1. Thermochemistry
AH Chemistry, Unit 2(c)

2. Introduction
Thermochemistry concerns the study of energy changes in chemical reactions.
First Law of Thermodynamics: “Energy is conserved”.
Hess’s law states that “the overall reaction enthalpy is the sum of the reaction enthalpies of each of step of a reaction”.

3. Calorimetry
The term used to describe the quantitative determination of the change in heat energy which occurs during a chemical reaction.
A calorimeter is used to measure this energy change.

4. “Standard” enthalpy changes
Both reactants and products are considered in their most stable state at 1 atmosphere pressure and a specified temperature (usually 298K).

8. Standard Enthaply of Formation
Refers to the enthalpy change which occurs when one mole of a substance is prepared from its elements in their standard state.

9. C(s) + O2(g) ———> CO2(g) H2(g) + ½O2(g) ———> H2O(l)
Examples
C(s) + O2(g) ———> CO2(g)
H2(g) ½O2(g) ———> H2O(l)
2C(s) + ½O2(g) H2(g) ———> C2H5OH(l)
Notes
Only ONE MOLE of product on the RHS of the equation.
Elements In their standard states have zero enthalpy of formation.
Carbon is usually taken as the graphite allotrope.

14. Using enthalpies of formation
Hº =   Hf (products) –   Hf (reactants)
Example
Calculate the standard enthalpy of reaction at 298K for the complete combustion of methane.
Substance
Enthalpy of formation (kJ mol-1)
Carbon dioxide
Water
Methane
-394
-286
-74.9
kJ mol l-1

15. Bond Enthalpies

21. Born-Haber cycles

22. What are they?
A thermochemical cycle applied to the formation of an ionic crystal.
Used to calculate the enthalpy of lattice formation, which cannot be determined directly by experiment.
Standard Enthalpy of Lattice Formation: the enthalpy change which occurs when one mole of an ionic crystal is formed from the ions of their gaseous states, under standard conditions.

23. Born-Haber cycle Na+ (g) + e- + Cl (g) H first ionisation energy
H first electron affinity
Na (g) + Cl (g)
Na+ (g) + Cl- (g)
H bond enthalpy
H lattice
Na (g) + ½ Cl2 (g)
H atomisation
Na (s) + ½ Cl2 (g)
H formation
NaCl (s)

24. Apply Hess’s Law: enthalpy H NaCl (s) Na+ (g) + Cl- (g) H lattice
Na (s) + ½ Cl2 (g)
H formation
H atomisation
Na (g) + ½ Cl2 (g)
Na (g) + Cl (g)
Na+ (g) + e- + Cl (g)
H first ionisation energy
H first electron affinity
Apply Hess’s Law:
HatmNa + HatmCl + H1st IE + H1st EA + Hlattice =
Hformation

25. Born-Haber Cycles: applying Hess’s Law
HatmNa + HatmCl + H1st IE + H1st EA + Hlattice =
Hformation
Rearrange to find the lattice energy:
Hlattice =
Hformation - (HatmNa + HatmCl + H1st IE + H1st EA)
So Born-Haber cycles can be used to calculate a measure of ionic bond strength based on experimental data.

26. Definitions Standard enthalpy of atomisation of an element:
The energy required to produce one mole of isolated gaseous atoms from the element in its standard state.
Electron affinity:
The enthalpy change for the process of adding one mole of electrons to one mole of isolated atoms in the gaseous state.

35. Enthalpy of solution
Energy must be supplied to break up the ionic lattice: ∆Hlatt
Energy is released when the free ions form bonds with water molecules: ∆Hhyd
Hydration enthalpy:
The energy released when one mole of individual ions become hydrated.

37. Illustrate the hydration enthalpy for:
A sodium ion
An sulphate ion