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Homework Problems
Chapter 11 Homework Problems: 4, 14, 25, 34, 50, 52, 56, 59, 68, 74,
76, 78, 83, 84, 87, 104, 106, 108, 118
Skip sections 11-10, 11-11, and 11-13.
CHAPTER 11
Liquids, Solids, and Intermolecular Forces
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.
Dipole Moment and Miscibility
Miscibility refers to the ability of one liquid to mix with another
without forming two separate liquid phases. As a general rule (for
reasons discussed in Chapter 12) polar liquids will mix well with polar
liquids, and nonpolar liquids with nonpolar liquids. A polar and a
nonpolar liquid will generally not mix to form a homogeneous solution.
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.
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
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. 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
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.
Solvation - The close association of solvent molecules
with solute molecules or
ions.
Hydration - Solvation when
the solvent is water.
To summarize the types of 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.
responsible for the solubility of some ionic compounds in water.
It is
Properties of Liquids
There are several general properties of liquids.
Viscosity - Resistance of a liquid to flow.
Surface tension - Resistance of a liquid to spreading out.
Generally speaking, the stronger the intermolecular forces, the
larger the values for viscosity and surface tension.
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
Enthalpy and Entropy Changes For Phase Transitions
Note
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
Vapor Pressure
The vapor pressure (pvap) of a liquid or solid is equilibrium partial
pressure of the substance in the gas phase above the liquid or solid.
Experimentally, it is found that
the vapor pressure of a liquid increases
as temperature increases. This is not
surprising, as the molecules in the liquid
phase have a higher average energy as
temperature increases.
Clausius-Clapeyron Equation
The dependence of vapor pressure on temperature is well
described by the Clausius-Clapeyron equation (which may be derived).
ln(pvap) = - Hvap + constant
RT
This may be rearranged to give
ln(p2/p1)= - Hvap 1
R
1
T2 T1
p1 is the vapor pressure at T1
p2 is the vapor pressure at T2
slope = - Hvap/R
The Clausius-Clapeyron equation assumes that the value for
Hvap is constant. In fact, the value decreases slightly with increasing
temperature.
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
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
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).
Classification of Crystalline Solids
Crystalline solids can be classified into three groups.
1) Ionic solids – Composed of cations and anions; held together
by ionic bonding
2) Molecular solids – Composed of molecules; held together by
weak intermolecular forces (dipole-dipole, hydrogen bonding, or
dispersion)
3) Atomic solids – Composed of atoms
a) Nonbonding – Composed of noble gas atoms. Held together
by dispersion forces,
b) Metallic – Composed of metal atoms, held together by
metallic bonding.
c) Network covalent – Composed of atoms where every atom
is attached to other atoms in the solid through a network of covalent
bonds.
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
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
Atomic Solids
An atomic solid is a solid substance composed of individual
atoms (as opposed to ions or molecules).
There are three types of atomic solids
Nonbonding atomic solid – An atomic solid where the atoms are
held together by dispersion forces. The only nonbonding atomic solids
are the solid forms of the noble gases (He, Ne, Ar, …)
Metallic atomic solid – The solid form of a metal. Consists of
individual metal atoms held together by a “sea” of electrons in metallic
bonding.
Network covalent atomic solid – A solid where every atom is
connected to every other atom by a network of covalent bonds. In a
sense a network covalent atomic solid is a single “supermolecule”
diamond
Nonbonding Atomic Solid
The only nonbonding atomic solids are the solid forms of the
noble gases. Since these substances do not form covalent bonds with
other noble gase atoms the forces holding nonbonding atomic solids
together are weak dispersion forces.
Properties
Extremely
low
melting
point
Extremely low boiling point
Poor conductors of heat and
electricity in solid state
Kr, Tfus = - 157. C, Tvap = - 152. C
Metallic Atomic 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
Network Covalent Atomic Solids
A network covalent atomic 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
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)