1. © 2006 Brooks/Cole - Thomson Inter- molecular Forces Have studied INTRA molecular forces— the forces holding atoms together to form molecules. Now turn to forces between molecules — INTER molecular forces. Forces between molecules, between ions, or between molecules and ions.

2. © 2006 Brooks/Cole - Thomson Ion-Ion Forces for comparison of magnitude Na + —Cl - in salt These are the strongest forces. Lead to solids with high melting temperatures. NaCl, mp = 800 o C MgO, mp = 2800 o C

3. © 2006 Brooks/Cole - Thomson Covalent Bonding Forces for comparison of magnitude C–H, 413 kJ/mol C=C, 610 kJ/mol C–C, 346 kJ/mol CN, 887 kJ/mol

4. © 2006 Brooks/Cole - Thomson Attraction Between Ions and Permanent Dipoles Water is highly polar and can interact with positive ions to give hydrated ions in water.

5. © 2006 Brooks/Cole - Thomson Attraction Between Ions and Permanent Dipoles Water is highly polar and can interact with positive ions to give hydrated ions in water.

6. © 2006 Brooks/Cole - Thomson Attraction Between Ions and Permanent Dipoles Many metal ions are hydrated. This is the reason metal salts dissolve in water.

7. © 2006 Brooks/Cole - Thomson Attraction between ions and dipole depends on ion charge and ion-dipole distance. Measured by ∆H for M n+ + H 2 O --> [M(H 2 O) x ] n+ -1922 kJ/mol -405 kJ/mol -263 kJ/mol Attraction Between Ions and Permanent Dipoles

8. © 2006 Brooks/Cole - Thomson Dipole-Dipole Forces Such forces bind molecules having permanent dipoles to one another.

9. © 2006 Brooks/Cole - Thomson Dipole-Dipole Forces Influence of dipole-dipole forces is seen in the boiling points of simple molecules. CompdMol. Wt.Boil Point N 2 28-196 o C CO28-192 o C Br 2 16059 o C ICl16297 o C

10. © 2006 Brooks/Cole - Thomson Hydrogen Bonding A special form of dipole-dipole attraction, which enhances dipole-dipole attractions. H-bonding is strongest when X and Y are N, O, or F

11. © 2006 Brooks/Cole - Thomson Hydrogen Bonding in H 2 O H-bonding is especially strong in water because the O—H bond is very polarthe O—H bond is very polar there are 2 lone pairs on the O atomthere are 2 lone pairs on the O atom Accounts for many of water’s unique properties.

12. © 2006 Brooks/Cole - Thomson Hydrogen Bonding in H 2 O Ice has open lattice-like structure. Ice density is < liquid. And so solid floats on water. Snow flake: www.snowcrystals.com

13. © 2006 Brooks/Cole - Thomson Hydrogen Bonding in H 2 O Ice has open lattice-like structure. Ice density is < liquid and so solid floats on water. One of the VERY few substances where solid is LESS DENSE than the liquid.

14. © 2006 Brooks/Cole - Thomson Hydrogen Bonding in H 2 O H bonds ---> abnormally high specific heat capacity of water (4.184 J/gK) This is the reason water is used to put out fires, it is the reason lakes/oceans control climate, and is the reason thunderstorms release huge energy.

15. © 2006 Brooks/Cole - Thomson Hydrogen Bonding H bonds leads to abnormally high boiling point of water.

16. © 2006 Brooks/Cole - Thomson Boiling Points of Simple Hydrogen-Containing Compounds Active Figure 13.8

17. © 2006 Brooks/Cole - Thomson Hydrogen Bonding in Biology H-bonding is especially strong in biological systems — such as DNA. DNA — helical chains of phosphate groups and sugar molecules. Chains are helical because of tetrahedral geometry of P, C, and O. Chains bind to one another by specific hydrogen bonding between pairs of Lewis bases. —adenine with thymine —adenine with thymine —guanine with cytosine —guanine with cytosine

