1. Chapter 13 “States of Matter” Pre-AP Chemistry Charles Page High School Stephen L. Cotton

2. Bellringer #25 (Apr 5 th, 2011) What is KINETIC energy? How is kinetic energy different from POTENTIAL energy? What is KINETIC energy? How is kinetic energy different from POTENTIAL energy?

3. Mini Quiz 13.1 Answer the questions below on a looseleaf sheet of paper using your worksheet from yesterday. Answer the questions below on a looseleaf sheet of paper using your worksheet from yesterday. When does a vacuum exist? When does a vacuum exist? How does atmospheric pressure result? How does atmospheric pressure result? How does atmospheric pressure decrease? How does atmospheric pressure decrease? When could particles of all substances stop moving? When could particles of all substances stop moving?

4. Agenda (Apr 5 th, 2011) Bellringer #25 Bellringer #25 Homework Check/Homework Collection Homework Check/Homework Collection 13.1 Objectives 13.1 Objectives Review Section 13.1 Review Section 13.1 Homework: Cornell Notes 13.2 Homework: Cornell Notes 13.2

5. Section 13.1 The Nature of Gases OBJECTIVES: OBJECTIVES: Describe the assumptions of the “kinetic theory” as it applies to gases. Describe the assumptions of the “kinetic theory” as it applies to gases.

6. Section 13.1 The Nature of Gases OBJECTIVES: OBJECTIVES: Interpret gas pressure in terms of kinetic theory. Interpret gas pressure in terms of kinetic theory.

7. Section 13.1 The Nature of Gases OBJECTIVES: OBJECTIVES: Define the relationship between Kelvin temperature and average kinetic energy. Define the relationship between Kelvin temperature and average kinetic energy.

8. Section 13.1 The Nature of Gases Kinetic refers to motion Kinetic refers to motion Potential energy refers to an object ABOUT to move Potential energy refers to an object ABOUT to move The energy an object has because of it’s motion is called kinetic energy The energy an object has because of it’s motion is called kinetic energy The kinetic theory states that the tiny particles in all forms of matter are in constant motion! The kinetic theory states that the tiny particles in all forms of matter are in constant motion!

9. Section 13.1 The Nature of Gases (Kinetic Theory) Three basic assumptions of the kinetic theory as it applies to gases: Three basic assumptions of the kinetic theory as it applies to gases: #1. Gas is composed of particles- usually molecules or atoms #1. Gas is composed of particles- usually molecules or atoms Small, hard spheres Small, hard spheres Insignificant volume; relatively far apart from each other Insignificant volume; relatively far apart from each other No attraction or repulsion between particles No attraction or repulsion between particles

10. Section 13.1 The Nature of Gases (Kinetic Theory) #2. Particles in a gas move rapidly in constant random motion #2. Particles in a gas move rapidly in constant random motion Move in straight paths, changing direction only when colliding with one another or other objects Move in straight paths, changing direction only when colliding with one another or other objects Average speed of O 2 in air at 20 o C is an amazing 1700 km/h! Average speed of O 2 in air at 20 o C is an amazing 1700 km/h!

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12. Section 13.1 The Nature of Gases (Kinetic Theory) #3. Collisions are perfectly elastic- meaning kinetic energy is transferred without loss from one particle to another- the total kinetic energy remains constant #3. Collisions are perfectly elastic- meaning kinetic energy is transferred without loss from one particle to another- the total kinetic energy remains constant

13. Section 13.1 The Nature of Gases Gas Pressure – defined as the force exerted by a gas per unit surface area of an object Gas Pressure – defined as the force exerted by a gas per unit surface area of an object Due to: a) force of collisions, and b) number of collisions Due to: a) force of collisions, and b) number of collisions No particles present? Then there cannot be any collisions, and thus no pressure – called a vacuum No particles present? Then there cannot be any collisions, and thus no pressure – called a vacuum

14. Section 13.1 The Nature of Gases Atmospheric pressure results from the collisions of air molecules with objects Atmospheric pressure results from the collisions of air molecules with objects Decreases as you climb a mountain because the air layer thins out as elevation increases Decreases as you climb a mountain because the air layer thins out as elevation increases Barometer is the measuring device for atmospheric pressure, which is dependent upon weather & altitude Barometer is the measuring device for atmospheric pressure, which is dependent upon weather & altitude

