1. Chapter 11 Intermolecular Forces, Liquids, and Solids
Chemistry, The Central Science, 11th edition
Theodore L. Brown, H. Eugene LeMay, Jr., and Bruce E. Bursten
Chapter 11 Intermolecular Forces, Liquids, and Solids
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2. Overview of Chapter 11 – Intermolecular Forces
Review/On your own:
compare liquids and solids with gases in the context of the KMT
describe the three intermolecular forces and apply this knowledge to predictions about substances’ properties, such as melting and boiling points, solubilities, surface tension, etc.
explain what vapor pressure is, and how it differs between contained and uncontained containers
interpret heating curves, vapor pressure graphs, and phase diagrams
calculate enthalpy changes associated with phase changes
relate vapor pressure to volatility of a liquid and the intermolecular forces
interpret vapor pressure graphs
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3. Overview of Chapter 11 – Intermolecular Forces
More depth:
interpret phase diagrams
describe the three intermolecular forces and apply this knowledge to predictions about substances’ properties, such as melting and boiling points, solubilities, surface tension, etc.
NOTHING ON SOLIDS (11.7 and 11.8)
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4. States of Matter Shape Volume Density Compressible? Mass? Solid Liquid
Gas
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5. States of Matter
The fundamental difference between states of matter is the distance between particles.
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6. States of Matter
Because in the solid and liquid states particles are closer together, we refer to them as condensed phases.
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7. 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.
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8. Intermolecular Forces
The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together.
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9. Intermolecular Forces
They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.
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10. Intermolecular Forces
These intermolecular forces as a group are referred to as van der Waals forces.
Note: Sometimes scientists use the term van der Waals force to refer to dispersion forces only. Be sure to describe the interaction as thoroughly as possible to clearly distinguish amongst them.
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11. van der Waals Forces
London dispersion forces (weakest, present in all molecules)
Dipole-dipole interactions
Hydrogen bonding
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12. Ion-Dipole Interactions
Ion-dipole interactions (a fourth type of force), are important in solutions of ions.
The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.
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13. 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.
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14. Dipole-Dipole Interactions
The more polar the molecule, the higher is its boiling point.
Note: There is a stronger relationship between the strength of IMFs and boiling points, more so than with melting points.
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15. 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.
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16. 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.
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17. 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.
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18. London Dispersion Forces
London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.
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19. 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.
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20. 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.
Note the differences in boiling points
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21. Factors Affecting London Forces
The strength of dispersion forces tends to increase with increased molecular weight.
BECAUSE larger atoms have larger electron clouds which are easier to polarize.
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22. Which Have a Greater Effect
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.
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23. 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.
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24. 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.
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25. 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, small and can approach an electronegative atom very closely, making it much easier to interact with it.
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26. Sample Exercise 11.1 (p. 450)
In which of the following substances is hydrogen bonding likely to play an important role in determining physical properties: methane (CH4), hydrazine (H2NNH2), methyl fluoride (CH3F), or hydrogen sulfide (H2S)?
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27. Sample Exercise 11.1 (p. 450)
In which of the following substances is hydrogen bonding likely to play an important role in determining physical properties: methane (CH4), hydrazine (H2NNH2), methyl fluoride (CH3F), or hydrogen sulfide (H2S)?
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28. Practice Exercise 1 (11.1)
Which of the following substances is most likely to be a liquid at room temperature?
Formaldehyde, H2CO
Fluoromethane, CH3F
Hydrogen cyanide, HCN
Hydrogen peroxide, H2O2
Hydrogen sulfide, H2S
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29. Practice Exercise 1 (11.1)
Which of the following substances is most likely to be a liquid at room temperature?
Formaldehyde, H2CO
Fluoromethane, CH3F
Hydrogen cyanide, HCN
Hydrogen peroxide, H2O2
Hydrogen sulfide, H2S
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30. Practice Exercise 2 (11.1)
In which of the following substances is significant hydrogen bonding possible: methylene chloride (CH2Cl2), phosphine (PH3), hydrogen peroxide (HOOH), or acetone (CH3COCH3)?
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31. Practice Exercise 2 (11.1)
In which of the following substances is significant hydrogen bonding possible: methylene chloride (CH2Cl2), phosphine (PH3), hydrogen peroxide (HOOH), or acetone (CH3COCH3)?
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32. Summarizing Intermolecular Forces
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33. Sample Exercise 11.2 (p. 454)
List the substances BaCl2, H2, CO, HF, and Ne in order of increasing boiling points.
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34. Sample Exercise 11.2 (p. 454)
List the substances BaCl2, H2, CO, HF, and Ne in order of increasing boiling points.
H2 < Ne < CO < HF < BaCl2
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35. Practice Exercise 1 (11.2)
List the substances Ar, Cl2, CH4, and CH3COOH in order of increasing strength of intermolecular attractions.
CH4 < Ar < CH3COOH < Cl2
Cl2 < CH3COOH < Ar < CH4
CH4 < Ar < Cl2 < CH3COOH
CH3COOH < Cl2 < Ar < CH4
Ar < Cl2 < CH4 < CH3COOH
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36. Practice Exercise 1 (11.2)
List the substances Ar, Cl2, CH4, and CH3COOH in order of increasing strength of intermolecular attractions.
