2. Learning objective 2.16:
The student is able to explain the properties (phase, vapor pressure, viscosity, etc.) of small and large molecular compounds in terms of the strengths and types of intermolecular forces.

3. Distinguish between intermolecular and intramolecular attractions
Define:
Sublimation
Deposition
Condensation
Evaporation
Melting
Freezing
Freezing point
Boiling point
Vapor pressure
Viscosity
Surface tension
ΔH of fusion
ΔH of vaporization
ΔH of sublimation
Polar
Nonpolar
Dipole-dipole forces
Ion-dipole forces
Hydrogen “bonding”
London dispersion forces
Distinguish between intermolecular and intramolecular attractions
Put a list of compounds in order of increasing melting point, boiling point, and vapor pressure
Use and label the parts of a phase diagram
Use the Clausius-Clapeyron equation to relate temperature to vapor pressure of a substance

4. Solid, Liquid, or Gas
What are three factors determine whether a substance is a solid, a liquid, or a gas:
The attractive intermolecular forces between particles that tend to draw the particles together.
Temperature: The kinetic energies of the particles (atoms, molecules, or ions) that make up a substance. Kinetic energy tends to keep the particles moving apart.
Pressure: pressure is increased or decreased as the volume of a closed container changes

5. Types of Attractive Forces
There are several types of attractive intermolecular forces:
Ionic
Ion-dipole forces
Dipole-dipole forces
Hydrogen bonding
Induced-dipole forces
Ion-induced
Dipole-induced
London dispersion forces
All of the intermolecular forces that hold a liquid together are called cohesive forces.

6. Ionic Bonds Electrons are transferred
Electronegativity differences are
generally greater than 1.7
The formation of ionic bonds is
always exothermic!

7. Coulomb’s Law Q is the charge. r is the distance between the centers.
If charges are opposite, E is negative
exothermic
Same charge, positive E, requires energy to bring them together.
endothermic

8. Ion-Dipole Forces
An ion-dipole force is an attractive force that results from the electrostatic attraction between an ion and a neutral molecule that has a dipole.
Most commonly found in solutions. Especially important for solutions of ionic compounds in polar liquids.
Ion-dipole attractions become stronger as either the charge on the ion increases, or as the magnitude of the dipole of the polar molecule increases.

9. Dipole-Dipole Forces
Dipole-dipole forces are attractive forces between the positive end of one polar molecule and the negative end of another polar molecule.
They are much weaker than ionic or covalent bonds and have a significant effect only when the molecules involved are close together (touching or almost touching).

10. Hydrogen Bonding
Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen
Hydrogen bonding between ammonia and water

11. Hydrogen Bonding in DNA
Thymine hydrogen bonds to Adenine T A

12. Hydrogen Bonding in DNA
Cytosine hydrogen bonds to Guanine C G

13. Ion-Induced Dipole Forces
Induced dipole forces result when an ion or a dipole induces a dipole in an atom or a molecule with no dipole. These are weak forces.
Ion-Induced Dipole Forces
An ion-induced dipole attraction is a weak attraction that results when the approach of an ion induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species.

14. Dipole-Induced Dipole Forces
A dipole-induced dipole attraction is 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.

15. London Dispersion Forces
The temporary separations of charge that lead to the London force attractions are what attract one nonpolar molecule to its neighbors.
London forces increase with the size of the molecules.
Fritz London

16. London Dispersion Forces

17. London Forces in Hydrocarbons

18. Boiling point as a measure of intermolecular attractive forces

19. Relative Magnitudes of Forces
The types of bonding forces vary in their strength as measured by average bond energy.
Ionic bonds
Strongest
Weakest
*Ion-dipole interactions
*Hydrogen bonding (12-16 kcal/mol )
Dipole-dipole interactions (2-0.5 kcal/mol)
Ion induced dipole interactions
Induced Dipole-dipole interactions
London forces (less than 1 kcal/mol)

21. CH4 CaCl2 NaHCO3 CH3Cl C12H22O11 NH3 H2S NH4Cl
Identify the predominant intermolecular forces present in the solids of each of the following substances:
dipole-dipole
ion-dipole
hydrogen bonding
London dispersion
dipole-induced dipole
ionic
ion-induced dipole
CH4
CaCl2
NaHCO3
CH3Cl
C12H22O11
NH3
H2S
NH4Cl

22. CH4 (d) CaCl2 (f) NaHCO3 (f) CH3Cl (a) C12H22O11 (c) NH3 (c) H2S (a)
Identify the predominant intermolecular forces present in the solids of each of the following substances:
dipole-dipole
ion-dipole
hydrogen bonding
London dispersion
dipole-induced dipole
ionic
ion-induced dipole
CH4 (d)
CaCl2 (f)
NaHCO3 (f)
CH3Cl (a)
C12H22O11 (c)
NH3 (c)
H2S (a)
NH4Cl (f)

23. SO2 and CH4 MgCl2 and C3H8 C4H10 and BF3 C12H22O11 and C3H8
Identify the intermolecular forces present between the two substances listed:
dipole-dipole
ion-dipole
hydrogen bonding
London dispersion
dipole-induced dipole
ionic
ion-induced dipole
SO2 and CH4
MgCl2 and C3H8
C4H10 and BF3
C12H22O11 and C3H8
NBr3 and H2S
H2O and C12H22O11
NBr3 and H2O
CO2 and CCl4

24. SO2 and CH4 (e) MgCl2 and C3H8 (g) C4H10 and BF3 (d) NBr3 and H2S (a)
Identify the strongest intermolecular forces present between the two substances listed:
dipole-dipole
ion-dipole
hydrogen bonding
London dispersion
dipole-induced dipole
ionic
ion-induced dipole
SO2 and CH4 (e)
MgCl2 and C3H8 (g)
C4H10 and BF3 (d)
C12H22O11 and C3H8 (e)
NBr3 and H2S (a)
H2O and C12H22O11 (c)
NBr3 and H2O (c)
CO2 and CCl4 (d)