15. Measuring Pressure The first device for measuring atmospheric pressure was developed by Evangelista Torricelli during the 17 th century. The device was called a “barometer” Baro = weight Meter = measure Torricelli

16. Section 13.1 The Nature of Gases The SI unit of pressure is the pascal (Pa) The SI unit of pressure is the pascal (Pa) At sea level, atmospheric pressure is about 101.3 kilopascals (kPa) At sea level, atmospheric pressure is about 101.3 kilopascals (kPa) Older units of pressure include millimeters of mercury (mm Hg), and atmospheres (atm) – both of which came from using a mercury barometer Older units of pressure include millimeters of mercury (mm Hg), and atmospheres (atm) – both of which came from using a mercury barometer You WILL see atmospheres (atm) again in a later unit You WILL see atmospheres (atm) again in a later unit

17. Section 13.1 The Nature of Gases Mercury Barometer – Fig. 13.2, page 386 – a straight glass tube filled with Hg, and closed at one end; placed in a dish of Hg, with the open end below the surface Mercury Barometer – Fig. 13.2, page 386 – a straight glass tube filled with Hg, and closed at one end; placed in a dish of Hg, with the open end below the surface At sea level, the mercury would rise to 760 mm high at 25 o C- called one standard atmosphere (atm) At sea level, the mercury would rise to 760 mm high at 25 o C- called one standard atmosphere (atm)

18. An Early Barometer The normal pressure due to the atmosphere at sea level can support a column of mercury that is 760 mm high.

19. Section 13.1 The Nature of Gases For gases, it is important to relate measured values to standards For gases, it is important to relate measured values to standards Standard values are defined as a temperature of 0 o C and a pressure of 101.3 kPa, or 1 atm Standard values are defined as a temperature of 0 o C and a pressure of 101.3 kPa, or 1 atm This is called Standard Temperature and Pressure, or STP This is called Standard Temperature and Pressure, or STP

20. Section 13.1 The Nature of Gases What happens when a substance is heated? Particles absorb energy! What happens when a substance is heated? Particles absorb energy! Some of the energy is stored within the particles- this is potential energy, and does not raise the temperature Some of the energy is stored within the particles- this is potential energy, and does not raise the temperature Remaining energy speeds up the particles (increases average kinetic energy)- thus increases temperature Remaining energy speeds up the particles (increases average kinetic energy)- thus increases temperature

21. Section 13.1 The Nature of Gases The particles in any collection have a wide range of kinetic energies, from very low to very high- but most are somewhere in the middle, thus the term average kinetic energy is used The particles in any collection have a wide range of kinetic energies, from very low to very high- but most are somewhere in the middle, thus the term average kinetic energy is used The higher the temperature, the wider the range of kinetic energies The higher the temperature, the wider the range of kinetic energies

22. Section 13.1 The Nature of Gases An increase in the average kinetic energy of particles causes the temperature to rise. An increase in the average kinetic energy of particles causes the temperature to rise. As it cools, the particles tend to move more slowly, and the average K.E. declines. As it cools, the particles tend to move more slowly, and the average K.E. declines. Is there a point where they slow down enough to stop moving? Is there a point where they slow down enough to stop moving?

23. Section 13.1 The Nature of Gases The particles would have no kinetic energy at that point, because they would have no motion The particles would have no kinetic energy at that point, because they would have no motion Absolute zero (0 K, or –273 o C) is the temperature at which the motion of particles theoretically ceases Absolute zero (0 K, or –273 o C) is the temperature at which the motion of particles theoretically ceases This has never been reached, but about 0.5 x 10 -9 K has been achieved This has never been reached, but about 0.5 x 10 -9 K has been achieved

24. Section 13.1 The Nature of Gases The Kelvin temperature scale reflects a direct relationship between temperature and average kinetic energy The Kelvin temperature scale reflects a direct relationship between temperature and average kinetic energy Particles of He gas at 200 K have twice the average kinetic energy as particles of He gas at 100 K Particles of He gas at 200 K have twice the average kinetic energy as particles of He gas at 100 K

25. Section 13.1 The Nature of Gases Solids and liquids differ in their response to temperature Solids and liquids differ in their response to temperature However, at any given temperature the particles of all substances, regardless of their physical state, have the same average kinetic energy However, at any given temperature the particles of all substances, regardless of their physical state, have the same average kinetic energy What happens to the temperature of a substance when the average kinetic energy of its particles decreases? What happens to the temperature of a substance when the average kinetic energy of its particles decreases?

