1. Chapter 11 Intermolecular Forces, Liquids and Solids
CHEMISTRY The Central Science 9th Edition
Chapter 11 Intermolecular Forces, Liquids and Solids
David P. White
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2. A Molecular Comparison of Liquids and Solids
Physical properties of substances understood in terms of kinetic molecular theory:
Gases are highly compressible, assumes shape and volume of container:
Gas molecules are far apart and do not interact much with each other.
Liquids are almost incompressible, assume the shape but not the volume of container:
Liquids molecules are held closer together than gas molecules, but not so rigidly that the molecules cannot slide past each other.
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3. A Molecular Comparison of Liquids and Solids
Solids are incompressible and have a definite shape and volume:
Solid molecules are packed closely together. The molecules are so rigidly packed that they cannot easily slide past each other.
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4. A Molecular Comparison of Liquids and Solids
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5. A Molecular Comparison of Liquids and Solids
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6. A Molecular Comparison of Liquids and Solids
Converting a gas into a liquid or solid requires the molecules to get closer to each other:
cool or compress.
Converting a solid into a liquid or gas requires the molecules to move further apart:
heat or reduce pressure.
The forces holding solids and liquids together are called intermolecular forces.
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7. Intermolecular Forces
The covalent bond holding a molecule together is an intramolecular forces.
The attraction between molecules is an intermolecular force.
Intermolecular forces are much weaker than intramolecular forces (e.g. 16 kJ/mol vs. 431 kJ/mol for HCl).
When a substance melts or boils the intermolecular forces are broken (not the covalent bonds).
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8. Intermolecular Forces
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9. Intermolecular Forces
Ion-Dipole Forces
Interaction between an ion and a dipole (e.g. water).
Strongest of all intermolecular forces.
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10. Intermolecular Forces
Dipole-Dipole Forces
Dipole-dipole forces exist between neutral polar molecules.
Polar molecules need to be close together.
Weaker than ion-dipole forces.
There is a mix of attractive and repulsive dipole-dipole forces as the molecules tumble.
If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity.
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11. Intermolecular Forces
Dipole-Dipole Forces

12. Intermolecular Forces
Dipole-Dipole Forces
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13. Intermolecular Forces
London Dispersion Forces
Weakest of all intermolecular forces.
It is possible for two adjacent neutral molecules to affect each other.
The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom).
For an instant, the electron clouds become distorted.
In that instant a dipole is formed (called an instantaneous dipole).
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14. Intermolecular Forces
London Dispersion Forces
One instantaneous dipole can induce another instantaneous dipole in an adjacent molecule (or atom).
The forces between instantaneous dipoles are called London dispersion forces.
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15. Intermolecular Forces
London Dispersion Forces
Polarizability is the ease with which an electron cloud can be deformed.
The larger the molecule (the greater the number of electrons) the more polarizable.
London dispersion forces increase as molecular weight increases.
London dispersion forces exist between all molecules.
London dispersion forces depend on the shape of the molecule.
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16. Intermolecular Forces
London Dispersion Forces
The greater the surface area available for contact, the greater the dispersion forces.
London dispersion forces between spherical molecules are lower than between sausage-like molecules.
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17. Intermolecular Forces
London Dispersion Forces

18. Intermolecular Forces
London Dispersion Forces
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19. Intermolecular Forces
Hydrogen Bonding
Special case of dipole-dipole forces.
By experiments: boiling points of compounds with H-F, H-O, and H-N bonds are abnormally high.
Intermolecular forces are abnormally strong.
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20. Intermolecular Forces
Hydrogen Bonding
H-bonding requires H bonded to an electronegative element (most important for compounds of F, O, and N).
Electrons in the H-X (X = electronegative element) lie much closer to X than H.
H has only one electron, so in the H-X bond, the + H presents an almost bare proton to the - X.
Therefore, H-bonds are strong.
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21. Hydrogen Bonding

22. Hydrogen Bonding

23. Intermolecular Forces
Hydrogen Bonding
Hydrogen bonds are responsible for:
Ice Floating
Solids are usually more closely packed than liquids;
Therefore, solids are more dense than liquids.
Ice is ordered with an open structure to optimize H-bonding.
Therefore, ice is less dense than water.
In water the H-O bond length is 1.0 Å.
The O…H hydrogen bond length is 1.8 Å.
Ice has waters arranged in an open, regular hexagon.
Each + H points towards a lone pair on O.
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24. Intermolecular Forces
Hydrogen Bonding

