1. Intermolecular Forces
Chapter 11

2. Intermolecular Forces: Introduction
Forces between separate molecules and dissolved ions (not bonds)
“Van der Waals Forces”
15% as strong as covalent or ionic bonds

3. Intermolecular Forces: Introduction
Low temperature – strong
High temperature – kinetic energy of motion overcomes the IMF
Boiling point is a good indicator
Stronger IMF = higher boiling point
Weaker IMF = lower boiling point

4. Ion-Ion
Full to Full Charge
Ion-Dipole
Full to Partial Charge
Hydrogen Bonds
Partial to Partial Charge (H involved)
Dipole-Dipole
Partial to Partial Charge (No H)
London Dispersion Forces
Non-polar to Non-Polar

5. Predict what type of IMF would form between:
Br2 and I2
KCl and water
Water and ammonia
Two SO2 molecules
NaCl ions in a crystal

6. Type 1: Ion-Ion Forces Full Charges to full charges
High melting points (ionic solids)
Ex: Melting Point
Na ~98 oC
NaCl ~800 oC

7. Type 2: Ion-Dipole Forces
Full Charges to partial charges
Very Strong

10. Type 3: Hydrogen Bonds
Stronger than dipole-dipole that do not have hydrogen (no inner electrons, strong + pull)
Generally involves hydrogen and O, N or F
R-H · · · · O-R
R-H · · · · N-R
R-H · · · · F-R

11. Miscibility

12. Miscibility “Like dissolves like”
Substances that can hydrogen bond dissolve in one another.

13. Glucose and other sugars

14. Water Beading

15. Ice

17. Ice

18. DNA

19. DNA is TWO molecules that are hydrogen bonded (like a zipper)

20. Boiling Point Generally increases with increasing molar mass
H2O unusually high - H-bonding

21. Type 4: Dipole-Dipole Slightly weaker IMF
Involve d+ and d- charges other than those in hydrogen bonding

22. Draw Lewis Dot Structures to explain the following boiling points
MM (amu) BP (oC)
CH3CH2CH
CH3CHO
CH3CN

23. Type 5: London Dispersion Forces
Very weak IMF
Caused by temporary imbalances in electrons

24. London Forces: Inorganic Molecules
More electrons, more chance for temporary dipole
Boiling Point Table
Halogen
Molar Mass
BP(oC)
Noble Gas
Atomic Mass
F2 (g)
38.0
-188 He
4.0
-268
Cl2 (g)
71.0
-35 Ne
20.2
-246
Br2 (l)
159.8 59 Ar
39.9
-186
I2 (s)
253.8
185 Kr
83.8
-152

25. Explain the differences in boiling point between Cl2 (-35oC) and Ar (-186oC)

26. London Forces: Organic Compounds
The longer the carbon chain, the higher the London Dispersion Forces (the higher the melting point and boiling point)
Chainlike molecules greater London Forces than “bunched up” molecules (branched)

28. Ex 1
Rank the following compounds in terms of increasing melting points: NaCl, CF4, CH3OH.

29. Ex 2
Separate the following compounds by whether they have dipole-dipole attractions (including H-bonding) or London Forces. Which should have the highest dipole-dipole attraction? Which should have the strongest London Force? Br2, Ne, HCl, N2, HF

30. Ex 3
Rank the following in order of increasing boiling point: BaCl2, H2, CO, HF, Kr

31. Ex 4
Rank the following in order of increasing boiling point: N2, KBr, O2, CH3CH2OH, HCN

32. Properties of Liquid Viscosity – resistance of a liquid to flow
Oil is more viscous than water
Water has H-bonds (stronger)
Oil has London forces (weaker, but there are many more of them, long carbon chain)

33. 40 W oil 10 W oil

34. Miscibility

35. Surface Tension

36. Surface Tension

37. Capillary Action Water is attracted to the glass
Mercury more attracted to itself

38. Heating Curves Changes of state do not have a temperature change.
Melting/Freezing
Boiling/Condensing
A glass of soda with ice will stay at 0oC until all of the ice melts.
Graph “flattens out” during changes of state

39. Heating Curves Steam heats up Temperature (oC) Boiling Water warms up
Melting
Ice warms up
Heat (Joules)

40. Heating Curve
No phase change is occurring (heating ice, water, or steam):
q = mCpDT
Melting or boiling:
q = mLf or q = mLv

41. Lf and Lv Latent Heat - heat for phase changes. No temperature change.
Lf –latent heat of freezing/melting
Lv –latent heat of boiling/condensing

42. Heating Curves Temperature (oC) Use q = mLv Boiling Use q = mLf
Melting
Use q = mCpDT
Heat (calories)

43. Important Values Substance Cp Steam 2.01 J/goC Water 4.18 J/goC
Ice J/goC
Latent Heat of fusion (water) Lf = J/g
Latent Heat of vaporization(water) Lv = J/g

