1. Intermolecular Forces and
Liquids and Solids
Chapter 11
Copyright © The McGraw-Hill Companies, Inc.  Permission required for reproduction or display.

2. A phase is a homogeneous part of the system in contact with other parts of the system but separated from them by a well-defined boundary.
2 Phases
Solid phase - ice
Liquid phase - water
11.1

3. Kinetic-Molecular Description
Intermolecular forces are attractive forces between molecules
Intramolecular forces hold atoms together in a molecule. (covalent bond)
Intermolecular vs Intramolecular
41 kJ to vaporize 1 mole of water (inter)
930 kJ to break all O-H bonds in 1 mole of water (intra)
“Measure” of intermolecular force
boiling point
melting point
DHvap
DHfus
DHsub
Generally, intermolecular forces are much weaker than intramolecular forces.(only about 15% as strong)
11.2

4. Types of Intermolecular Forces (IMF)
• Should actually be called Interparticulate Forces (molecules, ions, and/or atoms)
• Ion - ion forces
• Ion-dipole forces
• Dipole-dipole forces
• Dispersion Forces
• Hydrogen Bonds
Ion - ion forces: (lattice energy-ionic compound)
• Remember Chapter 9
• Force depends on the charge on the ions and
the distance separating the ions
• (STRONG FORCES)
E = k
Q+Q- r

5. Intermolecular Forces
1. Ion-Dipole Forces
Attractive forces between an ion and a polar molecule
Example: ions in solution
Ion-Dipole Interaction
11.2

6. Higher charge, smaller size strong interaction
The strength of the interaction depends on the charge and size of the ion and on the magnitude of the dipole moment and size of the molecule.
Water molecules
Higher charge, smaller size strong interaction
11.2

7. Van der Waals Forces
• Dipole-Dipole
• Temporary dipoles (Dispersion forces)
– Ion induced
– Dipole induced
– Instantaneous dipole (London Dispersion forces)
• Hydrogen bonds

8. Intermolecular Forces
2. Dipole-Dipole Forces (permanent dipole moment)
Attractive forces between polar molecules
Orientation of Polar Molecules in a Solid
11.2

9. Intermolecular Forces
3. Temporary dipoles (Dispersion forces)
What attractive force occurs in nonpolar substances?
Attractive forces that arise as a result of temporary dipoles induced in atoms or molecules
– Ion induced
– Dipole induced
– Instantaneous dipole
ion-induced dipole interaction
The electron distribution of atom is distorted by the force exerted by the ions or polar molecules.
dipole-induced dipole interaction
11.2

10. Polarizability increases with: greater number of electrons
The likelihood of a dipole moment being induced depends not only on the charge of the ion or the strength of the ddipole but also on the polarizability of the atom or molecules.
Polarizability is the ease with which the electron distribution in the atom or molecule can be distorted.
Polarizability increases with:
greater number of electrons
more diffuse electron cloud
Dispersion forces usually increase with molar mass (more electrons), or size of the atom.
Melting point increases as the number of electrons in the molecule increase.
11.2

11. London Dispersion Force (Dispersion force)
Polarizability allows gases containing atoms or nonpolar molecules (e.g.He,N2) to condense.
Induced dipoles interacting with each other. This type of interaction produces dispersion forces, which arise as a result of temporary dipoles induced in atoms or molecules, is responsible for the condensation of nonpolar gases.
**Dispersion force exists between all species.
At any instant it is likely that the atom has a dipole momnet created by a specific positions of electrons. This dipole moment is called instantaneous dipole.

12. What type(s) of intermolecular forces exist between each of the following molecules?
HBr
HBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules.
CH4
CH4 is nonpolar: dispersion forces. S O
SO2
SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules.
11.2

13. Worked Example 11.1

14. Intermolecular Forces
4. Hydrogen Bond
The hydrogen bond is a special dipole-dipole interaction between the hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. A H … B or
A & B are N, O, or F
11.2

15. Hydrogen Bonds:
• Strength of H bonds: up to 40 kJ/mol
• Lots of H bonds = strong
• compare with strength of typical covalent bonds:
250 kJ/mole)

16. Why is the hydrogen bond considered a “special” dipole-dipole interaction?
Decreasing molar mass
Decreasing boiling point
11.2

17. Worked Example 11.2

19. Strong intermolecular forces
Properties of Liquids
Surface tension is the amount of energy required to stretch or increase the surface of a liquid by a unit area.
Strong intermolecular forces
High surface tension
The intermolecular attraction tend to pull molecules into liquids and cause the surface to tighten like a plastic film.
11.3

