2. Chapter 11
Section 1 Notes
2/1/16

3. Chapter 11 States of Matter
Section 1 States and State Changes
States of Matter
Most substances can be in three states: solid, liquid, and gas.
Solid Particles Have Fixed Positions
The particles in a solid are very close together and have an orderly, fixed arrangement.
Solid particles can vibrate only in place and do not break away from their fixed positions.
Solids have fixed volumes and shapes.

4. Mercury in Three States
Chapter 11
Section 1 States and State Changes
Mercury in Three States

5. States of Matter, continued
Chapter 11
Section 1 States and State Changes
States of Matter, continued
Liquid Particles Can Move Easily Past One Another
The particles in a liquid are very close together and have a random arrangement.
Liquid particles have enough energy to be able to move past each other readily, which allows liquids to flow.
Liquids have fixed volumes but can flow to take the shape of the lower part of a container.

6. States of Matter, continued
Chapter 11
Section 1 States and State Changes
States of Matter, continued
Liquid Forces Lead to Surface Wetting and Capillary Action
Liquid particles can have cohesion, attraction for each other.
Liquid particles can also have adhesion, attraction for particles of solid surfaces.
The balance of cohesion and adhesion determines whether a liquid will wet a solid surface.
The forces of adhesion and cohesion will pull water up a narrow glass tube, called a capillary tube.

7. Comparing Cohesion and Adhesion
Chapter 11
Section 1 States and State Changes
Comparing Cohesion and Adhesion

8. States of Matter, continued
Chapter 11
Section 1 States and State Changes
States of Matter, continued
Gas Particles Are Essentially Independent
The particles in a gas are very far apart and have a random arrangement.
The attractive forces between particles in a gas do not have a great effect, so the particles move almost independently of one another.
The shape, volume, and density of an amount of gas change depending on the size and shape of the container.

9. Chapter 11
Section 1 States and State Changes
Solid, Liquid and Gas

10. Chapter 11 Changing States
Section 1 States and State Changes
Changing States
Most substances can undergo six changes of state: freezing, melting, evaporation, condensation, sublimation, and deposition.
Temperature, Energy, and State
Generally, adding energy to a substance will increase the substance’s temperature.
But after a certain point, adding more energy will cause a substance to experience a change of state instead of a temperature increase.

11. Changes of State (copy into notes)
Chapter 11
Section 1 States and State Changes
Changes of State (copy into notes)

12. Changing States, continued
Chapter 11
Section 1 States and State Changes
Changing States, continued
Liquid Evaporates to Gas
Energy is required to separate liquid particles. They gain energy when they collide with each other.
If a particle gains a large amount of energy, it can leave the liquid’s surface and join gas particles.
Evaporation is the change of state from liquid to gas. Evaporation is an endothermic process.
Boiling point is the temperature and pressure at which a liquid turns into vapor.

13. Changing States, continued
Chapter 11
Section 1 States and State Changes
Changing States, continued
Gas Condenses to Liquid
When gas particles no longer have enough energy to overcome the attractive forces between them, they go into the liquid state.
Condensation is the change of state from a gas to a liquid. Condensation is an exothermic process.
Condensation can take place on a cool night, causing water vapor in the air to form dew on plants.

14. Vaporization and Condensation
Chapter 11
Visual Concepts
Vaporization and Condensation

15. Changing States, continued
Chapter 11
Section 1 States and State Changes
Changing States, continued
Solid Melts to Liquid
As a solid is heated, the particles vibrate faster and faster in their fixed positions.
At a certain temperature, some of the molecules have enough energy to break out of their fixed positions.
Melting is the change of state from solid to liquid. Melting is an endothermic process.
Melting point is the temperature and pressure at which a solid becomes a liquid.

16. Changing States, continued
Chapter 11
Section 1 States and State Changes
Changing States, continued
Liquid Freezes to Solid
As a liquid is cooled, the movement of particles becomes slower and slower.
At a certain temperature, the particles are pulled together into the fixed positions of the solid state.
Freezing is the change of state from a liquid to a solid. Freezing is an exothermic process.
Freezing point is the temperature at which a substance freezes.

