1. Chapter 9 Intermolecular Forces, Liquids, and Solids

2. States of Matter
The fundamental difference between states of matter is the distance between particles.

3. States of Matter
Because in the solid and liquid states particles are closer together, we refer to them as condensed phases.

4. The States of Matter
The state a substance is in at a particular temperature and pressure depends on two antagonistic entities:
The kinetic energy of the particles
The strength of the attractions between the particles

5. Intermolecular Forces
The attractions between molecules (intermolecular) are not nearly as strong as the intramolecular attractions that hold compounds together.

6. Intermolecular Forces
They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.

7. Intermolecular Forces
These intermolecular forces as a group are referred to as van der Waals forces.

8. Van der Waals Forces London dispersion forces
Dipole-dipole interactions
Hydrogen bonding

9. Ion-Dipole Interactions
A fourth type of force, ion-dipole interactions are an important force in solutions of ions.
The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents!

10. London Dispersion Forces
While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom.

11. London Dispersion Forces
At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.

12. London Dispersion Forces
Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1.

13. London Dispersion Forces
London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.

14. London Dispersion Forces
These forces are present in all molecules, whether they are polar or nonpolar.
The tendency of an electron cloud to distort in this way is called polarizability.

15. Factors Affecting London Forces
The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane).
This is due to the increased surface area in n-pentane and greater polarizability.

16. Factors Affecting London Forces
The strength of dispersion forces tends to increase with increased molecular weight or increasing number of electrons.
Larger atoms have larger electron clouds, which are easier to polarize.

17. Dipole-Dipole Interactions
Molecules that have permanent dipoles are attracted to each other.
The positive end of one is attracted to the negative end of the other and vice-versa.
These forces are only important when the molecules are close to each other.

18. Dipole-Dipole Interactions
The more polar the molecule, the higher is its boiling point.

19. Which Have a Greater Effect: Dipole-Dipole Interactions or Dispersion Forces?
If two molecules are of comparable size and shape, dipole-dipole interactions will likely be the dominating force.
If one molecule is much larger than another, dispersion forces will likely determine its physical properties.

20. How Do We Explain This?
The nonpolar series (SnH4 to CH4) follow the expected trend.
The polar series follows the trend from H2Te through H2S, but water is quite an anomaly.

21. Hydrogen Bonding
The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong!
We call these interactions hydrogen bonds.

22. Hydrogen Bonding
Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine.
Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

23. Summarizing Intermolecular Forces

26. Intermolecular Forces Affect Many Physical Properties
The strength of the attractions between particles can greatly affect the properties of a substance or solution.

27. occurs across the surface and below the surface
The molecular basis of surface tension.
the net vector
for attractive
forces is downward
hydrogen bonding
occurs across the surface
and below the surface
hydrogen bonding
occurs in three
dimensions

28. Viscosity Resistance of a liquid to flow is called viscosity.
It is related to the ease with which molecules can move past each other.
Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

29. Surface Tension
Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

30. stronger cohesive forces
Shape of water or mercury meniscus in glass.
capillarity
stronger cohesive forces
adhesive forces
H2O Hg

31. Phase Changes

32. Energy Changes Associated with Changes of State
Heat of Fusion: Energy required to change a solid at its melting point to a liquid.

33. Energy Changes Associated with Changes of State
Heat of Vaporization: Energy required to change a liquid at its boiling point to a gas.

34. Energy Changes Associated with Changes of State
The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other.
The temperature of the substance does not rise during the phase change.

35. Vapor Pressure
At any temperature, some molecules in a liquid have enough energy to escape.
As the temperature rises, the fraction of molecules that have enough energy to escape increases.

36. Vapor Pressure
The effect of temperature on the distribution of molecular speed in a liquid.

37. Vapor Pressure
As more molecules escape the liquid, the pressure they exert increases.

38. Liquid-gas equilibrium.
The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate.

39. Vapor Pressure
The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate.
kEtOH=K[EtOH]= Ptequil

40. Vapor Pressure
The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure.
The normal boiling point is the temperature at which its vapor pressure is 760 torr.

41. A linear plot of vapor pressure- temperature relationship.
Vapor pressure as a function of temperature and intermolecular forces.

42. The Clausius-Clapeyron Equation

43. Phase Diagrams
Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.

44. Phase Diagrams The AB line is the liquid-vapor interface.
It starts at the triple point (A), the point at which all three states are in equilibrium.

45. Phase Diagrams
It ends at the critical point (B); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.

46. Phase Diagrams
Each point along this line is the boiling point of the substance at that pressure.

47. Phase Diagrams The AD line is the interface between liquid and solid.
The melting point at each pressure can be found along this line.

