1. Intermolecular Forces and
Liquids and Solids
Chapter 11
Copyright © The McGraw-Hill Companies, Inc.  Permission required for reproduction or display.

2. States of Matter
The fundamental difference between states of matter is the distance between particles.

3. States of Matter
Because in the solid and liquid states particles are closer together, we refer to them as condensed phases.

4. The States of Matter
The state a substance is in at a particular temperature and pressure depends on two antagonistic entities:
The kinetic energy of the particles
The strength of the attractions between the particles

5. A phase is a homogeneous part of the system in contact with other parts of the system but separated from them by a well-defined boundary.
2 Phases
Solid phase - ice
Liquid phase - water
11.1

6. Review: Bonding
Bonds are formed by the interaction of electrons between elements
Two types of chemical bonds
Ionic bond – transfer electrons from the atom of one element to another
Covalent bond – sharing of electrons between atoms of elements

7. Ions and Ionic Compounds
Ionic compounds are made up of charged particles called ions.
Ions are formed when an atom gains or loses an electron
A binary ionic compound is the simplest type of ionic compound made of two elements
Often occurs when a metal reacts with a nonmetal
All binary ionic compounds are solid

8. Covalent bonding Occurs when elements share electrons
Usually occurs between nonmetals

9. Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms. This occurs because the electrons spend more time in the vicinity of one atom or another. H F
electron rich
region
e- poor
e- rich
electron poor
region F H d+ d-
9.5

10. Electronegativity - relative, F is highest
Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.
Electronegativity - relative, F is highest H F
9.5

11. Which of the following molecules have a dipole moment?
H2O, CO2, SO2, and CH4 O H S O
dipole moment
polar molecule
dipole moment
polar molecule C H C O
no dipole moment
nonpolar molecule
no dipole moment
nonpolar molecule
10.2

12. Intermolecular Forces
Intermolecular forces are attractive forces between molecules.
Intramolecular forces hold atoms together in a molecule.
“Measure” of intermolecular force
boiling point
melting point
DHvap
DHfus
DHsub
Generally, intermolecular forces are much weaker than intramolecular forces.
11.2

13. Intermolecular Forces
Dipole-Dipole Forces
hydrogen bonding
Dipole-induced dipole
Ion-dipole Forces
Dispersion Forces (van der Waals Forces, London Dispersion Forces)

14. Intermolecular Forces
Dipole-Dipole Forces
Attractive forces between polar molecules
Orientation of Polar Molecules in a Solid
11.2

15. Dipole-induced Dipole Forces
A weak attraction that results when a polar molecule induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species
A nonpolar molecule with no dipole moment. The dot indicates the nucleus.
When a polar molecule approaches a nonpolar molecule the nonpolar molecule develops a dipole

16. Intermolecular Forces
Ion-Dipole Forces
Attractive forces between an ion and a polar molecule
Ion-Dipole Interaction
11.2

17. Strength of an ionic bond
Ions with higher charges attract each other more strongly than ions with lower charges.
The smaller the distance between the charges the stronger the attraction.

18. 11.2

19. Intermolecular Forces
Dispersion Forces
Attractive forces that arise as a result of temporary dipoles induced in atoms or molecules
The separation of the negative and positive charges in a nonpolar molecule is due to the proximity of an ion or a polar molecule
ion-induced dipole interaction
dipole-induced dipole interaction
11.2

20. Dispersion forces are also known as London forces
These forces are present in all molecules, whether they are polar or nonpolar.
The tendency of an electron cloud to distort in this way is called polarizability.
The strength of dispersion forces tends to increase with increased molecular weight.
Larger atoms have larger electron clouds, which are easier to polarize.

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

22. Intermolecular Forces Affect Many Physical Properties
The strength of the attractions between particles can greatly affect the properties of a substance or solution.

23. Dispersion Forces
London Forces - London showed that the magnitude of this attractive interaction is directly proportional to the polarizability of the atom or molecule
van der Waals forces - attractive or repulsive force between molecules (or between parts of the same molecule) other than those due to covalent bonds or to the electrostatic interaction of ions with one another or with nonpolar molecules
Fritz London interpreted temporary dipoles in 1930
Johannes Diderik van der Waals

24. Intermolecular Forces
Dispersion Forces Continued
Polarizability is the ease with which the electron distribution in the atom or molecule can be distorted.
Polarizability increases with:
greater number of electrons
more diffuse electron cloud
Dispersion forces usually increase with molar mass.
11.2

25. What type(s) of intermolecular forces exist between each of the following molecules?
HBr
HBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules.
CH4
CH4 is nonpolar: dispersion forces. S O
SO2
SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules.
11.2

26. Dipole-dipole forces are attractive forces between
Hydrogen Bond
The hydrogen bond is a special dipole-dipole interaction between the hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom.
IT IS NOT A BOND.
Dipole-dipole forces are attractive forces between
polar molecules

27. The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong.
We call these interactions hydrogen bonds.
Hydrogen bonding arises in part
from the high electronegativity of
nitrogen, oxygen, and fluorine.