26. Bellringer #26 (Apr 6 th, 2011) How does pressure occur? How does pressure occur? Name something that can INCREASE pressure? Name something that can INCREASE pressure?

27. Bellringer #27 (Apr 11 th, 2011) How would you compare the INTERMOLECULAR attractions of a liquid to the INTERMOLECULAR attractions of a gas?

28. Agenda (Apr 6 th, 2011) Bellringer #27 Bellringer #27 13.2 Cornell Note Check 13.2 Cornell Note Check Write 13.2 Objectives Write 13.2 Objectives Discuss 13.2: The Nature of Liquids Discuss 13.2: The Nature of Liquids Homework: 13.3 Cornell Notes Homework: 13.3 Cornell Notes

29. Section 13.2 The Nature of Liquids OBJECTIVES: OBJECTIVES: Identify factors that determine physical properties of a liquid. Identify factors that determine physical properties of a liquid.

30. Section 13.2 The Nature of Liquids OBJECTIVES: OBJECTIVES: Define “evaporation” in terms of kinetic energy. Define “evaporation” in terms of kinetic energy.

31. Section 13.2 The Nature of Liquids OBJECTIVES: OBJECTIVES: Describe the equilibrium between a liquid and its vapor. Describe the equilibrium between a liquid and its vapor.

32. Section 13.2 The Nature of Liquids OBJECTIVES: OBJECTIVES: Identify the conditions at which boiling occurs. Identify the conditions at which boiling occurs.

33. Section 13.2 The Nature of Liquids Liquid particles are also in motion. Liquid particles are also in motion. Liquid particles are free to slide past one another Liquid particles are free to slide past one another Gases and liquids can both FLOW, as seen in Fig. 13.5, p.390 Gases and liquids can both FLOW, as seen in Fig. 13.5, p.390 However, liquid particles are attracted to each other, whereas gases are not However, liquid particles are attracted to each other, whereas gases are not

34. Section 13.2 The Nature of Liquids Particles of a liquid spin and vibrate while they move, thus contributing to their average kinetic energy Particles of a liquid spin and vibrate while they move, thus contributing to their average kinetic energy But, most of the particles do not have enough energy to escape into the gaseous state; they would have to overcome their intermolecular attractions with other particles But, most of the particles do not have enough energy to escape into the gaseous state; they would have to overcome their intermolecular attractions with other particles

35. Section 13.2 The Nature of Liquids The intermolecular attractions also reduce the amount of space between particles of a liquid The intermolecular attractions also reduce the amount of space between particles of a liquid Thus, liquids are more dense than gases Thus, liquids are more dense than gases Increasing pressure on liquid has hardly any effect on it’s volume Increasing pressure on liquid has hardly any effect on it’s volume

36. Section 13.2 The Nature of Liquids Increasing the pressure also has little effect on the volume of a solid Increasing the pressure also has little effect on the volume of a solid For that reason, liquids and solids are known as the condensed states of matter For that reason, liquids and solids are known as the condensed states of matter

37. Section 13.2 The Nature of Liquids The conversion of a liquid to a gas or vapor is called vaporization The conversion of a liquid to a gas or vapor is called vaporization When this occurs at the surface of a liquid that is not boiling, the process is called evaporation When this occurs at the surface of a liquid that is not boiling, the process is called evaporation Some of the particles break away and enter the gas or vapor state; but only those with the minimum kinetic energy needed to move into the gaseous state Some of the particles break away and enter the gas or vapor state; but only those with the minimum kinetic energy needed to move into the gaseous state

