1. Intermolecular Forces, Liquids, and Solids
Seneca Valley SHS
AP Chemistry
Chapter 10 1 1 1 1

2. Solids, Liquids & Gases, A Comparison
The fundamental difference between states of matter is the distance between particles.

3. Solids, Liquids & Gases, A Comparison
Because in the solid and liquid states particles are closer together, we refer to them as condensed phases.
© 2009, Prentice-Hall, Inc.

4. Intermolecular Forces
Chapters 8 & 9 – Bonding (AKA intramolecular forces)
The forces holding solids and liquids together are called intermolecular forces.
Intermolecular forces are much weaker than chemical bonds. (< 15% as strong)
Example: energy needed to vaporize water: 40.7 kJ compared to energy needed to break O-H bond in water: 934 kJ

5. Intermolecular Forces
When a substance melts or boils the intermolecular forces are broken (not the covalent bonds).
When a substance condenses intermolecular forces are formed.
Intermolecular forces in ice
Intermolecular forces between water molecules

6. Importance of Intermolecular Forces
Intermolecular forces of attraction are strong enough to control physical properties such as:
Boiling & melting points
Vapor pressures
Viscosity
Surface tension
Collectively, the intermolecular forces being discussed are called van der Waal forces.
London dispersion
Dipole-Dipole
Hydrogen bonding

7. Intermolecular Forces – Dipole-Dipole
Polar molecules 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 (like in solid & liquid states).

8. Intermolecular Forces – Dipole-Dipole
In molecules of equal mass & size, the strength of the IMF increase with increasing polarity.

9. Intermolecular Forces
London Dispersion Forces
Weakest of all intermolecular forces.
Only IMF between noble gases and nonpolar molecules.
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).
Relatively weak and short-lived

10. Intermolecular Forces
London Dispersion Forces
One instantaneous dipole can induce another instantaneous dipole in an adjacent molecule (or atom).

11. Intermolecular Forces - London Dispersion
Strength of dispersion forces are directly related to:
Surface area available for contact
London dispersion forces between spherical molecules are lower than between sausage-like molecules.
Polarizability
Polarizability - ease with which an electron cloud can be deformed.
Polarizability increases with the number of electrons in the molecule.
Pi bonding also enhances the polarizability of a molecule.

12. Intermolecular Forces
London Dispersion Forces
Nonpolar Molecule
Boiling Point (°C, 1 atm)
# electrons
Comparison of strength H2
-253 O2
-183
Cl2
-34
Br2 59

13. Intermolecular Forces – How to rank IMF
Dispersion forces exist between ALL molecules.
When comparing relative strengths of intermolecular attractions, remember:
When molecules have similar molecular weights and shapes, dispersion forces are approximately equal. The relative strength of IMF will be measured by strengths of dipole-dipole interactions (which one is more polar).
When molecules vary drastically in molecular weights, dispersion forces are the decisive factor in determining which substance has stronger IMF (which one is larger).

14. Intermolecular Forces

15. Intermolecular Forces
Hydrogen Bonding
Special case of dipole-dipole forces when H is bonded to N, O, or F.
By experiments: boiling points of compounds with H-F, H-O, and H-N bonds are abnormally high.
Hydrogen bonds are abnormally strong but still weaker than chemical bonds (5-25 kJ/mol versus kJ/mol)

16. Intermolecular Forces
Hydrogen Bonding
Hydrogen bonds are responsible for:
Unique properties of water

17. Intermolecular Forces
Hydrogen Bonding
Hydrogen bonds are responsible for:
Protein Structure
Protein folding is a consequence of H-bonding.
DNA Transport of Genetic Information

18. Intermolecular Forces
Comparing Intermolecular Forces

19. 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.

20. Some Properties of Liquids
Surface Tension
Bulk molecules (those in the liquid) are equally attracted to their neighbors. 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.
This is what leads to a meniscus!

21. Some Properties of Liquids
Capillary Action
The rise of liquids up very narrow tubes
Helps water and dissolved nutrients move upward through plants

22. Bonding in Solids Solids can be classified into two categories
Amorphous – considerable disorder in their structures
Example: Glass
Crystalline – highly regular arrangement of atoms usually represented by a lattice
Amorphous
Crystalline

23. Bonding in Crystalline Solids

24. 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 at low temperatures.
Here is the structure of ice. Notice how the hydrogen bonds become rigid forming a lattice linking each molecule of water together.

25. Bonding in Solids Covalent Network Solids – C & SiO2
Atoms held together by covalent bonds in large networks.
Examples: diamond, graphite, quartz (SiO2), silicon carbide (SiC), and boron nitride (BN).
Diamond crystal structure: Each atom here is C.

26. Bonding in Solids Covalent Network Solids Crystalline SiO2 (quartz)
Each one of these structures is bonded together with regular, repeating covalent bonds.

27. Bonding in Solids Ionic Solids Ions are at the lattice points.
Electrostatic attraction keeps them frozen into this position.

28. Bonding in Solids Metals
Atoms held together by non-directional covalent bonds in metallic crystals.

29. Bonding in Solids Metals – Electron Sea Model
Atoms in a metal are arranged in a regular manner and vibrate about fixed positions
The outermost electrons move freely, forming a ‘sea of electrons’ enveloping the positive metal ions.
The metal ions are attracted to and hence held together by the ‘sea of electrons’ – these constititute metallic bonding.
The movements of the electrons are random under normal conditions.
However, when a potential difference is applied across the metal, the electrons move towards the direction of the positive pole.

30. Bonding in Solids Metals – Properties
As a result of this non-directional bonding, metals are:
Great conductors of heat & electricity
Ductile (able to drawn into wire)
Malleable (able to pounded into a sheet)
Able to form alloys easily
Substance that contains a mixture of two elements and has metallic properties

31. Bonding in Solids Metal Alloys Two types
Substitutional – some of the metal atoms are replaced by other metal atoms of similar size
Examples include brass (copper & zinc), sterling silver (silver & copper), pewter (tin, copper, bismuth & antimony) and plumber’s solder (tin & antimony)
Interstitial – some of the holes between the metallic crystal are filled with small atoms
Example would be steel (iron mixed with carbon)

32. Bonding in Solids
Metal Alloys

33. Phase Changes
Surface molecules are only attracted inwards towards the bulk molecules.
Sublimation: solid  gas.
Vaporization: liquid  gas.
Melting or fusion: solid  liq.
Deposition: gas  solid.
Condensation: gas  liquid.
Freezing: liquid  solid.
Energy Changes Accompanying Phase Changes
Sublimation: Hsub > 0 (endo)
Vaporization: Hvap > 0 (endo)
Melting or Fusion: Hfus > 0 (endo)
Deposition: Hdep < 0 (exo)
Condensation: Hcon < 0 (exo)
Freezing: Hfre < 0 (exo)

34. Phase Changes Energy Changes Accompanying Phase Changes
All phase changes are possible under the right conditions (e.g. water sublimes when snow disappears without forming puddles).
The sequence
heat solid  melt  heat liquid  boil  heat gas
is endothermic.
cool gas  condense  cool liquid  freeze  cool solid
is exothermic.

35. Phase Changes
Energy Changes Accompanying Phase Changes

36. 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.

37. 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.

38. 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.

39. Vapor Pressure Volatility, Vapor Pressure, and Temperature
If equilibrium is never established then the liquid evaporates.
Volatile substances evaporate rapidly.
The higher the temperature, the higher the average kinetic energy, the faster the liquid evaporates.

40. 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.
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.
Normal boiling point is the boiling point at 760 mmHg (1 atm).