44. Heating Curves: Example 1
How much energy must be removed to cool grams of water at 20.0oC to make ice at
–10.0oC?

45. Heating Curves: Example 1
Temperature (oC)
Melting (q=mLf)
Water cools (q=mCpDT)
Ice cools (q=mCpDT)
Heat

46. Heating Curves: Example 1
Cooling the water
q = mCpDT = (100 g)(4.18 J/goC)(0oC-20oC)
q = J (8.36 kJ) (ignore the negative sign for now)
Freezing the water
q = mLf = (100.0 g)(334.7 J/g)= kJ
3. Cooling the ice down to –10.0oC
q = mCpDT = (100 g)(2.09 J/goC)(-10oC-0oC)
q = 2.09 kJ (we will ignore the negative sign for now)
8.36 kJ kJ kJ= kJ

47. Ex 2
How much energy is needed to convert 18.0 grams of ice at -25oC to steam at 125oC? ANS: 56.0 kJ

48. Ex 3
How much energy must be used to convert grams of water from steam at 110.0oC to ice at -25.0oC ? (309 kJ)

49. Vapor Pressure Pressure of a gas above a liquid caused by that liquid
Temperature – measure of the average kinetic energy of molecules
At any given moment, some molecules have enough energy to escape
EX: Even cold water will evaporate

52. Volatility Volatile – liquids that evaporate easily
Acetone
Often weak intermolecular forces
Boiling point – point at which the vapor pressure of a liquid = vapor pressure of the atmosphere
Normal Boiling Point – vapor pressure = 1 atm
Steam pressure cookers – Forces water to boil at a higher temperature
High Altitude – water boils at a lower temperature

53. Four Types of Solids Molecular Solids (single molecules)
Covalent Network Solids (one large molecule)
Ionic Solids
Metallic Solids

54. 1. Molecular Solids
Held together by IMF
Ice
Plastics

55. 2. Covalent Network Solids
Basically one big molecule
Held together by covalent bonds
Diamond
Graphite
Quartz (SiO2)

57. 3. Ionic Solids Held together by electrostatic attraction (Ion:Ion)
Usually crystalline (unit cells)

58. 4. Metallic Solids Atoms share electrons very freely
Positive nuclei in a “sea of electrons”
Electrons held loosely
Conducts electricity
Photoelectric effect
Malleable and ductile

59. Rank by boiling point (low to high)
Rank by boiling point (low to high). Below each, tell me which IMF is important: CH3OH Cl2 N2 CH3Cl CH3CH2CH2CH3 CH3CH2CH2OH Draw Lewis Dots for these CH3OCH3 CH3CH2CH3

60. How much energy must be removed to convert 100.0 grams of steam at 120.0oC to ice at -5.00oC?

61. (100. 0g)(2. 01 J/gK)(20. 0oC) = 4. 02 kJ (100. 0g)(2259. 4 J/g) = 225
(100.0g)(2.01 J/gK)(20.0oC) = 4.02 kJ (100.0g)( J/g) = kJ (100.0g)(4.18 J/gK)(100.0oC) = 41.8 kJ (100.0g)(334.7 J/g) = kJ (100.0g)(2.09 J/gK)(5.0oC) = 1.05 kJ kJ

62. The heat of vaporization of ammonia is 23. 35 kJ/mol
The heat of vaporization of ammonia is kJ/mol. How many grams of ammonia must evaporate to freeze grams of water initially at 10.00oC? Assume the ice stays at 0oC when frozen, and is used to cool a cup of too hot tea. (Ans: 27.4 g)

63. Water qcool (100)(4. 18)(10) = 4. 18 kJ qfreeze (100)(334. 7) = 33
Water qcool (100)(4.18)(10) = 4.18 kJ qfreeze (100)(334.7) = kJ kJ Ammonia qv = mLv m =qv/Lv m = kJ/23.35 kJ/mol = 1.61 mol mass = (1.61 mol)(17.0 g/mol) = 27.4 g

64. a) H-bonding, (b)London (c) Ion-dipole (d) dipole-dipole
a) H-bonding, (b)London (c) Ion-dipole (d) dipole-dipole. Ion-dipole and h-bonding are stronger
10. a) Solids = attractive forces (IMF) win
Liquids = Balance
Gases = Kinetic energy wins
b) Increasing T increase KE, eventually overcoming IMF
c) High pressure forces gas molecules clsoe together and IMF’s can win

65. a) Distance greater in liquid state
b) More movement, more volume, lower density
14. Overall, net forces are attractive
16.a) CH3OH has h-bonding, CH3SH does not
b) Xe is heavier, greater London Forces
c) Cl2 more polarizable than Kr
d) Acetone has dipole-dipole forces
18. a) True b) False c) False d) True
20. a) Br2 b) C5H11SH c) CH3CH2CH2Cl

66. 22. Propyl alcohol is longer and more polarizable 24
22. Propyl alcohol is longer and more polarizable 24. a) HF has hydrogen bonds, HCl dipole/dipole b) CHBr3 higher molar mass, more dispersion c) ICl has dipole-dipole, Br2 only dispersion 26.a) Dispersion, C8H18 higher boiling point b) C3H8(dispersion) CH3OCH3 (dip-dip) c) HOOH (h-bonding) HSSH (dip-dip) d) NH2NH2 (h-bonding) CH3CH3 (dispersion) 32. H2NNH2, HOOH, H2O can all h-bond

67. a) Exo b) Endo c) Endo d) Exo
g CCl2F2 kJ