecting Lattice Enthalpy and Hydration (OCR A-Level Chemistry A): Revision Notes

Factors Affecting Lattice Enthalpy and Hydration

Introduction to ionic compounds

Ionic compounds have several characteristic properties that arise from their giant ionic lattice structure. These compounds typically display high melting and boiling points, are soluble in polar solvents like water, and conduct electricity when molten or dissolved in aqueous solution.

However, there is actually considerable variation in the physical properties of different ionic compounds. Some ionic compounds melt easily with a Bunsen burner flame, whilst others require such extreme temperatures that they can line the interior of industrial furnaces. Similarly, while many ionic compounds dissolve readily in water, others are completely insoluble.

The key to understanding these differences lies in examining two important energy quantities: lattice enthalpy and hydration enthalpy. The relative magnitudes of these values help explain the variety of melting points, boiling points, and solubility trends observed in ionic compounds.

Factors affecting lattice enthalpy

Lattice enthalpy is the enthalpy change when one mole of an ionic compound is separated into gaseous ions. The magnitude of lattice enthalpy depends primarily on two factors: ionic size and ionic charge.

Effect of ionic size on lattice enthalpy

The size of ions has a significant impact on lattice enthalpy values. This can be clearly demonstrated by comparing Group 1 chlorides, where the cation size increases down the group whilst the anion remains constant.

When we examine the data for sodium chloride, potassium chloride, and rubidium chloride, several clear trends emerge:

• As we move down Group 1, the ionic radius of the cation increases ($\text{Na}^+ < \text{K}^+ < \text{Rb}^+$)
• The electrostatic attraction between the positive and negative ions decreases as they become further apart
• The lattice enthalpy becomes less negative (less exothermic): $\Delta_\text{lattice}H^\circ = -786 \text{ kJ mol}^{-1}$ for NaCl compared to $-689 \text{ kJ mol}^{-1}$ for RbCl
• The melting point decreases: $801°\text{C}$ for NaCl compared to $718°\text{C}$ for RbCl

Effect of ionic charge on lattice enthalpy

Ionic charge has an even more dramatic effect on lattice enthalpy than ionic size. This can be illustrated by comparing sodium oxide ($\text{Na}_2\text{O}$) with calcium oxide ($\text{CaO}$), where the cations have similar sizes but different charges.

When ionic charge increases:

• The electrostatic attraction between ions increases significantly
• The lattice enthalpy becomes more negative (more exothermic)
• The melting point increases substantially

Combined effects across Period 3

When we examine trends across Period 3 of the periodic table, both ionic size and charge change simultaneously, creating either reinforcing or opposing effects.

For the cations across Period 3 ($\text{Na}^+$, $\text{Mg}^{2+}$, $\text{Al}^{3+}$), we observe two supporting effects:

• Increasing charge provides greater electrostatic attraction
• Decreasing ionic size allows ions to pack more closely, increasing attraction

Both effects work together to create increasingly negative lattice enthalpies across the period.

For the anions ($\text{S}^{2-}$, $\text{Cl}^-$), we encounter two opposing effects:

• Increasing charge would increase attraction
• Increasing ionic size decreases attraction

Predicting melting points

The magnitude of lattice enthalpy provides a reliable indicator of an ionic compound's melting point. Metal oxides such as $\text{MgO}$, $\text{Al}_2\text{O}_3$, and $\text{ZrO}_2$ have extremely exothermic (very negative) lattice enthalpies and correspondingly very high melting points. These stable metal oxides are used as refractory materials for lining furnaces and industrial equipment.

Lattice enthalpy is the primary indicator of melting point, but other structural factors, such as how ions pack together in the lattice, must also be considered when making precise predictions.

Factors affecting hydration

Hydration enthalpy is the enthalpy change when one mole of gaseous ions dissolves in water and becomes surrounded by water molecules. Hydration enthalpies follow similar trends to lattice enthalpies, being affected by both ionic size and ionic charge.

Effect of ionic size on hydration enthalpy

The size of an ion significantly influences its hydration enthalpy. This trend can be observed by examining Group 1 cations, where ionic radius increases down the group.

For the Group 1 cations:

• $\text{Na}^+$: $\Delta_\text{hyd}H^\circ = -406 \text{ kJ mol}^{-1}$
• $\text{K}^+$: $\Delta_\text{hyd}H^\circ = -320 \text{ kJ mol}^{-1}$
• $\text{Rb}^+$: $\Delta_\text{hyd}H^\circ = -296 \text{ kJ mol}^{-1}$

As ionic radius increases down the group:

• The attraction between the ion and surrounding water molecules decreases
• The hydration enthalpy becomes less negative (less exothermic)

Effect of ionic charge on hydration enthalpy

Ionic charge has a pronounced effect on hydration enthalpy, similar to its effect on lattice enthalpy.

When ionic charge increases:

• The attraction between the ion and water molecules increases dramatically
• The hydration enthalpy becomes significantly more negative (more exothermic)

Predicting solubility

The solubility of an ionic compound in water depends on the balance between lattice enthalpy and hydration enthalpy. When an ionic compound dissolves, the ionic lattice must first be broken apart (requiring energy equal to the lattice enthalpy), and then the separated ions become hydrated (releasing energy equal to the hydration enthalpy).

Predicting dissolution:

If the sum of the hydration enthalpies (for both cation and anion) is larger in magnitude than the lattice enthalpy, the overall enthalpy change of solution will be exothermic and the compound should dissolve.

Remember!

Exam focus checklist

✓ Be able to explain how ionic size affects lattice enthalpy and hydration enthalpy with reference to electrostatic attraction

✓ Be able to explain how ionic charge affects lattice enthalpy and hydration enthalpy

✓ Understand that both factors can work together (supporting effects) or against each other (opposing effects)

✓ Be able to use lattice enthalpy values to predict relative melting points of ionic compounds

✓ Understand the energy changes involved when an ionic compound dissolves in water

✓ Remember that solubility predictions require consideration of both enthalpy and entropy factors