1. Accelerated Chemistry
LIQUIDS AND SOLIDS

2. Chapter Goals Kinetic-Molecular Description of Liquids and Solids
Intermolecular Attractions and Phase Changes
The Liquid State
Viscosity
Surface Tension
Capillary Action
Evaporation
Vapor Pressure
Boiling Points and Distillation
Heat Transfer Involving Liquids

3. Chapter Goals The Solid State Melting Point
Heat Transfer Involving Solids
Sublimation and the Vapor Pressure of Solids
Phase Diagrams (P versus T)
Amorphous Solids and Crystalline Solids
Structures of Crystals
Bonding in Solids
Band Theory of Metals
Synthesis Question

4. Kinetic-Molecular Description of Liquids and Solids
Solids and liquids are condensed states.
The atoms, ions, or molecules in solids and liquids are much closer to one another than in gases.
Solids and liquids are highly incompressible.
Liquids and gases are fluids.
They easily flow.
The intermolecular attractions in liquids and solids are strong.

5. Kinetic-Molecular Description of Liquids and Solids
Schematic representation of the three common states of matter.
gas
liquid
solid
cool
heat

6. Kinetic-Molecular Description of Liquids and Solids
If we compare the strengths of interactions among particles and the degree of ordering of particles, we see that
Gases< Liquids < Solids
Miscible liquids are soluble in each other.
Examples of miscible liquids:
Water dissolves in alcohol.
Gasoline dissolves in motor oil.

7. Kinetic-Molecular Description of Liquids and Solids
Immiscible liquids are insoluble in each other.
Two examples of immiscible liquids:
Water does not dissolve in oil.
Water does not dissolve in cyclohexane.

8. Intermolecular Attractions and Phase Changes
There are four important intermolecular attractions.
This list is from strongest attraction to the weakest attraction.
Ion-ion interactions
The force of attraction between two oppositely charged ions is governed by Coulomb’s law.

9. Intermolecular Attractions and Phase Changes
Coulomb’s law determines:
The melting and boiling points of ionic compounds.
The solubility of ionic compounds.
: Arrange the following ionic compounds in the expected order of increasing melting and boiling points.
NaF, CaO, CaF2

10. Intermolecular Attractions and Phase Changes

11. Intermolecular Attractions and Phase Changes
Hydrogen bonding
Consider H2O a very polar molecule.
2 Factors
Great polarity of bond
Close approach
of dipoles
Effects Boiling Points
Strong H bonds

12. Intermolecular Attractions and Phase Changes
Hydrogen bonding
Consider H2O a very polar molecule.

13. Intermolecular Attractions and Phase Changes

14. Intermolecular Attractions and Phase Changes
Dipole-dipole interactions
Consider BrF a polar molecule.

15. Intermolecular Attractions and Phase Changes
London Forces are very weak.
They are the weakest of the intermolecular forces.
This is the only attractive force in nonpolar molecules.
Consider Ar as an isolated atom.

16. Intermolecular Attractions and Phase Changes
In a group of Ar atoms the temporary dipole in one atom induces other atomic dipoles.

17. Intermolecular Attractions and Phase Changes
Similar effects occur in a group of I2 molecules.
The effect is shown in this movie.

18. The Liquid State Viscosity Viscosity is the resistance to flow.
For example, compare how water pours out of a glass compared to molasses, syrup or honey.
Oil for your car is bought based on this property.
10W30 or 5W30 describes the viscosity of the oil at high and low temperatures.

19. The Liquid State
An example of viscosity of two liquids.
One

20. The Liquid State Surface Tension
Surface tension is a measure of the unequal attractions that occur at the surface of a liquid.
The molecules at the surface are attracted unevenly.