38. Section 13.2 The Nature of Liquids A liquid will also evaporate faster when heated A liquid will also evaporate faster when heated Because the added heat increases the average kinetic energy needed to overcome the attractive forces Because the added heat increases the average kinetic energy needed to overcome the attractive forces But, evaporation is a cooling process But, evaporation is a cooling process Cooling occurs because those with the highest energy escape first (the warmest molecules) Cooling occurs because those with the highest energy escape first (the warmest molecules)

39. Section 13.2 The Nature of Liquids Particles left behind have lower average kinetic energies; thus the temperature decreases Particles left behind have lower average kinetic energies; thus the temperature decreases Similar to removing the fastest runner from a race- the remaining runners have a lower average speed Similar to removing the fastest runner from a race- the remaining runners have a lower average speed Evaporation helps to keep our skin cooler on a hot day. Why? Evaporation helps to keep our skin cooler on a hot day. Why?

40. Section 13.2 The Nature of Liquids Evaporation of a liquid in a closed container is somewhat different Evaporation of a liquid in a closed container is somewhat different Fig. 13.6b on page 391 shows that no particles can escape into the outside air Fig. 13.6b on page 391 shows that no particles can escape into the outside air When some particles do vaporize, these collide with the walls of the container producing vapor pressure (think of a tea kettle) When some particles do vaporize, these collide with the walls of the container producing vapor pressure (think of a tea kettle)

41. Section 13.2 The Nature of Liquids Eventually, some of the particles will return to the liquid, or condense Eventually, some of the particles will return to the liquid, or condense After a while, the number of particles evaporating will equal the number condensing- the space above the liquid is now saturated with vapor After a while, the number of particles evaporating will equal the number condensing- the space above the liquid is now saturated with vapor A dynamic equilibrium exists A dynamic equilibrium exists Rate of evaporation = rate of condensation Rate of evaporation = rate of condensation

42. Section 13.2 The Nature of Liquids We now know the rate of evaporation from an open container increases as heat is added We now know the rate of evaporation from an open container increases as heat is added The heating allows larger numbers of particles at the liquid’s surface to overcome the attractive forces The heating allows larger numbers of particles at the liquid’s surface to overcome the attractive forces Heating allows the average kinetic energy of all particles to increase Heating allows the average kinetic energy of all particles to increase

43. Section 13.2 The Nature of Liquids The boiling point (bp) is the temperature at which the vapor pressure of the liquid is just equal to the external pressure on the liquid The boiling point (bp) is the temperature at which the vapor pressure of the liquid is just equal to the external pressure on the liquid Bubbles form throughout the liquid, rise to the surface, and escape into the air Bubbles form throughout the liquid, rise to the surface, and escape into the air

44. Section 13.2 The Nature of Liquids Since the boiling point is where the vapor pressure equals external pressure, the bp changes if the external pressure changes Since the boiling point is where the vapor pressure equals external pressure, the bp changes if the external pressure changes Normal boiling point- defined as the bp of a liquid at a pressure of 101.3 kPa (or standard pressure) Normal boiling point- defined as the bp of a liquid at a pressure of 101.3 kPa (or standard pressure)

45. Section 13.2 The Nature of Liquids Normal bp of water = 100 o C Normal bp of water = 100 o C However, in Denver = 95 o C, since Denver is 1600 m above sea level and average atmospheric pressure is about 85.3 kPa (Recipe adjustments?) However, in Denver = 95 o C, since Denver is 1600 m above sea level and average atmospheric pressure is about 85.3 kPa (Recipe adjustments?) In pressure cookers, which reduce cooking time, water boils above 100 o C due to the increased pressure In pressure cookers, which reduce cooking time, water boils above 100 o C due to the increased pressure

47. Section 13.2 The Nature of Liquids Turning down the source of external heat drops the liquid’s temperature below the boiling point Turning down the source of external heat drops the liquid’s temperature below the boiling point Supplying more heat allows particles to acquire enough KE to escape- the temperature does not go above the boiling point, the liquid only boils at a faster rate Supplying more heat allows particles to acquire enough KE to escape- the temperature does not go above the boiling point, the liquid only boils at a faster rate

48. - Page 394 Questions: a. 60 o C b. about 20 kPac. about 30 kPa

49. Bellringer #28 (Apr 12 th, 2011) Rank in INCREASING order the INTERMOLECULAR FORCES of solids, liquids, and gases. Rank in INCREASING order the INTERMOLECULAR FORCES of solids, liquids, and gases. Rank in INCREASING ORDER the kinetic energy of solids, liquids, and gases. Rank in INCREASING ORDER the kinetic energy of solids, liquids, and gases.

