Intermolecular Forces, Gases, and Liquids

1. Intermolecular Forces, Gases, and Liquids Ch.13

2. Gases • Kinetic-Molecular Theory says molecules/atoms separated • Little, if any, interactions • Not so in solids and liquids • Examples: • Big difference in volume between liquids & solids and gases • Gases compressible, liqs & solids not

3. Intermolecular Forces • Various electrostatic forces that attract molecules in solids/liqs • Much weaker than ionic forces • About 15% (or less) that of bond energies • Remember, ionic bonds extremely powerful • Boiling pt of NaCl = 1465 °C!

4. Intermolecular Forces • Reason behind importance of knowing about IMF: • 1) b.p. & m.p. and heats of vaporization (lg) and fusion (sl) • 2) solubility of gases, liquids, and solids • 3) determining structures of biochemicals (DNA, proteins)

5. Remember dipole moments? • Dipole moment = product of magnitude of partial charges (+/-) & their distance of separation • = (1 Debye = 3.34 x 10-30 C x m) • Important in IMF

6. Ion-dipole: Ionization in aqueous medium (water) • 1) stronger attraction if ion/dipole closer • Li+ vs. Cs+ in water • 2) higher ion charge, stronger attraction • Be2+ vs. Li+ in water • 3) greater dipole, stronger attraction • Dissolved salt has stronger attraction to water than methanol

8. Solvation energy • Or, enthalpy of hydration (if water) = energy of ionization in aq. media • Water molecules surround both ions • Example: • Take hydration energies of G I metal ions • Exothermicity decreases as you go down the column • Cations become larger • Easier to dissociate

9. Permanent dipoles • Positive end of one molecule attracted to negative end of other • For ex: HCl • Dipole-dipole attractions • Cmpds that exhibit greater d-d attractions have higher b.p., and Hvap • Polar cmpds exhibit greater d-d attractions than non-polar cmpds • NH3 vs. CH4 •  equivalent molar masses (g/mol): 17 vs. 16, respectively • Boiling points: -33°C vs. -162°C, respectively

10. Hydrogen Bonding • A type of “super” dipole-dipole interaction • Interaction between e--rich atom connected to H entity & another H attached to e—rich atom • e--rich atom = O, F, N • Density water > than ice • Opposite of almost every other substance • Inordinately high heat capacity of water • High surface tension • Insects walk on water • Concave meniscus

11. Hydrogen Bonding • Boiling pts. of H2O, HF, and NH3much higher

12. Surface Tension • Outer molecules interact with surface, while inner interact with other molecules • It has a “skin” • Skin toughness = surface tension • Energy required to break through surface • Smaller surface area reason that water drops spherical

13. Capillary Action • When water goes up a small glass tube • Due to polarity of Si-O bonding with water • Adhesive forces > cohesive forces of water • Creates a chain or bridge • Pulls water up tube • Limited by balancing gravity with adhesive/cohesive forces • Thus, water has a concave meniscus

14. Mercury • Forms a convex meniscus • Doesn’t “climb” a glass tube • Due to cohesive forces > adhesive forces

15. Viscosity • Hydrogen-bonding increases viscosity • But large non-polar liquids like oil have: • 1) large unwieldy molecules w/greater intermolecular forces • 2) greater ability to be entangled w/one another • Did you ever hear the expression, “You’re as slow as molasses in January”?