18. © 2006 Brooks/Cole - Thomson Portion of a DNA chain Double helix of DNA

19. © 2006 Brooks/Cole - Thomson Base-Pairing through H-Bonds

20. © 2006 Brooks/Cole - Thomson Base-Pairing through H-Bonds

21. © 2006 Brooks/Cole - Thomson FORCES INVOLVING INDUCED DIPOLES How can non-polar molecules such as O 2 and I 2 dissolve in water? The water dipole INDUCES a dipole in the O 2 electric cloud. Dipole-induced dipole

22. © 2006 Brooks/Cole - Thomson FORCES INVOLVING INDUCED DIPOLES Solubility increases with mass the gas

23. © 2006 Brooks/Cole - Thomson FORCES INVOLVING INDUCED DIPOLES Process of inducing a dipole is polarizationProcess of inducing a dipole is polarization Degree to which electron cloud of an atom or molecule can be distorted in its polarizability.Degree to which electron cloud of an atom or molecule can be distorted in its polarizability.

24. © 2006 Brooks/Cole - Thomson IM FORCES — INDUCED DIPOLES Consider I 2 dissolving in ethanol, CH 3 CH 2 OH. O H -  +  I-I R -  +  O H +  -  R The alcohol temporarily creates or INDUCES a dipole in I 2.

25. © 2006 Brooks/Cole - Thomson FORCES INVOLVING INDUCED DIPOLES Formation of a dipole in two nonpolar I 2 molecules. Induced dipole-induced dipole The induced forces between I 2 molecules are very weak, so solid I 2 sublimes (goes from a solid to gaseous molecules).

26. © 2006 Brooks/Cole - Thomson FORCES INVOLVING INDUCED DIPOLES The magnitude of the induced dipole depends on the tendency to be distorted. Higher molec. weight ---> larger induced dipoles. MoleculeBoiling Point ( o C) MoleculeBoiling Point ( o C) CH 4 (methane) - 161.5 CH 4 (methane) - 161.5 C 2 H 6 (ethane)- 88.6 C 2 H 6 (ethane)- 88.6 C 3 H 8 (propane) - 42.1 C 3 H 8 (propane) - 42.1 C 4 H 10 (butane) - 0.5 C 4 H 10 (butane) - 0.5

27. © 2006 Brooks/Cole - Thomson Boiling Points of Hydrocarbons Note linear relation between bp and molar mass. CH 4 C2H6C2H6C2H6C2H6 C3H8C3H8C3H8C3H8 C 4 H 10

28. © 2006 Brooks/Cole - Thomson Intermolecular Forces Summary

29. © 2006 Brooks/Cole - Thomson Measuring Equilibrium Vapor Pressure Liquid in flask evaporates and exerts pressure on manometer. Active Fig. 13.17

30. © 2006 Brooks/Cole - Thomson Equilibrium Vapor Pressure Active Figure 13.18

31. © 2006 Brooks/Cole - Thomson Liquids Equilibrium Vapor Pressure FIGURE 13.18: VP as a function of T. 1. The curves show all conditions of P and T where LIQ and VAP are in EQUILIBRIUM 2. The VP rises with T. 3. When VP = external P, the liquid boils. This means that BP’s of liquids change with altitude. This means that BP’s of liquids change with altitude. 4. If external P = 760 mm Hg, T of boiling is the NORMAL BOILING POINT 5. VP of a given molecule at a given T depends on IM forces. Here the VP’s are in the order C 2 H 5 H 5 C 2 H H 5 C 2 H Hwateralcoholether increasing strength of IM interactions extensive H-bondsH-bonds dipole- dipole O O O

32. © 2006 Brooks/Cole - Thomson Liquids HEAT OF VAPORIZATION is the heat req’d (at constant P) to vaporize the liquid. LIQ + heat ---> VAP Compd.∆H vap (kJ/mol) IM Force H 2 O40.7 (100 o C)H-bonds SO 2 26.8 (-47 o C)dipole Xe12.6 (-107 o C)induced dipole

33. © 2006 Brooks/Cole - Thomson Phase Diagrams

34. © 2006 Brooks/Cole - Thomson TRANSITIONS BETWEEN PHASES Section 13.10 Lines connect all conditions of T and P where EQUILIBRIUM exists between the phases on either side of the line. (At equilibrium particles move from liquid to gas as fast as they move from gas to liquid, for example.)