28. Examples of the Intermolecular Forces

29. Strong intermolecular forces
Properties of Liquids
Surface tension is the amount of energy required to stretch or increase the surface of a liquid by a unit area.
Strong intermolecular forces
High surface tension
11.3

30. Properties of Liquids
Cohesion is the intermolecular attraction between like molecules
Adhesion is an attraction between unlike molecules
Cohesion < Adhesion
Cohesion > Adhesion
Adhesion
attracted to glass
Cohesion
attracted to each other
11.3

31. High Viscosity = Slow liquid Flow
Properties of Liquids
Viscosity is a measure of a fluid’s resistance to flow.
Resistance of a liquid to flow is called viscosity.
It is related to the ease with which molecules can move past each other.
Viscosity increases with stronger intermolecular forces and decreases with higher temperature.
High Viscosity = Slow liquid Flow
11.3

32. Ice is less dense than water
Water is a Unique Substance
Maximum Density
40C
Density of Water
Ice is less dense than water
11.3

33. Phase Changes
Transformation from one phase to another when energy is added or removed
Physical changes characterized by changes in molecular order

34. Greatest
Order
Least
T2 > T1
Evaporation
Condensation
11.8

35. Vapor Pressure
The vapor pressure of any substance increases non-linearly with temperature according to the Clausius-Clapeyron relation.
The atmospheric pressure boiling point of a liquid is the temperature where the vapor pressure equals the ambient atmospheric pressure. With any incremental increase in that temperature, the vapor pressure becomes sufficient to overcome atmospheric pressure and lift the liquid to form bubbles inside the bulk of the substance.
bp of water
99°C
210°F
Death Valley
86 m
bp of water
69°C
156.2°F
Mt. Everest
8850 m

36. The boiling point is the temperature at which the (equilibrium) vapor pressure of a liquid is equal to the external pressure.
The normal boiling point is the temperature at which a liquid boils when the external pressure is 1 atm.
11.8

37. Vapor Pressure
The higher the vapor pressure of a liquid at a given temperature, the lower the normal boiling point (i.e., the boiling point at atmospheric pressure) of the liquid.

38. Vapor Pressure Versus Temperature
Molar heat of vaporization (Hvap) is the energy required to vaporize 1 mole of a liquid at its boiling point.
ln P = -
Hvap RT
+ C
Clausius-Clapeyron Equation
P = (equilibrium) vapor pressure
T = temperature (K)‏
R = gas constant (8.314 J/K•mol)‏
Vapor Pressure Versus Temperature
11.8

39. H2O (s) H2O (l)‏
The melting point of a solid or the freezing point of a liquid is the temperature at which the solid and liquid phases coexist in equilibrium
Melting
Freezing
11.8

40. Molar heat of fusion (Hfus) is the energy required to melt 1 mole of a solid substance at its freezing point.
11.8

41. Heating Curve
11.8

42. Sublimation and Deposition
Sublimation – Process in which molecules go directly from the solid into the vapor phase.
Deposition – Process where the molecules make the transition from vapor to solid directly.
Naphthalene
Camphor

43. H2O (s) H2O (g)‏
Molar heat of sublimation (Hsub) is the energy required to sublime 1 mole of a solid.
Sublimation
Deposition
Hsub = Hfus + Hvap
( Hess’s Law)‏
11.8

44. The critical temperature (Tc) is the temperature above which the gas cannot be made to liquefy, no matter how great the applied pressure. This is the highest temperature a substance can exist as a liquid.
The critical pressure (Pc) is the minimum pressure that must be applied to bring about liquefaction at the critical temperature.
11.8

45. A phase diagram summarizes the conditions at which a substance exists as a solid, liquid, or gas.
Phase Diagram of Water
Triple point -The only condition under which all 3 phases can be equilibrium with on another.
11.9

46. Phase Diagram of Carbon Dioxide
At 1 atm
CO2 (s) CO2 (g)‏
11.9

47. Effect of Increase in Pressure on the Melting Point
of Ice and the Boiling Point of Water
11.9