1. Chapter 11 Intermolecular Attractive Forces, Liquids, and Solids
John D. Bookstaver
St. Charles Community College
St. Peters, MO
 2006, Prentice Hall, Inc.
“Tweaked”
by Robert Hernandez
Seabreeze Highschool
2015

2. What conditions are best for an “Ideal” gases?
HIGH P, LOW T
HIGH P, HIGH T
LOW P, LOW T
LOW P, HIGH T ?
Remember Mercury!

3. States of Matter
Draw in the circles representing solid and liquid in your notes. g s l
gas
liquid
solid
The fundamental difference between states of matter is the distance between particles.

4. The States of Matter 1. The KE of the particles and…
The state at a particular temperature and pressure depends on two opposite qualities:
1. The KE of the particles and…
2. The strength of attractions between them

5. The States of Matter (property review)
Gas
Liquid
Solid
Assumes SHAPE of container
Yes No
Assumes VOLUME of container
No…*
Compressible?
Flows?
Diffusion speed?
Fast
Slow
Zero
Fill in the boxes in your notes, correct the mistakes, when revealed
* Assume the shape of the portion of the container they occupy.

6. Intermolecular Attractive Forces
The attractions between molecules are not nearly as strong as the intramolecular attractions (bonds) holding compounds together.

7. Intermolecular Attractive Forces
This means that – if stressed…
It’s the INTERmolecular bonds that will sever first.. 7

8. Intermolecular Attractive Forces
They are, however, strong enough to determine physical properties such as;
boiling and melting points,
vapor pressure,
viscosity.

9. Intermolecular Attractive Forces
These intermolecular forces as a group are referred to as van der Waals forces.

10. 1. London dispersion forces 2. Dipole-dipole interactions
van der Waals Forces
1. London dispersion forces
2. Dipole-dipole interactions
3. Hydrogen bonding
4. Ion-dipole
(non-van der Waals, but is IMF)

11. 1. London Dispersion Forces
Electrons in this helium atom repel each other…
It does happen that they occasionally wind up on the same side of the atom.

12. 1. London Dispersion Forces
Instantaneous dipole
At that instant, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.

13. 1. London Dispersion Forces
Induced dipole d- d+ d- d+
Instantaneous dipole
Another helium atom nearby…
would have an induced dipole in it…
..as the electrons on the left side of the INSTANTANEOUS repel the electrons in the cloud on the INDUCED.

14. 1. London Dispersion Forces
Induced dipole
Instantaneous dipole
London dispersion forces, are attractions between an instantaneous dipole and an induced dipole.

15. 1. London Dispersion Forces
These forces are present in all molecules, whether they are polar or nonpolar.
(stronger) (weaker)
The tendency of an electron cloud to distort in this way is called polarizability.

16. 1. London Dispersion Forces
< d- d+ d- d+ D D
The greater the distance between the poles, the stronger the dispersion forces.

17. Factors Affecting London Forces
The strength of dispersion forces tends to increase with increased molar mass b/c…
larger atoms have larger electron clouds, which are more polarizable.

18. Factors Affecting London Forces
The shape of the molecule affects the strength of dispersion forces:
long, skinny molecules… (like pentane)
tend to have stronger dispersion forces than short, fat ones (like neopentane).
This is due to the increased surface area in n-pentane…
Cylinders have greater surface area..
..than spheres…
SA > SA

19. 2. Dipole-Dipole Interactions
Molecules that have permanent dipoles are attracted to each other.
The positive end of one is attracted to the negative end of the other and vice-versa.
Like charges repulse one another d+ d- d+ d- d+ d- d+ d- d+ d- d+ d-

20. 2. Dipole-Dipole Interactions
As with London Dispersion Forces, the greater the distance between the poles, the stronger the dispersion forces.

21. 2. Dipole-Dipole Interactions
The more polar the molecule, the higher is its boiling/melting point.

22. Which Have a Greater Effect: Dipole-Dipole Interactions or Dispersion Forces?
If two molecules are of comparable size and shape, dipole-dipole interactions will be more significant.
If one molecule is much larger than another, dispersion forces will be more significant in determining its physical properties.
mass = ~
H2S  34.08 g/mol
SiH g/mol
mass
<
H2S  34.08 g/mol
H2Se g/mol

23. How Do We Explain This? And then there’s this freak! Boiling Points
Polar molecules are “stickier”…
Boiling Points
Nonpolar molecules merely rely on size.

24. 3. Hydrogen Bonding
The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong.
H-bonding arises in part from high electronegativity
of N, O, and F.

25. 1st Image of Hydrogen Bonding
In May, Felix Fischer and colleagues at the University of California Berkeley in the US used AFM to image molecules before and after a chemical transformation. The remarkable images showed the formation of covalent bonds in a cyclisation reaction.

