2. 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
3. Intermolecular Forces
11.2
Intermolecular forces are attractive forces between molecules.
Intramolecular forces hold atoms together in a molecule.
Intermolecular vs Intramolecular
• 41 kJ to vaporize 1 mole of water (inter)
• 930 kJ to break all O-H bonds in 1 mole of water (intra)
Generally,
intermolecular
forces are much
weaker than
intramolecular
forces.
“Measure” of intermolecular force
boiling point
melting point
DHvap
DHfus
DHsub
7. Intermolecular Forces
Dispersion Forces
Attractive forces that arise as a result of temporary
dipoles induced in atoms or molecules
11.2
ion-induced dipole interaction
dipole-induced dipole interaction
9. Intermolecular Forces
Dispersion Forces Continued
11.2
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.
10. S
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.
SO2
SO2 is a polar molecule: dipole-dipole forces. There are
also dispersion forces between SO2 molecules.
11.2
11. Intermolecular Forces
Hydrogen Bond
11.2
The hydrogen bond is a special dipole-dipole interaction
between they hydrogen atom in a polar N-H, O-H, or F-H bond
and an electronegative O, N, or F atom.
A H…B A H…A
or
A & B are N, O, or F
13. Why is the hydrogen bond considered a
“special” dipole-dipole interaction?
Decreasing molar mass
Decreasing boiling point
11.2
14. 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
15. Properties of Liquids
Cohesion is the intermolecular attraction between like molecules
11.3
Adhesion is an attraction between unlike molecules
Adhesion
Cohesion
16. Properties of Liquids
Viscosity is a measure of a fluid’s resistance to flow.
11.3
Strong
intermolecular
forces
High
viscosity
18. A crystalline solid possesses rigid and long-range order. In a
crystalline solid, atoms, molecules or ions occupy specific
(predictable) positions.
An amorphous solid does not possess a well-defined
arrangement and long-range molecular order.
A unit cell is the basic repeating structural unit of a crystalline
solid.
Unit Cell
lattice
point
Unit cells in 3 dimensions 11.4
At lattice points:
• Atoms
• Molecules
• Ions
21. An amorphous solid does not possess a well-defined
arrangement and long-range molecular order.
A glass is an optically transparent fusion product of inorganic
materials that has cooled to a rigid state without crystallizing
Crystalline
quartz (SiO2)
Non-crystalline
quartz glass 11.7
23. The equilibrium vapor pressure is the vapor pressure
measured when a dynamic equilibrium exists between
condensation and evaporation
H2O (l) H2O (g)
Rate of
condensation
Rate of
evaporation
=
Dynamic Equilibrium
11.8
25. Molar heat of vaporization (DHvap) is the energy required to
vaporize 1 mole of a liquid at its boiling point.
ln P = -
DHvap
RT
+ C
Clausius-Clapeyron Equation
P = (equilibrium) vapor pressure
T = temperature (K)
R = gas constant (8.314 J/K•mol)
11.8
Vapor Pressure Versus Temperature
26. 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
27. The critical temperature (Tc) is the temperature above which
the gas cannot be made to liquefy, no matter how great the
applied pressure.
The critical pressure
(Pc) is the minimum
pressure that must be
applied to bring about
liquefaction at the
critical temperature.
11.8
29. Melting
11.8
Freezing
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
30. Molar heat of fusion (DHfus) is the energy required to melt
1 mole of a solid substance at its freezing point.
11.8
34. Sublimation
11.8
Deposition
H2O (s) H2O (g)
Molar heat of sublimation
(DHsub) is the energy required
to sublime 1 mole of a solid.
DHsub = DHfus + DHvap
( Hess’s Law)
35. A phase diagram summarizes the conditions at which a
substance exists as a solid, liquid, or gas.
Phase Diagram of Water
11.9
37. 11.9
Effect of Increase in Pressure on the Melting Point
of Ice and the Boiling Point of Water
Editor's Notes
Phase changes, transformations from one phase to
another, occur when energy (usually in the form of heat) is added or removed from
a substance.
What happens at the molecular level during evaporation? In the beginning, the
traffi c is only one way: Molecules are moving from the liquid to the empty space. Soon
the molecules in the space above the liquid establish a vapor phase. As the concentration
of molecules in the vapor phase increases, some molecules return to the liquid
phase, a process called condensation.
The critical phenomenon of sulfur hexafl uoride. (a) Below the critical temperature the clear liquid phase is visible. (b) Above
the critical temperature the liquid phase has disappeared. (c) The substance is cooled just below its critical temperature. The fog
represents the condensation of vapor. (d) Finally, the liquid phase reappears.
The graph is divided into three
regions, each of which represents a pure phase. The line separating any two regions
indicates conditions under which these two phases can exist in equilibrium.
a) The phase diagram of water. Each solid line between two phases specifi es the conditions
of pressure and temperature under which the two phases can exist in equilibrium. The point
at which all three phases can exist in equilibrium (0.006 atm and 0.01°C) is called the triple
point.
b) This phase diagram tells us that increasing the pressure on ice lowers its melting
point and that increasing the pressure of liquid water raises its boiling point.
The phase diagram of carbon
dioxide. Note that the solidliquid
boundary line has a
positive slope. The liquid phase
is not stable below 5.2 atm, so
that only the solid and vapor
phases can exist under
atmospheric conditions.
What would happen if melting and boiling were carried
out at some other pressure? Figure 12.32(b) shows clearly that increasing the pressure
above 1 atm will raise the boiling point and lower the melting point. A decrease in
pressure will lower the boiling point and raise the melting point
