1. What is energy? Two types: Potential and Kinetic Page 2
The ability to do work
The ability to transfer heat
Two types: Potential and Kinetic
Page 2

2. Potential Energy Stored Energy Energy due to
Position
Chemical Bonds
Nuclear
Position: boulder at the top of the hill
Chemical: tank of gas, hamburger
Nuclear: atomic bomb, nuclear reactor

3. Kinetic Energy Active Energy Energy of Motion
Electromagnetic waves, (ex. Light, Xrays)
Heat
Motion
Electrical current,
A moving truck has the ability to flatten you - do work on you!

4. Kinetic and Potential Energy

5. Electromagnetic Spectrum

6. First Law of Thermodynamics
Energy can neither be created nor destroyed, but may change from one form to another.
Page 2 bottom

7. Entropy – 2nd Law Entropy is the amount of disorder in a system
Entropy always increases over time (in the absence of an input of outside energy)
Example: cleaning up your room

8. Page 3

10. Page 4

11. Exothermic vs Endothermic

12. Chemical OR Physical changes can be exothermic or endothermic.
Definition
Stored energy
Energy of motion
Energy is
Absorbed
Released
Temperature
+∆H
Chemical -∆T
Physical +∆T
-∆H
Chemical +∆T
Physical -∆T
Type of Energy Conversion
Example
2H2O + energy  2H2 + O2
Energy is on the left
2H2 + O2  2H2O + energy
Energy is on the right

14. Regents Question: 06/02 #64-66
A hot pack contains chemicals that can be activated to produce heat. A cold pack contains chemicals that feel cold when activated.
Based on energy flow, state the type of chemical change that occurs in a hot pack.
Exothermic
A cold pack is placed on an injured leg. Indicate the direction of the flow of energy between the leg and the cold pack.
From the leg to the cold pack (Hot to Cold)
What is the Law of Conservation of Energy? Describe how the Law of Conservation of Energy applies to the chemical reaction that occurs in the hot pack.
Energy cannot be created nor destroyed. It can only be changed from one form to another. The heat produced in the hot pack was stored in the chemical bonds.

15. Page 5
Measuring Energy

16. Energy
There are two units which are commonly used:
Calories (c):
amount of energy it takes to raise one gram of water one degree Celsius
Joules (J):
4.18 Joules = 1 calorie
Metric system is most commonly used in chemistry

17. Criteria
Heat
Temperature
Similar/Different
Energy
Kinetic
Motion of Molecules
Both are about Motion
Quantitative Aspect
How fast molecules are moving. Measured by temperature
Kelvin and Celsius
Celsius based on properties of water
Kelvin based on Celsius
Definition
A form of Kinetic Energy that involves movement of molecules
The measurement of how fast a molecule is moving
Temperature is a measurement of Kinetic Energy
Examples

18. Page 6 Converting between Celsius and Kelvin Reference Table
Why is it not out already?
Temperature
Towards bottom
K = ◦C + 273

19. Heat and Temperature
Temperature measures the average speed of the atoms
Heat is the amount of kinetic energy of the atoms

20. Page 7

21. To convert between Kelvin and Celsius use K=◦C+273
J Deutsch 2003

23. Page 8

24. Phases of Matter- Page 9 Ice Ice Ice Regular Irregular Irregular
Minimal
Moderate
Fast

25. Page 10

26. Graph page 11 onto page 13. Be sure to have an appropriate scale
Graph page 11 onto page 13. Be sure to have an appropriate scale. Circle the points and connect them.

27. F D E B C A

28. A change in phase is a change in Potential Energy, not Kinetic Energy
Boiling Point
Potential energy changes, so temperature doesn’t
Melting Point

30. Energy and phase changesAB
solid warms up (KE/PE constant)

32. Energy and phase changesAB
solid warms up (KE/PE constant) BC
solid melts (KE constant/PE)

34. Energy and phase changesAB
solid warms up (KE/PE constant) BC
solid melts (KE constant/PE) CD
liquid warms up (KE/PE constant)

36. Energy and phase changesAB
solid warms up (KE/PE constant) BC
solid melts (KE constant/PE) CD
liquid warms up (KE/PE constant) DE
liquid boils (KE constant/PE)

38. Energy and phase changesAB
solid warms up (KE/PE constant) BC
solid melts (KE constant/PE) CD
liquid warms up (KE/PE constant) DE
liquid boils (KE constant/PE) EF
gas warms (KE/PE constant)

40. Regents Question: 06/02 #28
As ice melts at standard pressure, its temperature remains at 0°C until it has completely melted. Its potential energy
(1) decreases
(2) increases
(3) remains the same þ
J Deutsch 2003

41. Regents Question: 08/02 #54
A sample of water is heated from a liquid at 40°C to a gas at 110°C. The graph of the heating curve is shown in your answer booklet.
a On the heating curve diagram provided in your answer booklet, label each of the following regions:
Liquid, only Gas, only Phase change
Liquid Only
Gas Only
Phase change
J Deutsch 2003

42. Regents Question: cont’d
b For section QR of the graph, state what is happening to the water molecules as heat is added.
c For section RS of the graph, state what is happening to the water molecules as heat is added.
They move faster, their temperature increases.
Their intermolecular bonds are breaking, their potential energy is increasing.
J Deutsch 2003

43. Regents Question: 01/02 #47
What is the melting point of this substance?
(1) 30°C (3) 90°C
(2) 55°C (4) 120°C þ
J Deutsch 2003

44. Graph page 14 onto page 16. Be sure to have an appropriate scale
Graph page 14 onto page 16. Be sure to have an appropriate scale. Circle the points and connect them.

45. A B C D E F

46. Energy and phase changesAB
Gas cools down (KE/PE constant) BC
Gas condenses (KE constant/PE ) CD
liquid cools down (KE  /PE constant) DE
liquid freezes (KE constant/PE ) EF
Solid cools down (KE  /PE constant)

47. Pages 17-18

50. Page 19
How do we calculate amount of heat,(Q), if we are not given a graphic?

51. 3 equations for Q Q = mCT Q = mHf Q = mHv
Have to figure out which one to use for a given problem.
Depends which section of heating curve.
Look for hints in the problem.

52. Calculating Heat Transferred
Simple system: Pure substance in a single phase. To calculate heat gained or lost, use:
Q = mCT
Q = amount of heat transferred
m = mass of substance
C = specific heat capacity of the substance (Table B).
T = temperature change = Tfinal – Tinitial

53. Q = mCT Final temperature Ending temperature Temperature changed
From ____ to ____
Water
Temperature changed
Temperature increased
Temperature decreased
Initial / Start temperature

54. Amount of energy required to convert 1 gram of a pure substance from the solid to the liquid phase at the melting point.
Heat of Fusion Hf

55. Calculating Energy Change at Phase Change
Q = mHf
Use this equation to calculate energy changes for phase changes between ice & liquid water at 0C.

56. Q = mHf
Ice
Freezing
Melting
At 0C (for H2O)
At constant temperature

57. Amount of energy required to convert 1 gram of a pure substance from the liquid to the gas phase at the boiling point.
Heat of Vaporization Hv

58. Calculating Energy Change at Phase Change
Q = mHv
Use this equation to calculate energy changes for phase changes between steam & liquid water at 100C.

59. Q = mHv Steam Boiling Condensation At 100C (for H2O)
At constant temperature