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Homework Problems
Chapter 11 Homework Problems: 2, 14, 16, 18, 22, 28, 32, 36, 46, 56,
60, 64, 70, 75, 84, 103
CHAPTER 11
Intermolecular Forces and the Physical Properties
of Liquids and Solids
Solids, Liquids, Gases
The three common states of matter are solids, liquids, and gases.
Their properties can be summarized as follows
state
properties
solid
definite volume and definite shape
liquid
definite volume but indefinite shape
gas
indefinite volume and indefinite shape
Compressibility
Compressibility refers to the decrease in the volume occupied by
a substance when the pressure applied on the substance increases.
Because the particles making up solids and liquids are in close contact
with one another, these two phases are not easily compressed (to a first
approximation solids and liquids are incompressible). Since gases are
mostly empty space, they are highly compressible. This also means that
solids and liquids will be much higher density than gases.
Effect of Temperature and Pressure on Phase
The phase of a substance depends on both temperature and
pressure. Generally speaking, substances go from solid to liquid to gas
as temperature increases.
It is often possible to condense a gas into a liquid or solid by
increasing pressure. We will discuss this further later in the chapter.
Intermolecular Forces
Intermolecular forces are the forces that exists between different
molecules or particles. We are more concerned with long range
attractive forces and will ignore short range repulsive forces.
Ion-ion - The attractive force acting between cations and anions.
These are strong, and are found in substances where ionic bonding
occurs..
Dipole-Dipole Forces
Dipole-dipole - The attractive force acting between polar
molecules. The attraction is between the partial positive charge (+) on
one molecule and the partial negative charge (-) on a different molecule.
Generally speaking, the larger the partial positive and negative charges
the stronger the dipole-dipole attraction.
Dipole-Dipole Forces and Boiling Point
When molecules have strong intermolecular attractive forces it
takes more energy to overcome those attractive forces. One way of
seeing this is in the boiling point for a substance. Generally speaking,
the stronger the dipole-dipole attraction between molecules the higher
the boiling point, particularly for substances
with approximately the
same molecular mass.
Hydrogen Bonding
Hydrogen bonding - A particularly strong form of dipole-dipole
attractive force. It is the attractive force that exists between a hydrogen
atom bonded to an N, O, or F atom and lone pair electrons on a different
N, O, or F atom (often an atom in a different molecule).
Evidence For Hydrogen Bonding
One effect of hydrogen bonding is to raise the boiling point of a
liquid. This occurs because it requires more energy (and so a higher
temperature) to break apart strong attractive forces between molecules
than it does to break apart weak attractive forces between molecules.
substance
boiling point
hydrogen bonding?
H2O
100.0 C
yes
H2S
- 60.7 C
no
H2Se
- 41.5 C
no
H2Te
- 4.4 C
no
Boiling Points for Binary Hydrogen Compounds
London Dispersion Forces
London dispersion forces - The attractive force that is due to the
formation of instantaneous dipoles in a molecule. These instantaneous
dipoles arise from the random motion of the electrons in the molecule.
Strength of London Dispersion Forces
London dispersion forces are present in all molecules, but are the
only intermolecular force present in nonpolar molecules. The strength of
London dispersion forces is approximately proportional to the number of
electrons, and so to the size of the molecule. Therefore, as a general rule,
the larger the molecule the stronger the London dispersion forces.
substance boiling point
He
- 268.6 C
Ne
- 245.9 C
Ar
- 185.7 C
Kr
- 152.3 C
Xe
- 107.1 C
Induced Dipole and Polarizability
London dispersion forces are closely related to a second property
of molecules, called polarizability. The polarizability of a molecule
refers to the extent to which the electron cloud distribution is distorted
when a positive or negative charge is brought close to the molecule. In
general, large molecules, and molecules containing atoms whose valence
electrons are far away from the atomic nucleus are more polarizable than
other molecules. The more polarizable a molecule the stronger the
London dispersion forces it experiences.
Ion-Dipole Forces
Ion-dipole - The attractive force between an ion and a polar
molecule. Responsible for the dissolution of some ionic substances in
polar liquids such as water. In general, the smaller the ion and the larger
the charge the stronger the ion-dipole attractive force.
Solvation - The close association of
solvent molecules with solute molecules
or ions.
Hydration - Solvation when the solvent
is water.
Summary of the Types of Intermolecular Forces
Ion – ion. Forces between cations and anions.
Dipole – dipole. Forces between molecules with a permanent
dipole moment. This category includes hydrogen bonding, a particularly
strong type of dipole – dipole force.
London dispersion forces. Due to random movement of
electrons. All particles have this type of force, but it is most important in
molecules with no permanent dipole moment.
Mixed forces: Ion – dipole is the most important, but this also
includes ion – induced dipole and dipole – induced dipole. It is
responsible for the solubility of some ionic compounds in water. We will
discuss these and related forces in detail when discussing solution
formation.
