1. Why is water more stable as a solid below 0 o C but as a liquid above it?

2. What controls why only some things happen?

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

4. TYPES of ENERGY KineticPotential MechanicalGravitational ThermalElectrostatic ElectricalChemical Radiant

5. Dropping a rock? 1.Gravitational  Mechanical 2.Gravitational  Thermal 3.Gravitational  Gravitational

6. Using a flashlight? 1.Thermal  Radiant 2.Chemical  Thermal 3.Chemical  Radiant 4.Electrostatic  Radiant

7. Driving a Car?

8. One turn of a giant windmill produces 0.26 kWh of electricity. How many kJ is this?

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

10. Dissolution of NH 4 NO 3 HCl(aq) + NaOH(aq)  H 2 O(l) + NaCl(aq)

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

12. Example 1: 5 g wood at 0 o C + 5 g wood at 100 o C Example 2: 10 g wood at 0 o C + 5 g wood at 100 o C Example 3: 5 g copper at 0 o C + 5 g copper at 100 o C Example 4: 5 g wood at 0 o C + 5 g copper at 100 o C Clicker Choices: 1: 0 o C 2: 33 o C 3: 50 o C 4. 67 o C 5: 100 o C 6: other Temperature Changes from Heat Exchange Bill, go to the next slide. You know you want to.

13. Temperature Changes from Heat Exchange 5 g wood at 0 o C + 5 g wood at 100 o C 1.0 o C 2.33 o C 3.50 o C 4.67 o C 5.100 o C 6.other

14. Temperature Changes from Heat Exchange 10 g wood at 0 o C + 5 g wood at 100 o C 1.0 o C 2.33 o C 3.50 o C 4.67 o C 5.100 o C 6.other

15. Temperature Changes from Heat Exchange 5 g copper at 0 o C + 5 g copper at 100 o C 1.0 o C 2.33 o C 3.50 o C 4.67 o C 5.100 o C 6.other

16. Temperature Changes from Heat Exchange 5 g wood at 0 o C + 5 g copper at 100 o C 1.0 o C 2.33 o C 3.50 o C 4.67 o C 5.100 o C 6.other

17. Conclusion I: What happens to thermal (heat) energy? When objects of different temperature meet: Warmer object cools Cooler object warms Thermal energy is transferred q warmer = -q cooler

18. Conclusion II: What controls the magnitude of an object’s temperature change?

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

20. Quantitative: Calculating Heat Exchange: Specific Heat Capacity

21. Calculate the specific heat capacity of copper: Calculate the specific heat capacity of wood:

22. Experiment! What is the specific heat capacity of this 106.1-g hunk of metal?

23. Which has the smallest specific heat capacity? 1.Wood 2.Copper 3.Silver 4.Water

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

26. Heat transfer between substances:

27. Conceptually Easy Example with Annoying Algebra: If we mix 250 g H 2 O at 95 o C with 50 g H 2 O at 5 o C, what will the final temperature be?

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

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

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

31. 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 But what’s really happening?

32. Energy and Phase Changes: Quantitative Treatment Melting: Heat of Fusion (DH fus ) for Water: 333 J/g Boiling: Heat of Vaporization (DH vap ) for Water: 2256 J/g

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

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

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

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

37. Experiment! Reaction of HCl + HNO 3

38. Calorimetry Experiment N 2 H 4 + 3 O 2  2 NO 2 + 2 H 2 O Energy released = E absorbed by water + E absorbed by calorimeter E water = E calorimeter = Total E =  H = energy/moles = 0.500 g N 2 H 4 600 g water 420 J/ o C