1. Liquids & Aqueous solutions

2. Ideal Gas Equation R = 0.08206 atm L/mol K
Ideal gas equation predicts gas behavior

3. A Model for Gas Behavior
Ideal gas law describes what gases do, but not why.
Kinetic Molecular Theory of Gases (KMT): model that explains gas behavior.
developed in mid-1800s
based on concept of an ideal or perfect gas

4. Ideal gas Tiny particles in constant, random, straight-line motion
Molecules collide w/ each other & w/ walls of container
Gas molecules are points; gas volume is empty space between molecules
Molecules independent of each other (no attractive or repulsive forces between them).

5. Molecular motion & temperature
Moving molecule has kinetic energy
KE depends on mass (m) and speed (u)
Temperature (in K) proportional to average molecular KE
At higher T, average speed higher
At lower T, average speed lower
At T = 0, speed = 0 (molecules stop moving)

6. Different gases at same temperature
All have same average KE (same temperature)
Heavier gases are slower; lighter gases are faster

7. Molecular motion and pressure
Molecules colliding with container → gas pressure
What if there are more molecules?
More collisions → higher pressure

8. Molecular motion and pressure
Molecules colliding with container → gas pressure
What if the container is smaller?
More collisions → higher pressure

9. Molecular motion and pressure
Molecules colliding with container → gas pressure
What if the molecules are moving faster?
Harder, more frequent collisions → higher pressure

10. Molecular motion and volume
Moving molecules fill the container
Light molecules escape faster, heavy molecules more slowly
Large spaces between molecules allow gas to be compressed

11. KMT & liquids Ideal gas remains a gas when cooled, even to 0 K
Real gases condense to liquid state when cooled
How do we explain condensation?
Pressure
(atm)
Temperature (K)
Ideal gas pressure decreases steadily & becomes zero at absolute zero
Real gas pressure decreases abruptly to zero when gas condenses to liquid

12. Condensation KMT ignores attractions between gas molecules
Gas molecules are too far apart & too fast for attractions to act
BUT attractive forces do exist between all molecules!
At low enough T, attractions overcome kinetic energy & molecules stick together to form a liquid

13. Vaporization
At liquid surface, faster molecules have enough kinetic energy to escape (evaporate)
As higher-energy molecules leave liquid, average kinetic energy of liquid decreases
The temperature of liquid decreases (evaporative cooling)

14. What if the vapor can’t escape?
(a) Molecules evaporate
(b) Some vapor molecules return to liquid
(c) Rates of evaporation & condensation equal
When rate of vaporization = rate of condensation, system has reached dynamic equilibrium
vapor pressure at equilibrium is constant
vp is balance of KE (temperature) & attractions
vp affected only by temperature
liquids with high vp at room temp are volatile

15. Vapor pressure always increases as temperature increases
Vapor pressure curves
Vapor pressure always increases as temperature increases

16. Evaporation & boiling

17. Evaporation & boiling Evaporation occurs only at surface
As temperature increases, evaporation increases
At some point, evaporation begins to occur throughout the liquid instead of just at the surface: boiling!
Vapor bubbles form throughout the liquid
Bubbles rise to surface, burst, and release vapor

18. Boiling point
Boiling begins when liquid’s vapor pressure matches atmospheric pressure
Temperature at which this occurs is boiling point
b.p. at standard pressure is normal boiling point
When atmospheric pressure is lower, bp is lower
When atmospheric pressure is higher, bp is higher

19. Intermolecular attractions
Attractive forces exist between all atoms/molecules
Strength of attractions indicated by boiling point
When comparing two substances,
Low b.p. ⇒ weaker intermolecular attractions
High b.p. ⇒ stronger intermolecular attractions

20. Trends in boiling points

21. Intermolecular attractions
For molecules of similar structure,
boiling point increases as molar mass increases
intermolecular attractive forces increase as molar mass increases

22. Then there’s hydrogen . . .

23. The O–H bond in water is very polar, and the atoms are very small
The dipoles are close together, so their attraction is very strong
An H atom is covalently bonded (red-white) to its own O and weakly bonded (dotted line) to the neighboring O
This weak bond to a neighboring O is called a hydrogen bond

24. Hydrogen bonding
Hydrogen bonding occurs only between molecules containing N–H, O–H, and F–H bonds
Hydrogen bonding is much stronger than ordinary intermolecular attractions ⇒ very high boiling points for their mass
Hydrogen bonds are not as strong as covalent bonds (15-40 kJ/mol, vs >150 kJ/mol)

25. Boiling point & attractions
CH3-CH2-CH3
CH3-O-CH3
CH3-CH2-NH2
CH3-CH2-OH
molar mass
44 g/mol
45 g/mol
46 g/mol
normal bp
231 K
248 K
290 K
352 K
Similar masses⇒ predict similar boiling points
1st molecule is nonpolar
2nd molecule is polar, a little stickier: higher bp
3rd & 4th molecules have hydrogen bonding between molecules, much stickier: much higher bp

26. Heating curve
Add energy
Temperature

27. Heating curve Temperature (g) boiling condensing (l) melting freezing
Add energy
Temperature
(s)
melting
freezing
boiling
condensing
(l)
(g)
point
melting/freezing
point

28. Heating curve Melting & boiling are ENDOTHERMIC
Freezing & condensing are EXOTHERMIC

29. Changing temperature changes KE (#1, 3, 5)
Changing state changes potential energy (#2, 4)
energy to melt or freeze = heat of fusion (∆Hfusion)
energy to vaporize or condense = heat of vaporization (∆Hvaporization)

30. Energy in phase changes