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

2. Kinetic Theory are evidence of this. uKinetic theory says that molecules are in constant motion. uPerfume molecules moving across the room

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

4. Ë The particles are in constant random motion. H Move in straight lines until they bounce off each other or the walls. Ì All collisions are perfectly elastic

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

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

7. Kinetic Energy % ofMolecules% ofMolecules High temp. Low temp.

8. Kinetic Energy % ofMolecules% ofMolecules High temp. Low temp. Few molecules have very high kinetic energy

9. Kinetic Energy % ofMolecules% ofMolecules High temp. Low temp. Average kinetic energies are temperatures

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

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

12. Pressure u Pressure is the result of collisions of the molecules with the sides of a container. u A vacuum is completely empty space - it has no pressure. u Pressure is measured in units of atmospheres (atm). u It is measured with a device called a barometer.

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

14. Barometer uAt one atmosphere pressure a column of mercury 760 mm high. uA second unit of pressure is mm Hg u1 atm = 760 mm Hg 760 mm 1 atm Pressure

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

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

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

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

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

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

21. Evaporation u Molecules at the surface break away and become gas. u Only those with enough KE escape u Evaporation is a cooling process. u It requires heat. u 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. / 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. Dynamic equilibrium

24. / 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 Dynamic equilibrium

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

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

27. Kinetic energy % of Molecules% of Molecules T1T1 Energy needed to overcome intermolecular forces

28. Kinetic energy % of Molecules% of Molecules T2T2 u At higher temperature more molecules have enough energy u Higher vapor pressure.

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

30. Changing the Boiling Point u Lower the pressure (going up into the mountains). u Lower external pressure requires lower vapor pressure. u Lower vapor pressure means lower boiling point. u Food cooks slower.

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

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

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

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

35. Cubic

36. Body-Centered Cubic

37. Face-Centered Cubic

38. Phase Changes Solid Liquid Gas Melting Vaporization CondensationFreezing

39. Liquid Sublimation Melting Vaporization Condensation Solid Freezing Gas

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

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

42. Water and Ice Ice Water and Steam Steam -20 0 20 40 60 80 100 120 0 40120 220760800 Heating Curve for Water

43. Water and Ice Ice Water and Steam Steam -20 0 20 40 60 80 100 120 0 40120 220760800 Heating Curve for Water Heat of Vaporization

44. Water and Ice Ice Water and Steam Steam -20 0 20 40 60 80 100 120 0 40120 220760800 Heating Curve for Water Heat of Fusion

45. Water and Ice Ice Water and Steam Steam -20 0 20 40 60 80 100 120 0 40120 220760800 Heating Curve for Water Slope = Specific Heat Steam Water Ice

46. Water and Ice Ice Water and Steam Steam -20 0 20 40 60 80 100 120 0 40120 220760800 Heating Curve for Water Both Water and Steam

47. Water and Ice Ice Water and Steam Steam -20 0 20 40 60 80 100 120 0 40120 220760800 Heating Curve for Water Ice and Water

48. Water and Ice Ice Water and Steam Steam -20 0 20 40 60 80 100 120 0 40120 220760800 Heating Curve for Water

49. Calcualting Energy u Three equations  Heat = specific heat x mass x  T u Heat = heat of fusion x mass u 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 u Heat of vaporization= 540 u Heat of fusion = 80 cal/g

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

52. Examples  How much heat does it take to heat 12 g of ice at -6  C to 25  C water?  How much heat does it take to heat 35 g of ice at 0  C to steam at 150  C?