Chapter 22 Lattice Enthalpy And Entropy

The enthalpy change that accompanies the formation of 1 mol of an ionic compound from its gaseous ions under standard conditions e.g K+ (g) + Cl - (g) –/ KCl(s)

The enthalpy change that accompanies the formation of 1 mol of gaseous atoms from the element in its standard state under standard conditions e.g 1/2 Cl2 (g) —/ Cl(g)

It will change state from liquid to gas, and change from diatomic to an atom Might be seen as two different energy changes or combined together

1/2 I2 (s) —/ I (g) e.g for MgI2, you will need to multiple this number by two, as even thought it is one iodine molecule, we look at the number of atoms produced?

The enthalpy change that accompanies the addition of 1 electron to each atom in 1 mole of gaseous atoms to form 1 mol of gaseous 1- ions e.g Cl(g) + e- —/ Cl- (g)

The first electron affinity is exothermic the negative electron is attracted to the positive nucleus The second electron affinity are endothermic as we are adding a negative electrons to a negative ion, so energy is needed

Mg2+ (g) + 2e- + 2Cl (g) First and second IE —————- 2 x EA Cl Mg2+(g) + 2Cl-(g) ----------- Mg(g) + 2 Cl (g) Atomisation 2 x Cl and 1 x Mg ------------ Mg (s) + Cl2(g) MgCl2 (s)

Multiply enthalpy changes by number of moles in formula Seperate lines for each stage (don’t combine atomisation of both elements) Start from the beginning of the arrow you want to calculate

The enthalpy change the accompanies 1 mol of a solute dissolving in a solvent under standard conditions MgCl2 (s) + aq —// Mg2+ (aq) + 2Cl- (aq)

The enthalpy change that accompanies 1 mol of gaseous ions dissolving to form 1 mol of aqueous ions under standard conditions e.g Cl- (g) + aq —/ Cl- (aq)

1) The ionic lattice breaks down 2) The solvent molecules (e.g water) surround the ions

Gaseous ions —-Lattice enthalpy —- Ionic lattice —- Hydration —–Solution enthalpy —– Aqueous ions

Down the group, the ionic radius increases whilst the charge of the ion is the same This decreases the charge density of the ion There is also more electron shielding This reduces the electrostatic attraction between the oppositely charged ions, so a less exothermic enthalpy

Greater charge of the ions Smaller ionic radius and similar electron shielding So the ions have a greater charge density, and so have a greater electrostatic forces of attraction So more exothermic lattice enthalpies

The atomic radius is increasing and the electron shielding greater This means they have a smaller charge density So the ion-dipole interactions between water and the ion is lesser, so a less exothermic enthalpy

Greater charge of the ions Smaller ionic radius and similar electron shielding So the ions have a greater charge density Greater ion-dipole interactions with water, so a more exothermic enthalpy

When the sum of the hydration enthalpies is greater that the lattice enthalpy But also the arrangement of the lattice, shape of the molecules and entropy need to be taken into account

A measure of the dispersal of energy within a chemical system, a measure of disorder There is a natural tendency for energy to spread out

Gases > Liquids > Solids Gases are the most disordered

Entropy products - Entropy reactants If positive, the entropy has increased, more random

Formation of gases increases entropy More moles of gas, more disordered (Even if both are contain gases) e.g N2O4 –/ 2NO2 Entropy products greater than reactants

Whether a reaction is possible at a certain temperature

The overall energy of a reaction, taking into account enthalpy and entropy change

Gc= Enthalpyc - Temperature(Entropyc) KJ/mol K KJ Often entropy is in joules, so must be converted first