17. What process occurs as a gas turns into a liquid?
Chapter 11
Section 1 States and State Changes
Warm-up 2/2/16
What process occurs as a gas turns into a liquid?
Is the above an endothermic or exothermic process?
What process occurs as a liquid turns into a solid?
Is the above process an endothermic or exothermic process?

18. Chapter 11
Visual Concepts
Freezing

19. Changing States, continued
Chapter 11
Section 1 States and State Changes
Changing States, continued
Solid Sublimes to Gas
The particles in a solid are constantly vibrating. Some particles have higher energy than others.
Particles with high enough energy can escape from the solid.
Sublimation is the change of state from solid to gas.
Sublimation is an endothermic process.

20. Changing States, continued
Chapter 11
Section 1 States and State Changes
Changing States, continued
Gas Deposits to Solid
Molecules in the gaseous state become part of the surface of a crystal.
When a substance changes state from a gas to a solid, the change is often called deposition.
Deposition is an exothermic process.

21. Comparing Sublimation and Deposition
Chapter 11
Visual Concepts
Comparing Sublimation and Deposition

22. Warm-Up 2/3/16
What is the specific heat of aluminum if the temperature of a 28.4 g sample of aluminum is increased by 8.1 oC when 207 J of heat is added?
What is the specific heat of silicon if the temperature of a 4.11 g sample of silicon is increased by 3.8 oC when 11.1 J of heat is added?
How much heat must be added to a 8.21 g sample of gold to increase its temperature by 6.2 oC? The specific heat of gold is 0.13 J/goC.

23. Warm-Up 2/4/16
If 40.5 J of heat is added to a 15.4 g sample of silver, how much will the temperature increase by? The specific heat of silver is J/g*C.

24. Warm-Up 2/5/16
1. The specific heat of ethanol is 2.46 J/g oC. Find the heat required to raise the temperature of 193 g of ethanol from 19oC to 35oC.
2. To what temperature will a 50.0 g piece of glass raise if it absorbs 5275 joules of heat and its specific heat capacity is 0.50 J/g°C? The initial temperature of the glass is 20.0°C.

25. Warm-Up 2/9/16
Which type of bond is the strongest?
What determines bond strength?
How is bond strength related to melting/boiling points?

26. Chapter 11
Section 2 Notes
2/9/16

27. Comparing Ionic and Covalent Compounds
Chapter 11
Section 2 Intermolecular Forces
Comparing Ionic and Covalent Compounds
It takes energy to overcome the forces holding particles together.
Thus, it takes energy to cause a substance to go from the liquid to the gaseous state.
The boiling point of a substance is therefore a good measure of the strength of the forces that hold the particles together.
Melting point also relates to attractive forces between particles.

28. Chapter 11 Comparing Ionic and Covalent Compounds, continued
Section 2 Intermolecular Forces
Comparing Ionic and Covalent Compounds, continued
Most covalent compounds melt at lower temperatures than ionic compounds do.

29. Chapter 11 Comparing Ionic and Covalent Compounds, continued
Section 2 Intermolecular Forces
Comparing Ionic and Covalent Compounds, continued
Oppositely Charged Ions Attract Each Other
Ionic substances generally have much higher forces of attraction than covalent substances.
For small ions, attractions between ions of opposite charge hold the ions tightly in a crystal lattice.
These attractions are overcome only by heating to very high temperatures.

30. Chapter 11 Comparing Ionic and Covalent Compounds, continued
Section 2 Intermolecular Forces
Comparing Ionic and Covalent Compounds, continued
Oppositely Charged Ions Attract Each Other, continued
If the ions are larger, then the distances between them are larger and the forces are weaker.
Thus, ionic compounds with small ions have high melting points.

31. Chapter 11 Comparing Ionic and Covalent Compounds, continued
Section 2 Intermolecular Forces
Comparing Ionic and Covalent Compounds, continued
Intermolecular Forces Attract Molecules to Each Other
Intermolecular forces are the forces of attraction between molecules of covalent compounds.
Intermolecular forces include dipole-dipole forces and London dispersion forces.

