Lattice Energy

This lesson covers: 

1. Definitions of enthalpy change and standard conditions
2. Types of enthalpy change
3. Lattice energy and factors affecting it
4. Factors affecting electron affinity
5. Trend in electron affinity down groups 16 and 17

Defining enthalpy change

Enthalpy change (ΔH) is the heat energy transferred in a reaction at constant pressure. Its units are kilojoules per mole (kJ mol^-1).

• ΔH^⦵ indicates that the enthalpy change was measured under standard conditions of 298 K and 100 kPa.
• Exothermic reactions release heat energy to the surroundings and have a negative ΔH value.
• Endothermic reactions absorb heat energy from the surroundings and have a positive ΔH value.

Types of enthalpy change

There are many types of enthalpy change, each referring to a specific chemical or physical process:

1. Enthalpy change of formation (ΔH^⦵_f) - The enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions. For example:

2C_(s) + 3H_2(g) + 1⁄2O_2(g) ➔ C₂H₅OH_(l)

ΔH^⦵_f values are usually exothermic as energy is released when bonds form between the elements to make the compound.

1. Enthalpy change of atomisation (ΔH^⦵_at) - The enthalpy change when 1 mole of gaseous atoms is formed from the element in its standard state under standard conditions. For example:

1⁄2Cl_2(g) ➔ Cl_(g)

ΔH^⦵_at values are always endothermic as energy must be supplied to break the bonds holding the atoms in the element together.

1. First ionisation energy (ΔH^⦵_IE1) - The enthalpy change when 1 mole of gaseous 1+ ions is formed from 1 mole of gaseous atoms. For example:

Mg_(g) ➔ Mg⁺_(g) + e^-

ΔH^⦵_IE1 values are always endothermic as energy is needed to overcome the electrostatic attraction between the nucleus and the electron being removed.

1. Second ionisation energy (ΔH^⦵_IE2) - The enthalpy change when 1 mole of gaseous 2+ ions is formed from 1 mole of gaseous 1+ ions. For example:

Mg⁺_(g) ➔ Mg²⁺_(g) + e^-

ΔH^⦵_IE2 values are always endothermic as even more energy is required to remove an electron from a positively charged ion.

1. First electron affinity (ΔH^⦵_ea1) - The enthalpy change when 1 mole of gaseous 1- ions is formed from 1 mole of gaseous atoms. For example:

O_(g) + e^- ➔ O^-_(g)

ΔH^⦵_ea1 values are usually exothermic as the attraction between the nucleus and the incoming electron releases energy.

1. Second electron affinity (ΔH^⦵_ea2) - The enthalpy change when 1 mole of gaseous 2- ions is formed from 1 mole of gaseous 1- ions. For example:

O^-_(g) + e^- ➔ O^2-_(g)

ΔH^⦵_ea2 values are always endothermic as energy must be supplied to overcome the repulsion between the negative ion and the second incoming electron.

1. Bond enthalpy (ΔH^⦵_diss) - The enthalpy change when 1 mole of a particular covalent bond in the gaseous state is broken. For example:

Cl_2(g) ➔ 2Cl_(g)

ΔH^⦵_diss values are always endothermic as energy is required to overcome the attraction between the bonded atoms and break the bond.

Lattice energy measures ionic bond strength

Lattice energy (ΔH^⦵_latt), also known as lattice enthalpy, is a measure of the strength of the electrostatic forces holding ions together in an ionic lattice. It is defined as the enthalpy change when one mole of a solid ionic compound is formed from its gaseous ions under standard conditions (298 K, 100 kPa).

For example:  Na⁺_(g) + Cl^-_(g) ➔ NaCl_(s)   ΔH^⦵_latt = −787 kJ mol^-1

Lattice energy values are always negative (exothermic) because energy is released when the oppositely charged ions come together to form the solid lattice.

The more negative the lattice energy, the stronger the ionic bonding in the compound. For instance, MgO has a more negative lattice energy (-3,791 kJ mol^-1) than NaCl (-787 kJ mol^-1), indicating that MgO has stronger ionic bonds.

Factors affecting lattice enthalpy

The lattice enthalpy of an ionic compound depends on two key factors.

1. Ionic charge

• Ions with higher charges experience stronger electrostatic attractions than ions with lower charges.
• This leads to more energy being released when the lattice forms, resulting in a more negative lattice enthalpy.
• For example, MgCl₂ has a much more negative lattice enthalpy (-2,526 kJ mol^-1) than NaCl (-787 kJ mol^-1) because the Mg²⁺ ion has a higher charge than the Na⁺ ion, resulting in stronger electrostatic attractions in the lattice.

2. Ionic radius

• Smaller ions have a higher charge density and can pack more closely together in the lattice.
• This increases the strength of the electrostatic attractions between ions.
• Consequently, compounds with smaller ions tend to have more negative lattice enthalpies.
• For example, LiCl has a more negative lattice enthalpy (-853 kJ mol^-1) than NaCl (-787 kJ mol^-1) because the Li⁺ ion is smaller than the Na⁺ ion, allowing closer packing in the lattice.

So in general, compounds with small, highly charged ions have the most negative lattice enthalpies.

Factors affecting electron affinity

The factors influencing electron affinity in groups 16 and 17 are similar to those affecting first ionisation energy:

1. Nuclear charge - A higher nuclear charge creates a stronger electrostatic attraction between the nucleus and an incoming electron. This results in more energy being released when the electron is gained.
2. Atomic radius - In atoms with a larger atomic radius, an incoming electron is farther from the nucleus, experiencing a weaker electrostatic attraction. This leads to less energy being released when the electron is added.
3. Electron shielding - As the number of inner electron shells increases, the shielding effect on an incoming electron becomes greater. This reduces the attraction between the nucleus and the added electron, resulting in a smaller release of energy.

Trends in electron affinity

For groups 16 and 17, electron affinities generally become less negative moving down the group. However, the first element in each group (oxygen and fluorine) is an exception to this trend.

The table below illustrates the electron affinities for elements in these groups:

Several factors contribute to this trend:

1. Nuclear charge - As you move down a group, the increasing number of protons in the nucleus leads to a stronger attraction for added electrons. This would be expected to result in more negative electron affinities.
2. Atomic radius - However, the increasing atomic radius down a group means that the added electron is farther from the nucleus, reducing the electrostatic attraction. This factor tends to make electron affinities less negative.
3. Electron shielding - Additionally, the presence of more inner electron shells in larger atoms shields the outer electrons from the nucleus, further decreasing the attraction for an added electron.

The combined effects of increased atomic radius and electron shielding outweigh the influence of greater nuclear charge, resulting in less negative (less exothermic) electron affinities down the group.

The anomalies of oxygen and fluorine

Oxygen and fluorine are exceptions to the trend, having less negative electron affinities than their respective periods' trends would suggest. This anomaly arises from their extremely small atomic radii, which lead to greater electron-electron repulsion within the atoms. This repulsion reduces the attraction between the incoming electron and the nucleus, resulting in less energy being released when an oxygen or fluorine atom gains an electron compared to sulfur or chlorine respectively.