1. Thermodynamics

2. Energy is neither created or destroyed during chemical or physical changes, but it is transformed from one form to another. Euniverse = 0

3. Mechanical Gravitational Thermal Electrostatic Electrical Chemical
TYPES of ENERGY
Kinetic Potential
Mechanical Gravitational
Thermal Electrostatic
Electrical Chemical
Radiant
Energy Conversion Examples:
1. dropping a rock
2. using a flashlight
3. driving a car

5. Endo: heat added to system Exo: heat released by system
SYSTEMS and SURROUNDINGS
System: The thing under study
Surroundings: Everything else in the universe
Energy transfer between system and surroundings:
Endo: heat added to system Exo: heat released by system

6. HEAT: What happens to thermal (heat) energy?
Three possibilities:
Warms another object
Causes a change of state
Is used in an endothermic reaction

7. Temperature Changes from Heat Exchange
Example 1: 5 g wood at 0 oC g wood at 100 oC
Example 2: 10 g wood at 0 oC g wood at 100 oC
Example 3: 5 g copper at 0 oC g copper at 100 oC
Example 4: 5 g wood at 0 oC g copper at 100 oC
Clicker Choices:
1: 0 oC 2: 33 oC 3: 50 oC oC 5: 100 oC 6: other

8. What happens to thermal (heat) energy?
When objects of different temperature meet:
Warmer object cools
Cooler object warms
Thermal energy is transferred
qwarmer = -qcooler

9. Quantitative: Calculating Heat Exchange: Specific Heat Capacity

10. Specific Heat Capacity
The energy required to heat one gram of a substance by 1 oC.
Usefulness: #J transferred = S.H. x #g x T
How much energy is used to heat 250 g water from 17 oC to 100 oC?

12. What happens to thermal (heat) energy?
When objects of different temperature meet:
Warmer object cools
Cooler object warms
Thermal energy is transferred
qwarmer = -qcooler
specific heat x mass x T = specific heat x mass x T
warmer object cooler object

13. Heat transfer between substances:

14. Conceptually Easy Example with Annoying Algebra:
If we mix 250 g H2O at 95 oC with 50 g H2O at 5 oC,
what will the final temperature be?

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

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

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

18. 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

19. 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

20. 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)

21. 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)

22. Enthalpy Change and Chemical Reactions
DH = energy needed to break bonds – energy released forming bonds Example: formation of water:
DH = [498 + (2 x 436)] – [4 x 436] kJ = -374 kJ

23. Enthalpy Change and Chemical Reactions
DH is usually more complicated, due to solvent and
solid interactions.
So, we measure DH experimentally.
Calorimetry
Run reaction in a way that the heat exchanged
can be measured. Use a “calorimeter.”

24. 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

25. Hess’s Law
If reactions can be “added”
so can their H values.