35. © 2006 Brooks/Cole - Thomson Phase Diagram for Water Solid phase Liquid phase Gas phase

36. © 2006 Brooks/Cole - Thomson Phase Equilibria — Water Solid-liquid Gas-Liquid Gas-Solid

37. © 2006 Brooks/Cole - Thomson Triple Point — Water At the TRIPLE POINT all three phases are in equilibrium.

38. © 2006 Brooks/Cole - Thomson Phases Diagrams— Important Points for Water T(˚C)P(mmHg) Normal boil point 100 760 Normal freeze point0 760 Triple point 0.0098 4.58

39. © 2006 Brooks/Cole - Thomson Critical T and P Above critical T no liquid exists no matter how high the pressure. As P and T increase, you finally reach the CRITICAL T and P At P < 4.58 mmHg and T < 0.0098 ˚C solid H 2 O can go directly to vapor. This process is called SUBLIMATION This is how a frost-free refrigerator works.

40. © 2006 Brooks/Cole - Thomson Critical T and P COMPDT c ( o C)P c (atm) H 2 O374218 CO 2 3173 CH 4 -8246 Freon-1211241 (CCl 2 F 2 ) Notice that T c and P c depend on intermolecular forces.

41. © 2006 Brooks/Cole - Thomson Solid-Liquid Equilibria In any system, if you increase P the DENSITY will go up. Therefore — as P goes up, equilibrium favors phase with the larger density (or SMALLER volume/gram). Liquid H 2 OSolid H 2 O Liquid H 2 OSolid H 2 O Density1 g/cm 3 0.917 g/cm 3 cm 3 /gram11.09

42. © 2006 Brooks/Cole - Thomson Solid-Liquid Equilibria Raising the pressure at constant T causes water to melt. The NEGATIVE SLOPE of the S/L line is unique to H 2 O. Almost everything else has positive slope. Solid H 2 O Liquid H 2 O P T 760 mmHg 0 °C Normal freezing point The behavior of water under pressure is an example of LE CHATELIER’S PRINCIPLE At Solid/Liquid equilibrium, raising P squeezes the solid. It responds by going to phase with greater density, i.e., the liquid phase.

43. © 2006 Brooks/Cole - Thomson CO 2 Phase Diagram

44. © 2006 Brooks/Cole - Thomson CO 2 Phases Separate phases Increasing pressure More pressure Supercritcal CO 2 Figure 13.41

45. © 2006 Brooks/Cole - Thomson Dissolving An Ionic Solid Active Figure 14.9

46. © 2006 Brooks/Cole - Thomson Energetics of the Solution Process Figure 14.8

47. © 2006 Brooks/Cole - Thomson Energetics of the Solution Process If the enthalpy of formation of the solution is more negative that that of the solvent and solute, the enthalpy of solution is negative. The solution process is exothermic !

48. © 2006 Brooks/Cole - Thomson Colligative Properties On adding a solute to a solvent, the props. of the solvent are modified. Vapor pressure decreasesVapor pressure decreases Melting point decreasesMelting point decreases Boiling point increasesBoiling point increases Osmosis is possible (osmotic pressure)Osmosis is possible (osmotic pressure) These changes are called COLLIGATIVE PROPERTIES. They depend only on the NUMBER of solute particles relative to solvent particles, not on the KIND of solute particles.

49. © 2006 Brooks/Cole - Thomson Concentration Units MOLE FRACTION, X For a mixture of A, B, and C WEIGHT % = grams solute per 100 g solution MOLALITY, m

50. © 2006 Brooks/Cole - Thomson Calculating Concentrations Dissolve 62.1 g (1.00 mol) of ethylene glycol in 250. g of H 2 O. Calculate mol fraction, molality, and weight % of glycol.

51. © 2006 Brooks/Cole - Thomson Understanding Colligative Properties To understand colligative properties, study the LIQUID- VAPOR EQUILIBRIUM for a solution.