Properties of Liquids
There are several properties of liquids that are related to the
strength of the intermolecular forces acting between molecules.
Surface tension - Resistance of a liquid to spreading out.
Viscosity - Resistance of a liquid to flow.
Generally speaking, the stronger the intermolecular forces, the
larger the values for viscosity and surface tension.
Surface Tension
Surface tension is the resistance of a liquid to spreading out.
A molecule within the liquid feels forces from all of the
surrounding molecules. However, at the surface of the liquid a molecule
is attracted back to the liquid by the molecules below. This attractive
force tends to cause liquids to arrange themselves to minimize their
surface area (the reason small droplets of water form beads.
Cohesion and Adhesion
Cohesion refers to the attractive force between molecules of the
same type. A second term, adhesion, refers to the attractive force
between different types of molecules, such as between molecules of a
liquid and those making up the surface of the container. In a narrow tube
the competition between these forces affects the shape of the liquid
meniscus, the curved surface of the liquid.
Capillary Action
When adhesive forces are stronger than cohesive forces, a liquid
will be drawn up into a narrow diameter tube, by a process called
capillary action.
Note that the narrower the diameter of the tube the higher the level of
the liquid drawn into the tube. For
liquids like mercury, where adhesive
forces are larger than cohesive forces, a
reverse process occurs.
Viscosity
Viscosity is the resistance of a liquid to flow. Liquids with strong
intermolecular attractive forces will flow more slowly than those where
such forces are small.
The device at right is a viscometer. It
can be used to measure the viscosity of a liquid.
Vapor Pressure
Vapor pressure is defined as the equilibrium partial pressure of
vapor above a pure solid or gas. The vapor pressure of substance
increases as temperature increases.
The normal boiling point for a substance corresponds to the
temperature at which the liquid and vapor pressures are at equilibrium
and the vapor pressure is 1.00 atm.
Clausius-Clapeyron Equation
Based on experiment, the vapor pressure above a liquid is found
to obey a simple equation, called the Clausius-Clapeyron equation
ln(p) = - Hvap + C
RT
Based on this equation we expect a plot of ln(p) vs 1/T to give a straight
line with slope m = - Hvap/R.
The above equation may be used to find a second useful
expression
ln(p2/p1) = - (Hvap/R) [ (1/T2) – (1/T1) ]
where p1 is the vapor pressure at temperature T1
p2 is the vapor pressure at temperature T2
Clausius-Clapeyron Equation (Example)
slope = - Hvap/R
Example: The normal boiling point for water occurs at T = 100.0 C.
The enthalpy of vaporization for water is Hvap = 40.67 kJ/mol. Based
on this information estimate the vapor pressure of water at T = 20.0 C.
Example: The normal boiling point for water occurs at T = 100.0 C.
The enthalpy of vaporization for water is Hvap = 40.67 kJ/mol. Based
on this information estimate the vapor pressure of water at T = 20.0 C.
Recall that one form of the Clausius-Clapyron equation is:
ln(p2/p1)= - Hvap
R
1 _ 1
T2 T1
Example: The normal boiling point for water occurs at T = 100.0 C.
The enthalpy of vaporization for water is Hvap = 40.67 kJ/mol. Based
on this information estimate the vapor pressure of water at T = 20.0 C.
Let
T1 = 100.0 C = 373. K ; p1 = 1.00 atm
T2 = 20.0 C = 293. K
ln(p2/p1) = - 40670. J/mol
(8.314 J/mol.K)
1
1
293. K
373. K
= - 3.581
So (p2/p1) = e-3.581 = 0.0279
p2 = (0.0279) p1 = (0.0279) (1.00 atm) = 0.0279 atm = 21. torr
Solids
Solids can be divided into two general categories
Crystalline solid - Has a regular arrangement of the particles
making up the solid (a crystal structure). Four main types exist: ionic,
molecular, covalent, and metallic solids.
Amorphous solid Does not have a regular
arrangement of the particles making up the solid
(no regular crystal structure).
Crystal Structure
For crystalline solids, the crystal structure of a solid substance
refers to the arrangement of the particles making up the solid. This is
often given in terms of the unit cell, the smallest part of the crystal that
can be used to construct the entire crystal.
Examples of Crystal Structures
Examples of different types of crystal structures are given below.
Note that information about crystal structures is determined using
techniques such as X-ray diffraction or electron diffraction.
Types of Crystalline Solids
Crystalline solids can be classified into four groups.
1) Ionic – Composed of cations and anions; held together by ionic
bonding.
2) Covalent – Composed of atoms where every atom is attached
to other atoms in the solid through a network of covalent bonds.
3) Molecular – Composed of molecules; held together by weak
intermolecular forces (dipole-dipole, hydrogen bonding, or dispersion)
4) Metallic – Composed of metal atoms, held together by metallic
bonding.
