1. Energy Energy: the ability to do work –Potential Energy: stored energy –Kinetic Energy: energy of motion Heat: –Energy associated with motion of particles –Units: joule (J) kilojoule (kJ) calorie (cal) –The energy needed to raise the temperature of 1 g of water by 1˚C 1 = 4.184 J (exact number)

2. Energy in Chemical Reactions Activation energy: energy needed for reaction to occur –“hill” we must climb over Heat of reaction: –Amount of heat absorbed or released during a reaction –Exothermic: heat is released –Endothermic: heat is absorbed 2

3. Identify each reaction as Ex) exothermic or En) endothermic. A. N 2 + 3H 2 2NH 3 + 22 kcal B. CaCO 3 + 133 kcalCaO + CO 2 C. 2SO 2 + O 2 2SO 3 + heat 3

4. Calculation using Heat of Reaction The reaction occurring in a cold pack is as follows: 4 Is this reaction exothermic or endothermic? If 10. g of ammonium nitrate is contained in the pack, how much heat can be absorbed or released if it reacts completely? Convert to moles! Moles  energy? Energy is a reactant!

5. In the reaction N 2 (g) + O 2 (g) 2NO(g) ΔH = 43.2 kcal N 2 (g) + O 2 (g) + 43.2 kcal 2NO(g) If 15.0 g NO are produced, how many kcal were absorbed? 5

6. Specific Heat (SH) The amount of heat needed to raise the temperature of 1 g of a substance by 1 °C 6 How much energy is needed to raise the temperature of 20.0 g of iron by 14.5 °C?

7. A hot-water bottle contains 750 g of water at 65°C. If the water cools to body temperature (37°C), how many calories of heat could be transferred to sore muscles? 7

8. Energy and Nutrition & Calculation Nutritional Calorie = Cal = 1 kcal= 1000 cal 8 Carbohydrate Fat (lipid) Protein Caloric Values: food water

9. Entropy –Entropy Describes degree of disorder –Increasing disorder helps drive reactions –Changes in disorder are indicated by ΔS 9

10. Gibbs Free Energy 10 Used to determine if reaction is spontaneous or not Takes into consideration heat of reaction, temperature and entropy Negative ΔG is spontaneous and positive ΔG is nonspontaneous

11. States of Matter Matter: anything that occupies space and has mass Physical states: Solid Liquid Gas (vapor) 11 Shape? Volume? Arrangement of particles? Interaction between particles? Movement of particles? Attractive forces between particles hold substances together Dipole-dipole Hydrogen bonds Dispersion Intermolecular forces

12. Intermolecular Forces Dipole-dipole: –Polar molecules act like magnets –Ex. HCl and HCl Hydrogen bonds: –Strongest dipole-dipole interaction –Only occur between: Hydrogen and… Nitrogen, Oxygen or Fluorine –Ex. Water Dispersion: –Very weak force –Nonpolar molecules (have brief polar moment) 12

13. Changes of State 13

14. Melting/Freezing A substance Is melting when it changes from a solid to a liquid. Is freezing when it changes from a liquid to a solid. Such as water has a freezing (melting) point of 0°C. 14

15. Sublimation Occurs when a solid changes directly to a gas. Is typical of dry ice, which sublimes at -78  C. Takes place in frost-free refrigerators. Is used to prepare freeze-dried foods for long- term storage. 15

16. Evaporation/Condensation Water Evaporates when molecules on the surface gain sufficient energy to form a gas. Condenses when gas molecules lose energy and form a liquid. At boiling, All the water molecules acquire enough energy to form a gas. Bubbles appear throughout the liquid. 16

17. The heat of fusion Is the amount of heat released when 1 gram of liquid freezes (at its freezing point). Is the amount of heat needed to melt 1 gram of a solid (at its melting point). For water (at 0°C) is 80. cal 1 g water 17 Changes of State Equations

18. The heat of vaporization is the amount of heat Absorbed to vaporize 1 g of a liquid to gas at the boiling point. Released when 1 g of a gas condenses to liquid at the boiling point. Boiling Point of Water = 100°C Heat of Vaporization (water) = 540 cal 1 g water = 540 cal / g 18

19. Heating and Cooling Curves

20. Ice Cream 20