16. Dipole/Induced Dipole Forces • Polar entities induce dipole in nonpolar species like O2 • O2 can now dissolve in water • If not, fishes in trouble! • Process called “polarization” • Generally, higher molar mass, greater polarizability of molecule • Why? • (larger the species, more likely e- held further away  easier to polarize)

17. Polarizability

18. Induced dipole/induced dipole forces • Non-polar entities can cause temporary dipoles between other non-polar entities •  causing intermolecular attractions • Pentane, hexane, etc. • The higher the molar mass, the greater the intermolecular attractions • N-pentane has greater interactions than neo-pentane • Latter’s smaller area for interactions • I2 has a higher Hvap & b.p. than other halogens • cause nonpolar substances to condense to liquids • and to freeze into solids • (when the temperature is lowered sufficiently) • Also called: London Dispersion Forces

19. Intermolecular Bonding Compared • Strength • Strongest: Ion-dipole • Strong: Dipole-dipole (incl. H-bonding) • Less strong: dipole/induced-dipole • Least strong: induced-dipole/induced-dipole (London dispersion forces) • Keep in mind  a compound can have more than one of the above!

20. Problem • Rank the following in order of increasing boiling point and explain why: • NH3, CH4, and CO2

21. Properties of Liquids • (l)  (g) • Vaporization = endothermic • Condensation = exothermic • Boiling • Why do we have bubbles?

22. Vapor Pressure • Leave a bottle of water open…. • Will evaporate • Keep the lid on…. • can have equilibrium between liquid and gas • Equilibrium vapor pressure/vapor pressure • Measure of tendency of molecules to vaporize at given temp.

23. What does this graph tell us?

24. Volatility • Ability of liquid to evaporate • Higher the vapor pressure, greater the volatility • Are polar cmpds or non-polar cmpds of equal molecular mass more volatile?

25. Clausius-Clapeyron Equation • Calculates ∆Hvap • What is this an equation for? • What are the variables? • C = constant unique to cmpd • R = ideal gas constant • 8.314472 J/molK

26. Clausius-Clapeyron Equation • Or, if given two pts.:

27. Clausius-Clapeyron Problem • Methanol has a normal boiling point of 64.6°C and a heat of vaporization of 35.2 kJ/mol. What is the vapor pressure of methanol at 12.0°C? • Does the answer make sense? • Would water have a higher heat of vaporization? • Why? • Heat of vaporization of water = 40.65 kJ/mol

28. Boiling Point • Bp  temp. at which vapor pressure = external (atmospheric pressure) • At higher elevations atmospheric pressure is lower • Thus, water boils at less than 100 °C

29. Critical Temperature and Pressure • As temp. rises so does vapor pressure, but not infinitely • At the critical point liq/gas interface disappears • Critical temp/pressure • Tc/Tp • Gives supercritical fluid • Density of a liq • Viscosity of gas • H2O: • Tc = 374 °C • Tp = 217.7 atm! • Normal earth pressure  1 atm

30. Supercritical fluid • CO2 used in decaffeinating coffee • Read about it on page 614

31. Phase diagram • Gives info on phase states of a substance at varying pressures and temperatures

32. Deciphering a phase diagram • Triple point • Where all 3 states coexist • Curves denote existence of two states • Fusion (solid & liq) • Vaporization (liq & gas) • Sublimation (solid & gas) • Off the lines • Single state

33. Water’s phase diagram • Graph explains why water boils at lower temps at higher altitudes (next slide) • If you apply increasing pressure (const. T of 0°C) to ice will it convert to water? • Solid-liquid line has negative slope • It’s the opposite of most species • Why?

34. Sublimation • Going from solid to gas without going through the liquid state • Enthalpy of sublimation • Hsublimation • Iodine & dry ice (solid CO2) sublimate • Opposite of sublimation • Deposition (gs) • Iodine demo

35. CO2’s Phase Diagram • Explains sublimation • How? • Why is it called “dry ice”?

36. Iodine’s Phase Diagram: But does it really sublimate?

37. Problem • The normal melting and boiling points of xenon are -112°C and -107°C, respectively. • Its triple point is a -121°C and 0.371 atm and its critical point is at 16.6°C and 57.6 atm. • a) Sketch the phase diagram for Xe, showing the axes, the four points given above, and indicating the area in which each phase is stable. • b) If Xe gas is cooled under an external pressure of 0.131 atm, will it undergo condensation or deposition?