1. Energy Conservation and Transfer Chm 2.1.1

2. States of Matter Solid KMT –particles packed tightly together –high attraction –Lowest energy of all states –Volume: definite –Shape: definite –Density: High (incompressible) –Types: Crystalline and Amorphous

3. States of Matter Liquid KMT –particles are “fluid”, –less attraction –Higher energy than in solids –Volume: definite –Shape: not definite –Density: relatively high

4. States of Matter Gases KMT –particles are spread apart –very low or no attraction –Highest energy –Volume and Shape not definite –Density: very low –able to be compressed

5. States of Matter

6. Phase Changes of Matter Endothermic –Melting (liquefying) –Boiling –Vaporization (Evaporation) –Sublimation Exothermic –Freezing (solidifying) –Condensing –Deposition

7. Temperature Measures “speed” or Kinetic Energy of particles (T) Measured in ºC or Kelvin (K) K = ºC + 273 Heat measures energy absorbed or released (q) Measured in joules (J) or kilojoules (kJ)

8. Equilibrium in Changes of State Equilibrium: a dynamic condition in which two opposing changes occur at equal rates, at the same time in a closed system

10. Equilibrium Vaporization liquid + heat energy  vapor liquid + heat energy  vaporCondensation vapor  liquid + heat energy At Equilibrium Liquid + Heat Energy ↔ Vapor

11. Vapor Pressure Vapor Pressure = vapor molecules at equilibrium exert pressure Effects of Temperature of Vapor Pressure –Increase Temperature = increases VP –Decrease Temperature = decrease VP –Boiling = when VP of liquid is = to the atmospheric pressure

12. Phase Diagram for H 2 O

13. Phase Diagram for CO 2

14. Heating Curves

15. Heat Calculations Specific heat: –The amount of energy needed to raise 1g of a substance 1°C or 1K –Cp = J/g°C or J/gK –Measured under conditions of constant pressure –All substances have their own specific heat –Cp = q/mΔT

16. Calculating Heat Calculating Heat from Cp –q = mCpΔT Ex: Calculate the amount of heat absorbed by a 6.0g sample of aluminum when the temperature is increased from 500°C to 550°C.

17. The heating curve above gives data for water when it is cooled. How much heat energy is released from 4.0g of water when it is cooled from 100° to 0°

18. Heat of Fusion The amount of energy required/released when melting/freezing 1g of water at 0°C Hf = 334 J/g q = mHf Ex: how much energy will be absorbed by a 3.0g ice cube at 0°C?

19. Heat of Vaporization The amount of energy absorbed or released when vaporizing/condensing 1g of water at 100°C Hv = 2,260 J/g q = mHv Example: How much energy is absorbed when 9.36 g of water is vaporized?

20. Law of Conservation of Energy In a closed system –Energy is cannot be created or destroyed only transferred. –heat lost in a system is gained by the surroundings –Heat gained in a system is lost by the surroundings –q lost = q gained

21. An ice cube while melting gained 25 J of energy. How much energy was lost by its surroundings? A hot piece of metal is put into a cup of water. What will happen to the metal? What will happen to the water? Copper was heated and placed in a cup of water at room temperature (25°C). The water’s temperature rose 30°C. How much energy did the copper loose?