32. Chapter 11 Dipole-Dipole Forces
Section 2 Intermolecular Forces
Dipole-Dipole Forces
Dipole-Dipole Forces Affect Melting and Boiling Points
Dipole-dipole forces are interactions between polar molecules.
When molecules are very polar, the dipole-dipole forces are very significant.
The more polar the molecules are, the higher the boiling point of the substance.

33. Chapter 11
Section 2 Intermolecular Forces
Dipole-Dipole Forces

34. Chapter 11 Hydrogen Bonds
Section 2 Intermolecular Forces
Hydrogen Bonds
A hydrogen bond is a dipole-dipole force occurring when a hydrogen atom that is bonded to a highly electronegative atom of one molecule is attracted to two unshared electrons of another molecule.
In general, compounds with hydrogen bonding have higher boiling points than comparable compounds.

35. Hydrogen Bonds, continued
Chapter 11
Section 2 Intermolecular Forces
Hydrogen Bonds, continued
Hydrogen Bonds Form with Electronegative Atoms
Strong hydrogen bonds can form with a hydrogen atom that is covalently bonded to very electronegative atoms in the upper-right part of the periodic table: nitrogen, oxygen, and fluorine.
Hydrogen Bonds Are Strong Dipole-Dipole Forces
The combination of the large electronegativity difference (high polarity) and hydrogen’s small size accounts for the strength of the hydrogen bond.

36. Chapter 11
Section 2 Intermolecular Forces
Hydrogen Bonding

37. Hydrogen Bonds, continued
Chapter 11
Section 2 Intermolecular Forces
Hydrogen Bonds, continued
Hydrogen Bonding Explains Water’s Unique Properties
Each water molecule forms multiple hydrogen bonds, so the intermolecular forces in water are strong.
The angle between the two H atoms is 104.5°.
When water forms ice, the ice crystals have large amounts of open space. Thus, ice has a low density.
Water is unusual in that its liquid form is denser than its solid form.

38. Chapter 11
Section 2 Intermolecular Forces
Ice and Water

39. London Dispersion Forces
Chapter 11
Section 2 Intermolecular Forces
London Dispersion Forces
A substance with weak attractive forces will be a gas because there is not enough attractive force to hold molecules together as a liquid or a solid.
However, many nonpolar substances are liquids.
What forces of attraction hold together nonpolar molecules and atoms?
London dispersion forces are the dipole-dipole force resulting from the uneven distribution of electrons and the creation of temporary dipoles.

40. London Dispersion Force
Chapter 11
Section 2 Intermolecular Forces
London Dispersion Force

41. London Dispersion Forces, continued
Chapter 11
Section 2 Intermolecular Forces
London Dispersion Forces, continued
Properties Depend on Types of Intermolecular Force
The differences in the properties of the substances are related to the differences in the types of forces that act within each substance.
Nonpolar molecules can experience only London dispersion forces.
Polar molecules experience both dipole-dipole forces and London dispersion forces.

42. Warm-Up 2/10/16
In a spontaneous reaction, what is happening to the entropy?
What is the process called in which a liquid becomes a solid?
What is the equation to solve for ∆Hreaction if ∆Hproducts and ∆Hreactants are known?

43. Warm-Up 2/11/16
What is the exact definition of specific heat?
A solid with what type of bond will have the highest boiling point?
Would it require more energy to raise the temperature of 10 grams of mercury 10 degrees Celsius or 5 grams of mercury 5 degrees Celsius?

44. Warm-Up 2/12/16
What is enthalpy?
In what state does pressure effect the entropy of a substance?
What is a hydrogen bond?
In a polar covalent molecule what types of intermolecular forces can occur?

