T16: Lattice energies

date:Archy de Berker

Ionic/ Covalent continuumWe can calculate the lattice energies of compounds based upon the size of the ions, the charges of the ions and the solid structure. However, in doing this we assume ionic purity.

These are theoretical lattice energies. They are usually close to the real thing, which we find by using the Born-Haber cycle.

However, sometimes their is a discrepancy between the two, indicating that our assumptions are incorrect; the bonding is not purely ionic, with a polarised covalent bond there instead. This results in stronger lattices.

Fajan’s rules: There will be an increasing tendency towards covalency with:

» Smaller positive ions » Larger negative ions » Ions with more charge

Rules in Born-Haber cycles

Multiply up ∆H atm

& Emn

So for Cl2, the enthalpy of atomisation and the enthalpy of ionisa-

tion should be counted twice: look at the number of moles of ions, not of original atoms.

Multple positive charges = multiple Emns

So for Mg2+ for instance, use Em1 + Em

2.

Multiple negative charges = multiple Eaff

So O2 use E

aff O (g) but also E

aff O- (g).

Ionisation energiesAre always endothermic.

Electron affinitiesAre usually exothermic

Atomisation energiesAre always endothermic

Lattice energiesAre always exothermic

Lattice energies & reactivityThe lattice energy of a compound can help us to predict whether a reaction will happen; for instance, Sodium is ex-tremely reactive towards Chlorine because the lattice energy is extremely exothermic; because the lattice energy of NaCl is very negative.

Why?

1. The size of the ionsThe larger the ions, the farther apart their charges, and the weaker the lattice.

2. The charges on the ionsThe more charge, the better; stronger attractions lead to stronger lattices.

3. The structure of the solidAffects the distance, and therefore the attraction, between the ions.

Periodicity of period 3 chlorides & oxides

This is due to the increasing charges and consequently

decreasing cation size as we move along the period.

This causes decreasing melting points along the period,

since ionic compounds have far stronger bonds than the

weak intermolecular forces of molecular compounds.

Hydrolysis of chloridesMetal chlorides dissolve easily, because of the ionic charac-ter.

Non-metal chlorides (and aluminium), however, don’t dissolve.

instead hydrolysing.

This is the break up of water into H+ and OH -, and often pro-

duces hydrochloric acid.

Hydrolysis of oxidesBoth metal and non-metal oxides react with water in a hydroly-sis reaction.

Metal oxides tend to form hydroxides, which are alkaline, whilst non-metals often form acids.

However, aluminium is an amphoteric compound; it is capa-

ble of reaction with both acids and alkalis.

structures

date:Archy de Berker

Giant Ionic Lattice

Giant just means containing an unpredictable number of parti-cles; unlike a molecular compounds such as, H

2O

, which always

contains 2H, 1O.

The more charge, the stronger the lattice.

e.g KCl (770) has a lower melting point than NaCl (801)

The smaller the ions, the stronger the lattice.

e.g RbCl (715) has a lower melting point than NaCl (801)

The more negative the lattice energy, the stronger the lattice is, and the higher the melting point

Molecular Structures

With molecular forces, the intramolecular bonds are largely/wholly covalent. It is, however, the intermolecular forces which are most important, and these are merely weak VdW, diople-dipole or hydrogen bonds.

Therefore, to melt a molecular compound no covalent or ionic bonds are broken, just intermolecular forces.

MORE COVALENT = LOWER MELTING POINT