1. Johannes Diderik van der Waals
Van der Waals Forces
Johannes Diderik van der Waals

2. Polarity Separation of charge
An asymmetrical difference in electronegativity along a bond or in a molecule

3. Circle the polar molecules. Label + and -Al Cl S      

4. Non-Polar 104.5o C. __________ molecules are symmetrical
 D. What is the bond angle in H2O? _______
 E. The motion of particles in these phases: Solid Liquid Gas
104.5o

5. Van der Waals Forces
Small, weak interactions between molecules

6. Van der Waals Forces Intermolecular: between molecules (not a bond)
Intramolecular: bonds within molecules (stronger)

7. What is being attracted?
+ attracted to -
 electrostatic attraction
e- s of one atom to another atom’s nucleus e- + + e-

8. Evidence of VDW Forces?
Non-polar molecules can form gases, liquids and solids.
Ex: CO2 C O C O C O

9. 3 Types of Van der Waals Forces
1)    dipole-dipole
2)    dipole-induced dipole
3) dispersion

10. Dipole-Dipole
Two polar molecules align so that + and - are matched (electrostatic attraction)
Ex: ethane (C2H6) vs. fluromethane (CH3F)

11. Fluoromethane (CH3F) – boiling point = 194.7 K polar or non-polar?
H H
H C F H C F - 
Dipole-Dipole
Ethane (C2H6) – boiling point = K
polar or non-polar?
H H H H
H C C H H C C H
NOT Dipole-Dipole

12. Try This:
Draw two KBr molecules and draw their dipole-dipole interactions with a dashed line. Br K

13. What does to “induce” mean?
To cause or bring about
Ex:
Induced vomiting
Induced labor
Induced coma

14. Dipole-Induced Dipole
A dipole can induce (cause)
a temporary dipole to form in a
non-polar molecule
The molecules then line up
to match + and - charges

15. Ar Example H Cl +    INDUCED DIPOLE non-polar A DIPOLE
(it’s polar)
Dipole – Induced Dipole
(weak and short-lived)

16. Draw CO2 (aq) What does (aq) mean? dissolved in WATER
So…draw CO2 (g) in H2O (l) O H C O C O      

17. Where is CO2 (aq) seen? Carbonated water CO2 is not very soluble…
1 CO2 in 1000 H2O molecules

18. Dispersion Forces A temporary dipole forms in a non-polar molecule…
which leads to…
a temporary dipole to form in ANOTHER non-polar molecule
Dispersion is the ONLY intermolecular attraction that occurs between non-polar molecules

19. (weakest and very short-lived)
Dispersion Forces e- e- e- e- e- e-
Cl-Cl e- e-
Cl-Cl e- e-   e- e-   e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e-
INDUCED
DIPOLE
TEMPORARY
DIPOLE
non-polar
non-polar
Dispersion
(weakest and very short-lived)

20. Tokay Gecko: Dispersion Forces!

21. Review Dipole – Dipole between two polar molecules
Dipole – Induced Dipole
b/w a polar & a non-polar molecule
Dispersion
between two non-polar molecules

22. Hydrogen Bonding STRONGEST Intermolecular Force!!
A special type of dipole-dipole attraction
Bonds form due to the polarity of water
Draw 3 H2O molecules in your notes
Ice
Liquid

23. Hydrogen Bonding con’t
Hydrogen bonds keep water in the liquid phase over a wider range of temperatures than is found for any other molecule of its size

24. Hydrogen bonds account for the high boiling point of water

25. Expansion of Ice
Ice expands when water freezes compared to most substances that contract when freezing
Ice bomb video

26. Denisty vs Temperature of H2O
4 oC—max density of water – liquid!
Solid
Ice
Liquid
water

27. Hexagonal Ice

28. Halos, Sundogs, & Pillars are caused by hexagonal ice crystals

29. Ponds Freezing
Solid water (ice) has a lower density than liquid water

31. Why is this good?
Ponds freeze from the top down, insulating the water below and keeping it from freezing solid
Without this, ponds would freeze solid and thaw more slowly

32. Surface Tension
Enhancement of the intermolecular attractive forces at the surface

33. Evidence
Lab:
Dixie cup
Penny
Capillary tube
needle

34. What causes surface tension?
The cohesive forces between molecules are shared with all neighboring atoms.
Since the surface has no neighboring atoms above, they exhibit stronger attractive forces for their neighbors next to and below them

35. Surface tension is a result of cohesive intermolecular forces

36. How many drops can you get on a penny?
Water?
TTE?
Why is there a difference???
Water has strong Hydrogen Bonds and TTE has weaker intermolecular forces

37. How is surface tension affected by soap?
Breaks the surface tension!

38. Capillary Rise
glass
gravity
H2O Hg
Water rises up the capillary tube because there are unbalanced forces between the water and glass and the water and gravity

39. Which is larger? Adhesion or Cohesion?
Adhesion: attraction between H2O (Hg) & glass
Cohesion: attraction of H2O (Hg) molec. to each other
Adhesion > Cohesion
Cohesion > Adhesion

40. Do other liquids exhibit capillary rise?
As long as they are attracted to glass and have enough cohesion

41. IM forces and interactions between liquids and surfaces
Cohesion > Adhesion
Liquid “Beads”
on Surface
Cohesion < Adhesion
Liquid “Wets”
the Surface

42. Evaporation
Diagram the distribution of kinetic energy at a temperature
5oC
25oC
75oC
# particles
low KE
ave KE
high KE