45. Chapter 11
Section 3
2/12/16

46. Enthalpy, Entropy, and Changes of State
Chapter 11
Section 3 Energy of State Changes
Enthalpy, Entropy, and Changes of State
Enthalpy is the total energy of a system.
Entropy measures a system’s disorder.
The energy added during melting or removed during freezing is called the enthalpy of fusion.
Particle motion is more random in the liquid state, so as a solid melts, the entropy of its particles increases. This increase is the entropy of fusion.

47. Chapter 11
Visual Concepts
Entropy

48. Gibbs Energy and State Changes, continued
Chapter 11
Section 3 Energy of State Changes
Gibbs Energy and State Changes, continued
Pressure Can Affect Change-of-State Processes
Boiling points are pressure dependent because pressure has a large effect on the entropy of a gas.
When a gas is expanded (pressure is decreased), its entropy increases because the degree of disorder of the molecules increases.
Thus, boiling points change as atmospheric pressure changes with changes in elevation.

49. Gibbs Energy and State Changes, continued
Chapter 11
Section 3 Energy of State Changes
Gibbs Energy and State Changes, continued
Pressure Can Affect Change-of-State Processes
Liquids and solids are almost incompressible.
Therefore, changes of atmospheric pressure have little effect on the entropy of substances in liquid or solid states.
Also, ordinary changes in pressure have essentially no effect on melting and freezing.
Thus, melting and freezing points hardly change at all as elevation changes.

50. Chapter 11
Section 4 Phase Equilibrium
Chapter 11
Section 4

51. Chapter 11 Two-Phase Systems
Section 4 Phase Equilibrium
Two-Phase Systems
A phase is a region that has the same composition and properties throughout.
For example, ice water is a system that has a solid phase and a liquid phase.

52. Two-Phase Systems, continued
Chapter 11
Section 4 Phase Equilibrium
Two-Phase Systems, continued
Equilibrium Involves Constant Interchange of Particles
A dynamic equilibrium exists when particles are constantly moving between two or more phases yet no net change in the amount of substance in either phase takes place.
When you cap a bottle of rubbing alcohol, the liquid and gas are at equilibrium.
That is, the rate of evaporation equals the rate of condensation.

53. Chapter 11 Phase Diagrams
Section 4 Phase Equilibrium
Phase Diagrams
A substance’s state depends on temperature and that pressure affects state changes.
A phase diagram is a graph that shows the relationship between the physical state of a substance and the temperature and pressure of the substance.

54. Phase Diagrams, continued
Chapter 11
Section 4 Phase Equilibrium
Phase Diagrams, continued
A phase diagram has three lines:
One line shows the liquid-gas equilibrium.
Another line shows the liquid-solid equilibrium
A third line shows the solid-gas equilibrium.
The three lines meet at the triple point, the temperature and pressure at which the three states of a substance coexist at equilibrium.

55. Chapter 11
Section 4 Phase Equilibrium
Phase Diagram for H2O

56. Phase Diagrams, continued
Chapter 11
Section 4 Phase Equilibrium
Phase Diagrams, continued
Phase Diagrams Relate State, Temperature, and Pressure
For any given point (x, y) on the phase diagram for water, you can see in what state water will be.
For example, at the coordinates 363 K (x = 90°C) and standard pressure (y = kPa), the point falls in the region labeled “Liquid.”

57. Phase Diagrams, continued
Chapter 11
Section 4 Phase Equilibrium
Phase Diagrams, continued
Gas-Liquid Equilibrium
Line AD shows liquid-vapor equilibrium.
As you follow line AD upwards, the vapor pressure is increasing, so the density of the vapor increases and the density of the liquid decreases.
At a temperature and pressure called the critical point the liquid and vapor phases are identical.
Above this point, the substance is called a supercritical fluid.

58. Phase Diagrams, continued
Chapter 11
Section 4 Phase Equilibrium
Phase Diagrams, continued
Solid-Liquid Equilibrium
Line AC shows solid-liquid equilibrium.
Only solid is present to the left of AC.
Water is an unusual substance: the solid is less dense than the liquid.
If the pressure is increased at point F, at constant temperature, water will melt. Thus, the line AC has a slightly negative slope, which is very rare in phase diagrams of other substances.