21. The Liquid State
Floating paper clip demonstration of surface tension.
TWO

22. The Liquid State Capillary Action
Capillary action is the ability of a liquid to rise (or fall) in a glass tube or other container

23. The Liquid State
Cohesive forces are the forces that hold liquids together.
Adhesive forces are the forces between a liquid and another surface.
Capillary rise implies that the:
Adhesive forces > cohesive forces
Capillary fall implies that the:
Cohesive forces > adhesive forces

24. Mercury exhibits a capillary fall.
The Liquid State
Water exhibits a capillary rise.
Mercury exhibits a capillary fall.
Water
Mercury

25. The Liquid State
Capillary action also affects the meniscus of liquids.

26. The Liquid State Evaporation
Evaporation is the process in which molecules escape from the surface of a liquid and become a gas.
Evaporation is temperature dependent.
Insert Fig Have it appear (wipe left) after the last sentence.

27. The Liquid State Vapor Pressure
Vapor pressure is the pressure exerted by a liquid’s vapor on its surface at equilibrium.
Vapor Pressure (torr) and boiling point for three liquids at different temperatures.
0oC 20oC 30oC normal boiling point
diethyl ether oC
ethanol oC
water oC
What are the intermolecular forces in each of these compounds?
You do it!

28. Vapor Pressure as a function of temperature.
The Liquid State
Vapor Pressure as a function of temperature.

29. Boiling Points and Distillation
The Liquid State
Boiling Points and Distillation
The boiling point is the temperature at which the liquid’s vapor pressure is equal to the applied pressure.
The normal boiling point is the boiling point when the pressure is exactly 1 atm.
Distillation is a method we use to separate mixtures of liquids based on their differences in boiling points.

30. The Liquid State Distillation
Distillation is a process in which a mixture or solution is separated into its components on the basis of the differences in boiling points of the components.
Distillation is another vapor pressure phenomenon.

31. Heat Transfer Involving Liquids
The Liquid State
Heat Transfer Involving Liquids
From earlier
How much heat is released by 2.00 x 102 g of H2O as it cools from 85.0oC to 40.0oC? The specific heat of water is J/goC.
You have done this

32. The Liquid State

33. The Liquid State
Molar heat capacity is the amount of heat required to raise the temperature of one mole of a substance 1.00 oC.
The molar heat capacity of ethyl alcohol, C2H5OH, is 113 J/moloC. How much heat is required to raise the T of 125 g of ethyl alcohol from 20.0oC to 30.0oC?
1 mol C2H5OH = 46.0 g

34. The Liquid State

35. The Liquid State
The calculations we have done up to now tell us the energy changes as long as the substance remains in a single phase.
Next, we must address the energy associated with phase changes.
For example, solid to liquid or liquid to gas and the reverse.
Heat of Vaporization is the amount of heat required to change 1.00 g of a liquid substance to a gas at constant temperature.
Heat of vaporization has units of J/g.
Heat of Condensation is the reverse of heat of vaporization, phase change from gas to liquid.

36. Molar heat of vaporization or DHvap Molar heat of condensation
The Liquid State
Molar heat of vaporization or DHvap
The DHvap is the amount of heat required to change 1.00 mole of a liquid to a gas at constant temperature.
DHvap has units of J/mol.
Molar heat of condensation
The reverse of molar heat of vaporization is the heat of condensation.

37. The Liquid State

38. STUFF Molar Heat of Fusion ice 6.02 kJ/mol
Molar Heat of Vaporization liquid to vapor kJ/mol
Specific Heat of water
Ice J/g C
Liquid J/g C
Steam 2.0 j/g C

39. The Liquid State
How many joules of energy must be absorbed by 5.00 x 102 g of H2O at 50.0oC to convert it to steam at 120oC? The molar heat of vaporization of water is 40.6 kJ/mol and the molar heat capacities of liquid water and steam are 75.3 J/mol oC and 36.4 J/mol oC, respectively.
You do it!

40. The Liquid State
Next, let’s calculate the energy required to boil the water.
Finally, let’s calculate the heat required
to heat steam from 100 to 120oC.

41. The Liquid State
The total amount of energy for this process is the sum of the 3 pieces we have calculated

42. The Liquid State
If 45.0 g of steam at 140oC is slowly bubbled into 450 g of water at 50.0oC in an insulated container, can all the steam be condensed?

