1. Temperature can change Phase change can occur Reaction can take place
When Heat is Transferred …
Temperature can change
Phase change can occur
Reaction can take place

2. Thermal Energy and Phase Changes
First: What happens?

3. Thermal Energy and Phase Changes
First: What happens?

4. Thermal Energy and Phase Changes
First: What happens?

5. But what’s really happening?
Warming:
Molecules move more rapidly
Kinetic Energy increases
Temperature increases
Melting/Boiling:
Molecules do NOT move more rapidly
Temperature remains constant
Intermolecular bonds are broken
Chemical potential energy (enthalpy) increases

6. Energy and Phase Changes: Quantitative Treatment
Melting:
Heat of Fusion (DHfus) for Water: 333 J/g
Boiling:
Heat of Vaporization (DHvap) for Water: 2256 J/g

7. Total Quantitative Analysis
Convert 40.0 g of ice at –30 oC to steam at 125 oC
Warm ice: (Specific heat = 2.06 J/g-oC)
Melt ice:
Warm water (s.h. = 4.18 J/g-oC)

8. Total Quantitative Analysis
Convert 40.0 g of ice at –30 oC to steam at 125 oC
Boil water:
Warm steam (s.h. = 1.92 J/g-oC)
Total Energy:

9. Enthalpy Change and Chemical Reactions
DH = energy needed to break bonds – energy released forming bonds Example: formation of water:
DH = ?

10. Energy Diagrams

11. Chemical Reactions and Enthalpy Change
C3H8(g) + 5 O2(g)  3 CO2(g) + 4 H2O(g) H = kJ
What does this mean?
How much energy is released by a 468-g tank of propane?

12. Chemical Reactions and Enthalpy Change
C3H8(g) + 5 O2(g)  3 CO2(g) + 4 H2O(g) H = kJ
How much energy is released by a 468-g tank of propane?
How much water (starting at 20 oC) can this boil?

13. Return to Calorimetry Constant Pressure gives H
Constant Volume give E
Count the calorimeter, or not…

14. Bomb Calorimetry Experiment
N2H4 + 3 O2  2 NO2 + 2 H2O
Energy released = E absorbed by water +
E absorbed by calorimeter
Ewater =
Ecalorimeter =
Total E =
H = energy/moles =
0.500 g N2H4
600 g water
420 J/oC

15. Hess’s Law
If you can add the reactions, you can add the H’s.

16. Hess’s Law
Enthalpy is a
State Function.

17. Rules for manipulating Reaction Equations
Reverse reactions: Change sign of H
Multiple by a constant: Multiply sign of H by that constant

18. Hess’s Law: Your Lab
Mg(s) + ½ O2(g)  MgO(s) H = ???

19. Standard Enthalpy of Formation
Formation Reaction: Reaction that forms one mole of a compound from its elements in their natural states.
K2CO3(s)??
Standard enthalpy of formation: Enthalpy change on forming one mole of a compound from its elements under standard conditions.