CH4 < Ar < CH3COOH < Cl2
Cl2 < CH3COOH < Ar < CH4
CH4 < Ar < Cl2 < CH3COOH
CH3COOH < Cl2 < Ar < CH4
Ar < Cl2 < CH4 < CH3COOH
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37. Practice Exercise 2 (11.2)
Identify the intermolecular forces present in the following substances, and
b) select the substance with the highest boiling point:
CH3CH3, CH3OH, CH3CH2OH.
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38. Practice Exercise 2 (11.2)
Identify the intermolecular forces present in the following substances,
CH3CH3, Only dispersion
CH3OH, CH3CH2OH dispersion and H-bonds
b) select the substance with the highest boiling point:
CH3CH2OH
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39. Intermolecular Forces Affect Many Physical Properties
The strength of the attractions between particles can greatly affect the properties of a substance or solution.
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40. 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.
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41. Surface Tension
Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.
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42. Phase Changes
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43. Energy Changes Associated with Changes of State
The heat of fusion is the energy required to change a solid at its melting point to a liquid.
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44. Energy Changes Associated with Changes of State
The heat of vaporization is defined as the energy required to change a liquid at its boiling point to a gas.
Note how much more energy is needed to vaporize a substance than to melt it.
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45. 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 a phase change.
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46. Sample Exercise 11.3 (p. 460)
Calculate the enthalpy change upon converting 1.00 mol of ice at -25oC to water vapor (steam) at 125oC under a constant pressure of 1 atm. The specific heats of ice, water, and steam are 2.03 J/g-K, 4.18 J/g-K, and 1.84 J/g-K, respectively. For H2O, DHfus = 6.01 kJ/mol, and DHvap = kJ/mol.
Note: multi-step problem, tedious, but not hard
(56.0 kJ)
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47. Sample Exercise 11.3 (p. 460) 0.91 kJ 6.01 kJ  melting 7.52 kJ
40.7 kJ  vaporization
0.83 kJ
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48. Practice Exercise 1 (11.3)
What information about water is needed to calculate the enthalpy change for converting 1 mol H2O(g) at 100oC to H2O(l) at 80oC?
Heat of fusion
Heat of vaporization
Heat of vaporization and specific heat of H2O(g)
Heat of vaporization and specific heat of H2O(l)
Heat of fusion and specific heat of H2O(l)
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49. Practice Exercise 1 (11.3)
What information about water is needed to calculate the enthalpy change for converting 1 mol H2O(g) at 100oC to H2O(l) at 80oC?
Heat of fusion
Heat of vaporization
Heat of vaporization and specific heat of H2O(g)
Heat of vaporization and specific heat of H2O(l)
Heat of fusion and specific heat of H2O(l)
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50. Practice Exercise 2 (11.3)
What is the enthalpy change during the process in which g of water at 50.0oC is cooled to ice at -30.0oC?
(Use the specific heats and enthalpies for phase changes given in Sample Exercise 11.4.)
(-60.4 kJ)
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51. 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.
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52. Vapor Pressure
As more molecules escape the liquid, the pressure they exert increases.
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53. Vapor Pressure
The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate.
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54. 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.
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55. Sample Exercise 11.4 (p. 464)
Use this figure to estimate the boiling point of diethyl ether under an external pressure of 0.80 atm.
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56. Sample Exercise 11.4 (p. 464)
Use this figure to estimate the boiling point of diethyl ether under an external pressure of 0.80 atm.
27oC
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57. Practice Exercise 1 (11.4)
In the mountains, water in an open container will boil when
its critical temperature exceeds room temperature
its vapor pressure equals atmospheric pressure
its temperature is 100oC
enough energy is supplied to break covalent bonds
none of these is correct
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58. Practice Exercise 1 (11.4)
In the mountains, water in an open container will boil when
its critical temperature exceeds room temperature
its vapor pressure equals atmospheric pressure
its temperature is 100oC
enough energy is supplied to break covalent bonds
none of these is correct
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59. Practice Exercise 2 (11.4)
At what external pressure will ethanol have a boiling point of 60oC?
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60. Practice Exercise 2 (11.4)
At what external pressure will ethanol have a boiling point of 60oC?
about 340 torr (0.45 atm)
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61. Phase Diagrams
Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.
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62. Phase Diagrams The circled line is the liquid-vapor interface.
It starts at the triple point (T), the point at which all three states are in equilibrium.
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63. Phase Diagrams
It ends at the critical point (C); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.
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64. Phase Diagrams
Each point along this line is the boiling point of the substance at that pressure.
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65. Phase Diagrams
The circled line in the diagram below is the interface between liquid and solid.
The melting point at each pressure can be found along this line.
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66. Phase Diagrams
Below the triple point the substance cannot exist in the liquid state.
Along the circled line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line.
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67. 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.
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68. 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.
Why?
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69. Phase Diagram of Carbon Dioxide
Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 sublimes at normal pressures.