25. Intermolecular Forces

26. Some Properties of Liquids
Viscosity
Viscosity is the resistance of a liquid to flow.
A liquid flows by sliding molecules over each other.
The stronger the intermolecular forces, the higher the viscosity.
Surface Tension
Bulk molecules (those in the liquid) are equally attracted to their neighbors.
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27. Some Properties of Liquids
Viscosity
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28. Surface Tension

29. Some Properties of Liquids
Surface Tension
Surface molecules are only attracted inwards towards the bulk molecules.
Therefore, surface molecules are packed more closely than bulk molecules.
Surface tension is the amount of energy required to increase the surface area of a liquid.
Cohesive forces bind molecules to each other.
Adhesive forces bind molecules to a surface.
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30. Some Properties of Liquids
Surface Tension
Meniscus is the shape of the liquid surface.
If adhesive forces are greater than cohesive forces, the liquid surface is attracted to its container more than the bulk molecules. Therefore, the meniscus is U-shaped (e.g. water in glass).
If cohesive forces are greater than adhesive forces, the meniscus is curved downwards.
Capillary Action: When a narrow glass tube is placed in water, the meniscus pulls the water up the tube.
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31. Phase Changes
Surface molecules are only attracted inwards towards the bulk molecules.
Sublimation: solid  gas.
Vaporization: liquid  gas.
Melting or fusion: solid  liquid.
Deposition: gas  solid.
Condensation: gas  liquid.
Freezing: liquid  solid.
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32. Phase Changes

33. Phase Changes Energy Changes Accompanying Phase Changes
Sublimation: Hsub > 0 (endothermic).
Vaporization: Hvap > 0 (endothermic).
Melting or Fusion: Hfus > 0 (endothermic).
Deposition: Hdep < 0 (exothermic).
Condensation: Hcon < 0 (exothermic).
Freezing: Hfre < 0 (exothermic).
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34. Phase Changes Energy Changes Accompanying Phase Changes
Generally heat of fusion (enthalpy of fusion) is less than heat of vaporization:
it takes more energy to completely separate molecules, than partially separate them.
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35. Phase Changes

36. Phase Changes Energy Changes Accompanying Phase Changes
All phase changes are possible under the right conditions.
The sequence
heat solid  melt  heat liquid  boil  heat gas
is endothermic.
cool gas  condense  cool liquid  freeze  cool solid
is exothermic.
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37. Phase Changes Heating Curves
Plot of temperature change versus heat added is a heating curve.
During a phase change, adding heat causes no temperature change.
These points are used to calculate Hfus and Hvap.
Supercooling: When a liquid is cooled below its melting point and it still remains a liquid.
Achieved by keeping the temperature low and increasing kinetic energy to break intermolecular forces.
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38. Prentice Hall © 2003
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39. Phase Changes Critical Temperature and Pressure
Gases liquefied by increasing pressure at some temperature.
Critical temperature: the warmest temperature for liquefaction of a gas using pressure. Above this temperature, additional pressure only causes compression.
Critical pressure: pressure required for liquefaction at the critical temperature.
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40. Phase Changes Critical Temperature and Pressure Prentice Hall © 2003
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41. Vapor Pressure Explaining Vapor Pressure on the Molecular Level
Some of the molecules on the surface of a liquid have enough energy to escape the attraction of the bulk liquid.
These molecules move into the gas phase.
As the number of molecules in the gas phase increases, some of the gas phase molecules strike the surface and return to the liquid.
After some time the pressure of the gas will be constant at the vapor pressure.
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42. Vapor Pressure Explaining Vapor Pressure on the Molecular Level
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43. Vapor Pressure Explaining Vapor Pressure on the Molecular Level
Dynamic Equilibrium: the point when as many molecules escape the surface as strike the surface.
Vapor pressure is the pressure exerted when the liquid and vapor are in dynamic equilibrium.
Volatility, Vapor Pressure, and Temperature
If equilibrium is never established then the liquid evaporates.
Volatile substances evaporate rapidly.
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44. Vapor Pressure Volatility, Vapor Pressure, and Temperature
The higher the temperature, the higher the average kinetic energy, the faster the liquid evaporates.
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45. Vapor Pressure
Volatility, Vapor Pressure, and Temperature

