1. Intermolecular Forces
(a) Particles in solid (b) Particles in liquid (c) Particles in gas

2. Properties of Solids, Liquids, and Gases
Property Solid Liquid Gas
Shape Has definite shape Takes the shape of Takes the shape
the container of its container
Volume Has a definite volume Has a definite volume Fills the volume of
High densities High densities the container
Low densities
Bonding Ionic, Metallic, Covalent Covalent Covalent
Arrangement of Fixed, very close Random, close Random, far apart,
Particles Crystalline or amorphous Collisions
Interactions Between Very strong forces: Strong forces: Essentially none
Particles (i.e. Melting point, (i.e. Boiling point,
malleability, ductility, Surface Tension,
conductivity…) Viscosity, Vapor
pressure…)

3. Liquid Surface Tension H2O(l) Water molecules minimize
their surface area (“skin”)
molecules at surface
interact only with molecules
in the interior of liquid
surface molecules
subjected to inward force,
so surface is under tension
surface tension increases
with increasing intermolecular
forces
some writing from Kotz (PowerPoints online)
H2O(l) Water
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 31

4. Liquid Viscosity H2O(l) Water resistance of a liquid to flow
greatest in substances with
strong intermolecular forces,
which hinder flow
longer molecules higher
viscosity than shorter ones
some writing from Kotz (PowerPoints online)
H2O(l) Water
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 31

5. Acetone

6. Gasoline (Hexane)

7. Corn Syrup

8. Motor Oil

9. Molasses

10. Intermolecular Forces Dispersion (London Force)
(a) Interaction of any two
atoms or molecules.
Electrons unevenly
distributed. Creates
instantaneous (temporary
dipole). Polarization
increases with size.
(b) interaction of many
dipoles. WEAKEST forces!
- +
- + + -
Attraction
Repulsion
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 442

11. Intermolecular Forces Dipole-Dipole
(a) Interaction of two
polar molecules. Polar
molecules have permanent
dipoles from electronegativity
difference. Higher melting and
boiling points due to
stronger IM forces.
(b) interaction of many
dipoles in a liquid.
- +
- + + -
Attraction
Repulsion
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 442

12. Intermolecular Forces Hydrogen Bonding
Strong intermolecular forces of attraction between molecules containing fluorine, oxygen, or nitrogen bonded to hydrogen
Results from large
electronegativity
difference and small
atomic size of hydrogen
Hydrogen
bonds O H
Chemical
Bonds O H
Chemical
Bonds
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 442

13. Boiling Points of Covalent Hydrides
H2O
H2Te
H2Se
SnH4
H2S
Group 16 Hydrogen Compounds
Compound Molar Mass Melting Point Boiling Point H fusion H vapor
(oC) (oC) (cal/mol) (cal/mol)
H2O
H2S
H2Se
H2Te
GeH4
SiH4
CH4 50
100

14. Evaporation (Vaporization)
Molecules must have sufficient energy to break IM forces.
Molecules at the surface break away and become gas (“volatility”).
Only molecules with enough KE escape.
Breaking IM forces absorbs energy.
Evaporation is endothermic.
Rate of evaporation increases with
increasing surface area, increasing
temperature, and weaker IM forces

15. Condensation Forming IM forces from gas to liquid
Condensation is exothermic because energy is released.
Dynamic equilibrium: rate of vaporization
equals rate of condensation (gas molecules above liquid becomes constant).
Vapor pressure: partial pressure
of gas in dynamic equilibrium with liquid
Vapor pressure increases
with increasing temperature
and weaker IM forces

16. Boiling point: temperature at which the vapor pressure of a liquid is equal to the pressure above it. Normal boiling point is boiling point at atmospheric pressure
Microscopic view of a liquid near its surface
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 446

17. Formation of a bubble is opposed by the pressure of the atmosphere
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 452

18. Effect of Pressure on Boiling Point
Boiling Point of Water at Various Locations
Location
Feet above sea level
Patm (kPa)
Boiling Point (C)
Top of Mt. Everest, Tibet
29,028
32.0 70
Top of Mt. Denali, Alaska
20,320
45.3 79
Top of Mt. Whitney, California
14,494
57.3 85
Leadville, Colorado
10,150
68.0 89
Top of Mt. Washington, N.H.
6,293
78.6 93
Boulder, Colorado
5,430
81.3 94
Madison, Wisconsin
900
97.3 99
New York City, New York 10
101.3
100
Death Valley, California
-282
102.6
100.3