48. Phase Diagrams Below A the substance cannot exist in the liquid state.
Along the AC line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line.

49. Phase Diagram of Water
Note the high critical temperature and critical pressure:
These are due to the strong van der Waals forces between water molecules.

50. Phase Diagram of Water The slope of the solid–liquid line is negative.
This means that as the pressure is increased at a temperature just below the melting point, water goes from a solid to a liquid.

51. Phase Diagram of Carbon Dioxide
Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 sublimes at normal pressures.

52. Phase Diagram of Carbon Dioxide
The low critical temperature and critical pressure for CO2 make supercritical CO2 a good solvent for extracting nonpolar substances (such as caffeine).

53. CO2 as a critical fluid!

54. Iodine subliming. test tube with ice iodine solid iodine vapor

55. Phase diagrams for CO2 and H2O.

56. Solids We can think of solids as falling into two groups:
Crystalline—particles are in highly ordered arrangement.

57. Solids
Amorphous—NO regular geometric order in the arrangement of particles at the molecular level.

58. Crystalline Solids
Because of the regular order in a crystalline solid, we can focus on the three dimensional repeating pattern or arrangement called the unit cell.

59. Crystalline Solids
There are several types of basic arrangements in crystals. In the cubic class we have primitive, body-centered, and face-centered unit cells.

60. Crystalline Solids
We can determine the formula of an ionic solid by determining how many ions of each element fall within the unit cell.

61. Characteristics of the Major Types of Crystalline Solids
Type of
Solids
Interparticle Forces
Physical Behavior
Particles
Examples (mp,0C)
Atomic
Atoms
Dispersion
Soft, very low mp, poor thermal & electrical conductors
Group 8A(18)
[Ne-249 to Rn-71]
Molecular
Molecules
Dispersion, dipole-dipole, H bonds
Fairly soft, low to moderate mp, poor thermal & electrical conductors
Nonpolar - O2[-219], C4H10[-138], Cl2
[-101], C6H14[-95]
Polar - SO2[-73], CHCl3[-64], HNO3[-42], H2O[0.0]
Ionic
Positive & negative ions
Ion-ion
attraction
Hard & brittle, high mp, good thermal & electrical conductors when molten
NaCl [801]
CaF2 [1423]
MgO [2852]
Metallic
Atoms
Metallic bond
Soft to hard, low to very high mp, excellent thermal and electrical conductors, malleable and ductile
Na [97.8]
Zn [420]
Fe [1535]
Network
Atoms
Covalent bond
Very hard, very high mp, usually poor thermal and electrical conductors

62. Closest Packing of Identical Non-penetrating Spheres
layer a
layer b
hexagonal closest packing
cubic closest packing
layer a
layer c
closest packing of first and second layers
abab… (74%)
hexagonal unit cell
abcabc… (74%)
expanded side views
face-centered unit cell

63. Ionic Solids What are the formulas for these compounds?
In ionic crystals, ions pack themselves so as to maximize the attractions and minimize repulsions between the ions.
What are the formulas for these compounds?
(a) Green: chlorine; Gray: cesium
(b) Yellow: sulfur; Gray: zinc
(c) Green: calcium; Gray: fluorine
(a)
(b)
(c)
CsCl
ZnS
CaF2

64. Covalent-Network and Molecular Solids
Diamonds are an example of a covalent-network solid in which atoms are covalently bonded to each other.
They tend to be hard and have high melting points.

65. Covalent-Network and Molecular Solids
Graphite is an example of a molecular solid in which atoms are held together with van der Waals forces.
They tend to be softer and have lower melting points.

66. Metallic Solids
Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces.
In metals, valence electrons are delocalized throughout the solid.

67. The band of molecular orbitals in lithium metal.

68. Electrical conductivity in a conductor, semiconductor, and insulator.

69. Crystal structures and band representations of doped semiconductors.

70. The levitating power of a superconducting oxide.
rare earth magnet
superconducting ceramic disk
liquid nitrogen

71. Structures of two typical molecules that form liquid crystals.

72. Some Uses of Modern Ceramics and Ceramic Mixtures
Applications
SiC, Si3N4, TiB2, Al2O3
Whiskers(fibers) to strength Al and other ceramics
Si3N4
Car engine parts; turbine rotors for “turbo” cars; electronic sensor units
Si3N4, BN, Al2O3
Supports or layering materials(as insulators) in electronic microchips
SiC, Si3N4, TiB2, ZrO2, Al2O3, BN
Cutting tools, edge sharpeners(as coatings and whole devices), scissors, surgical tools, industrial “diamond”
BN, SiC
Armor-plating reinforcement fibers(as in Kevlar composites)
ZrO2, Al2O3
Surgical implants(hip and knee joints)

