5.2.1 (a,b) Lattice enthalpy, Born–Haber and related enthalpy cycles

(a) explanation of the term lattice enthalpy (formation of 1 mol of ionic lattice from gaseous ions, ∆H_LE) and use as a measure of the strength of ionic bonding in a giant ionic lattice

{Definition required.}

(b) use of the lattice enthalpy of a simple ionic solid (i.e. NaCl, MgCl₂) and relevant energy terms for:

(i) the construction of Born–Haber cycles

(ii) related calculations

{Relevant energy terms: enthalpy change of formation, ionisation energy, enthalpy change of atomisation and electron affinity.}

{Definition required for first ionisation energy and enthalpy change of formation only}

You'll have learned definitions for the enthalpy change of formation and combustion in Year 12

The lattice enthalpy follows the same pattern:

"The enthalpy change when 1 mole of an ionic compound forms from its gaseous ions under standard conditions"

More specifically, this definition is Lattice Formation enthalpy

And so should be exothermic.

But, were you given Lattice Dissociation enthalpy - it would be the same but endothermic.

Sodium Chloride

Step one is forming one mole of the required ionic substance from its elements, in their standard states and in their usual formula.

In this case that means we must use Chlorine as Cl₂ - even though it would be more convenient not to - because Chlorine doesn't exist naturally as individual Cl atoms.

That means that we only need half a mole of Chlorine.

Remember that we're not allowed to set up the reaction as 2Na + Cl₂--> 2 NaCl - because the definition says we must form one mole of substance.

We can assume that all formation enthalpies are exothermic because they always are in A level questions even if not in reality!

Next we start atomising the reactant elements.

It doesn't matter at all which one you choose to atomise first because, in reality, if you heated the mixture the processes would both start happening at the same time. Doing them sequentially is theoretical

In this case we have turned the one mole of Sodium metal (as a solid) into one mole of sodium ions in the gaseous state.

The A level examiner will want to see the state symbols.

The next step is either to atomise the Chlorine or to ionise the Sodium by removing an electron.

The example below is for ionising Sodium

This energy change is the first ionisation energy (Δ _{1st IE}H) that you learned about in Year 12

Having dealt with the Sodium, this example now atomises the Chlorine

This is the same process as breaking the Cl-Cl bond

But the Atomisation Energy (Δ _AH) of Chlorine makes 1 mole of Cl atoms

While the bond dissociation energy of Chlorine breaks one mole of Cl-Cl bonds and so makes 2 moles of Cl atoms.

As a result - Bond dissociation energy of Chlorine = 2 x Atomisation energy of chlorine.

Examiners sometimes give you the bond dissociation, sometimes the Atomisation energy.

So, be careful!

Now the Chloride ion must be made by gaining the electron that we removed from the Sodium

This is not the ionisation energy - that's for losing electrons.

This is the Electron Affinity Δ _EAH

Now that we have Sodium and Chlorine ions they can combine to make Sodium Chloride

This is Δ _LEH - the Lattice Enthalpy

Hess' Law states that enthalpy change is independent of route taken.

Δ _fH (NaCl) = Δ _atH (Na) + Δ _ieH (Na) + Δ _atH (Cl) + Δ _EAH (Cl) + Δ _LEH (NaCl)

You'll be given all but one of these values and asked to calculate the missing one.