20. Adhesion<cohesion Adhesion>cohesion
Properties of Liquids
One example of surface tension is capillary action.
Cohesion is the intermolecular attraction between like molecules
Adhesion is an attraction between unlike molecules
Adhesion
Adhesion<cohesion
Adhesion>cohesion
mercury
Cohesion
water
11.3

21. Strong intermolecular forces
Properties of Liquids
Viscosity is a measure of a fluid’s resistance to flow.
Strong intermolecular forces
High viscosity
CH2-OH
CH-OH
11.3

22. Trapping prevails-denser
The structure and property of water
Water is a Unique Substance
Maximum Density
40C
Density of Water
Thermal expansion
Ice is less dense than water
Trapping prevails-denser
The highly ordered 3-D structure prevents the molecules from getting too close to one another
11.3

23. Solid---crystal structure
A crystalline solid possesses rigid and long-range order. In a crystalline solid, atoms, molecules or ions occupy specific (predictable) positions. (e.g. NaCl)
An amorphous solid does not possess a well-defined arrangement and long-range molecular order. (e.g. glass)
Structures of crystalline solids:
A unit cell is the basic repeating structural unit of a crystalline solid.
At lattice points:
Atoms
Molecules
Ions
lattice
point
Unit Cell
Unit cells in 3 dimensions
11.4

24. Figure 11.15

25. 11.4

26. 11.4

27. 11.4

28. Corner atom edge atom face centered atom
Shared by 2 unit cells
Shared by 8 unit cells
Shared by 4 unit cells
11.4

29. 1 atom/unit cell
2 atoms/unit cell
4 atoms/unit cell
(8 x 1/8 = 1)
(8 x 1/8 + 1 = 2)
(8 x 1/8 + 6 x 1/2 = 4)
11.4

30. 11.4

31. How to use unit cell to calculate formula:
• how many cells touch the atom in question
position in unit cell fraction
• center 1
• face 1/2
• edge 1/4
• corner 1/8

32. 4 atoms/unit cell in a face-centered cubic cell
When silver crystallizes, it forms face-centered cubic cells. The unit cell edge length is 409 pm. Calculate the density of silver.
d = m V
V = a3
= (409 pm)3 = 6.83 x cm3
4 atoms/unit cell in a face-centered cubic cell
107.9 g
mole Ag x
1 mole Ag
6.022 x 1023 atoms x
m = 4 Ag atoms
= 7.17 x g
d = m V
7.17 x g
6.83 x cm3 =
= 10.5 g/cm3
11.4

33. Worked Example 11.3a

34. Worked Example 11.3b

35. Close-packing and coordination number
• hexagonal close-packing = 12
– ABABAB-
• cubic close-packing = 12
– Face centered cubic
– ABCABCABC-
• Body centered cubic = 8
• simple cube = 6

36. Figure 11.20

37. X-ray diffraction by Crystals
X-ray diffraction and Bragg equation are used to locate atoms.
X-ray diffraction refers to the scattering of X rays by the units of a crystalline solid.
11.5

38. Extra distance =
BC + CD =
2d sinq
= nl
(Bragg Equation)
11.5

39. X rays of wavelength nm are diffracted from a crystal at an angle of Assuming that n = 1, what is the distance (in pm) between layers in the crystal?
nl = 2d sin q
n = 1
q =
l = nm = 154 pm nl
2sinq =
1 x 154 pm
2 x sin14.17
d =
= pm
11.5

40. Types of Cyrstals
• Ionic
• NaCl
• Covalent
• Network covalent dolics
• SiO2 and diamond
• Molecular
• Sugar
• Metallic
• Copper

41. Types of Crystals Ionic Crystals
Lattice points occupied by cations and anions
Held together by electrostatic attraction
Hard, brittle, high melting point
Poor conductor of heat and electricity
CsCl
ZnS
CaF2
11.6

42. Types of Crystals Covalent Crystals Lattice points occupied by atoms
Held together by covalent bonds
Hard, high melting point
Poor conductor of heat and electricity
carbon
atoms
sp2
sp3
diamond
graphite
allotrope
11.6

43. Types of Crystals Molecular Crystals
Lattice points occupied by molecules
Held together by intermolecular forces
Soft, low melting point
Poor conductor of heat and electricity
11.6

44. Cross Section of a Metallic Crystal
Types of Crystals
Metallic Crystals
Lattice points occupied by metal atoms
Held together by metallic bonds
Soft to hard, low to high melting point
Good conductors of heat and electricity
Cross Section of a Metallic Crystal
nucleus &
inner shell e-
mobile “sea”
of e-
11.6

