Lattice Enthalpy

The enthalpy change that accompanies the formation of one mole of an ionic compound from its gaseous ions under standard conditions, 25 degrees celsius and 1 atmosphere

K+(aq) + Claq) -> KCl(s)

It is exothermic-energy is released when the ionic bonds are formed

A more negative lattice enthalpy means a stronger lattice/stronger ionic bonds

It is not measured directly-it is impossible to form one mole of an ionic lattice from gaseous ions experimentally

It is calculated using Hess' law, using a Born-Haber cycle

Standard Enthalpy of Formation: the enthalpy change that takes place when one mole of a compound is formed from its constituent elements in their standard states under standard conditions

Standard Enthalpy Change of Atomisation: the enthalpy change that takes place when one mole of gaseous ions forms from the element in its standard state

First Ionisation Energy: the enthalpy change accompanying the removal of one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions

Second Ionisation Energy: the enthalpy change accompanying the removal of one electron from each ion in one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions

First Electron Affinity: the enthalpy change accompanying the addition of one electron to each atom in one mole of gaseous atoms to form one mole of gaseous 1- ions

Second Electron Affinity: the enthalpy change that accompanies the addition of one electron to each ion in one mole of gaseous 1- ions to form one mole of gaseous 2- ions

The enthalpy change that takes place when one mole of a compound is completely dissolved in water under standard conditions

Can be exothermic or endothermic

KCl(aq) + aq -> K+(aq) + Cl-(aq)

When a solid dissolves, the ionic lattice breaks down into gaseous ions, which are then hydrated.

The crystal lattice is broken down and the ions are separated.

Energy is required to overcome the attractive forces between oppositely charged ions. This is the opposite of producing the lattice enthalpy. The enthalpy change involved in breaking the ionic lattice is the same magnitude as the lattice enthalpy but with the opposite sign.

The enthalpy change that takes place when one mole of isolated gaseous ions is dissolved in water, forming one mole of aqueous ions under standard conditions

Gaseous ions bond with water molecules

The positively charged ions are attracted to the slightly negative oxygen

The negatively charged ions are attracted to the slightly positive hydrogen

The lattice enthalpy of an ionic solid is calculated using the enthalpy changes of hydration of the gaseous ions and the enthalpy changes of solution of the ionic solid

Lattice Enthalpy of the ionic solid + Enthalpy of Solution of the ionic solid = Enthalpy of hydration of one of the gaseous ions + Enthalpy of hydration of the other gaseous ion

Lattice Enthalpy

Ionic Radius

Smaller ions pack closely together and so have stronger forces of attraction

Larger ions are further apart and so have weaker forcesof attraction

As ionic radius increases, attraction between ions decreases, so lattice enthalpy becomes less negative

Ionic Charge

The higher the charge on an ion, the more negative the lattice enthalpy

Greater attraction between positive and negative ions

Ionic radius decreases, the ions get closer together, producing more attraction

Enthalpy Change of Hydration

Ionic Radius

As ionic radius decreases, enthalpy of hydration becomes more exothermic

Small ions exert more attraction over water molecules

More energy released 

Ionic Charge

The greater the charge. the more negative the enthalpy of hydration

Ions with a greater charge have a greater attraction for water molecules

Positive ions attracted to the slightly negative oxygen

Negative ions attracted to the slightly positive oxygen