19. Heating Curves Temperature Change Phase Change Heat Capacity
change in KE (molecular motion)
depends on heat capacity
Phase Change
change in PE (molecular arrangement)
temperature remains constant
Heat Capacity
energy required to raise the temp of 1 gram of a substance by 1°C (q = mC∆T)
Courtesy Christy Johannesson

20. Heating Curve for Water
vaporization E
gas D
100
condensation C
liquid
melting
Temperature (oC)
Melting, freezing, vaporization, condensation, sublimation, and deposition are six common phase changes.
Note: The temperature of a substance does not change during a phase change. B A
freezing
solid
Heat added
LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 487

21. Heating Curves Gas - KE  Boiling - PE  Liquid - KE  Melting - PE 
140
Gas - KE 
120
100
Boiling - PE  80 60 40
Liquid - KE 
Temperature (oC) 20
Melting - PE 
-20
-40
Solid - KE 
-60
-80
-100
Energy

22. Heating Curves Heat of Fusion (Hfus) Heat of Vaporization (Hvap)
energy required to melt 1 gram of a substance at its melting point. Breaking intermolecular forces in the solid. Water = 6.02 kJ/mol M.P.)
Heat of Vaporization (Hvap)
energy required to vaporize 1 gram of a substance at its boiling point.
usually larger than Hfus (requires complete separation of molecules)
higher temperatures = lower (Hvap)
Water = 40.7 kJ/mol B.P.)
Courtesy Christy Johannesson

23. Calculating Energy Changes - Heating Curve for Water
140
DH = mol x DHvap
120
DH = mol x DHfus
100 80 60
∆H = mCgasDT 40
Temperature (oC) 20
∆H = mCliquidDT
-20
-40
-60
∆H = mCsolidDT
-80
-100
Energy

24. Specific Heat Capacities
Tro's "Introductory Chemistry", Chapter 3

25. Energy Changes Accompanying Phase Changes
Gas
Vaporization
Condensation
Sublimation
Deposition
Energy of system
Liquid
Melting
Freezing
Solid
Brown, LeMay, Bursten, Chemistry 2000, page 405

26. Phase Diagram Water
Show the phases of a substance at different temperatures and pressures.

27. Phase Diagram of CO2

28. Solids
Ionic – Giant lattice of ions, strong electrostatic attractions between metal and nonmetal, strong bonds, high melting points, crystalline, poor conductors except in molten and aqueous states (NaCl, KBr…)
Discrete Covalent – Groups of atoms covalently bonded by sharing electrons between two nonmetals/metalloids, weaker bonds than ionic, IMFs with relatively low melting points (except very large polymers), soft, poor conductors (plastics, wax…)
Giant Covalent – Massive structures of atoms bonded through network covalent bonds, very strong bonds with very high melting points, hard/rigid, poor conductors (graphite (carbon), diamond (carbon), SiO2, Si…)
Metallic – Closely packed array of atoms or ions with “sea” of free moving electrons, all metals and alloys (mixtures), stronger bonds than covalent but weaker than ionic, high melting points, malleable/ductile/shiny, very good conductors (Cu, Na…)

33. Molecular Structure of Ice
Hydrogen
bonding
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 455

36. Natural Polymers
Cellulose is a macromolecule composed of individual sugar molecules (glucose) that are bonded together to give molecular weights in the millions.
Cellulose is the basis for cotton and rayon fibers as well as the structural support in plants
Starch
Cellulose
RNA and DNA
Natural Polymers = silk, cotton, starch, sand, and asbestos, as well as the incredibly complex polymers known as RNA (ribonucleic acid) and DNA
Chitin
Cotton
Natural Rubber 37

37. Synthetic Polymers “Everyday Polys” Polypropylene Polystyrene
Polyvinyl Chloride
Polyester
Polyethylene
Polyvinyl chloride
Polypropylene
Polystyrene
Polyethylene