13. Energetics II (LATTICE ENTHALPY)

The enthalpy when 1 mole of an ionic solid is formed from its constituent gaseous ions, under standard conditions Eg: Na+ (g) + Cl- (g) —> NaCl (s) Lattice enthalpy of formation is always exothermic Units are KJmol-1

Larger charge smaller size of the ion = more exothermic lattice enthalpy as the ionic bond is stronger

The amount of energy absorbed when 1 mole of an ionic compound is split apart into free gaseous ions This value is always endothermic

The enthalpy change when 1 mol of a gaseous atoms is formed from its element in its standard state Always endothermic Eg: Na (s) —> Na(g) Cl2 (g) —> 1/2Cl2 (g)

The energy change when 1 mole of electrons is added to 1 mole of gaseous atom Eg: Cl(g) + e- —> Cl- (g) Note, we do not use 1/2 Cl2 as that is a molecule. the definition explicitly states we need gaseous atom

First electron affinity is usually exothermic Second electron affinity for adding an electron to a negatively charged ion (e.g. O⁻ to O²⁻), and is endothermic due to repulsion

If theoretical and experimental lattice energy values are similar, so compound is almost 100% ionic If the theoretical and experimental values are different then compound has some covalent character Theoretical = perfect ionic model Experimental lattice energy = normally more exothermic than theoretical

Cation- talk about it being more polarizing on the anion (due to more higher charge) Anion - talk about it being more polarizable (due to larger ionic radius) These factors mean more polarisation, so more covalent character meaning the experimental value is negative, meaning greater difference in the values

–> MgCl2 has a higher lattice ie more exothermic lattice enthlpy than CaCl2 –> This is because the magensium ion has a smaller ionic radius for the same 2+ charge (leading to greater charge density) than calcium ions –> So the attraction between magensium ions and chloride ions is stronger than between calcium ions and chloride ions

• Ions are perfectly spherical. • Charge is distributed evenly throughout the ion. • Attractions are purely electrostatic (with no covalent character).

Lattice enthalpy = enthalpy of formation - sum of all other enthalpies Remember to check if there is a 2nd IE OR 2nd EA Do you have to x2 any value? Rearrange to find unknown value

Enthalpy change when 1 mole of an ionic solid dissolves in an amount of water large enough so that the dissolved ions are well separated and do not interact with each other Eg: NaCl(s) –> Na+(aq) + Cl-(aq) Endothermic or exothermic

Enthalpy change when 1 mole of a gaseous ions become hydrated (dissolved in water) Eg: Na+(g) + Cl-(g) + aq –> Na+ (aq) + Cl- (aq) Exothermic

An ionic compound will dissolve if the enthalpy change of solution is negative. The enthalpy change of hydration is greater than the energy needed to break apart lattice

The entropy change is big enough to overcome a small enthalpy change of solution value

ΔHsolution​= +ΔHlattice + ΔHhydration​ So if the lattice enthalpy has a negative sign, make it positive for this equation Don’t change the ΔHhydration​ sign

Lattice dissociation is endothermic whilst hydration of ions is exothermic, therefore the dissolving of a magnesium sulphate is exothermic because the enthalpy of hydration is greater in magnitude than the lattice energy

Lattice enthalpy change of formation = exothermic, units are Kjmol-1 Enthalpy change of formation = exothermic Atomisation = endothermic 1st electron affinity - exothermic, 2nd electron affinity = endothermic ionisation energy - endothermic lattice enthalpy of disassociation - endothermic Enthalpy of hydration - exothermic Enthalpy of solution - exo or endo, needs to be exo to happen