52. © 2006 Brooks/Cole - Thomson P solvent = X solvent P o solvent Understanding Colligative Properties VP of H 2 O over a solution depends on the number of H 2 O molecules per solute molecule. P solvent proportional to X solvent P solvent proportional to X solvent VP of solvent over solution = (Mol frac solvent)(VP pure solvent) RAOULT’S LAW RAOULT’S LAW

53. © 2006 Brooks/Cole - Thomson P A = X A P o A Raoult’s Law An ideal solution is one that obeys Raoult’s law. Because mole fraction of solvent, X A, is always less than 1, then P A is always less than P o A. The vapor pressure of solvent over a solution is always LOWERED!

54. © 2006 Brooks/Cole - Thomson Raoult’s Law Assume the solution containing 62.1 g of glycol in 250. g of water is ideal. What is the vapor pressure of water over the solution at 30 o C? (The VP of pure H 2 O is 31.8 mm Hg; see App. E.)

55. © 2006 Brooks/Cole - Thomson Changes in Freezing and Boiling Points of Solvent For a 2-component system where A is the solvent and B is the soluteFor a 2-component system where A is the solvent and B is the solute ∆PA = VP lowering = XBPoA ∆PA = VP lowering = XBPoA VP lowering is proportional to mol frac solute!VP lowering is proportional to mol frac solute! For very dilute solutions, ∆PA = KmolalityB where K is a proportionality constant.For very dilute solutions, ∆PA = KmolalityB where K is a proportionality constant. This helps explain changes in melting and boiling points.This helps explain changes in melting and boiling points.

56. © 2006 Brooks/Cole - Thomson Vapor Pressure Lowering Figure 14.14

57. © 2006 Brooks/Cole - Thomson The boiling point of a solution is higher than that of the pure solvent.

58. © 2006 Brooks/Cole - Thomson Elevation of Boiling Point Elevation in BP = ∆T BP = K BP m (where K BP is characteristic of solvent)

59. © 2006 Brooks/Cole - Thomson Change in Boiling Point Dissolve 62.1 g of glycol (1.00 mol) in 250. g of water. What is the BP of the solution? K BP = +0.512 o C/molal for water (see Table 14.3).

60. © 2006 Brooks/Cole - Thomson Change in Freezing Point The freezing point of a solution is LOWER than that of the pure solvent. FP depression = ∆T FP = K FP m Pure water Ethylene glycol/water solution

61. © 2006 Brooks/Cole - Thomson Lowering the Freezing Point Water with and without antifreezeWhen a solution freezes, the solid phase is pure water. The solution becomes more concentrated.

62. © 2006 Brooks/Cole - Thomson Calculate the FP of a 4.00 molal glycol/water solution. K FP = -1.86 o C/molal (Table 14.4) Freezing Point Depression

63. © 2006 Brooks/Cole - Thomson How much NaCl must be dissolved in 4.00 kg of water to lower FP to -10.00 o C?. Freezing Point Depression

64. © 2006 Brooks/Cole - Thomson How much NaCl must be dissolved in 4.00 kg of water to lower FP to -10.00 o C?. Freezing Point Depression

65. © 2006 Brooks/Cole - Thomson Boiling Point Elevation and Freezing Point Depression ∆T = Kmi ∆T = Kmi A generally useful equation i = van’t Hoff factor = number of particles produced per formula unit. CompoundTheoretical Value of i glycol1 NaCl2 CaCl 2 3

66. © 2006 Brooks/Cole - Thomson Osmosis The semipermeable membrane allows only the movement of solvent molecules. Solvent molecules move from pure solvent to solution in an attempt to make both have the same concentration of solute. Driving force is entropy

67. © 2006 Brooks/Cole - Thomson Osmosic Pressure, ∏ Equilibrium is reached when pressure — the OSMOTIC PRESSURE, ∏ — produced by extra solution counterbalances pressure of solvent molecules moving thru the membrane. ∏ = cRT (c is conc. in mol/L) Osmotic pressure

68. © 2006 Brooks/Cole - Thomson Osmosis and Living Cells

69. © 2006 Brooks/Cole - Thomson Reverse Osmosis Water Desalination Water desalination plant in Tampa

70. © 2006 Brooks/Cole - Thomson Osmosis Calculating a Molar Mass Dissolve 35.0 g of hemoglobin in enough water to make 1.00 L of solution. ∏ measured to be 10.0 mm Hg at 25 ˚C. Calculate molar mass of hemoglobin.