Unit 3: Phase Changes & Energy

1. Unit 3: Phase Changes & Energy

2. PERSPECTIVE The universe is made up of: • The system – the thing that you are studying • The surroundings- everything else

3. What is happening? • Thermal Energy - determined by the movement of the molecules and the potential energy of the arrangement of molecules; sum of all the energy in a sample of matter. • Temperature- average kinetic energy of the system. • Heat - energy transferred from a warmer substance to a colder one by the collisions of molecules.

4. Wait– there’s more. • Thermal Equilibrium-when two objects or systems reach the same temperature and stop exchanging energy through heat. • Heat Transfer- the exchange of thermal energy through conduction (direct contact), convection (in fluids, less dense rise, more dense sink) or radiation (waves).

5. Warming/Cooling Curves • show how a substance’s temperature changes as energy is added (warming) or removed (cooling) • temperature remains constant during a phase change

6. Heating/Cooling Curves Change of Temperature Change of Phase

8. Specific heat • Quantity of heat needed to raise 1 gram of a substance 1° C; unit is J/g°C • Since each substance has a unique chemical composition, each substance has a unique specific heat. • Substances with a low specific heat make good thermal conductors. • Substances with a high specific heat make good thermal insulators.

9. Specific Heat of H2O This means that it takes 4.18 Joules of energy to increase the temperature of 1 gram of liquid water by 1°C.

10. Calculating Heat evolved (removed from the system) or absorbed (added to the system) • FORMULA: q= c x m x ΔT • q= heat (J) • c = specific heat (J/g°C) • m = mass (g) • ΔT = change in temp. (°C)

11. Warming Curve - Water

12. Specific Heat Practice Problem #1 Calculate the amount of heat, in Joules, needed to warm 250 g of water from 25°C to 95°C.

13. Practice Problem #2 How much heat is lost when 50.0 grams of Al is cooled from 130°C to 62°C? The specific heat of Al is 0.897 J/g°C

14. *K classes only* Energy unit conversions: 1 calorie (cal) = 4.18 Joules 1 kilocalorie = 1 Calorie = 1000 calories Practice: convert your answer from pp#2 into calories and Calories

15. Change of state and heat • heat of fusion- amount of heat needed to melt 1 gram of a substance at its melting point • Hf copper = 205 J/g • Hfwater = 334 J/g • q = mHf • heat of vaporization- amount of heat needed to boil 1 gram of a substance at its boiling point • Hv water = 2260 J/g • q= mHv

16. Warming Curve - Water

17. Change of state Practice Problem #1 • Calculate the amount of heat, in Joules, needed to melt 70.0g of copper at its melting point.

18. Change of state Practice Problem #2 • Calculate the heat required, in Joules, to change 250g of water at 100°C to steam at 100°C.

19. Change of state Practice Problem #3 • Calculate the amount of heat needed to change 20g of ice at -10.0°C to water at 80.0°C.

20. Heat Loss= Heat Gain • What is the specific heat of an unknown metal if a 50g piece of it at 175 oC is dropped into 100g of water at 25 oC. The final temperature of the system is 52 oC.

21. Warming Graph

22. Phase Changes http://hyperphysics.phy-astr.gsu.edu/hbase/kinetic/vappre.html • When energy is added or removed from a system, one phase can change to another.

23. PHASE CHANGES THAT REQUIRE ENERGY (ENDOTHERMIC) • Melting- the process of going from a solid to a liquid. This process requires energy to break the “inter-particular” (between particles) forces that hold the solid together. The stronger the forces, the higher the melting point.

24. PHASE CHANGES THAT REQUIRE ENERGY (ENDOTHERMIC) • Vaporization- The process of changing from a liquid to a gas. • When this occurs only at the surface, the process is called evaporation.

25. PHASE CHANGES THAT REQUIRE ENERGY (ENDOTHERMIC) • Evaporation is how our body controls its temperature. For sweat to evaporate, energy is required. This energy comes from your body. This loss of energy leaves us feeling cooler. • A substance that evaporates very easily is said to be volatile.

26. EVAPORATION • As temperature increases the amount of evaporation increases. If evaporation is taking place in a closed container, the evaporated particles exist as a vapor that will exert a pressure on the liquid called vapor pressure.

27. VAPOR PRESSURE Two factors affect vapor pressure: Temperature - increase temp, increase vapor pressure Forces between the particles - the weaker the force, the higher the vapor pressure

29. BOILING POINT • Boiling Point- the temperature where the vapor pressure of a liquid is equal to atmospheric pressure. • A vapor pressure curve can be used to determine the boiling point. • “normal” BP is the bp at standard pressure

30. PHASE CHANGES THAT REQUIRE ENERGY (ENDOTHERMIC) • Sublimation- Changing directly from the solid phase to the gas phase. Dry Ice and iodine are 2 examples.

31. PHASE CHANGES THAT RELEASE ENERGY (EXOTHERMIC) • Freezing- Removing heat from liquid molecules causes them to slow down and form intermolecular bonds. • How does the melting point and freezing point of the same substance compare?

32. PHASE CHANGES THAT RELEASE ENERGY (EXOTHERMIC) • Condensation- going from a gas to a liquid. Bonds are forming. Energy must be released. • Deposition- going directly from a gas to a solid. Snowflakes are an example.

33. Phase Changes

34. Reading a phase diagram • CRITICAL POINT - the point at which critical temperature and critical pressure meet • Above the critical temperature, a liquid cannot be formed by an increase in pressure, though a solid may be formed under sufficient pressure. • The critical pressure is the vapor pressure at the critical temperature. • TRIPLE POINT - the temperature and pressure at which the three phases (gas, liquid, and solid) of that substance coexist in thermodynamic equilibrium

35. Phase diagram - water

36. Phase diagram –CO2