43. Which molecules will evaporate?
This lowers the total kinetic energy (temperature) of the entire system
Only high energy molecules can vaporize
# particles
low KE
ave KE
high KE

44. Boiling P atm P atm Pvap P atm Pvap = Patm Pvap t = 5 min t = 0 min

45. Boiling Boiling occurs when Vapor Pressure = Barometric Pressure
When Vapor Pressure = 760 mmHg, Boiling Point = 100oC

46. Evaporation Questions
Why do we sweat?
breaking water’s bonds has a cooling effect
high energy molecules are lost

47. 2. Why does water stay cool in clay containers?
Since clay is porous, high energy molecules escape leaving lower temperature water
When the water added to the sand evaporates in the Pot-in-Pot Cooler, it pulls heat from the smaller pot, keeping vegetables cool.
Refrigeration for the other 90% 47

48. 3. Why can liquid water change to vapor at room temperature?
High energy molecules escape
Evaporation occurs at all temperatures
# particles
low KE
ave KE
high KE

49. 4. Define vapor pressure
Force of particles leaving a liquid
Pressure of molecules in their bubbles
Can solids have a vapor pressure?
Yes! Solid Gas
Ex: ice, dry ice, plastics

50. 5. What is the difference between evaporation and boiling?
Evaporation: occurs at any temperature; high energy molecules escape
Boiling: occurs when atmospheric pressure = vapor pressure

51. Volatile Substances Easily evaporate Weak attractive forces
Low boiling point
High vapor pressure

52. Non-volatile substances
Do not easily evaporate
Strong attractive forces
High boiling point
Low vapor pressure

53. Equilibrium A + B C + D Forward Reaction Reverse Reaction
Rate of forward reaction =
Rate of reverse reaction

54. Dynamic Equilibrium Acetone (l) Acetone (g)
Reaction looks like it has stopped,
but is dynamic at the molecular level

55. What conditions are necessary for equilibrium?
Closed System
Rate of fwd rxn = rate of rev rxn
Constant temp, pressure, color
Both reactants and products are present (but not necessarily equal)

56. Henri Louis Le Chatlier (1850-1936)
Inventor of acetylene torch
Professor of Industrial Chemistry and Metallurgy
Instrumental in the development of cement and Plaster of Paris

57. LeChatlier’s Principle
When a stress is applied
to a system at equilibrium,
the system will respond
to partially undo the stress
Add Reactant, Add Product, Remove Reactant, Remove Product, Add Heat, Increase Pressure,…

58. Predicting adjustments
produced
Haber process
N H NH3 + energy
used
produced
used
Add energy
System wants?
Shift?
Amount of
N2 and H2?
Amount of NH3?
Remove NH3
Use energy
Produce NH3

59. 2 H+ + 2 CrO42- Cr2O72- + H2O produced used produced used Add HCl
System wants?
Shift?
Color?
Add NaOH
ORANGE
(Add H+)
Use H+
(Use H+)
Produce H+
YELLOW H+ H+ H+
Na+

60. 2 H+ + 2 CrO42- Cr2O72- + H2O Add H+ X = CrO4-2 O = Cr2O7-2 x x xx x x
x x xox xx
x x x xx xxxx
o x x x x
x x x x
ox o oo o
o o ooo o
o o x o oo oo
oo o ooo oo
Add OH-

61. 2 NO2 N2O4 + energy DARKER produced used produced used Add Heat
System wants?
Shift?
Color?
Remove Heat
Increase Pressure
DARKER
Use Heat
LIGHTER
Produce Heat
LIGHTER
Decrease Pr.

62. H2O (l) + energy H2O (g) produced used produced used Add Heat
System wants?
Shift?
Observation?
Remove Heat
Decrease Pressure
Increase Pressure
Use Heat
Evaporation
Produce Heat
Condensation
Evaporation
Increase Pr.
Decrease Pr.
Condensation

63. How Do Pressure Cookers Work?
Pressure cookers increase the pressure above the water so that water boils at a ________ temperature and cooks food ________
HIGHER
QUICKER

64. Lab Practice Problem NaCl Na+ + Cl-
a) Which direction would the reaction shift if MgCl2 (Mg2+ and Cl-) were added to the system above? Explain.
b) What would happen to the amount of NaCl if Cl- were removed from the system? Explain.
Cl-
Cl-
Na+
Cl-
Cl-
Cl-
Cl-
Na+
Cl-
Na+
Na+
Cl-
Na+
NaCl
NaCl
NaCl
NaCl
NaCl
NaCl

65. Phase Changes Temperature (oC) 105 KE PE 100 KE PE KE
KE
Where is there a KE?
Where is there a PE?
- 5
Time

66. Terms Melting Point Temp when substances changes from l  s
Boiling point
Temp when substance changes from l  g
KE—
where there is a change in temperature
PE—
where there’s a phase change
(constant temp)

67. Calculations
Calculate the amount of heat needed to raise the temperature of 100 ml of water from 15oC to 65oC.
Q = mcT
Q = (100g)(1 cal/goC)(50oC)
Q = 5000 cal

68. = 3. Calculate the amount of heat needed to melt 100 g of ice.
REMEMBER: Heat of Fusion = 80 cal/g
80 cal
x cal
x = 8000 cal =
1 g
100 g

69. = 2. Calculate the amount of heat needed to boil 100 ml of water.
HEAT OF VAPORIZATION = 540 cal/g
540 cal
x cal
x = 54,000 cal =
1 g
100 g