27. No, really, what IS a liquid??!!

28. What Is a Liquid?
A liquid is a state of matter in which a sample of matter:
is made up of very small particles (atoms, molecules, and/or ions).
flows and can change its shape.
is not easily compressible and maintains a relatively fixed volume.
The particles that make up a liquid:
are close together with no regular arrangement,
vibrate, move about, and slide past each other.
This bottle contains both liquid bromine [Br2(l), the darker phase at the bottom of the bottle] and gaseous bromine [Br2(g), the lighter phase above the liquid]. The circles show microscopic views of both liquid bromine and gaseous bromine.

29. More Properties of a Liquid
Surface Tension: The resistance to an increase in its surface area (polar molecules, liquid metals).
Capillary Action: Spontaneous rising of a liquid in a narrow tube.

30. Even More Properties of a Liquid
Viscosity: Resistance to flow
High viscosity is an
indication of strong
intermolecular forces

31. Evaporation
Evaporation is the change of a liquid to a gas.
Microscopic view of a liquid.
Microscopic view after evaporation.
When a liquid is heated sufficiently or when the pressure on the liquid is decreased sufficiently, the forces of attraction between molecules do not prevent them from moving apart, and the liquid evaporates to a gas.
Example: The sweat on the outside of a cold glass evaporates when the glass warms.
Example: Gaseous carbon dioxide is produced when the valve on a CO2 fire extinguisher is opened and the pressure is reduced.

32. Condensation
Condensation is the change from a vapor to a condensed state (solid or liquid).
When a gas is cooled sufficiently or, in many cases, when the pressure on the gas is increased sufficiently, the forces of attraction between molecules prevent them from moving apart, and the gas condenses to either a liquid or a solid.
Example: Water vapor condenses and forms liquid water (sweat) on the outside of a cold glass or can.
Example: Liquid carbon dioxide forms at the high pressure inside a CO2 fire extinguisher.
Microscopic view of a gas.
Microscopic view after condensation.

34. Vapor Pressure
The vapor pressure of a liquid is the equilibrium pressure of a vapor above its liquid (or solid)
The vapor pressure of a liquid varies with its temperature, as the following graph shows for water. The line on the graph shows the boiling temperature for water.
As the temperature of a liquid or solid increases its vapor pressure also increases. Conversely, vapor pressure decreases as the temperature decreases.

35. Factors That Affect Vapor Pressure
Types of Molecules: the types of molecules that make up a solid or liquid determine its vapor pressure. If the intermolecular forces between molecules are:
relatively strong, the vapor pressure will be relatively low.
relatively weak, the vapor pressure will be relatively high.
substance
vapor pressure at 25oC
diethyl ether
0.7 atm
bromine
0.3 atm
ethyl alcohol
0.08 atm
water
0.03 atm
Surface Area: the surface area of the solid or liquid in contact with the gas has no effect on the vapor pressure.

36. Temperature Dependence of Vapor Pressures
The vapor pressure above the liquid varies exponentially with changes in the temperature.
The Clausius-Clapeyron equation shows how the vapor pressure and temperature are related. It can be written as:

37. Clausius – Clapeyron Equation
A straight line plot results when ln P vs. 1/T is plotted and has a slope of Hvap/R.
Clausius – Clapeyron equation is true for any two pairs of points.
Write the equation for each and combine to get:

38. Using the Clausius – Clapeyron Equation
Boiling point - the temperature at which the vapor pressure of a liquid is equal to the pressure of the external atmosphere.
Normal boiling point - the temperature at which the vapor pressure of a liquid is equal to atmospheric pressure (1 atm).
E.g. Determine normal boiling point of chloroform if its heat of vaporization is 31.4 kJ/mol and it has a vapor pressure of mmHg at 25.0°C.
E.g.2. The normal boiling point of benzene is 80.1°C; at 26.1°C it has a vapor pressure of mmHg. What is the heat of vaporization?
334 K
33.0 kJ/mol

39. Phase Transitions Melting: change of a solid to a liquid.
H2O(s)  H2O(l)
H2O(l)  H2O(s)
H2O(l)  H2O(g)
H2O(g)  H2O(l)
H2O(s)  H2O(g)
H2O(g)  H2O(s)
Melting: change of a solid to a liquid.
Freezing: change a liquid to a solid.
Vaporization: change of a liquid to a gas.
Condensation: change of a gas to a liquid.
Sublimation : Change of solid to gas
Deposition: Change of a gas to a solid.

40. Water phase changes
Temperature remains constant during a phase change.
Energy

41. Energy of Heat and Phase Change
Heat of vaporization: heat needed for the vaporization of a liquid.
H2O(l) H2O(g) DH = 40.7 kJ
Heat of fusion: heat needed for the melting of a solid.
H2O(s) H2O(l) DH = 6.02 kJ
Temperature does not change during the change from one phase to another.
E.g. Start with a solution consisting of 50.0 g of H2O(s) and 50.0 g of H2O(l) at 0°C. Determine the heat required to heat this mixture to 100.0°C and evaporate half of the water.
130 kJ

42. Phase Diagrams
Triple point- Temp. and press. where all three phases co-exist in equilibrium.
Critical temp.- Temp. where substance must always be gas, no matter what pressure.
Critical pressure- vapor pressure at critical temp.
Critical point- point where system is at its critical pressure and temp.

43. Phase changes by Name

44. Water

45. Carbon dioxide

46. Carbon

47. Sulfur