Chapter 10

1. Chapter 10 Kinetic Theory

2. States of Matter Definite Volume? Definite Shape? Expansion w/ heat Com-pressible? Small Expans. Solid YES YES NO Small Expans. Liquid NO NO YES Large Expans. Gas NO NO YES

3. Kinetic Theory • Kinetic theory says that molecules are in constant motion. • Perfume molecules moving across the room are evidence of this.

4. The Kinetic Theory of GasesMakes three assumptions about gases • A Gas is composed of particles • usually molecules or atoms • Considered to be hard spheres far enough apart that we can ignore their volume. • Between the molecules is empty space.

5. The particles are in constant random motion. • Move in straight lines until they bounce off each other or the walls. • All collisions are perfectly elastic They don’t stick together

6. The Average speed of an oxygen molecule is 1656 km/hr at 20ºC • The molecules don’t travel very far without hitting each other so they move in random directions.

7. Kinetic Energy and Temperature • Temperature is a measure of the Average kinetic energy of the molecules of a substance. • Higher temperature = faster molecules. • At absolute zero (0 K), all molecular motion would stop.

8. High temp. % of Molecules Low temp. • Kinetic Energy

9. High temp. Low temp. % of Molecules Few molecules have very high kinetic energy • Kinetic Energy

10. High temp. % of Molecules Low temp. Average kinetic energies are temperatures • Kinetic Energy

11. Temperature • The average kinetic energy is directly proportional to the temperature in Kelvin • If you double the temperature (in Kelvin) you double the average kinetic energy. • If you change the temperature from 300 K to 600 K the kinetic energy doubles.

12. Temperature • If you change the temperature from 300ºC to 600ºC the Kinetic energy doesn’t double. • 873 K is not twice 573 K

13. Pressure • Pressure is the result of collisions of the molecules with the sides of a container. • A vacuum is completely empty space - it has no pressure. • Pressure is measured in units of atmospheres (atm), mm Hg, or kiloPascals. • It is measured with a device called a barometer.

14. Barometer • At one atmosphere pressure a column of mercury 760 mm high. 1 atm Pressure Column of Mercury Dish of Mercury

15. Barometer • At one atmosphere pressure a column of mercury 760 mm high. • Other units for pressure • 1 atm = 760 mm Hg • 1 atm = 101.3 kPa 1 atm Pressure 760 mm

16. Avagadro’s Hypothesis • Equal volumes of gas at the same temperature and pressure have equal numbers of molecules. • That means ...

17. Avagadro’s Hypothesis • Has the same number of particles as .. 2 Liters of Helium 2 Liters of Oxygen

18. This is where we get the fact that 22.4 L =1 mole • Only at STP • 0ºC • 1 atm • This way we compare gases at the same temperature and pressure.

19. Think of it in terms of pressure. • The same pressure at the same temperature should require that there be the same number of particles. • The smaller particles must have a greater average speed to have the same kinetic energy.

20. Liquids • Particles are in motion. • Attractive forces between molecules keep them close together. • These are called intermolecular forces. • Inter = between • Molecular = molecules

21. Breaking intermolecular forces. • Vaporization - the change from a liquid to a gas below its boiling point. • Evaporation - vaporization of an uncontained liquid ( no lid on the bottle ).

22. Evaporation • Molecules at the surface break away and become gas. • Only those with enough KE escape • Evaporation is a cooling process. • It requires heat. • Endothermic.

23. Condensation • Change from gas to liquid • Achieves a dynamic equilibrium with vaporization in a closed system. • What is a closed system? • A closed system means matter can’t go in or out. (put a cork in it) • What the heck is a “dynamic equilibrium?”

24. Dynamic equilibrium • When first sealed the molecules gradually escape the surface of the liquid • As the molecules build up above the liquid some condense back to a liquid.

25. Dynamic equilibrium • As time goes by the rate of vaporization remains constant • but the rate of condensation increases because there are more molecules to condense. • Equilibrium is reached when…

26. Dynamic equilibrium Rate of Vaporization = Rate of Condensation • Molecules are constantly changing phase “Dynamic” • The total amount of liquid and vapor remains constant “Equilibrium”

27. Vaporization • Vaporization is an endothermic process - it requires heat. • Energy is required to overcome intermolecular forces • Responsible for cool earth. • Why we sweat.

28. Energy needed to overcome intermolecular forces % of Molecules T1 Kinetic energy

29. % of Molecules • At higher temperature more molecules have enough energy • Higher vapor pressure. T2 Kinetic energy

30. Boiling • A liquid boils when the vapor pressure = the external pressure • Normal Boiling point is the temperature a substance boils at 1 atm pressure. • The temperature of a liquid can never rise above it’s boiling point.

31. Changing the Boiling Point • Lower the pressure (going up into the mountains). • Lower external pressure requires lower vapor pressure. • Lower vapor pressure means lower boiling point. • Food cooks slower - why?

32. Changing the Boiling Point • Raise the external pressure (Use a pressure cooker). • Raises the vapor pressure needed. • Raises the boiling point. • Food cooks faster.

33. Solids • Intermolecular forces are strong • Can only vibrate and revolve in place. • Particles are locked in place - don’t flow. • Melting point is the temperature where a solid turns into a liquid.

34. The melting point is the same as the freezing point. • When heated the particles vibrate more rapidly until they shake themselves free of each other. • Ionic solids have strong intermolecular forces so a high mp. • Molecular solids have weak intermolecular forces so a low mp.

35. Crystals • A regular repeating three dimensional arrangement of atoms in a solid. • Most solids are crystals. • Amorphous solids lack an orderly internal structure. • Think of them as supercooled liquids.

36. Melting Vaporization Freezing Condensation Phase Changes Solid Gas Liquid

37. Sublimation Vaporization De-Sublimation Melting Solid Gas Liquid Freezing Condensation

38. Energy and Phase Change • Heat of vaporization - energy required to change one gram of a substance from liquid to gas. • Heat of condensation - energy released when one gram of a substance changes from gas to liquid. • For water: 540 cal/g

39. Energy and Phase Change • Heat of fusion - energy required to change one gram of a substance from solid to liquid. • Heat of solidification - energy released when one gram of a substance changes from liquid to solid. • For water: 80 cal/g

40. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800

41. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Heat of Vaporization

42. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Heat of Fusion

43. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Steam Water Slope = Specific Heat Ice

44. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Both Water and Steam

45. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Ice and Water

46. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800

47. Calculating Energy • Three equations • Heat = • heat x mass x DT • Heat = heat of fusion x mass • Heat = heat of vaporization x mass

48. Numbers to Know • For ice S.H. = 0.50 cal/g°C • For water S.H = 1 cal/g°C • For steam S.H. = 0.50 cal/g°C • Heat of vaporization= 540 cal/g • Heat of fusion = 80 cal/g

49. How to do it • The total heat = the sum of all the heats you have to use • Go in order • Use the graph Heat Ice Melt Ice Heat Water Boil Water Heat Steam + + + + Below 0°C At 0°C 0°C - 100°C At 100°C Above 100°C

50. Examples • How much heat does it take to heat 1.0g of ice at 0 °C to steam at 115 °C ? • How much heat does it take to heat 12 g of ice at -6°C to 25°C water?