Ionic Solids
An ionic solid is composed of cations and anions. Ionic solids
are held together by the strong electrostatic force of attraction that exists
between particles of opposite charge. Examples: NaCl, CaCO3, Al2N3,
FeCl3.
Properties
Hard and brittle
High melting point
High boiling point
Poor conductors of heat and
electricity in solid state
Good conductors of electricity when dissolved in water
NaCl, Tfus = 801 C, Tvap = 1413 C
Covalent Solids
A covalent solid (sometimes called a network covalent solid) is a
“supermolecule” in which every atom is connected to every other atom
through a network of covalent bonds. Because all of the atoms are
connected by covalent bonds, these substances are generally among the
hardest substances known. Exam-ples: C (diamond), C(graphite), SiO2
(quartz).
Properties
Extremely hard
Very high melting point
Very high boiling point
Poor conductors of heat and
electricity in solid state
Insoluble in water
diamond
C , Tfus > 3550 C, Tvap = 4827 C
Molecular Solids
A molecular solid is composed of molecules. Molecular solids
are held together by the weak van der Waals attractive forces (dipoledipole and London dispersion forces) that exist between molecules.
Examples: H2O, Ar, CS2, C10H8 (naphthalene), C6H12O6 (sugar).
Properties
Soft
Low melting point
Low boiling point
Poor conductors of heat and
electricity in solid state
Poor conductors of electriCS2, Tfus = - 111 C, Tvap = 46 C
city when dissolved in water
Metallic Solids
A metallic atomic solid represents the solid form for metals.
Metallic solids can be thought of as metal cations immersed in a sea of
loosely held valence electrons. Examples: K, Fe, Cu, Na, Pb.
Properties
Can be hard or soft
Low to high melting point
Low to high boiling point
Good conductors of heat
and electricity in solid state
Insoluble in water
Na, Tfus = 98 C, Tvap = 892 C
Phase Transitions
The conversion of a substance from one phase to another phase is
called a phase transition. Transitions can be caused both by adding heat
and by removing heat from a substance.
adding heat (H > 0)
removing heat (H < 0)
s   fusion (melting)
  s freezing
  g vaporization
g   condensation
s  g sublimation
g  s deposition
Recall that the enthalpy change for a phase transition is usually
reported at the normal transition temperature, that is, the temperature at
which the phase transition occurs when p = 1.00 atm.
Since enthalpy is a state function:
Hfreez = - Hfus
Hcond = - Hvap
Hdep = - Hsub
Relationships Among Phase Transitions
Based on Hess’ law, we would expect
Hsub = Hfus + Hvap
Thermodynamics of Phase Transitions
We can study the thermodynamics of phase transitions by finding
the heating curve for a substance. This is simply a plot of temperature
vs. amount of heat added, under conditions where the heat is added
slowly enough to maintain equilibrium.
Experimentally we expect to see two regions in the
heating curve. Normally the
temperature of the substance
will increase as heat is added.
However, at the temperature
where a phase transition
occurs the added heat will be
used to carry out the transition, and so temperature will
remain constant until the
phase transition is complete.
Sample Heating Curve
Phase Diagram
A phase diagram is a diagram indicating which phase or phases
are present at equilibrium as a function of pressure and temperature.
H2O
Important Features in a Phase Diagram
Phase boundaries - Indicate where two phases can exist simultaneously at equilibrium.
Triple point - Indicates a point where three phases can exist
simultaneously at equilibrium.
Normal melting point - Solid-liquid equilibrium at p = 1.00 atm.
Normal boiling point - Liquid-gas equilibrium at p = 1.00 atm.
Normal sublimation point - Solid-gas equilibrium at p = 1.00 atm.
(Note that substances will have either a normal melting and
normal boiling point, or a normal sublimation point, but not both.)
Critical point - Point below which a gas will undergo a phase
transition (g   or g  s) when compressed reversibly at constant
temperature. Above the critical point no such phase transition occurs. In
this region of the phase diagram a supercritical fluid is present.
Phase Diagram For CO2
CO2
End of Chapter 11
“Gibbs is perhaps the most brilliant person most people have
never heard of. Modest to the point of near-invisibility, he passed
virtually the whole of his life, apart from three years spent studying in
Europe, within a three-block area bounded by his house and the Yale
campus in New Haven, Connecticut. For his first ten years at Yale he
didn't even bother to draw a salary. (He had independent means.) From
1871, when he joined the university as a professor, to his death in 1903,
his courses attracted an average of slightly over one student a semester.”
- Bill Bryson A Short History of Nearly Everything
“Of all chemical bonds, hydrogen bonds are the weakest, the
most important, the least understood, and the hardest to measure.”
- John Emsley, “Science Watch” (2000)