43. The Liquid State

44. The Liquid State
In Denver the normal atmospheric pressure is 630 torr. At what temperature does water boil in Denver?

45. The Liquid State

46. The Liquid State Boiling Points of Various Kinds of Liquids
Gas MW BP(oC)

47. The Liquid State

48. The Liquid State

49. The Liquid State

50. The Liquid State

51. The Liquid State
At the molecular level what happens when a species boils?

52. The Solid State Normal Melting Point
The normal melting point is the temperature at which the solid melts (liquid and solid in equilibrium) at exactly 1.00 atm of pressure.
The melting point increases as the strength of the intermolecular attractions increase.

53. What experimental proof do you have?
The Solid State
Which requires more energy?
What experimental proof do you have?

54. Heat Transfer Involving Solids
Heat of Fusion
Heat of fusion is the amount of heat required to melt one gram of a solid at its melting point at constant temperature.
Heat of crystallization is the reverse of the heat of fusion.

55. Heat Transfer Involving Solids
Molar heat of fusion or Hfusion
The molar heat of fusion is the amount of heat required to melt a mole of a substance at its melting point.
The molar heat of crystallization is the reverse of molar heat of fusion

56. Heat Transfer Involving Solids
Here is a summary of the heats of transformation for water.

57. Heat Transfer Involving Solids
Calculate the amount of heat required to convert g of ice at -10.0oC to water at 40.0oC.
specific heat of ice is 2.09 J/goC

58. Heat Transfer Involving Solids

59. Sublimation and the Vapor Pressure of Solids
In the sublimation process the solid transforms directly to the vapor phase without passing through the liquid phase.
Solid CO2 or “dry” ice does this well.

60. Phase Diagrams (P versus T)
Phase diagrams are a convenient way to display all of the different phase transitions of a substance.
This is the phase diagram for water.

61. Phase Diagrams (P versus T)
Compare water’s phase diagram to carbon dioxide’s phase diagram.

62. Amorphous Solids and Crystalline Solids
Amorphous solids do not have a well ordered molecular structure.
Examples of amorphous solids include waxes, glasses, asphalt.
Crystalline solids have well defined structures that consist of extended array of repeating units called unit cells.
Crystalline solids display X-ray diffraction patterns which reflect the molecular structure.
The Bragg equation, detailed in the textbook, describes how an X-ray diffraction pattern can be used to determine the interatomic distances in crystals.

63. Structure of Crystals
Unit cells are the smallest repeating unit of a crystal.
As an analogy, bricks are repeating units for buildings.
There are seven basic crystal systems.

64. Structure of Crystals
We shall look at the three variations of the cubic crystal system.
Simple cubic unit cells.
The balls represent the positions of atoms, ions, or molecules in a simple cubic unit cell.

65. Structure of Crystals
In a simple cubic unit cell each atom, ion, or molecule at a corner is shared by 8 unit cells
Thus 1 unit cell contains 8(1/8) = 1 atom, ion, or molecule.

66. Structure of Crystals
Body centered cubic (bcc) has an additional atom, ion, or molecule in the center of the unit cell.
On a body centered cubic unit cell there are 8 corners + 1 particle in center of cell.
1 bcc unit cell
contains 8(1/8) + 1 = 2 particles.

67. Structure of Crystals
A face centered cubic (fcc) unit cell has a cubic unit cell structure with an extra atom, ion, or molecule in each face.

68. Structure of Crystals
A face centered cubic unit cell has 8 corners and 6 faces.
1 fcc unit cell contains
8(1/8) + 6(1/2) = 4 particles.