50. Agenda (Apr 12 th, 2011) Bellringer #28 Bellringer #28 Cornell Notes Check 13.3 Cornell Notes Check 13.3 Write 13.3 Objectives Write 13.3 Objectives Review 13.3 Review 13.3 Pg. 399 #17-20; pg. 407 #45-47; AND 13.4 Cornell Notes Pg. 399 #17-20; pg. 407 #45-47; AND 13.4 Cornell Notes

51. Section 13.3 The Nature of Solids OBJECTIVES: OBJECTIVES: Evaluate how the way particles are organized explains the properties of solids. Evaluate how the way particles are organized explains the properties of solids.

52. Section 13.3 The Nature of Solids OBJECTIVES: OBJECTIVES: Identify the factors that determine the shape of a crystal. Identify the factors that determine the shape of a crystal.

53. Section 13.3 The Nature of Solids OBJECTIVES: OBJECTIVES: Explain how allotropes of an element are different. Explain how allotropes of an element are different.

54. Section 13.3 The Nature of Solids Particles in a liquid are relatively free to move Particles in a liquid are relatively free to move Solid particles are not Solid particles are not Figure 13.10, page 396 shows solid particles tend to vibrate about fixed points, rather than sliding from place to place Figure 13.10, page 396 shows solid particles tend to vibrate about fixed points, rather than sliding from place to place

55. Section 13.3 The Nature of Solids Most solids have particles packed against one another in a highly organized pattern Most solids have particles packed against one another in a highly organized pattern Tend to be dense and incompressible Tend to be dense and incompressible Do not flow, nor take the shape of their container Do not flow, nor take the shape of their container Are still able to move, unless they would reach absolute zero Are still able to move, unless they would reach absolute zero

56. Section 13.3 The Nature of Solids When a solid is heated, the particles vibrate more rapidly as the kinetic energy increases When a solid is heated, the particles vibrate more rapidly as the kinetic energy increases The organization of particles within the solid breaks down, and eventually the solid melts The organization of particles within the solid breaks down, and eventually the solid melts The melting point (mp) is the temperature a solid turns to liquid The melting point (mp) is the temperature a solid turns to liquid

57. Section 13.3 The Nature of Solids At the melting point, the disruptive vibrations are strong enough to overcome the interactions holding them in a fixed position At the melting point, the disruptive vibrations are strong enough to overcome the interactions holding them in a fixed position Melting point can be reversed by cooling the liquid so it freezes Melting point can be reversed by cooling the liquid so it freezes Solid liquid Solid liquid

58. Section 13.3 The Nature of Solids Generally, most ionic solids have high melting points, due to the relatively strong forces holding them together Generally, most ionic solids have high melting points, due to the relatively strong forces holding them together Sodium chloride (an ionic compound) has a melting point = 801 o C Sodium chloride (an ionic compound) has a melting point = 801 o C Molecular compounds have relatively low melting points Molecular compounds have relatively low melting points

59. Section 13.3 The Nature of Solids Hydrogen chloride (a molecular compound) has a mp = -112 o C Hydrogen chloride (a molecular compound) has a mp = -112 o C Not all solids melt- wood and cane sugar tend to decompose when heated Not all solids melt- wood and cane sugar tend to decompose when heated Most solid substances are crystalline in structure Most solid substances are crystalline in structure

60. Section 13.3 The Nature of Solids In a crystal, such as Fig. 13.10, page 396, the particles (atoms, ions, or molecules) are arranged in a orderly, repeating, three- dimensional pattern called a crystal lattice In a crystal, such as Fig. 13.10, page 396, the particles (atoms, ions, or molecules) are arranged in a orderly, repeating, three- dimensional pattern called a crystal lattice All crystals have a regular shape, which reflects their arrangement All crystals have a regular shape, which reflects their arrangement

