1. States of Matter

2. States of Matter
Solid
Gas
Liquid

3. What causes the differences in solids, liquids, and gases?

4. Kinetic Molecular Theory
Describes behavior of matter in terms of particles in motion.
Makes assumptions of gas particles:
separated by empty space
particles are not attracted to each other
are in constant, random motion
collisions are elastic
kinetic energy determined by mass & velocity
KE = ½ mv2

5. I. Intermolecular Forces (between molecules) (Ch. 6, p.189-193)
Liquids & Solids
I. Intermolecular Forces (between molecules) (Ch. 6, p )
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6. A. Definition of Intermolecular Forces
Attractive forces between molecules.
Much weaker than chemical bonds within molecules.
a.k.a. van der Waals forces
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7. Intermolecular Forces
Attraction between
molecules

8. B. Types of IMF
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9. B. Types of IMF London Dispersion Forces View animation online.
C. Johannesson

10. Dispersion
Weak forces caused from temporary shifts in e- density

11. Polarity (Differing Electronegativities)
Eletronegativity Difference
Bond Character
> 1.7
ionic
0.4 – 1.7
polar covalent
< 0.4
Nonpolar covalent

12. B. Types of IMF Dipole-Dipole Forces + - View animation online.
C. Johannesson

13. Dipole-Dipole
Between partial positive area of one molecule with the partial negative area of another
Occurs in polar molecules

14. B. Types of IMF
Hydrogen Bonding
C. Johannesson

15. Hydrogen Bonding Special dipole-dipole
Between H and a highly electronegative atom (O, N, F)
SPECIAL NOTE – THE H atom has to be directly chemically bonded to an O, N, or F

16. Your DNA

17. C. Determining IMF CH2Cl2 polar = dispersion, dipole-dipole CH4
nonpolar = dispersion HF
H-F bond = dispersion, dipole-dipole, hydrogen bonding
C. Johannesson

18. II. Physical Properties A. Liquids vs. Solids
IMF Strength
Fluid
Density
Compressible
Diffusion
LIQUIDS
Stronger than in gases Y
high N
slower than in gases
SOLIDS
Very strong N
high
extremely slow
C. Johannesson

19. B. Liquid Properties Surface Tension
attractive force between particles in a liquid that minimizes surface area
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20. B. Liquid Properties Capillary Action
attractive force between the surface of a liquid and the surface of a solid
water
mercury
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21. C. Types of Solids Crystalline - repeating geometric pattern
covalent network
metallic
ionic
covalent molecular
Amorphous - no geometric pattern
decreasing
m.p.
C. Johannesson

22. C. Types of Solids
Ionic
(NaCl)
Metallic
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23. C. Types of Solids Covalent Molecular Covalent Network Amorphous (H2O)
(SiO2 - quartz)
Amorphous
(SiO2 - glass)
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24. III. Changes of State A. Phase Changes
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25. A. Phase Changes Evaporation
molecules at the surface gain enough energy to overcome IMF
Volatility
measure of evaporation rate
depends on temp & IMF
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26. A. Phase Changes temp volatility IMF volatility Boltzmann Distribution
# of Particles
volatility
IMF
volatility
Kinetic Energy
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27. A. Phase Changes Equilibrium
trapped molecules reach a balance between evaporation & condensation
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28. A. Phase Changes temp v.p. IMF v.p. Vapor Pressure
pressure of vapor above a liquid at equilibrium
v.p.
depends on temp & IMF
directly related to volatility
temp
temp
v.p.
IMF
v.p.
C. Johannesson

29. A. Phase Changes Patm b.p. IMF b.p. Boiling Point
temp at which v.p. of liquid equals external pressure
depends on Patm & IMF
Normal B.P. - b.p. at 1 atm
Patm
b.p.
IMF
b.p.
C. Johannesson

30. A. Phase Changes IMF m.p. Melting Point equal to freezing point
Which has a higher m.p.?
polar or nonpolar?
covalent or ionic?
polar
ionic
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31. A. Phase Changes Sublimation solid  gas
v.p. of solid equals external pressure
EX: dry ice, mothballs, solid air fresheners
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32. B. Heating Curves Gas - KE  Boiling - PE  Liquid - KE 
Melting - PE 
Solid - KE 
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33. B. Heating Curves Temperature Change change in KE (molecular motion)
depends on heat capacity
Heat Capacity
energy required to raise the temp of 1 gram of a substance by 1°C
“Volcano” clip -
water has a very high heat capacity
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34. B. Heating Curves Phase Change change in PE (molecular arrangement)
temp remains constant
Heat of Fusion (Hfus)
energy required to melt 1 gram of a substance at its m.p.
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35. B. Heating Curves Heat of Vaporization (Hvap)
energy required to boil 1 gram of a substance at its b.p.
usually larger than Hfus…why?
EX: sweating, steam burns, the drinking bird
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36. C. Phase Diagrams
Show the phases of a substance at different temps and pressures.
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37. Phase Diagrams
Triple point - The temperature and pressure at which the solid, liquid, and vapor phases of a pure substance can coexist in equilibrium.
*Be able to know what phase change occurs when pressure and/or temperature changes when looking a phase diagram.
C. Johannesson