Chapter 22: Enthalpy

• The enthalpy change when 1 mole of an ionic lattice is formed from its gaseous ions • ∆LEH⦵

• K+ (g) + Cl- (g) -> KCl (s) • Exothermic as bonds are being formed

• The more exothermic the lattice enthalpy, the stronger the ionic bonding in a giant ionic lattice

• It cannot be measured directly, so it is calculated indirectly using a type of Hess cycles called Born-Haber cycles, which break lattice enthalpy up into different steps

• Splitting the ionic compound into its elements in their standard states • Turning the elements in their standard states to gaseous atoms • Turning the metal atom into a cation • Turning the non-metal atom into an anion

• The standard enthalpy of formation • The enthalpy change that takes place when one mole of a compound is formed from its elements under standard conditions, with all reactants and products in their standard states • ΔfH⦵

• K (s) + 1/2 Cl2 (g) -> KCl (s) • It is exothermic for salts

• The standard enthalpy of atomisation • The enthalpy change that takes place when 1 mole of gaseous atoms is formed from the element in its standard state • ΔatH⦵

• K (s) -> K (g) • 1/2 Cl2 (g) -> Cl (g) • Endothermic because bonds are broken to form gaseous atoms

• Ionisation energy • First ionisation energy: the energy required for the removal of 1 mole of electrons from 1 mole of gaseous atoms of an element • ΔIE1H⦵

• K (g) -> K+ (g) + e- • Endothermic because energy is required to overcome the attraction between the positive nucleus and the negative electron

• Electron affinity • First electron affinity: the enthalpy change when 1 electron is added to each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1- ions • ΔEA1H⦵

• Cl (g) + e- -> Cl- (g)

• The first electron affinity is exothermic as the first electron is attracted by the nucleus more strongly than it is repelled by the other electrons • After this, the negative ion repels electrons, so the second and successive electron affinities are endothermic

• You draw a Born-Haber cycle with: • the cations, electrons and non-metal atoms at the top • the ionic solid at the bottom • on the right is the lattice enthalpy • on the left are all the other enthalpies with the exception of electron affinities • arrows point down for exothermic energy changes • arrows point up for endothermic energy changes • because of this, electron affinities have to go on the right above lattice enthalpy • Enthalpy of formation = all the enthalpy changes + lattice enthalpy (follow the arrows)

• If you have multiples of the atoms/ ions, do not forget to multiply the enthalpy changes • Remember to show all of the ionisation energies where there are more than one • If given values of bond dissociation enthalpy instead of enthalpy change of atomisation, remember to adapt for 1 mole of product • Second electron affinities

• It only works if the standard state of the element is gaseous and the element has covalent bonds • This is what the definition of average bond enthalpy specifies

• The ions are the highest level on the right • The first electron affinity is shown below this • Between the 2 levels on the left is the cation, electrons and non-metal atom • This way, there is an arrow down for the first electron affinity (exothermic) and an arrow up for the second (endothermic)

• Using the enthalpy changes of solution and hydration

• The enthalpy change that takes place when 1 mole of a solute dissolves in water • ΔsolH⦵

• The enthalpy change that takes place when 1 mole of gaseous ions form hydrated ions • ΔhydH⦵

• The lattice breaks up • The ions are separated and react with polar water molecules to form hydrated ions • These hydrated ions are held together by ion-dipole forces

• NaCl (s) -> Na+ (aq) + Cl- (aq) • It can be either endothermic or exothermic

• Na+ (g) -> Na+ (aq) • Cl- (g) -> Cl- (aq) • It is exothermic since ion-dipole forces are being formed