2. Unit: ENERGY

3. Topic 1: Types of Energy Energy = The ability to do work or produce heat Unit of energy = the Joule = J Two types: Potential Kinetic Lord Kelvin

4. Types of Energy 1. Potential Energy – stored energy Ex: food, batteries, chemicals not reacting

5. Types of Energy 2. Kinetic Energy – energy being used, motion Ex: heat, movement, fire, chemical reactions, using electricity

6. Average Kinetic Energy = Temperature

7. Types of Energy Law of Conservation of Energy: Energy cannot be created or destroyed during a chemical process; It is converted from one form of energy to another. Ex: Converted from Potential Energy to Kinetic Energy Ex: Battery to flashlight

8. Topic – Measurement of Heat Energy 1. Specific Heat – the amount of heat required to raise the temperature of a substance by 1º Celsius. Every substance has a different specific heat Which has a higher specific heat – water or sand? Water!

9. Water… Due to it’s high specific heat, water can absorb and release large amounts of heat The specific heat of water is on Table B 4.18 J/g ºC

10. Specific Heat Formula q = mCΔT q = heat (in joules) m = mass of the substance C = specific heat of the substance ΔT = the change in temperature (final – initial)

11. Sample Problem Sample Problem: How many joules are absorbed when 50.0g of water are heated from 30.2 ºC to 58.6 ºC?

12. Work: 1.List your variables: q = ? m = 50.0g C = 4.18 J/g x C ΔT = (58.6 ºC – 30.2 ºC) = 28.4 ºC

13. Answer: 2. Plug into the formula: q = mCΔT q = (50.0g)(4.18J/g x ºC)(28.4 ºC) q = 5936J q = 5.94 x 10 ³ J

14. Topic: Heating and Cooling Curves As you heat a substance, its particles speed up, its temperature rises, and kinetic energy increases. Eventually the temperature of that substance will level off. When would the temperature level off? During a phase change!

16. Cooling Curve

17. Heating and Cooling Curves When a phase change is occurring, and heat is being added at a constant rate, the temperature of the substance remains the same. Where does the heat energy go? It is used to break the bonds between the molecules What law does this demonstrate? Law of Conservation of Energy

19. Heating and Cooling Curves As a solid is heated at a constant rate, eventually some of the particles in the substance possess enough energy to break the bonds holding them in the solid phase. This is known as melting and is also called FUSION

20. Heating and Cooling Curves As heat is added during fusion, the temperature of the substance remains the same! The heat is absorbed by the substance in the form of potential energy. As time goes by, the amount of liquid increases and the amount of solid decreases.

21. Heating and Cooling Curves The amount of heat needed to needed to convert a solid to a liquid at its melting point is called the HEAT OF FUSION

22. Heating and Cooling Curves Once the solid has completely melted into a liquid, the temperature begins to rise again. As temperature increases, so does the kinetic energy of the substance. The temperature will rise until the boiling point is reached. Boiling is also known as vaporization.

23. Identify fusion and vaporization…

24. Question… Why does vaporization take longer than fusion? It requires more energy to turn a liquid into a gas than it does to turn a gas into a liquid.

25. Heating and Cooling Curves Once boiling begins, both liquid and gas phases are present and the temperature will remain the same until all of the liquid has vaporized. If heat is added to a substance in its gas phase, the temperature of the gas will begin to rise again.

26. Question: q = mCΔT when the temperature is changing. We use q = mCΔT when the temperature is changing. What about during a phase change??? There is no ΔT!

27. Enter… Heat of Fusion q = mCΔT During fusion (melting) there is not change in temperature so we can’t use q = mCΔT We use: Q = mH f Q = heat M = mass H f = heat of fusion

28. Problem: How many joules are required to melt 14.2 g of ice at 0°C? Q = mH f Q = (14.2g)(334J/g) Q =

29. I know what you’re thinking… What about vaporization?!?!?! We can’t use the same formula because it takes more energy to boil a liquid then to melt it. So……. Use Q = mH v Hv = Heat of vaporization

30. Question: How many grams of water will boil if 6,752J of energy are added to water at 100°C Q = mH v 6,752J = x(334J/g) X =

31. Topic: Exothermic vs. Endothermic Reactions

32. A. Exothermic Exothermic Reaction = heat energy is released into surroundings, therefore the surroundings get warmer Ex: Heat Pack NaOH + H 2 O  NaOH + Heat

33. Na 2 O 2 + Zn > Na 2 O + ZnO

34. B. Endothermic Endothermic Reaction = absorbs heat energy from its surroundings; therefore the surroundings get colder. Ex: Cold Pack

35. Endothermic Reactions of Barium Hydroxide and Ammonium Salts

36. C. Phase Changes A Phase change is ENDOTHERMIC if heat is absorbed Solid  Liquid  Gas Ex: Melting: Ice + Heat  Water Vaporization: Water + Heat  Steam

37. C. Phase Changes A Phase change is EXOTHERMIC if heat is given off Gas  Liquid  Solid Ex: Condensation: Steam  Water + Heat Freezing:Water  Ice + Heat