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70. Phase Diagram of Carbon Dioxide
The low critical temperature and critical pressure for CO2 make supercritical CO2 a good solvent for extracting nonpolar substances (like caffeine)
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71. Phase Diagram for Methane
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72. Sample Exercise 11.5 (p. 466)
Use the phase diagram for methane, CH4, shown in Figure to answer the following questions.
What are the approximate temperature and pressure of the critical point?
What are the approximate temperature and pressure of the triple point?
Is methane a solid, liquid, or gas at 1 atm and 0oC?
If solid methane at 1 atm is heated while the pressure is held constant, will it melt or sublime?
If methane at 1 atm and 0oC is compressed until a phase change occurs, in which state is the methane when the compression is complete?
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73. Sample Exercise 11.5 (p. 466)
Use the phase diagram for methane, CH4, shown in Figure to answer the following questions.
What are the approximate temperature and pressure of the critical point?
What are the approximate temperature and pressure of the triple point?
Is methane a solid, liquid, or gas at 1 atm and 0oC?
If solid methane at 1 atm is heated while the pressure is held constant, will it melt or sublime?
If methane at 1 atm and 0oC is compressed until a phase change occurs, in which state is the methane when the compression is complete?
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74. Practice Exercise 1 (11.5)
Based on the phase diagram for methane (Figure 11.30), what happens to methane as it is heated from -250 to 0oC at a pressure of 10-2 atm?
It sublimes at about -200oC.
It melts at about -200oC.
It boils at about -200oC.
It condenses at about -200oC.
It reaches the triple point at about -200oC.
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75. Practice Exercise 1 (11.5)
Based on the phase diagram for methane (Figure 11.30), what happens to methane as it is heated from -250 to 0oC at a pressure of 10-2 atm?
It sublimes at about -200oC.
It melts at about -200oC.
It boils at about -200oC.
It condenses at about -200oC.
It reaches the triple point at about -200oC.
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76. Practice Exercise 2 (11.5)
Using the phase diagram of methane to answer the following questions.
What is the normal boiling point of methane?
Over what pressure range does solid methane sublime?
Above what temperature does liquid methane not exist?
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77. Practice Exercise 2 (11.5)
Using the phase diagram of methane to answer the following questions.
What is the normal boiling point of methane?
Over what pressure range does solid methane sublime?
Above what temperature does liquid methane not exist?
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78. Solids We can think of solids as falling into two groups:
crystalline, in which particles are in highly ordered arrangement.
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79. Solids We can think of solids as falling into two groups:
amorphous, in which there is no particular order in the arrangement of particles.
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80. Attractions in Ionic Crystals
In ionic crystals, ions pack themselves so as to maximize the attractions and minimize repulsions between the ions.
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81. Crystalline Solids
Because of the ordered in a crystal, we can focus on the repeating pattern of arrangement called the unit cell.
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82. Crystalline Solids
There are several types of basic arrangements in crystals, like the ones depicted above.
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83. Crystalline Solids
We can determine the empirical formula of an ionic solid by determining how many ions of each element fall within the unit cell.
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84. Ionic Solids What are the empirical formulas for these compounds?
(a) Green: chlorine; Gray: cesium
(b) Yellow: sulfur; Gray: zinc
(c) Gray: calcium; Blue: fluorine
(a)
(b)
(c)
CsCl
ZnS
CaF2
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85. Types of Bonding in Crystalline Solids
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86. 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.
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87. 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.
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88. 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.
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89. Sample Integrative Exercise (p. 470)
The substance CS2 has a melting point of oC and a boiling
point of 46.3oC. Its density at 20oC is 1.26 g/cm3. It is highly
flammable.
What is the name of this compound?
List the intermolecular forces that CS2 molecules would have
with each other.
Write a balanced equation for the combustion of this
compound in air. (You will have to decide on the most likely
oxidation products.)
The critical temperature and pressure for CS2 are 552 K and
78 atm, respectively. Compare these values with those for
CO2 and discuss the possible origins of the differences.
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90. Sample Integrative Exercise (p. 470)
The substance CS2 has a melting point of oC and a boiling
point of 46.3oC. Its density at 20oC is 1.26 g/cm3. It is highly
flammable.
What is the name of this compound? Carbon disulfide
List the intermolecular forces that CS2 molecules would have
with each other. London dispersion forces
Write a balanced equation for the combustion of this
compound in air. (You will have to decide on the most likely
oxidation products.) most likely products will be SO2 and CO2
The critical temperature and pressure for CS2 are 552 K and
78 atm, respectively. Compare these values with those for
CO2 and discuss the possible origins of the differences.
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91. Sample Integrative Exercise (p. 470)
The critical temperature and pressure for CS2 are 552 K and
78 atm, respectively. Compare these values with those for
CO2 and discuss the possible origins of the differences.
The difference in critical temperatures is especially notable. The higher values for CS2 arise from the greater London dispersion attractions between the CS2 molecules compared with CO2. These greater attractions are due to the larger size of the sulfur compared to oxygen and therefore its greater polarizability.
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