Chapter 11: Liquids & Solids
Renee Y. Becker
Valencia Community College
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Kinetic Molecular Theory
Introduction
• Gases
– Gas particles act independent of one another
– Attractive forces are very weak
• Particles are free to move randomly
• Occupy whatever space available
Liquids and solids are different from gases in that
they have strong attractive forces between
particles
2
Polar Covalent Bonds
• Polar Covalent Bonds
– Form between a non-metal/non-metal of
different electronegativities
EN = ENA – ENB
When EN  1.7 ionic bond
When EN < 1.7 polar covalent bond
When EN < .5 non-polar covalent bond
3
Polar Molecules
• Polar Molecules
– Just as bonds can be polar, molecules as a
whole can be polar
– Net sum of individual bond polarities and lonepair contributions
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Dipole Moment
• Dipole moment, , (ionic and covalent)
– Measure of net molecular polarity
– The magnitude of the charge Q at either end of
the molecular dipole times the distance, r,
between the charges
 =Qxr
– Expressed in debyes, D, where 1 D = 3.336 x
10-30 coulomb meters
– Q = 1.6 x 10-19 C (electron charge)
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Dipole Moment
% Ionic Character = experimental  (100%)
calculated 
a high % IC means that the bond is similar to or is
an ionic bond
a low % IC means that it is more like a covalent
bond
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Chloromethane
Example 1
Cl
H
H
H
a) Calculate the dipole moment
b) Calculate the % ionic character of the bond
Experimentally measured dipole moment = 1.87 D
C-Cl bond distance = 178 pm = 178 x 10-12 m
If we assume that the contributions of the nonpolar C-H bonds
are small, then most of the chloromethane dipole moment
is due to the C-Cl bond
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Example 2
Hydrochloric acid
Cl
H
Calculate the % ionic charcater
1. Distance between atoms is 127 pm
2. Experimentally measured dipole
moment = 1.03 D
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Example 3
Tell which of the following compounds are likely to have
a dipole moment and show the direction of each.
a) SF6
b) CHCl3
c) CH2Cl2
d) CH2CH2
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Intermolecular Forces
• Van der Waals forces – intermolecular forces as a
whole, all are electrical in origin and result from
the mutual attraction of unlike charge or mutual
repulsion of like charges.
4 main types
• Dipole-dipole
• Ion-dipole
• Dispersion forces
• Hydrogen bonding
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Dipole-dipole
a) Neutral but polar molecules experience dipole-dipole forces as
a result of electrical interactions among dipoles on neighboring
molecules.
b) Forces can be attractive or repulsive, depending on the
orientation of the molecules.
c) These forces are weak 3-4 kJ/mol and only significant if
molecules are close
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HH
H
H
O
+
O Na O
O
H
H
H
H
H
O
H
Ion-dipole
Cl
H
O
H
Result of electrical interactions between an ion and the partial
charges on a polar molecule
b) Particularly important in aqueous solutions of ionic substances
such as NaCl, in which polar water molecules surround the ions
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London Dispersion Forces
a) Result from the motion of electrons
b) At any given time more electrons may be in a particular area of
the molecule
c) This gives the molecule an instantaneous dipole
d) This short lived dipole can affect the electron distribution in
neighboring molecules and induce temporary dipoles in them
e) More electrons a molecule has the stronger the dispersion forces
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Hydrogen Bonding
a)
Attractive interaction between a hydrogen atom bonded to a
very electronegative atom (O, N, F) and an unshared electron
pair on another electronegative atom
b)
Hydrogen bonds arise because O-H, N-H, and F-H bonds are
highly polar with partial positive charge on the hydrogen and
partial negative on the electronegative atom.
c)
Hydrogen has no core electrons to shield its nucleus and it is
small so it can be approached closely by other molecules
d)
The dipole-dipole attraction between the hydrogen and an
unshared electron pair on a nearby atom is usually strong
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Hydrogen Bonding
e) Water is able to form a vast 3D network of hydrogen bonds
because each H2O molecule has two hydrogens and two
electron pairs
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Intermolecular Forces
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Example 4
Identify the likely kinds of intermolecular forces in the
following
A) HCl
H Cl
H H
B) CH3CH3
H
H
H H
H
C) CH3NH2
H
H
N
H
H
D) Kr
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Example 5
Of the substances Ar, Cl2, CCl4 and HNO3 which
has:
a) The largest dipole-dipole forces?
b) The largest hydrogen-bond forces?
c) The smallest dispersion forces?