46. Vapor Pressure Vapor Pressure and Boiling Point
Liquids boil when the external pressure equals the vapor pressure.
Temperature of boiling point increases as pressure increases.
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48. Vapor Pressure Vapor Pressure and Boiling Point
Two ways to get a liquid to boil: increase temperature or decrease pressure.
Pressure cookers operate at high pressure. At high pressure the boiling point of water is higher than at 1 atm. Therefore, there is a higher temperature at which the food is cooked, reducing the cooking time required.
Normal boiling point is the boiling point at 760 mmHg (1 atm).
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49. Phase Diagrams
Phase diagram: plot of pressure vs. Temperature summarizing all equilibria between phases.
Given a temperature and pressure, phase diagrams tell us which phase will exist.
Any temperature and pressure combination not on a curve represents a single phase.
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50. Phase Diagrams Features of a phase diagram:
Triple point: temperature and pressure at which all three phases are in equilibrium.
Vapor-pressure curve: generally as pressure increases, temperature increases.
Critical point: critical temperature and pressure for the gas.
Melting point curve: as pressure increases, the solid phase is favored if the solid is more dense than the liquid.
Normal melting point: melting point at 1 atm.
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51. Phase Diagrams
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52. Phase Diagrams The Phase Diagrams of H2O and CO2 Prentice Hall © 2003
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53. Phase Diagrams The Phase Diagrams of H2O and CO2 Water:
The melting point curve slopes to the left because ice is less dense than water.
Triple point occurs at C and 4.58 mmHg.
Normal melting (freezing) point is 0C.
Normal boiling point is 100C.
Critical point is 374C and 218 atm.
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54. Phase Diagrams The Phase Diagrams of H2O and CO2 Carbon Dioxide:
Triple point occurs at -56.4C and 5.11 atm.
Normal sublimation point is -78.5C. (At 1 atm CO2 sublimes it does not melt.)
Critical point occurs at 31.1C and 73 atm.
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55. Structures of Solids Unit Cells
Crystalline solid: well-ordered, definite arrangements of molecules, atoms or ions.
Crystals have an ordered, repeated structure.
The smallest repeating unit in a crystal is a unit cell.
Unit cell is the smallest unit with all the symmetry of the entire crystal.
Three-dimensional stacking of unit cells is the crystal lattice.
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56. Structures of Solids
Unit Cells

57. Structures of Solids Unit Cells Three common types of unit cell.
Primitive cubic, atoms at the corners of a simple cube,
each atom shared by 8 unit cells;
Body-centered cubic (bcc), atoms at the corners of a cube plus one in the center of the body of the cube,
corner atoms shared by 8 unit cells, center atom completely enclosed in one unit cell;
Face-centered cubic (fcc), atoms at the corners of a cube plus one atom in the center of each face of the cube,
corner atoms shared by 8 unit cells, face atoms shared by 2 unit cells.
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58. Unit Cells