45. Crystal Structures of Metals
11.6

46. Types of Crystals
11.6

47. amorphous solid
An amorphous solid does not possess a well-defined arrangement and long-range molecular order.
A glass is an optically transparent fusion product of inorganic materials that has cooled to a rigid state without crystallizing
Crystalline
quartz (SiO2)
Non-crystalline
quartz glass
11.7

48. Liquid-vapor equilibrium
Greatest
Order
Least
Evaporation
Condensation
Liquid-vapor equilibrium
11.8

49. Liquid-vapor equilibrium
The equilibrium vapor pressure is the vapor pressure measured when a dynamic equilibrium exists between condensation and evaporation
H2O (l) H2O (g)
Rate of
condensation
evaporation =
Dynamic Equilibrium
Vapor pressure = the pressure of a vapor above its liquid in a closed container
11.8

50. Before
Evaporation
At Equilibrium
11.8

51. The boiling point is the temperature at which the (equilibrium) vapor pressure of a liquid is equal to the external pressure.
The normal boiling point is the temperature at which a liquid boils when the external pressure is 1 atm.
11.8

52. On top of a mountain
• Water boils at a lower temp. on top of a mountain than at sea level.
• Why? because the external pressure (atmospheric pressure) is less on top of a mountain.
In an autoclave
• An autoclave is used to sterilize medical instruments
• The pressure is often 2 atmospheres.
• Will water inside the autoclave boil at 100°C?

53. Vapor Pressure Versus Temperature
Molar heat of vaporization (DHvap) is the energy required to vaporize 1 mole of a liquid at its boiling point.
ln P = -
DHvap RT
+ C
Clausius-Clapeyron Equation
P = (equilibrium) vapor pressure
T = temperature (K)
R = gas constant (8.314 J/K•mol)
Vapor Pressure Versus Temperature
11.8

54. Worked Example 11.7

55. The critical temperature (Tc) is the temperature above which the gas cannot be made to liquefy, no matter how great the applied pressure.This is also the highest temperature at which a substance can exist as liquid.
The critical pressure (Pc) is the minimum pressure that must be applied to bring about liquefaction at the critical temperature.
Critical Temperature and Pressure:
• Remember: Average KE(gas) = kT
• If the KE of a gas is too large, the
association needed between particles to form a liquid will not happen
• T must be lower than critical temp. in order to liquify a gas -- regardless of the pressure.
11.8

56. Liquid-solid equilibrium
H2O (s) H2O (l)
The melting point of a solid or the freezing point of a liquid is the temperature at which the solid and liquid phases coexist in equilibrium
Melting
Freezing
11.8

57. Molar heat of fusion (DHfus) is the energy required to melt 1 mole of a solid substance at its freezing point.
11.8

58. Heating Curve
• Heat of fusion = energy needed to convert a solid to a liquid
• Heat of vaporization = energy needed to
convert a liquid to a gas
• Be able to draw a heating curve
11.8

59. Solid-vapor equilibrium
H2O (s) H2O (g)
Evaporation
Condensation
Molar heat of sublimation (DHsub) is the energy required to sublime 1 mole of a solid.
Sublimation
Deposition
DHsub = DHfus + DHvap
Melting
Freezing
( Hess’s Law)
11.8

60. Energy is needed to
• heat a solid
• heat a liquid
• heat a gas
• convert a solid to a liquid
• convert a liquid to a gas
• convert a solid to a gas

61. Phase Diagram of Water phase diagram
A phase diagram summarizes the conditions at which a substance exists as a solid, liquid, or gas.
Phase Diagram of Water
Solid line:
Two phases are in equilibrium
Triple point
All three phases are in equilibrium
11.9

62. Effect of Increase in Pressure on the Melting Point
of Ice and the Boiling Point of Water
11.9

63. Phase Diagram of Carbon Dioxide
At 1 atm
CO2 (s) CO2 (g)
11.9

64. Metallic bonding is the electromagnetic interaction between delocalized electrons, called conduction electrons, and the metallic nuclei within metals. Understood as the sharing of "free" electrons among a lattice of positively-charged ions (cations), metallic bonding is sometimes compared with that of molten salts; however, this simplistic view holds true for very few metals. In a more quantum-mechanical view, the conduction electrons divide their density equally over all atoms that function as neutral (non-charged) entities. Metallic bonding accounts for many physical properties of metals, such as strength, malleability, ductility, thermal and electrical conductivity, opacity, and luster.[1][2][3][4]