69. Bonding in Solids
Molecular Solids have molecules in each of the positions of the unit cell.
Molecular solids have low melting points, are volatile, and are electrical insulators.
Examples of molecular solids inlude:
water, sugar, carbon dioxide, benzene

70. Bonding in Solids
Covalent Solids have atoms that are covalently bonded to one another
Some examples of covalent solids are:
Diamond, graphite, SiO2 (sand), SiC

71. Bonding in Solids
Ionic Solids have ions that occupy the positions in the unit cell.
Examples of ionic solids include:
CsCl, NaCl, ZnS

72. Bonding in Solids
Metallic Solids may be thought of as positively charged nuclei surrounded by a sea of electrons.
The positive ions occupy the crystal lattice positions.
Examples of metallic solids include:
Na, Li, Au, Ag, ……..

73. Bonding in Solids Compound Melting Point (oC) ice 0.0 ammonia -77.7
Variations in Melting Points for Molecular Solids
What are the intermolecular forces in each solid?
Compound Melting Point (oC)
ice
ammonia
benzene, C6H
napthalene, C10H
benzoic acid, C6H5CO2H

74. Bonding in Solids Substance Melting Point (oC) sand, SiO2 1713
Variations in Melting Points for Covalent Solids
Substance Melting Point (oC)
sand, SiO
carborundum, SiC ~2700
diamond >3550
graphite

75. Bonding in Solids Variations in Melting Points for Ionic Solids
Compound Melting Point (oC)
LiF
LiCl
LiBr
LiI
CaF
CaCl
CaBr
CaI

76. Bonding in Solids Metal Melting Point (oC)
Variations in Melting Points for Metallic Solids
Metal Melting Point (oC) Na Pb Al Cu Fe W

77. Bonding in Solids
A group IVA element with a density of g/cm3 crystallizes in a face-centered cubic lattice whose unit cell edge length is 4.95 Å. Calculate the element’s atomic weight. What is the atomic radius of this element?

78. Bonding in Solids
Face centered cubic unit cells have 4 atoms, ions, or molecules per unit cell.
Problem solution pathway:
Determine the volume of a single unit cell.
Use the density to determine the mass of a single unit cell.
Determine the mass of one atom in a unit cell.
Determine the mass of 1 mole of these atoms

79. Use the density to determine the mass of a unit cell.
Bonding in Solids
Determine the volume of a single unit cell.
Use the density to determine the mass of a unit cell.

80. Bonding in Solids Determine the mass of one atom in the unit cell.
Determine the mass of one mole of these atoms.

81. Bonding in Solids
To determine an atomic radius requires some geometry.
For simple cubic unit cells:
The edge length = 2 radii

82. Bonding in Solids For face-centered cubic unit cells:
The face diagonal = 2 x edge length.
The diagonal length = 4 radii.

83. Bonding in Solids For body-centered cubic unit cells:
The body diagonal = 3 x edge length.
The diagonal length = 4 radii.

84. Bonding in Solids
Determine the diagonal length then divide by 4 to get the atomic radius.

85. Band Theory of Metals
Sodium’s 3s orbitals can interact to produce overlapping orbitals

86. Band Theory of Metals
The 3s orbitals can also overlap with unfilled 3p orbitals

87. Band Theory of Metals
Insulators have a large gap between the s and p bands.
Gap is called the forbidden zone.
Semiconductors have a small gap between the bands.

88. Synthesis Question
Maxwell House Coffee Company decaffeinates its coffee beans using an extractor that is 7.0 feet in diameter and 70.0 feet long. Supercritical carbon dioxide at a pressure of atm and temperature of 100.0oC is passed through the stainless steel extractor. The extraction vessel contains 100,000 pounds of coffee beans soaked in water until they have a water content of 50%.

89. Synthesis Question
This process removes 90% of the caffeine in a single pass of the beans through the extractor. Carbon dioxide that has passed over the coffee is then directed into a water column that washes the caffeine from the supercritical CO2. How many moles of carbon dioxide are present in the extractor?

90. Synthesis Question

91. Synthesis Question

92. Group Question
How many CO2 molecules are there in 1.0 cm3 of the Maxwell House Coffee Company extractor? How many more CO2 molecules are there in a cm3 of the supercritical fluid in the Maxwell House extractor than in a mole of CO2 at STP?