61. Section 13.3 The Nature of Solids The shape of a crystal depends upon the arrangement of the particles within it The shape of a crystal depends upon the arrangement of the particles within it The smallest group of particles within a crystal that retains the geometric shape of the crystal is known as a unit cell The smallest group of particles within a crystal that retains the geometric shape of the crystal is known as a unit cell

62. Section 13.3 The Nature of Solids There are three kinds of unit cells that can make up a cubic crystal system: There are three kinds of unit cells that can make up a cubic crystal system: 1. Simple cubic 1. Simple cubic 2. Body-centered cubic 2. Body-centered cubic 3. Face-centered cubic 3. Face-centered cubic 90 o angle

63. - Page 398

64. Section 13.3 The Nature of Solids Some solid substances can exist in more than one form Some solid substances can exist in more than one form Elemental carbon is an example, as shown in Fig. 13.13, page 399 Elemental carbon is an example, as shown in Fig. 13.13, page 399 1. Diamond, formed by great pressure 1. Diamond, formed by great pressure 2. Graphite, which is in your pencil 2. Graphite, which is in your pencil 3. Buckminsterfullerene (also called “buckyballs”) arranged in hollow cages like a soccer ball 3. Buckminsterfullerene (also called “buckyballs”) arranged in hollow cages like a soccer ball

65. Section 13.3 The Nature of Solids These are called allotropes of carbon, because all are made of pure carbon only, and all are solid These are called allotropes of carbon, because all are made of pure carbon only, and all are solid Allotropes are two or more different molecular forms of the same element in the same physical state Allotropes are two or more different molecular forms of the same element in the same physical state Not all solids are crystalline, but instead are amorphous Not all solids are crystalline, but instead are amorphous

66. Section 13.3 The Nature of Solids Amorphous solids lack an ordered internal structure Amorphous solids lack an ordered internal structure Rubber, plastic, and asphalt are all amorphous solids- their atoms are randomly arranged Rubber, plastic, and asphalt are all amorphous solids- their atoms are randomly arranged Another example is glass- substances cooled to a rigid state without crystallizing Another example is glass- substances cooled to a rigid state without crystallizing

67. Section 13.3 The Nature of Solids Glasses are sometimes called supercooled liquids Glasses are sometimes called supercooled liquids The irregular internal structures of glasses are intermediate between those of a crystalline solid and a free- flowing liquid The irregular internal structures of glasses are intermediate between those of a crystalline solid and a free- flowing liquid Do not melt at a definite mp, but gradually soften when heated Do not melt at a definite mp, but gradually soften when heated

68. Section 13.3 The Nature of Solids When a crystalline solid is shattered, the fragments tend to have the same surface angles as the original solid When a crystalline solid is shattered, the fragments tend to have the same surface angles as the original solid By contrast, when amorphous solids such as glass is shattered, the fragments have irregular angles and jagged edges By contrast, when amorphous solids such as glass is shattered, the fragments have irregular angles and jagged edges

69. Summary Chart State of Matter ArrangementMotion Intermolecular Forces Kinetic Energy Solid Very Close Particles, No Space Between Only Vibrate HighLow Liquid Some space between particles Slide, Vibrate, Spin MiddleMiddle Gas Mostly space between particles RandomNoneHigh

70. Bellringer #29 What happens to the particles in a solid when they get heated? What happens to the particles in a liquid when they get heated? What happens to the particles in a solid when they get heated? What happens to the particles in a liquid when they get heated?

71. Agenda Bellringer #29 Bellringer #29 Homework Collection/Homework Check Homework Collection/Homework Check Review 13.4 Review 13.4 Unit 13 Review Sheet Unit 13 Review Sheet Test FRIDAY!!!! Test FRIDAY!!!!

72. Section 13.4 Changes of State OBJECTIVES: OBJECTIVES: Identify the conditions necessary for sublimation. Identify the conditions necessary for sublimation.