73. Unit cells of some modern ceramic materials.
YBa2Cu3O7
SiC
silicon carbide BN
cubic boron nitride (borazon)
high temperature superconductor

74. Tools of the Laboratory
Diffraction of x-rays by crystal planes.

75. Tools of the Laboratory
Formation of an x-ray diffraction pattern of the
protein hemoglobin.

76. Tools of the Laboratory
Scanning tunneling micrographs.
gallium arsenide semiconductor
metallic gold

77. Problems

78. SAMPLE PROBLEM 1
Using the Clausius-Clapeyron Equation
PROBLEM:
The vapor pressure of ethanol is 115 torr at 34.90C. If DHvap of ethanol is 40.5 kJ/mol, calculate the temperature (in 0C) when the vapor pressure is 760 torr.
PLAN:
We are given 4 of the 5 variables in the Clausius-Clapeyron equation. Substitute and solve for T2.
SOLUTION:
34.90C = 308.0K 1 T2
308K - ln
760 torr
115 torr
-40.5 x103 J/mol
8.314 J/mol*K =
T2 = 350K = 770C

79. SAMPLE PROBLEM 2
Drawing Hydrogen Bonds Between Molecules
of a Substance
PROBLEM:
Which of the following substances exhibits H bonding? For those that do, draw two molecules of the substance with the H bonds between them.
(c)
(a)
(b)
PLAN:
Find molecules in which H is bonded to N, O or F. Draw H bonds in the format -B: H-A-.
SOLUTION:
(a) C2H6 has no H bonding sites.
(b)
(c)

80. SAMPLE PROBLEM 3
Predicting the Type and Relative Strength of
Intermolecular Forces
PROBLEM:
For each pair of substances, identify the dominant intermolecular forces in each substance, and select the substance with the higher boiling point.
(a) MgCl2 or PCl3
(b) CH3NH2 or CH3F
(c) CH3OH or CH3CH2OH
(d) Hexane (CH3CH2CH2CH2CH2CH3)
or 2,2-dimethylbutane
PLAN:
Use the formula, structure and Table 2.2 (button).
Bonding forces are stronger than nonbonding(intermolecular) forces.
Hydrogen bonding is a strong type of dipole-dipole force.
Dispersion forces are decisive when the difference is molar mass or molecular shape.

81. SAMPLE PROBLEM 3
Predicting the Type and Relative Strength of
Intermolecular Forces
continued
SOLUTION:
(a) Mg2+ and Cl- are held together by ionic bonds while PCl3 is covalently bonded and the molecules are held together by dipole-dipole interactions. Ionic bonds are stronger than dipole interactions and so MgCl2 has the higher boiling point.
(b) CH3NH2 and CH3F are both covalent compounds and have bonds which are polar. The dipole in CH3NH2 can H bond while that in CH3F cannot. Therefore CH3NH2 has the stronger interactions and the higher boiling point.
(c) Both CH3OH and CH3CH2OH can H bond but CH3CH2OH has more CH for more dispersion force interaction. Therefore CH3CH2OH has the higher boiling point.
(d) Hexane and 2,2-dimethylbutane are both nonpolar with only dispersion forces to hold the molecules together. Hexane has the larger surface area, thereby the greater dispersion forces and the higher boiling point.

82. SAMPLE PROBLEM 4
Determining Atomic Radius from Crystal Structure
PROBLEM:
Barium is the largest nonradioactive alkaline earth metal. It has a body-centered cubic unit cell and a density of 3.62 g/cm3. What is the atomic radius of barium?
(Volume of a sphere: V = 4/3pr3)
PLAN:
We can use the density and molar mass to find the volume of 1 mol of Ba. Since 68%(for a body-centered cubic) of the unit cell contains atomic material, dividing by Avogadro’s number will give us the volume of one atom of Ba. Using the volume of a sphere, the radius can be calculated.
density of Ba (g/cm3)
radius of a Ba atom
reciprocal divided by M
V = 4/3pr3
volume of 1 mol Ba metal
volume of 1 Ba atom
multiply by packing efficiency
volume of 1 mol Ba atoms
divide by Avogadro’s number

83. SAMPLE PROBLEM 4
Determining Atomic Radius from Crystal Structure
continued
SOLUTION:
1 cm3
3.62 g x
137.3 g Ba
mol Ba
Volume of Ba metal =
= 37.9 cm3/mol Ba
37.9 cm3/mol Ba
x 0.68
= 26 cm3/mol Ba atoms
26 cm3
mol Ba atoms x
mol Ba atoms
6.022x1023 atoms
= 4.3x10-23 cm3/atom
r3 = 3V/4p
= 2.2 x 10-8cm