59. Unit Cells

60. Structures of Solids
Unit Cells

61. Structures of Solids The Crystal Structure of Sodium Chloride
Two equivalent ways of defining unit cell:
Cl- (larger) ions at the corners of the cell, or
Na+ (smaller) ions at the corners of the cell.
The cation to anion ratio in a unit cell is the same for the crystal. In NaCl each unit cell contains same number of Na+ and Cl- ions.
Note the unit cell for CaCl2 needs twice as many Cl- ions as Ca2+ ions.
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62. Structures of Solids The Crystal Structure of Sodium Chloride
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63. Structures of Solids The Crystal Structure of Sodium Chloride
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64. Structures of Solids Close Packing of Spheres
Solids have maximum intermolecular forces.
Molecules can be modeled by spheres.
Atoms and ions are spheres.
Molecular crystals are formed by close packing of the molecules.
We rationalize maximum intermolecular force in a crystal by the close packing of spheres.
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65. Structures of Solids Close Packing of Spheres
When spheres are packed as closely as possible, there are small spaces between adjacent spheres.
The spaces are called interstitial holes.
A crystal is built up by placing close packed layers of spheres on top of each other.
There is only one place for the second layer of spheres.
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66. Structures of Solids Close Packing of Spheres
There are two choices for the third layer of spheres:
Third layer eclipses the first (ABAB arrangement). This is called hexagonal close packing (hcp);
Third layer is in a different position relative to the first (ABCABC arrangement). This is called cubic close packing (ccp).
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67. Structures of Solids Close Packing of Spheres Prentice Hall © 2003
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68. Structures of Solids Close Packing of Spheres
Each sphere is surrounded by 12 other spheres (6 in one plane, 3 above and 3 below).
Coordination number: the number of spheres directly surrounding a central sphere.
Hexagonal and cubic close packing are different from the cubic unit cells.
If unequally sized spheres are used, the smaller spheres are placed in the interstitial holes.
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69. Bonding in Solids There are four types of solid:
Molecular (formed from molecules) - usually soft with low melting points and poor conductivity.
Covalent network (formed from atoms) - very hard with very high melting points and poor conductivity.
Ions (formed form ions) - hard, brittle, high melting points and poor conductivity.
Metallic (formed from metal atoms) - soft or hard, high melting points, good conductivity, malleable and ductile.
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70. Bonding in Solids
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71. Bonding in Solids Molecular Solids
Intermolecular forces: dipole-dipole, London dispersion and H-bonds.
Weak intermolecular forces give rise to low melting points.
Room temperature gases and liquids usually form molecular solids and low temperature.
Efficient packing of molecules is important (since they are not regular spheres).
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72. Bonding in Solids Covalent-Network Solids
Intermolecular forces: dipole-dipole, London dispersion and H-bonds.
Atoms held together in large networks.
Examples: diamond, graphite, quartz (SiO2), silicon carbide (SiC), and boron nitride (BN).
In diamond:
each C atom has a coordination number of 4; each C atom is tetrahedral; there is a three-dimensional array of atoms.
Diamond is hard, and has a high melting point (3550 C).
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73. Bonding in Solids Covalent-Network Solids Prentice Hall © 2003
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74. Bonding in Solids Covalent-Network Solids In graphite
each C atom is arranged in a planar hexagonal ring;
layers of interconnected rings are placed on top of each other;
the distance between C atoms is close to benzene (1.42 Å vs Å in benzene);
the distance between layers is large (3.41 Å);
electrons move in delocalized orbitals (good conductor).
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75. Bonding in Solids Ionic Solids
Ions (spherical) held together by electrostatic forces of attraction.
There are some simple classifications for ionic lattice types.
NaCl structure
CsCl structure
Zinc blende structure
Fluorite structure
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76. Ionic Solids

77. Bonding in Solids Ionic Solids NaCl Structure
Each ion has a coordination number of 6.
Face-centered cubic lattice.
Cation to anion ratio is 1:1.
Examples: LiF, KCl, AgCl and CaO.
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78. Bonding in Solids Ionic Solids CsCl Structure
Cs+ has a coordination number of 8.
Primitive cubic lattice not body centered cubic
Different from the NaCl structure (Cs+ is larger than Na+).
Cation to anion ratio is 1:1.
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79. Bonding in Solids Ionic Solids Zinc Blende Structure
Typical example ZnS.
The S2- ions are placed in a tetrahedron around the Zn2+ ions.
Zn2+ ions have a coordination number of 4.
Another example: CuCl.
S2- ions adopt a fcc arrangement.
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80. Bonding in Solids Ionic Solids Fluorite Structure
Typical example CaF2.
Ca2+ ions in a fcc arrangement.
There are twice as many F- per Ca2+ ions in each unit cell.
Coordination number for Ca2+ is 8
Coordination number for F- is 4
Examples: BaCl2, PbF2.
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81. Bonding in Solids Metallic Solids
Metallic solids have metal atoms in hcp, fcc or bcc arrangements.
Coordination number for each atom is either 8 or 12.
Problem: the bonding is too strong for London dispersion and there are not enough electrons for covalent bonds.
Resolution: the metal nuclei float in a sea of electrons.
Metals conduct because the electrons are delocalized and are mobile.
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83. End of Chapter 11 Intermolecular Forces, Liquids and Solids
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