73. Section 13.4 Changes of State OBJECTIVES: OBJECTIVES: Describe how equilibrium conditions are represented in a phase diagram. Describe how equilibrium conditions are represented in a phase diagram.

74. Section 13.4 Changes of State OBJECTIVES: OBJECTIVES: Identify if a state change is endothermic or exothermic. Identify if a state change is endothermic or exothermic.

75. Section 13.4 Changes of State OBJECTIVES: OBJECTIVES: Identify the sections in a heating/cooling curve. Identify the sections in a heating/cooling curve.

76. Section 13.4 Changes of State Sublimation- the change of a substance from a solid directly to a vapor, without passing through the liquid state Sublimation- the change of a substance from a solid directly to a vapor, without passing through the liquid state Examples: iodine (Fig. 13.14, p. 401); dry ice (-78 o C); mothballs; solid air fresheners Examples: iodine (Fig. 13.14, p. 401); dry ice (-78 o C); mothballs; solid air fresheners

77. Section 13.4 Changes of State Sublimation is useful in situations such as freeze-drying foods- such as by freezing the freshly brewed coffee, and then removing the water vapor by a vacuum pump Sublimation is useful in situations such as freeze-drying foods- such as by freezing the freshly brewed coffee, and then removing the water vapor by a vacuum pump Also useful in separating substances - organic chemists use it separate mixtures and purify materials Also useful in separating substances - organic chemists use it separate mixtures and purify materials

78. Section 13.4 Changes of State The relationship among the solid, liquid, and vapor states (or phases) of a substance in a sealed container are best represented in a single graph called a phase diagram The relationship among the solid, liquid, and vapor states (or phases) of a substance in a sealed container are best represented in a single graph called a phase diagram Phase diagram- gives the temperature and pressure at which a substances exists as solid, liquid, or gas (vapor) Phase diagram- gives the temperature and pressure at which a substances exists as solid, liquid, or gas (vapor)

79. Section 13.4 Changes of State Fig. 13.15, page 403 shows the phase diagram for water Fig. 13.15, page 403 shows the phase diagram for water Each region represents a pure phase Each region represents a pure phase Line between regions is where the two phases exist in equilibrium Line between regions is where the two phases exist in equilibrium Triple point is where all 3 curves meet, the conditions where all 3 phases exist in equilibrium! Triple point is where all 3 curves meet, the conditions where all 3 phases exist in equilibrium!

80. Phase changes by Name Critical Point Temperature ( o C) Pressure (kPa)

81. - Page 403 Questions:

82. Exothermic vs. Endothermic When state changes occur a substance is either ABSORBING or RELEASING heat. When state changes occur a substance is either ABSORBING or RELEASING heat. Exothermic: when a substance/reaction RELEASES heat (ex: when liquid water freezes) Exothermic: when a substance/reaction RELEASES heat (ex: when liquid water freezes) Endothermic: when a substance/reaction ABSORBS heat (ex: when ice melts) Endothermic: when a substance/reaction ABSORBS heat (ex: when ice melts)

83. Endothermic or Exothermic? 1) Taking Ice Cream Out the Freezer 2) Lighting a Match 3) Morning Dew Forming on Grass 4) Boiling water

84. Heating Curve (shows ENDOTHERMIC changes)

85. Bellringer #30 (Apr 14 th, 2011) You take a glass of water outside on a HOT summer day. Someone calls you inside for awhile and you leave your glass of ice water outside. What types of phase changes have taken place with the glass of ice water? State if these changes are EXOTHERMIC or ENDOTHERMIC. You take a glass of water outside on a HOT summer day. Someone calls you inside for awhile and you leave your glass of ice water outside. What types of phase changes have taken place with the glass of ice water? State if these changes are EXOTHERMIC or ENDOTHERMIC.

87. Agenda (Apr 14 th, 2011) Bellringer #30 Bellringer #30 Questions about Review Sheet Questions about Review Sheet Continue Working on Review Sheet Continue Working on Review Sheet Extra Credit Option: pg.408 #50, 51, 52, 55, 56, 57, 58, 59, 63, 65 Extra Credit Option: pg.408 #50, 51, 52, 55, 56, 57, 58, 59, 63, 65