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Properties of Liquids
Viscosity
1. The measure of a liquids resistance to flow
2. Related to the ease with which individual molecules
move around in the liquid and thus to the
intermolecular forces present
3. Substances with small non-polar molecules have weak
intermolecular forces and low viscosities (free
flowing)
4. More polar substances have stronger intermolecular
forces and have higher viscosities
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Surface Tension
Properties of Liquids
1. The resistance of a liquid to spread out and increase its
surface area
2. Caused by differences in intermolecular forces
experienced by molecules at the surface and the
interior
3. Surface molecules feel attractive forces on only one
side and are drawn in toward the liquid
4. Interior molecules are drawn equally in all directions
5. Higher in liquids that have stronger intermolecular
forces
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Phase Changes
• Physical form changes but chemical identity does not
change
Fusion (melting)
Freezing
Vaporization
Condensation
Sublimation
solid  liquid
liquid  solid
liquid  gas
gas  liquid
solid  gas
Deposition
gas  solid
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Free energy change, G
Thermochemistry
1. All naturally occurring processes, every phase change
has a free-energy change
2. G =  H - T S
3. H enthalpy, heat flow, positive (from surrounding to
system, bond breaking takes energy), negative (from
system to surroundings, bond making)
4. S entropy, disorder, positive (ordered to disorderd),
negative (disordered to ordered)
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Thermochemistry
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Thermochemistry
Calculating the temperature at a phase change
1. G > 0 non-spontaneous G < 0 spontaneous
G = 0 equilibrium
2. Set G = 0
0 = H - TS and solve for T
T = H/S
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Heating Curve for H2O
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Heating Curve for H2O
Heating curve for H2O
E = molar heat capacity (T)
Energy to heat ice from -25C to 0C
Molar heat capacity of ice= 36.57 J/molC
E=
Energy to heat H2O from 0C to 100C
Molar heat capacity of liquid H2O = 75.4 J/ molC
E=
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Thermochemistry
Heat of fusion, Hfusion
The amount of energy required to overcome
enough intermolecular forces to convert a
solid into a liquid
Heat of Vaporization, Hvap
The amount of energy necessary to convert
a liquid into a gas
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Example 6
Chloroform has Hvap = 29.2 kJ/mol and
Svap = 87.5 J/K mol. What is the boiling
point of chloroform?
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Evaporation, Vapor Pressure, and Boiling Point
Vapor Pressure
1. The partial pressure of a gas in equilibrium and at
constant temperature with liquid
2. The pressure exerted by gaseous molecules above a
liquid
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Evaporation, Vapor Pressure, and Boiling Point
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Evaporation, Vapor Pressure, and Boiling Point
The higher the temperature and the lower the boiling
point of the substance the greater the fraction of
molecules in the sample that have sufficient
kinetic energy to break free from the surface of the
liquid and escape into the vapor.
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Evaporation, Vapor Pressure, and Boiling Point
Numerical value of Vapor Pressure depends on:
a)
Magnitude of intermolecular forces
The smaller the forces the higher the vapor
pressure, loosely held molecules escape
easily
b)
Temperature
The higher the temperature, the higher the
vapor pressure, larger fraction of molecules
have sufficient kinetic energy to escape
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Evaporation, Vapor Pressure, and Boiling Point
The Clausius-Clapeyron Equation
ln Pvap = - Hvap 1 + C
R
T
Y = m
x +b
m is the slope and b is the y-intercept
Where R is the gas constant 8.314 J/K mol
C is a constant characteristic of each specific
substance
T temperature in Kelvin
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Evaporation, Vapor Pressure, and Boiling Point
The Clausius-Clapeyron Equation
ln P2 = Hvap 1 - 1
P1
R T1
T2
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Evaporation, Vapor Pressure, and Boiling Point
• This equation makes it possible to calculate the heat of
vaporization of a liquid by measuring its vapor pressure
at several temperatures and then plotting the results to
obtain the slope
• Once the heat of vaporization and the vapor pressure at
one temperature are known, the vapor pressure of the
liquid at any other temperature can be calculated.
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Evaporation, Vapor Pressure, and Boiling Point
Normal boiling point
1. The temperature at which boiling occurs
when the pressure is exactly 1 atm.
2. Boiling point – when the vapor pressure of a
liquid rises to the point where it becomes equal
to the external pressure
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Example 7
The vapor pressure of ethanol at 34.7C is 100.0 mm Hg,
and the heat of vaporization of ethanol is 38.6 kJ/mol.
What is the vapor pressure of ethanol in mm Hg at
65.0C?
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Example 8
The normal boiling point of benzene is 80.1 C and the
heat of vaporization is 30.8 kJ/mol. What is the boiling
point of benzene (in C) on top of Mt. Everest where P
= 260 mm Hg?
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Phase Diagrams
• Shows which phase is stable at different
combinations of pressure and temperature.
39
Phase Diagrams
Triple Point: The only condition under which all three
phases can be in equilibrium with one another.
Critical Temperature (Tc): The temperature above which
the gas phase cannot be made to liquefy at any pressure.
Critical Pressure (Pc) : The minimum pressure required to
liquefy a gas at its critical temp.
Supercritical Fluid: Neither true liquid nor true gas
Normal boiling and melting point always at 1 atm
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Example 9
Can you label the following?
a) solid region
b) Liquid region
c) Gas region
d) Normal boiling point
e) Normal melting point
f) Triple point
g) Supercritical fluid region
h) Critical point, what is the
critical pressure and
temperature
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Types of solids
1. Molecular solid
-held together by intermolecular forces
-H2O(s), CO2(s)
2. Metallic solid
-positively charged atomic cores
surrounded by delocalized electrons
-Fe, Cu, Ag
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Types of solids
3. Ionic solid
-cations and anions held together by
electrical attraction of opposite charges
-NaCl
4. Covalent network solid
-atoms held together in large networks by
covalent bonds
-diamonds
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