27. 4. Ion-Dipole Interactions (non-vdW)
Ion-dipole interactions are an important force in solutions of ions.
They make it possible for ionic substances to dissolve in polar solvents.

28. Summarizing Intermolecular Forces
Stronger
(more significant)
Weaker
(less significant)
Ion Dipoles (NOT Van der Waals)
H-bonds (if H bonded to N, O, or F)
Dipole-dipole forces
(polar molecules, ion-dipole if ions)
London Dispersion forces
(nonpolar, induced dipoles)

29. Intermolecular Forces Affect Many Physical Properties
1. Cohesive/Adhesive Forces
2. Capillary Action
3. Surface Tension
4. Viscosity
5. Phase Changes
6. Vapor Pressure

30. 1. Cohesive/Adhesive Forces
Attraction to OTHER substances…
i.e.,
to glass surface
Cohesive
to self…

31. 2. Capillary Action
Attraction of water to glass… Forms the curved meniscus.

32. 2. Capillary Action
Attraction of water to glass INCREASES w/decreased radius…
Adhesive forces have less to pull.

33. 3. Surface Tension
the net inward force experienced by the molecules on the surface of a liquid.

34. 4. Viscosity resistance of a liquid to flow
increases with IMAF’s and decreases with higher temp.

35. 4. Viscosity The Pitch Drop Experiment… Begun 1927 35

36. 5. Phase Changes Energy of System Solid Liquid Gas Vaporization
Condensation
Sublimation
Deposition
Melting
Freezing

37. Energy Changes Associated with Changes of State
Heat of Fusion: energy required to change a solid at its melting point to a liquid.

38. Energy Changes Associated with Changes of State
Heat of Vaporization: energy required to change a liquid at its boiling point to a gas.
VERY
Heavy
Trends follow those we discussed
Polar
Large
nonpolar
Small
nonpolar

39. Rank these molecules from LOWEST to HIGHEST Heat of Fusion
_____ Large Nonpolar
_____ Polar
_____ Small nonpolar
_____ VERY heavy ? 2 3 1 4

40. Energy Changes Associated with Changes of State
The temp. does not change during the phase change.
Extra KE added is used in separating molecules further apart
Why? b/c energy added is used in separating molecules further apart

41. 6. Vapor Pressure
At any temperature, some molecules in a liquid have enough energy to escape.
As temperature increases, the fraction of molecules that have enough energy to escape increases.
Draw in the LOW and HIGH temperature curves on the graph in your notes
Fraction of Molecules
Minimum kinetic energy needed to escape liquid.
Low Temperature
Higher Temperature
Many more molecules can escape…
Kinetic Energy

42. 6. Vapor Pressure
As more molecules escape the liquid, the pressure they exert increases.
Dynamic Equilibrium:
vaporize/condense at same rate

43. 6. Vapor Pressure
The pressure exerted by a vapor in equilibrium with its solid or liquid phase.
So, in general… the HIGHER the boiling point, the LOWER the Vapor Pressure… Because there’s less “vapor”… Get it?
Water
At a given temperature, a liquid in a closed container will establish equilibrium with its own vapor. 43

44. 6. Vapor Pressure
boiling point: temperature at which liquid’s vapor pressure = atmospheric pressure.
The normal BP is temp. at which its vapor pressure is 1 atm.

45. Solids Fall into two groups:
1. Amorphous —no particular order in the arrangement of particles.
Melted quartz, cooled rapidly

46. Solids 2. Crystalline—particles are in highly ordered arrangement.
quartz

47. Covalent-Network Solids
Atoms are covalently bonded to each other.
Diamonds are an example of a covalent-network solid.
They tend to be hard and have high melting points.

48. Molecular Solids
Atoms are held together with van der Waals forces (IMAF’s).
Graphite is an example of a molecular solid
They tend to be softer and have lower melting points.

49. Ionic Solids +/- ions bonded by electrostatic attraction.
They tend to be hard and brittle and have higher melting points.

50. Metallic Solids “sea” of electrons
Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces.
In metals, valence electrons are delocalized throughout the solid. - -
Accounts for conductivity of metals
“sea” of electrons

51. Types of Bonding in Crystalline Solids
Type of Solid
Forces
Melting Point
Hard-ness
Cond.?
Molecular
IMAF
Low to High
Soft
Poor
Covalent
covalent
VERY high
VERY
Ionic
electrostatic
High
Hard, Brittle
Metallic
Good

52. Phase Diagrams
Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.

53. This means that Ice can be melted by applying pressure…
Phase Diagram of Water
This means that Ice can be melted by applying pressure… 53

54. Phase Diagram of Carbon Dioxide
Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 sublimes at normal pressures.