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

3.  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 3

4.  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). 4

5.  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) 5

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

7. 7  Molecules colliding with container → gas pressure  What if there are more molecules?  More collisions → higher pressure

8. 8  Molecules colliding with container → gas pressure  What if the container is smaller?  More collisions → higher pressure

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

10. 10  Moving molecules fill the container  Light molecules escape faster, heavy molecules more slowly  Large spaces between molecules allow gas to be compressed

11. 11  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) 0 0 Ideal gas pressure decreases steadily & becomes zero at absolute zero Real gas pressure decreases abruptly to zero when gas condenses to liquid

12. 12  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.  For every substance there are 2 opposing tendencies:  Kinetic energy of the molecules, which tend to make them move apart from each other (gas-like)  Attraction between molecules, which tends to make them stick together (liquid-like)  At any given temp., kinetic energy is the same for all molecules, so the attractive forces between molecules determines whether something is a liquid, solid, or gas.

14.  At same temp. (same KE), molecules are gases, liquids, solids.  This suggests that some molecules have stronger attractions between molecules.  How do you judge the strength of the attractive forces between molecules?  Look at the boiling point.  Low Boiling point = weak intermolecular forces (molecules are not sticky)  High Boiling point = strong intermolecular forces (molecules are sticky)

15.  Molecules of similar structure:  Boiling points increase as molar mass increases  Suggests that intermolecular attractions increase as molecular mass increase FormulaMolar Mass BPFormulaMolar Mass BP F2F2 38 g/mol85 KCH 4 16 g/mol111 K Cl 2 71 g/mol239 KC2H6C2H6 30 g/mol184 K Br 2 160 g/mol332 KC3H8C3H8 44 g/mol231 K I2I2 254 g/mol457 KC 4 H 10 58 g/mol273 K

16.  Predict which would have a higher boiling point, and why?  Na or K  F 2 or Br 2

17.  As you heat a solid, the temp. increases until the solid melts.  Temp. will remain constant until solid melts completely.  You observe the same pattern when a gas is cooled until it changes into a solid  Temp. will remain constant until gas changes into solid

18. 18 Add energy Temperature

19. 19 melting/freezing point Add energy Temperature (s) melting freezing boiling condensing (l) (g) boiling point

20. 20  Melting & boiling are ENDOTHERMIC  Freezing & condensing are EXOTHERMIC

21. 21  Changing temperature changes KE (#1, 3, 5)  Changing state changes potential energy (#2, 4)  energy to melt or freeze = heat of fusion (∆H fusion )  energy to vaporize or condense = heat of vaporization (∆H vaporization )

22. 22

23. 23

24. 24  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 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

25. 25  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)

26.  Heat of Fusion- change in energy when a solid substance melts  Heat of Vaporization- change in energy when liquid substance vaporizes (evaporates)

27.  Given that the heat of fusion (ice) is 6.0 kj/mol, how much energy is needed to melt an ice cube with a mass of 42 g?  Given that the heat of vaporization (water) is 41 kj/mol, how much energy is need to convert 24 g of water into steam?

28.  A quick look into the past…  At any given temp. molecules have the same avg. KE  Light molecules = fast  Heavy molecules = slow  Having the same avg. KE means some molecules are moving faster than avg., and some slower.

29.  Consider if you will…  A container of water at 25 o C (room temp.)  All the water molecules have the same avg. KE  Water is not hot, but a few fast molecules will leave and become water vapor.  Vapor molecules exert pressure = vapor pressure

30.  Now, consider this…  Water molecules can escape from open container, but not a closed container.  Energy is transferred by collisions w/ other vapor molecules.  Some vapor molecules will be slow enough to return to liquid.  Eventually, rate of evaporation = rate of condensing  This is called DYNAMIC EQUILIBRIUM

31.  Molecules continue to evaporate and condense.  # of vapor molecules and vapor pressure are constant.  A liquid in a closed container has a constant equilibrium vapor pressure.  Changing temp. changes equilibrium, and vapor pressure remains constant  at any given temp.

32.  Imagine this…  Open container and heat it.  KE increases, more molecules evaporate  Bubbles of water form throughout liquid, rise to surface and break… boiling begins

33.  As temp. increases, vapor pressure increases  A liquid boils when the vapor pressure matches the external pressure.  In an open container, external pressure is atmospheric pressure