Chapter 9- The States of Matter

1. Chapter 9- The States of Matter • Gases indefinite volume and shape, low density. • Liquids definite volume, indefinite shape, and high density. • Solids definite volume and shape, high density • Solids and liquids have high densities because their molecules are close together.

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

3. 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.

4. 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

5. 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.

6. 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.

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

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

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

10. 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.

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

12. 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). • It is measured with a device called a barometer.

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

14. Barometer • At one atmosphere pressure a column of mercury 760 mm high. • A second unit of pressure is mm Hg • 1 atm = 760 mm Hg 1 atm Pressure 760 mm

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

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

17. 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.

18. Think of it it 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.

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

20. 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 ).

21. 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.

22. 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?”

23. 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.

24. 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

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

26. Vaporization • Vaporization is an endothermic process - it requires heat. • Energy is required to overcome intermolecular forces • Responsible for cool earth. • Why we sweat. (Never let them see you.)

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

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

29. 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.

30. 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.

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

32. 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.

33. 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.

34. 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.

35. Cubic

36. Body-Centered Cubic

37. Face-Centered Cubic

38. Melting Vaporization Freezing Condensation Phase Changes Solid Gas Liquid

39. Sublimation Vaporization Condensation Melting Solid Gas Liquid Freezing Condensation

40. 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

41. 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

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

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 Heat of Vaporization

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 Heat of Fusion

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 Steam Water Slope = Specific Heat Ice

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 Both Water and Steam

47. 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

48. 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

49. Calcualting Energy • Three equations • Heat = specific heat x mass x DT • Heat = heat of fusion x mass • Heat = heat of vaporization x mass

50. 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 • Heat of fusion = 80 cal/g