1. Heat in Chemical Reactions and Processes
Thermochemistry Lecture 2

2. Enthalpy Symbol : H Units : J
Heat content of a system at constant pressure
Chemists are interested in how enthalpy changes (ΔH) during reactions and processes
Is heat absorbed or released
ΔHrxn = H products – H reactants

3. Enthalpy of Reactions If heat is absorbed If heat is released
ΔHrxn is positive meaning the Hproducts >Hreactants
Endothermic reaction
If heat is released
ΔHrxn is negative meaning the Hreactants > Hproducts
Exothermic reaction
Note: ΔHrxn = q

4. Enthalpy of Phase Changes
ΔHvap = heat required to vaporize one mole of a liquid
H2O (l)  H2O (g) ΔHvap = 40.7 KJ
Since condensation is opposite of vaporization
ΔHcond = KJ
ΔHfus = heat required to melt one mole of a solid
H2O (s)  H2O (l) ΔHvap = 6.01 KJ
Since freezing is opposite of melting
ΔHsolid = KJ

5. Practice Questions
Is the following endothermic or exothermic? Explain why. NH3 (g)  NH3 (l)

6. Practice Question 2
Is the following endothermic or exothermic? Explain why.
4 Fe (s) + 3 O2 (g)  2 Fe2O3 (s) kJ
ΔHrxn = KJ
Exothermic

7. Thermochemical equations
Balanced chemical equation
Includes:
physical state of reactants and products (s, l, g) ΔH
Example:
4 Fe (s) + 3 O2 (g)  2 Fe2O3 (s) ΔHrxn = KJ

8. Calculating Enthalpy Change
Some reactions are too slow to measure the enthalpy change
Therefore, chemists theoretically determine enthalpy
Hess’s Law: if you can add two or more thermochemical equations to produce a final equation for a reaction, then the sum of the enthalpy changes for the individual reactions is the enthalpy change for the final reaction

9. Applying Hess’s Law
Determine ΔH for the following reaction: 2 S(s) + 3O2(g)  2SO3(g) If: S(s) + O2(g)  SO2(g) ΔH = -297 KJ 2SO3(g)  O2(g) + 2SO2 ΔH = 198 KJ

10. Determine ΔH for the following reaction:
2H2O2(l)  O2(g) + 2H2O(l)
If:
2H2 (g) + O2(g)  2H2O(l) ΔH = -572 KJ
H2 (g) + O2(g)  H2O2(l) ΔH = -188 KJ

11. Standard enthalpy (heat) of formation
ΔH that accompanies the formation of one mole of the compound made from its constituents
Every free element in its standard state is assigned a ΔHf of exactly 0.0 KJ.
Example:
H2(g) + S(s)  H2S (g) ΔHf = -21 KJ

12. Practice Problem
Use the standard enthalpies of formation to calculate ΔHrxn for the combustion of methane. CH4(g) + O2(g)  CO2(g) + 2H2O(l)

13. Practice Problem
Use the standard enthalpies of formation to calculate ΔHrxn for the following reaction: CaCO3 (s)  CaO (s) + CO2 (g)