59. Phase Diagrams, continued
Chapter 11
Section 4 Phase Equilibrium
Phase Diagrams, continued
Solid-Gas Equilibrium
Line AB shows solid-vapor equilibrium.
If the pressure is decreased below the line AB, the solid will sublime.
Freeze-drying uses the process of sublimation of ice below the freezing point to dry foods.

60. Phase Diagrams, continued
Chapter 11
Section 4 Phase Equilibrium
Phase Diagrams, continued
Phase Diagrams Are Unique to a Particular Substance
Each pure substance has a unique phase diagram, although the general structure is the same.
The triple point is characteristic for each substance and serves to distinguish the substance from other substances.
Unlike boiling point, the melting point is affected little by changes in pressure, so the solid-liquid equilibrium line is nearly vertical in phase diagrams.

61. Chapter 11
Section 4 Phase Equilibrium
Phase Diagram for CO2

62. 1. What is the triple point?
Chapter 11
Section 4 Phase Equilibrium
Warm-Up 2/16/16
1. What is the triple point?
2. What is the change in state called going from a gas to a liquid?
3. What is the mass of a substance that absorbs 2450 joules of energy, with a specific heat of 1.5 J/g*C and a change in temperature of 25 degrees Celsius?
4. A 1.75 g sample at 75 °C is placed in 9.25 g of water. The temperature of the water increased from 22.2 °C to 26.5 ° C. Calculate the specific heat of the sample.
5. How can specific heat be used to identify a substance?

63. Phase Diagrams, continued
Chapter 11
Section 4 Phase Equilibrium
Phase Diagrams, continued
Heating Curve
Copy Heating Curve of Water

64. Chapter 11 Section 4 Continued 2/16/16
Section 4 Phase Equilibrium
Chapter 11
Section 4 Continued
2/16/16

65. Phase Diagrams, continued
Chapter 11
Section 4 Phase Equilibrium
Phase Diagrams, continued
Heating Curves
There are two main observations on the measured curve:
regions where the temperature increases as heat is added
plateaus where the temperature stays constant.

66. Phase Diagrams, continued
Chapter 11
Section 4 Phase Equilibrium
Phase Diagrams, continued
Heating Curves
Notice that, in general, the temperature goes up the longer the heating continues. However, there are two horizontal flat parts to the graph. These happen when there is a change of state. The plateaus are also called phase changes.

67. Phase Diagrams, continued
Chapter 11
Section 4 Phase Equilibrium
Phase Diagrams, continued
Heating Curves
The first change of state is melting (changing from a solid to a liquid). The temperature stays the same while a substance melts. For water, this temperature is 0°C because the melting point for water is 0°C.
The second change of state is boiling (changing from a liquid to a gas). The temperature stays the same while a substance boils. For water, this temperature is 100°C because the boiling point for water is 100°C

68. Phase Diagrams, continued
Chapter 11
Section 4 Phase Equilibrium
Phase Diagrams, continued
Heating Curves
Different substances have different melting points and boiling points, but the shapes of their heating curves are very similar. 
Heating curves show how the temperature changes as a substance is heated up. Cooling curves are the opposite. They show how the temperature changes as a substance is cooled down. Just like heating curves, cooling curves have horizontal flat parts where the state changes from gas to liquid, or from liquid to solid.

69. Identify the process that takes place during line segment BC of the heating curve.

70. Identify the process that takes place during line segment BC of the heating curve.
Melting

71. Heat energy goes into the vibration motion of the molecule, increasing its kinetic energy. Since temperature is the average kinetic energy, the temperature of the solid increases. The rate of temperature increase depends on the heat capacity of the solid.

72. Identify a line segment in which the average kinetic energy is increasing

73. Identify a line segment in which the average kinetic energy is increasing
AB, CD, or EF

74. What state of matter is the substance in at line CD?

75. What state of matter is the substance in at line CD?
Liquid

76. What state of matter is the substance in at line AB?

77. What state of matter is the substance in at line AB?
solid

78.
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