Detailed and comprehensive notes on unit 3.1.8 (thermodynamics) for AQA A level chemistry. Covers: enthalpy change definitions, born haber cycles, the perfect ionic model, entropy, Gibb’s free energy, enthalpies of solution and feasibility.
Content:
• Lattice enthalpy can be defined as either enthalpy of lattice dissociation or enthalpy of lattice
formation
• Born–Haber cycles are used to calculate lattice enthalpies using the following data:
o Enthalpy of formation
o Ionisation energy
o Enthalpy of atomisation
o Bond enthalpy
o Electron affinity
• Cycles are used to calculate enthalpies of solution for ionic compounds from lattice enthalpies and
enthalpies of hydration
Skills:
• Define each of the above terms and lattice enthalpy
• Construct born–haber cycles to calculate lattice enthalpies using these enthalpy changes
• Construct born–haber cycles to calculate one of the other enthalpy changes
• Compare lattice enthalpies from born–haber cycles with those from calculations based on a perfect
ionic model to provide evidence for covalent character in ionic compounds
• Define the term enthalpy of hydration
• Perform calculations of an enthalpy change using these cycles
Gibbs Free Energy Change and Entropy Change
Content:
• ∆H, whilst important, is not sufficient to explain feasible change
• The concept of increasing disorder (entropy change, ∆S)
• ∆S accounts for the above deficiency, illustrated by physical changes and chemical changes
• The balance between entropy and enthalpy determines the feasibility of a reaction given by the
relationship: ∆G = ∆H – T∆S (derivation not required)
• For a reaction to be feasible, the value of ∆G must be zero or negative
.
Skills:
• Calculate entropy changes from absolute entropy values
• Use the relationship ∆G = ∆H – T∆S to determine how ∆G varies with temperature
• Use the relationship ∆G = ∆H – T∆S to determine the temperature at which a reaction becomes
feasible
, Enthalpy Change Definitions
Enthalpy Change
• The heat energy transferred in a reaction at a constant pressure.
Bond Dissociation Enthalpy (Δ dissH):
• The bond dissociation enthalpy is the standard molar enthalpy change when one mole of a covalent
bond is broken into two gaseous atoms (or free radicals).
• E.g. Cl2 (g) ! 2Cl (g) [ΔdissH = +242 kJ mol-1]
• For diatomic molecules, ΔdissH will be 2x ΔatH, e.g.: ½ Cl2 (g) ! Cl (g) [ΔatH = +121 kJ mol-1]
Enthalpy of Formation (Δ fH):
• The energy transferred when 1 mole of the compound is formed from its elements under standard
conditions (298K and 100kpa), with all reactants and products being in their standard states.
• E.g. Na (s) + ½ Cl2 (g) ! NaCl (s) [ΔfH = - 411.2 kJ mol-1]
Enthalpy of Atomisation (Δ atH):
• The enthalpy change when 1 mole of gaseous atoms is formed from the element in its standard state.
• E.g. ½ O2 (g) ! O (g) [ΔatH = +249 kJ mol-1]
• The enthalpy change for a solid metal turning to a gaseous atom can be called the enthalpy of
sublimation and will be numerically the same as the enthalpy of atomisation.
First Ionisation Enthalpy (Δ IE1H):
• The enthalpy change required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1
mole of gaseous ions with a +1 charge.
• E.g. Mg (g) ! Mg+ (g) + e-
Second Ionisation Enthalpy (Δ IE2H):
• The enthalpy change to remove 1 mole of electrons from one mole of gaseous 1+ ions to produces
one mole of gaseous 2+ ions.
• E.g. Mg+ (g) ! Mg2+ (g) + e-
First Electron Affinity (Δ EA1H):
• The enthalpy change when 1 mole of gaseous atoms gain 1 mole of electrons to form 1 mole of
gaseous ions with a –1 charge
• E.g. O (g) + e- ! O- (g) [ΔEA1H = -141.1 kJ mol-1]
• The first electron affinity is exothermic for atoms that normally form negative ions. This is because
the ion is more stable than the atom, and there is an attraction between the nucleus and the electron.
Second Electron Affinity (Δ EA2H):
• The enthalpy change when one mole of gaseous 1- ions gains one electron per ion to produce
gaseous 2- ions.
• E.g. O- (g) + e- ! O2- (g) [ΔEA2H = +789 kJ mol-1]
• The second electron affinity can be endothermic because it takes energy to overcome the repulsive
force between a negative ion and an electron.
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